Moderné vzdelávanie pre vedomostnú spoločnosť/
Projekt je spolufinancovaný zo zdrojov EÚ
CHEMISTRY
Stavebná fakulta TU v Košiciach
doc. RNDr. Adriana Eštoková, PhD., prof. RNDr. Magdaléna Bálintová, PhD.
Táto publikácia vznikla za finančnej podpory z Európskeho sociálneho fondu
v rámci Operačného programu VZDELÁVANIE.
Prioritná os 1 Reforma vzdelávania a odbornej prípravy
Opatrenie 1.2 Vysoké školy a výskum a vývoj ako motory rozvoja
vedomostnej spoločnosti.
Názov projektu: Balík prvkov pre skvalitnenie a inováciu vzdelávania na
TUKE
ITMS 26110230070
Názov: Chemistry
Autori: doc. RNDr. Adriana Eštoková, PhD., prof. RNDr. Magdaléna Bálintová, PhD.
Vydavateľ: Technická univerzita v Košiciach
Rok: 2015
Vydanie: prvé
Počet výtlačkov: 25
Rozsah: 92 strán
Rukopis neprešiel jazykovou úpravou.
Za odbornú a obsahovú stránku zodpovedajú autori.
Content
1
ATOM STRUCTURE............................................................................................................... 6
1.1 Elemental particles ................................................................................................................. 6
1.2 Atomic structure models ......................................................................................................... 8
2
1.2.1
Bohr model ................................................................................................................ 9
1.2.2
Quantum mechanical model ................................................................................... 10
1.2.3
The electron configuration ...................................................................................... 12
CHEMICAL BONDING ......................................................................................................... 15
2.1 Covalent bond ...................................................................................................................... 17
2.1.1
Single and multiple covalent bonds ........................................................................ 17
2.1.2
Sigma and pi bonds ................................................................................................ 18
2.1.3
Polar and non-polar covalent bonds ....................................................................... 20
2.2 Metallic bonding ................................................................................................................... 20
2.2.1
Band theory ............................................................................................................. 21
2.3 Ionic bond ............................................................................................................................. 22
2.4 Van der Waals bond ............................................................................................................. 23
2.5 Hydrogen bonding ................................................................................................................ 24
3
STATE OF MATTER............................................................................................................. 25
3.1 Liquids .................................................................................................................................. 25
3.2 Solids .................................................................................................................................... 26
3.3 Gas27
3.4 Plasma ................................................................................................................................. 27
3.5 Phase changes..................................................................................................................... 28
3.5.1
Condensation and vaporisation .............................................................................. 28
3.5.2
Melting and freezing................................................................................................ 29
3.5.3
Sublimation ............................................................................................................. 30
3.6 Phase diagram ..................................................................................................................... 30
4
CHEMICAL REACTION ....................................................................................................... 32
4.1 Clasification of chemical reactions ....................................................................................... 32
4.1.1
Chemical reactions according to number of reaction phases ................................. 32
4.1.2
Chemical reactions according to type of changes .................................................. 32
4.2 Reaction mechanism ............................................................................................................ 33
4.2.1
Protolytical reakction ............................................................................................... 33
4.2.2
Oxidation-reduction reakction ................................................................................. 36
4.2.3
Precipitation reaction .............................................................................................. 37
4.3 Chemical kinetics ................................................................................................................. 37
4.3.1
5
Reaction rate ........................................................................................................... 37
WATER ................................................................................................................................. 42
5.1 Water structure ..................................................................................................................... 42
3
5.1.1
Hydrogen bonds in water ........................................................................................ 43
5.2 Water solutions..................................................................................................................... 44
5.2.1
Solubility .................................................................................................................. 44
5.2.2
Water hardness ....................................................................................................... 46
5.3 Autoionization of water and pH scale ................................................................................... 47
5.3.1
6
pH scale .................................................................................................................. 48
Heterogenous SYSTEMS .................................................................................................... 49
6.1 Disperse systems ................................................................................................................. 49
6.1.1
Molecular dispersion ............................................................................................... 50
6.1.2
Colloids ................................................................................................................... 50
6.1.3
Coarse dispersions ................................................................................................. 51
6.2 Adsorption ............................................................................................................................ 51
7
6.2.1
Adsorption at the gas/solid interface ....................................................................... 52
6.2.2
Interaction at the liquid /solid interface ................................................................... 54
BUILDING MATERIALS ....................................................................................................... 58
7.1 The important compounds used as building materials......................................................... 58
7.1.1
Silicon dioxide ......................................................................................................... 58
7.1.2
Silicates and Silicones; Polymeric Network Structures .......................................... 59
7.1.3
Carbonates and Hydrogen Carbonates .................................................................. 60
7.1.4
Calcium oxide - Lime .............................................................................................. 61
7.1.5
Calcium hydroxide .................................................................................................. 62
7.1.6
Magnesium oxide .................................................................................................... 62
7.1.7
Sulphate .................................................................................................................. 63
7.1.8
Aluminium oxide...................................................................................................... 64
7.1.9
Ferric oxide ............................................................................................................. 64
7.2 Portland cement – Manufacture, types properties, and specifications ................................ 64
8
7.2.1
Manufacture ............................................................................................................ 65
7.2.2
Composition of portland cement clinker.................................................................. 65
7.2.3
Technically important properties of portland cement .............................................. 68
7.2.4
Hydration of portland cement .................................................................................. 69
CONCRETE DETERIORATION ........................................................................................... 72
8.1 Concrete durability ............................................................................................................... 72
8.2 Concrete corrosion ............................................................................................................... 72
8.2.1
Leaching (dissolution of concrete) .......................................................................... 72
8.2.2
Acid attack .............................................................................................................. 73
8.2.3
Carbonation of concrete ......................................................................................... 73
8.2.4
Corrosion of concrete reinforcement ...................................................................... 73
8.2.5
Sulphate attack ....................................................................................................... 74
8.2.6
Freeze-thaw damage .............................................................................................. 76
8.2.7
Fire-resistance of concrete ..................................................................................... 76
4
9
METALS ................................................................................................................................ 78
9.1 Properties of metals ............................................................................................................. 79
9.2 Standard potential of metals ................................................................................................ 80
9.3 Metal production ................................................................................................................... 82
9.4 Iron 84
9.5 Non-ferrous metals ............................................................................................................... 85
9.5.1
Copper .................................................................................................................... 85
9.5.2
Aluminum ................................................................................................................ 86
9.5.3
Zinc ......................................................................................................................... 86
10 CORROSION OF METALS .................................................................................................. 88
10.1 Definition of corrosion ..................................................................................................... 88
10.2 Classification of corrosion ............................................................................................... 89
10.2.1 Dry corrosion ........................................................................................................... 89
10.2.2 Wet corrosion .......................................................................................................... 90
10.3 The factors affecting the rate of corrosion ...................................................................... 91
5
1 ATOM STRUCTURE
Chapter mission
This chapter presents an overview of knowledge about basic elemental particles in an atom.
Special attention is paid to the models representing the atom structure and to the electron
configuration.
Chapter objectives
After studying this chapter you should know to:

describe the basic elemental particles of atom,

explain current understanding of the atom structure,

understand the dependence between the electron configuration and chemical properties of
atoms.
1.1 Elemental particles
The atom is the smallest unit that defines the chemical elements and their isotopes. Every material
object or substance is made up of atoms. Everything that is solid, liquid, or gas is made up of atoms.
Although the word 'atom' comes from the Greek for indivisible, we now know that atoms are not the
smallest particles of matter. Atoms are made from smaller subatomic particles. Atoms consist of
electrons surrounding a nucleus.
Fig. 1.1 Illustration of the atomic nucleus and electrons surrounding
The nucleus is at the centre of the atom and contains the protons and neutrons. Protons and neutrons
are collectively known as nucleons. Neutrons are neutral, but protons and electrons are electrically
charged. Protons have a relative charge of +1, while electrons have a relative charge of -1.
The number of protons in an atom is called its atomic number. In the periodic table atoms are
arranged in atomic number order. The atomic number is also given the more descriptive name of
proton number. Number of protons + number of neutrons gives a mass number of the atom also called
the nucleon number. This information can be given simply in the form:
6
Protons and neutrons are responsible for most of the atomic mass. The mass of an electron is very
small (9.108 X 10
-28
grams). So, the atomic mass is the sum of the masses of protons and neutrons.
relative mass
relative charge
proton
1
+1
neutron
1
0
electron
1/1836
-1
It is the number of protons that determines the atomic number, e.g., H = 1. The number of protons in
an element is constant (e.g., H=1, Ur=92) but neutron number may vary, so mass number (protons +
neutrons) may vary. The number of neutrons in an atom can vary within small limits. For example,
there are three kinds of carbon atom
12
C,
13
C and
14
C. They all have the same number of protons, but
the number of neutrons varies.
protons
neutrons
mass number
carbon-12
6
6
12
carbon-13
6
7
13
carbon-14
6
8
14
These different atoms of carbon are called isotopes. The fact that they have varying numbers of
neutrons makes no difference whatsoever to the chemical reactions of the carbon. Isotopes are atoms
which have the same atomic number but different mass numbers.
The chemical properties of isotopes are the same, although the physical properties of some isotopes
may be different. Some isotopes are radioactive-meaning they "radiate" energy as they decay to a
more stable form, perhaps another element half-life: time required for half of the atoms of an element
to decay into stable form.
Fig. 1.2 Isotopes of carbon
Another example is hydrogen; with atomic number of 1 can have 0, 1, or 2 neutrons (Fig.1.3).
7
Fig. 1.3 Isotopes of hydrogen
Atoms are electrically neutral, and the positiveness of the protons is balanced by the negativeness of
the electrons. It follows that in a neutral atom:
number of electrons = number of protons
So, if an oxygen atom (atomic number = 8) has 8 protons, it must also have 8 electrons; if a chlorine
atom (atomic number = 17) has 17 protons, it must also have 17 electrons.
Electrons are arranged in energy levels or shells, and different energy levels can hold different
numbers of electrons. The electronic structure of an atom is a description of how the electrons are
arranged, which can be shown in a diagram or by numbers. There is a link between the position of an
element in the periodic table and its electronic structure.
1.2 Atomic structure models
The knowledge of structure of atoms has been developed over a long period of time (Fig. 1.4).
Fig. 1.4 Evolution of the atom structure models
8
There are two models of atomic structure in use today:

the Bohr model and

the quantum mechanical model.
Of these two models, the Bohr model is simpler and relatively easy to understand.
1.2.1 Bohr model
The Bohr Model is a planetary model in which the negatively-charged electrons orbit a small,
positively-charged nucleus similar to the planets orbiting the Sun (except that the orbits are not
planar). The gravitational force of the solar system is mathematically similar to the Coulomb (electrical)
force between the positively-charged nucleus and the negatively-charged electrons. The Bohr model
works well for very simple atoms such as hydrogen (which has 1 electron) but not for more complex
atoms.
Fig. 1.5 A planetary model of atom
1. The Bohr model shows that the electrons in atoms are in orbits of differing energy around the
nucleus (think of planets orbiting around the sun).
2. Bohr used the term energy levels (or shells) to describe these orbits of differing energy. He said
that the energy of an electron is quantized, meaning electrons can have one energy level or
another but nothing in between.
3. The energy level an electron normally occupies is called its ground state. But it can move to a
higher-energy, less-stable level, or shell, by absorbing energy. This higher-energy, less-stable
state is called the electron’s excited state.
4. After it’s done being excited, the electron can return to its original ground state by releasing the
energy it has absorbed, as shown in the diagram below.
9
Fig. 1.6 Return of electron from the excited state to its ground state
1.2.2 Quantum mechanical model
The quantum mechanical model is based on quantum theory, which says matter also has properties
associated with waves. According to quantum theory, it’s impossible to know the exact position and
momentum of an electron at the same time. This is known as the Uncertainty Principle. The quantum
mechanical model of the atom uses complex shapes of orbitals (sometimes called electron clouds),
volumes of space in which there is likely to be an electron. So, this model is based on probability
rather than certainty. Four numbers, called quantum numbers, were introduced to describe the
characteristics of electrons and their orbitals:

Principal quantum number: n

Angular momentum quantum number: l

Magnetic quantum number: m

Spin quantum number: s
The principal quantum number n describes the average distance of the orbital from the nucleus —
and the energy of the electron in an atom. It can have positive integer (whole number) values: 1, 2, 3,
4, and so on. The larger is the value of n the higher is the energy and the larger the orbital. Chemists
sometimes call the orbitals electron shells. The electron shells are labelled K, L, M, N, O, P, and Q;
going from innermost shell outwards (Fig. 1.7). These letters correspond to the n-values 1, 2, 3, 4, 5,
6, and 7. Electrons in outer shells have higher average energy and travel farther from the nucleus than
those in inner shells.
Fig. 1.7 The electron shells
10
The angular momentum quantum number “l” describes the shape of the orbital, and the shape is
limited by the principal quantum number n: The angular momentum quantum number l can have
positive integer values from 0 to n–1. For example, if the n value is 3, three values are allowed for l: 0,
1, and 2. Orbitals that have the same value of n but different values of l are called subshells. These
subshells are given different letters to help our distinguish them from each other. The following table
shows the letters corresponding to the different values of l.
Value of l
Letter
0
s
1
p
2
d
3
f
4
g
The following figure shows the shapes of the s, p, and d orbitals.
Fig. 1.8 Shape of the s, p, and d orbitals
There are two s orbitals — one for energy level 1 (1s) and the other for energy level 2 (2s). The s
orbitals are spherical with the nucleus at the center. Notice that the 2s orbital is larger in diameter than
the 1s orbital. In large atoms, the 1s orbital is nestled inside the 2s, just like the 2p is nestled inside the
3p.
The magnetic quantum number describes how the various orbitals are oriented in space. The value
of this number depends on the value of l. The values allowed are integers from –l to 0 to +l. For
example, if the value of l = 1 (p orbital), you can write three values for this number: –1, 0, and +1. This
means that there are three different p subshells for a particular orbital. The subshells have the same
energy but different orientations in space.
The second row of the figure (Fig. 1.9) shows how the p orbitals are oriented in space. Notice that the
three p orbitals correspond to magnetic quantum number values of –1, 0, and +1, oriented along the x,
11
y, and z axes. The third row of the figure (Fig. 1.9) illustrates five orientations of d orbitals in space and
in the fourth raw there are seven d orbitals orientations.
Fig. 1.9 Orientation of the s, p, and d orbitals in space
The spin quantum number describes the direction of the electron spinning in a magnetic field — either
clockwise or counterclockwise. Only two values are allowed: +1/2 or –1/2. For each subshell, there
can be only two electrons, one with a spin of +1/2 and another with a spin of –1/2.
1.2.3 The electron configuration
The electron configuration of an atom is the representation of the arrangement of electrons that are
distributed among the orbital shells and subshells. Commonly, the electron configuration is used to
describe the orbitals of an atom in its ground state, but it can also be used to represent an atom that
has ionized into a cation or anion by compensating with the loss of or gain of electrons in their
subsequent orbitals.
There are three rules for assigning electron orbitals:
1. Aufbau process: Understand how the energy level vary is the key to the Aufbau process, because
electrons tend to occupy the lowest energy level available.
2. The Pauli Exclusion Principle: This principle suggests that only two electrons with opposite spin
can occupy an atomic orbital. Stated another way, no two electrons have the same 4 quantum
numbers n, l, m, s. Pauli's exclusion principle can be stated in some other ways, but the idea is
12
that energy states have limit room to accommodate electrons. A state accepts two electrons of
different spins.
3. Hund's rule suggests that electrons prefer parallel spins in separate orbitals of subshells. This rule
guides us in assigning electrons to different states in each sub-shell of the atmic orbitals. In other
words, electrons fill each and all orbitals in the subshell before they pair up with opposite spins.
Many of the physical and chemical properties of elements can be correlated to their unique electron
configurations. The valence electrons, electrons on the outer most shell, become the determining
factor for the unique chemistry of the element. All atoms would like to attain electron configurations
like noble gases. That is, have completed outer shells. Atoms can form stable electron configurations
like noble gases by:

losing electrons

sharing electrons

gaining electrons.
For a stable configuration each atom must fill its outer energy level. In the case of noble gases that
means eight electrons in the last shell (with the exception of He which has two electrons). Atoms that
have 1, 2 or 3 electrons in their outer levels will tend to lose them in interactions with atoms that have
5, 6 or 7 electrons in their outer levels. Atoms that have 5, 6 or 7 electrons in their outer levels will tend
to gain electrons from atoms with 1, 2 or 3 electrons in their outer levels. Atoms that have 4 electrons
in the outer most energy level will tend neither to neither totally lose nor totally gain electrons during
interactions.
The electron configuration of each element is unique to its position on the periodic table (Fig. 1.10).
The energy level is determined by the period and amount of electrons by the atomic number of the
element. Orbitals on different energy levels are similar to each other, but they occupy different areas in
space. The 1s orbital and 2s orbital both have the characteristics of an s orbital (radial nodes,
spherical volume probabilities, can only hold two electrons, etc.) but as they are found in different
energy levels they occupy different spaces around the nucleus. Each orbital can be represented by
specific blocks on the periodic table.
The s-block is the region of the Alkali metals including Helium (groups 1 & 2), the d-block is the
Transition metals (groups 3 to 12), the p-block are the main group elements from group 13 to 18, and
the f-block are the Lanthanides and Actinides series.
13
Fig. 1.10 Electron configurations in the periodic table
Questions:
1. Characterize the major elemental particles of an atom.
2. Name the quantum numbers.
3. What determine the quantum numbers?
4. What kinds of atomic orbital do you know?
5. Write the arrangement of elements in the periodical table (metals, non-metals, metalloids).
References:
1. El Sair R.R.: Fundamentals of Chemistry. 2012, Romain Elsair & Ventus Publishing ApS, ISBN:
978-87-403-0105-2.
2. Nelson P.G.: Introduction to Inorganic Chemistry. 2011, Peter G. Nelson & Ventus Publishing
ApS, ISBN: 978-87-7681-732-9.
3. Beier S.P., Hede P.D Essentials of Chemistry. 2010, BookBoon, ISBN: 978-87-403-0322-3.
4. http://www.nyu.edu/pages/mathmol/textbook/atoms.html
14
2 CHEMICAL BONDING
Chapter mission
This chapter aims at an explanation of the chemical bonding principles and classification of
chemical bonds by their nature.
Chapter objectives
After studying this chapter you should know to:

classify the chemical bonds according to its nature,

explain the difference between the covalent and ionic bonding,

describe the kinds of covalent bonds,

explain a bonding among molecules and a metal bonding.
Chemical bonding is any of the interactions that account for the association of atoms into molecules,
ions, crystals, and other stable compounds. When atoms approach one another, their nuclei and
electrons interact and tend to distribute them in space in such a way that the total energy is lower than
it would be in any alternative arrangement. If the total energy of a group of atoms is lower than the
sum of the energies of the component atoms, they then bond together and the energy lowering is the
bonding energy (Fig. 2.1).
Fig. 2.1 Illustration of the bonding energy and bond length of bonding atoms
The strength of chemical bonds varies considerably; there are "strong bonds" such as covalent or ionic
bonds and "weak bonds" such as dipole–dipole interactions, the London dispersion force and
hydrogen bonding (Table 2.1).
15
Table 2.1 Examples of the bond length and bonding energy
Bond
Length (pm)
Energy (kJ/mol)
H–H
H–Cl
C–C
O=O
F–F
Hydrogen bond
…
O−H :O
74
127
154
121
142
160 to 200
200
436
432
348
498
158
5 to 30
21
According to the octet rule, atoms bond together to form molecules in such a way that each atom
participating in a chemical bond acquires an electron configuration resembling that of the noble gas
nearest it in the periodic table. Thus the outer shell of each bonded atom will contain eight electrons
(or two electrons for hydrogen and lithium).
There are two basic ways that the outer electrons of atoms can form bonds: by sharing electrons
between atoms or by transferring electrons from one atom to another. Thus, based on the principle of
formation, the chemical bonds can be divided into two groups:
1. Chemical bonds based on the electrons shared between atoms:
a. Covalent bond - electrons can be shared between neighbouring atoms.
b. Metal bond - electrons can be shared with all atoms in a material.
2. Chemical bonds based on the electrostatic attraction between two opposite charged particles
(ions, molecules):
a. Ionic bond - electrons can be transferred from one atom to another to form anion and
cation which are attracted each to other by Coulomb forces
b. Van der Waals bond – molecules are attracted each to other by Coulomb forces
c. Hydrogen bonding – special type of the dipole–dipole interactions.
There is a parameter which predicts the type of the chemical bonding between two atoms: the
electronegativity. The electronegativity is a chemical property that describes the tendency of an atom
or a functional group to attract electrons (or electron density) towards itself. Such an ability is high if
the ionization energy of the element is high (so that the atom is reluctant to give up electrons) and if its
electron affinity is also high (for then it is energetically favourable for it to acquire electrons). It follows
that atoms with high electronegativities are those in the upper right-hand corner of the periodic table,
close to fluorine (but excluding the noble gases). Such elements are likely to form anions when they
form compounds. Elements with low ionization energies (so that they readily give up electrons) and
low electron affinities (so that they have little tendency to acquire electrons) have low
electronegativities (i.e., they are electropositive) and occur at the lower left of the periodic table. Such
elements are likely to form cations during compound formation.
16
Prediction the type of the chemical bond between two atoms is based on the difference in the
electronegativity values of the atoms:

No electronegativity difference or very small difference (≤ 0.4) between two atoms leads to a
pure non-polar covalent bond.

A small electronegativity difference (≥ 0.4 but ≤ 1.7) leads to a polar covalent bond.

A large electronegativity difference (> 1.7) leads to an ionic bond.
2.1 Covalent bond
A covalent bond is a chemical bond that involves the sharing of electron pairs between atoms. The
stable balance of attractive and repulsive forces between atoms when they share electrons is known
as covalent bonding.
Hydrogen gas forms the simplest covalent bond in the diatomic hydrogen molecule. The halogens
such as chlorine also exist as diatomic gases by forming covalent bonds (Fig. 2). The nitrogen and
oxygen which makes up the bulk of the atmosphere also exhibits covalent bonding in forming diatomic
molecules.
Fig. 2.2 Covalent bonds examples
The atoms share a pair of electrons, and that pair is referred to as a bonding pair. The pairs of
electrons which do not participate in the bond have traditionally been called "lone pairs".
2.1.1 Single and multiple covalent bonds
A single bond is when two electrons (one pair of electrons) are shared between two atoms. It
represented by the two dots of the bonding pair, or by a single line which represents that pair (Fig. 3).
Although this form of bond is weaker and has a smaller density than a double bond and a triple bond,
it is the most stable because it has a lower level of reactivity meaning less vulnerability in losing
electrons to atoms that want to steal electrons.
17
Fig. 2.3 Single bond in a hydrogen molecule
A double bond is when two atoms share two pairs of electrons with each other (Fig. 4). It is depicted
by two horizontal lines between two atoms in a molecule. This type of bond is much stronger than a
single bond, but less stable; this is due to its greater amount of reactivity compared to a single bond.
Fig. 2.4 Double bond in an oxygen molecule
A triple bond is when three pairs of electrons are shared between two atoms in a molecule. It is the
least stable out of the three general types of covalent bonds.
Fig. 2.5 Triple bond in a nitrogen molecule
The single and multiple bonds differ not only by stability but by length as well (Fig. 6).
Fig. 2.6 Length of single and multiple bonds between carbon atoms
2.1.2 Sigma and pi bonds
Sigma bonds (σ bonds) are the strongest type of covalent chemical bond. They are formed by headon overlapping between atomic s and p orbitals. A σ-bond is symmetrical with respect to rotation about
the bond axis. Localization of electron density is symmetrical between the atom nuclei (Fig. 2.7).
Fig. 2.7 Localization of electron density between two atoms
18
Fig. 2.8 Sigma bonding by overlapping the s-s, s-p and p-p orbitals
Pi bond (π bond) is a covalent bond resulting from the formation of a molecular orbital by side-to-side
overlap of atomic orbitals p and d along a plane perpendicular to a line connecting the nuclei of the
atoms, denoted by the symbol π (Fig. 2.9).
Fig. 2.9 Pi bonding by overlapping the p-p, p-d and d-d orbitals
This sideways overlap also creates a molecular orbital, but of a different kind. In this one the electrons
aren't held on the line between the two nuclei, but above and below the plane of the molecule.
Localization of electron density in pi bond is above and below the plane of the molecule (Fig. 2.10).
Fig. 2.10 The sideways orbital overlap in pi bond
Multiple bonds are composed of one sigma bond together with pi or other bonds. A double bond has
one sigma plus one pi bond, and a triple bond has one sigma plus two pi bonds.
19
2.1.3 Polar and non-polar covalent bonds
Polarity refers to a separation of charge and can describe a bond or an entire molecule. Depending on
the relative electronegativities of the two atoms sharing electrons, there may be partial transfer of
electron density from one atom to the other. When the electronegativities are not equal, electrons are
not shared equally and partial ionic charges develop. Bonds connecting atoms of different
electronegativity are polar with a higher density of bonding electrons around the more electronegative
atom giving it a partial negative charge (designated as  ). The less electronegative atom has some of
-
+
its electron density taken away giving it a partial positive charge ( ).
The greater the electronegativity difference, the more ionic the bond is. Bonds that are partly ionic are
called polar covalent bonds. Nonpolar covalent bonds, with equal sharing of the bond electrons, arise
when the electronegativities of the two atoms are equal.
Nonpolar Covalent Bond

A bond between 2 nonmetal atoms that have the same electronegativity and therefore have
equal sharing of the bonding electron pair

Example: In H-H each H atom has an electronegativity value of 2.1, therefore the covalent
bond between them is considered nonpolar
Polar Covalent Bond

A bond between 2 nonmetal atoms that have different electronegativities and therefore have
unequal sharing of the bonding electron pair

Example: In H-Cl, the electronegativity of the Cl atom is 3.0, while that of the H atom is 2.1

This polarization of charge in the H-Cl bond is due to different electronegativities of chlorine
and hydrogen.
Fig. 2.11 Polarisation of the covalent bond in HCl molecule
2.2 Metallic bonding
Metallic bonding is the chemical bond characteristic of metals, in which mobile valence electrons are
shared among atoms in a usually stable crystalline structure. Metallic bonding occurs as a result of
electromagnetism and describes the electrostatic attractive force that occurs between conduction
electrons (in the form of an electron cloud of delocalize electrons) and positively charged metal ions. It
may be described as the sharing of free electrons among a lattice of positively charged ions (cations).
20
Fig. 2.12 Illustration of the valence electrons shared by metal atoms
Because the electrons move independently of the positive ions in a sea of negative charge, the metal
gains some electrical conductivity. It allows the energy to pass quickly through the electrons
generating a current. Conductivity of metals cn be explained through a band theory.
2.2.1 Band theory
Band theory, where the molecular orbitals of a solid become a series of continues energy levels, or a
band of electronic energy levels, can be used to explain the behavior or conductors, semiconductors
and insulators. Band theory models the behavior of electrons in solids by postulating the existence of
energy bands: valence and conduction bands (Figure 2.13).
Fig. 2.13 Energy bands
Conductors
In conductors, there is no band gap between their valence and conduction bands, since they overlap.
There is a continuous availability of electrons in these closely spaced orbitals.
Insulators
In insulators, the band gap between the valence band the the conduction band is so large that
electrons cannot make the energy jump from the valence band to the conduction band.
Semiconductors
21
Semiconductors have a small energy gap between the valence band and the conduction band.
Electrons can make the jump up to the conduction band, but not with the same ease as they do in
conductors.
2.3 Ionic bond
Ionic bond is formed from the electrostatic attraction between oppositely charged ions in a chemical
compound. Such a bond forms when the valence (outermost) electrons of one atom are transferred
permanently to another atom. The atom that loses the electrons becomes a positively charged ion
(cation), while the one that gains them becomes a negatively charged ion (anion).
Ionic bonding is the complete transfer of valence electron(s) between atoms. By losing those
electrons, these metals can achieve noble-gas configuration and satisfy the octet rule. Similarly,
nonmetals that have close to 8 electrons in its valence shell tend to readily accept electrons to achieve
its noble gas configuration. The electrostatic attractions between the positive and negative ions hold
the compound together. The predicted overall energy of the ionic bonding process, which includes the
ionization energy of the metal and electron affinity of the nonmetal, is usually positive, indicating that
the reaction is endothermic and unfavorable. However, this reaction is highly favorable because of
their electrostatic attraction. At the most ideal inter-atomic distance, attraction between these particles
releases enough energy to facilitate the reaction. Most ionic compounds tend to dissociate in polar
solvents because they are often polar. This phenomenon is due to the opposite charges on each ions.
Fig. 2.14 Formation of ions
The stable form of sodium chloride involves a very large number of NaCl units arranged in a lattice (or
regular arrangement) millions of atoms across. Because the lattice is rigid, this means that one gets a
solid: the ions do not move much one with respect to another. Also, because atoms are so small, even
a small crystal of salt will have billions of sodium chloride units in it. The ions are arranged so that
each positive (sodium) ion is close to many negative (chloride) ions, as shown on the following picture:
22
Fig. 2.15 Structure of NaCl
Elements in the first few columns of the periodic table have a few more electrons than predicted by the
octet rule: they therefore lose the electrons in the outermost shells fairly easily. For example, the alkali
metals (group I), such as sodium (Na) or potassium (K), which have, respectively, 11 and 19 electrons,
+
+
easily lose one electron to form monopositive ions, Na and K . These ions have 10 and 18 electrons,
respectively, so they are quite stable according to the octet rule.
Elements in the last few columns of the periodic table have one, two or three fewer electrons than
predicted by the octet rule: they therefore gain electrons fairly easily. For example, the halogens
(group VII), such as fluorine (F) or chlorine (Cl), which have, respectively, 9 and 17 electrons, easily
-
-
gain one electron to form mononegative ions, F or Cl . These ions have 10 and 18 electrons,
respectively. Likewise, elements in group II form doubly positive ions such as Mg
--
++
++
or Ca , and
--
elements in group VI form doubly negative ions such as O or S . All these ions obey the octet rule
and so are fairly stable.
At a simple level, a lot of importance is attached to the electronic structures of noble gases like neon
or argon which have eight electrons in their outer energy levels (or two in the case of helium). These
noble gas structures are thought of as being in some way a "desirable" thing for an atom to have. One
may well have been left with the strong impression that when other atoms react, they try to organize
things such that their outer levels are either completely full or completely empty.
2.4 Van der Waals bond
Van der Waals forces are driven by induced electrical interactions between two or more molecules
that are very close to each other. Van der Waals interaction is the weakest of all intermolecular
attractions between molecules. However, with a lot of Van der Waals forces interacting between two
objects, the interaction can be very strong. There are two kinds of Van der Waals forces: weak London
Dispersion Forces and stronger dipole-dipole forces. Even if a molecule is nonpolar, this displacement
of electrons causes a nonpolar molecule to become polar for a moment.
Since the molecule is polar, this means that all the electrons are concentrated at one end and the
molecule is partially negatively charged on that end. This negative end makes the surrounding
molecules have an instantaneous dipole also, attracting the surrounding molecules' positive ends. This
process is known as the London Dispersion Force of attraction.
The ability of a molecule to become polar and displace its electrons is known as the molecule's
"polarizability." The more electrons a molecule contains, the higher its ability to become polar.
23
Polarizability increases in the periodic table from the top of a group to the bottom and from right to left
within periods. This is because the higher the molecular mass, the more electrons an atom has. With
more electrons, the outer electrons are easily displaced because the inner electrons shield the
nucleus' positive charge from the outer electrons which would normally keep them close to the
nucleus.
2.5 Hydrogen bonding
A hydrogen bond is a weak type of force that forms a special type of dipole-dipole attraction which
occurs when a hydrogen atom bonded to a strongly electronegative atom exists in the vicinity of
another electronegative atom with a lone pair of electrons. These bonds are generally stronger than
ordinary dipole-dipole and dispersion forces, but weaker than true covalent and ionic bonds.
The hydrogen is attached directly to one of the most electronegative elements, causing the hydrogen
to acquire a significant amount of positive charge. Each of the elements to which the hydrogen is
attached is not only significantly negative, but also has at least one "active" lone pair. Lone pairs at the
2-level have the electrons contained in a relatively small volume of space which therefore has a high
density of negative charge. Lone pairs at higher levels are more diffuse and not so attractive to
positive things.
Fig. 2.16 Hydrogen bonding between a) water and b) HF molecules
Questions:
1. Characterise the covalent bond.
2. Name the types of covalent bond and give examples.
3. Define the term electronegativity.
4. Define the principle of the ionic bond formation.
5. Explain the polarisation of the covalent bond.
References:
1. El Sair R.R.: Fundamentals of Chemistry. 2012, Romain Elsair & Ventus Publishing ApS, ISBN:
978-87-403-0105-2.
2. Nelson P.G.: Introduction to Inorganic Chemistry. 2011, Peter G. Nelson & Ventus Publishing
ApS, ISBN: 978-87-7681-732-9.
3. Beier S.P., Hede P.D Essentials of Chemistry. 2010, BookBoon, ISBN: 978-87-403-0322-3.
24
3 STATE OF MATTER
Chapter mission
This chapter presents an overview of knowledge about the states of matter and their phase
changes. The phase rule and phase diagram of a single-component system are discussed as well.
Chapter objectives
After studying this chapter you should know to:

describe the differences in the states of matter,

explain the phase changes,

describe the phase diagram of a single-component sytem.
State of matter is one of the distinct forms that matter takes on. Four states of matter are observable in
everyday life: solid, liquid, gas, and plasma.
Fig. 3.1 States of matter
3.1 Liquids
The molecules that make up a liquid flow easily around one another. They are kept from flying apart
by attractive forces between them. Liquids assume the shape of their containers. The volume is
definite if the temperature and pressure are constant. The cohesive forces between liquid molecules
are responsible for the phenomenon known as surface tension. The molecules at the surface do not
have other like molecules on all sides of them and consequently they cohere more strongly to those
directly associated with them on the surface. This forms a surface "film" which makes it more difficult
to move an object through the surface than to move it when it is completely submersed.
In a sample of water, there are two types of molecules. Those that are on the outside, exterior, and
those that are on the inside, interior. The interior molecules are attracted to all the molecules around
them, while the exterior molecules are attracted to only the other surface molecules and to those
below the surface. This makes it so that the energy state of the molecules on the interior is much
25
lower than that of the molecules on the exterior. Because of this, the molecules try to maintain a
minimum surface area, thus allowing more molecules to have a lower energy state. This is what
creates what is referred to as surface tension. An illustration of this can be seen in Fig. 3.2.
Fig. 3.2 Interactions between interior and surface molecules
3.2 Solids
The molecules that make up a solid are arranged in regular, repeating patterns. They are held firmly in
place but can vibrate within a limited area. As a result, a solid has a stable, definite shape, and a
definite volume. Solids can only change their shape by force, as when broken or cut. The constituents
of a solid can be arranged in two general ways: they can form a regular repeating three-dimensional
structure called a crystal lattice, thus producing a crystalline solid, or they can aggregate with no
particular order, in which case they form an amorphous solid.
Fig. 3.3 Crystalline solid
Crystals tend to have relatively sharp, well-defined melting points because all the component atoms,
molecules, or ions are the same distance from the same number and type of neighbors; that is, the
regularity of the crystalline lattice creates local environments that are the same.
Amorphous solids have two characteristic properties. When cleaved or broken, they produce
fragments with irregular, often curved surfaces; and they have poorly defined patterns when exposed
to x-rays because their components are not arranged in a regular array. An amorphous, translucent
26
solid is called a glass. Almost any substance can solidify in amorphous form if the liquid phase is
cooled rapidly enough. Some solids, however, are intrinsically amorphous, because either their
components cannot fit together well enough to form a stable crystalline lattice or they contain
impurities that disrupt the lattice.
Fig. 3.4 Amorphous glass structure
3.3 Gas
The molecules that make up a gas fly in all directions at great speeds. They are so far apart that the
attractive forces between them are insignificant. Not only will a gas conform to the shape of its
container but it will also expand to fill the container. The typical distance between neighboring
molecules is much greater than the molecular size. A gas has no definite shape or volume, but
occupies the entire container in which it is confined.
3.4 Plasma
Like a gas, plasma does not have definite shape or volume. Unlike gases, plasmas are electrically
conductive, produce magnetic fields and electric currents, and respond strongly to electromagnetic
forces. Positively charged nuclei swim in a "sea" of freely-moving disassociated electrons, similar to
the way such charges exist in conductive metal. In fact it is this electron "sea" that allows matter in the
plasma state to conduct electricity.
Fig. 3.5 Plasma
27
3.5 Phase changes
A phase change is the transformation of a thermodynamic system from one phase or state of matter to
another one by heat transfer. The term is most commonly used to describe transitions between solid,
liquid and gaseous states of matter, and, in rare cases, plasma.
Fig. 3.6 Phase changes
3.5.1 Condensation and vaporisation
Condensation is the change of the physical state of matter from gas phase into liquid phase, and is the
reverse of vaporation. It can also be defined as the change in the state of water vapor to liquid water
when in contact with any surface.
Vaporization is a conversion of a substance from the liquid or solid phase into the gaseous (vapour)
phase. If conditions allow the formation of vapour bubbles within a liquid, the vaporization process is
called boiling. Direct conversion from solid to vapour is called sublimation.
Heat must be supplied to a solid or liquid to effect vaporization. If the surroundings do not supply
enough heat, it may come from the system itself as a reduction in temperature. The atoms or
molecules of a liquid or solid are held together by cohesive forces, and these forces must be
overcome in separating the atoms or molecules to form the vapour; the heat of vaporization is a direct
measure of these cohesive forces. Condensation of a vapour to form a liquid or a solid is the reverse
of vaporization, and in the process heat must be transferred from the condensing vapour to the
surroundings. The amount of this heat is characteristic of the substance, and it is numerically the
same as the heat of vaporization.
28
Fig. 3.7 Vaporisation and condensation processes
The boiling point of a liquid is the temperature at which its vapor pressure is equal to the pressure of
the gas above it. The normal boiling point of a liquid is the temperature at which its vapor pressure is
equal to one atmosphere (101,325 kPa).
Fig. 3.8 Vaporisation and boiling processes
3.5.2 Melting and freezing
Melting is a physical process that results in the phase transition of a substance from a solid to a liquid.
The internal energy of a substance is increased, typically by the application of heat or pressure,
resulting in a rise of its temperature to the melting point, at which the ordering of ionic or molecular
entities in the solid breaks down to a less ordered state and the solid liquefies. An object that has
melted completely is molten.
Freezing, or solidification, is a phase transition in which a liquid turns into a solid when its temperature
is lowered below its freezing point. For most substances, the melting and freezing points are the
same temperature; however, certain substances possess differing solid–liquid transition temperatures.
Fig. 3.9 Melting of ice
29
3.5.3 Sublimation
Sublimation is the transition of a substance directly from the solid to the gas phase without passing
through an intermediate liquid phase. Sublimation is an endothermic phase transition that occurs at
temperatures and pressures below a substance's triple point in its phase diagram. The reverse
process of sublimation is desublimation, or deposition.
Fig. 3.10 Sublimation of dry ice
3.6 Phase diagram
Phase diagram is a graphical representation of the physical states of a substance under different
conditions of temperature and pressure. A typical phase diagram has pressure on the y-axis and
temperature on the x-axis. As we cross the lines or curves on the phase diagram, a phase change
occurs. In addition, two states of the substance coexist in equilibrium on the lines or curves.
Fig. 3.11 Phase diagram of single-component sytem

Triple point – the point on a phase diagram at which the three states of matter: gas, liquid,
and solid coexist

Critical point – the point on a phase diagram at which the substance is indistinguishable
between liquid and gaseous states

Fusion(melting) (or freezing) curve – the curve on a phase diagram which represents the
transition between liquid and solid states
30

Vaporization (or condensation) curve – the curve on a phase diagram which represents the
transition between gaseous and liquid states

Sublimation (or deposition) curve – the curve on a phase diagram which represents the
transition between gaseous and solid states
The Phase Rule describes the possible number of degrees of freedom in a (closed) system at
equilibrium, in terms of the number of separate phases and the number of chemical constituents in the
system. It was deduced from thermodynamic principles by J. W. Gibbs in the 1870s.
The Degrees of Freedom [F] is the number of independent intensive variables (i.e. those that are
independent of the quantity of material present) that need to be specified in value to fully determine
the state of the system. Typical such variables might be temperature, pressure, or concentration.
A Phase [P] is a component part of the system that is immiscible with the other parts (e.g. solid, liquid,
or gas); a phase may of course contain several chemical constituents, which may or may not be
shared with other phases.
The Chemical Constituents [C] are simply the distinct compounds (or elements) involved in the
equations of the system. (If some of the system constituents remain in equilibrium with each other
whatever the state of the system, they should be counted as a single constituent.) The rule is:
F=C-P+2
(3.1)
For example:A system with one component and one phase (a balloon full of carbon dioxide, perhaps)
has two degrees of freedom: temperature and pressure, say, can be varied independently. If you have
two phases -- liquid and vapour for instance -- you lose a degree of freedom, and there is only one
possible pressure for each temperature. Add yet one more phase -- ice, water and water vapour in a
sealed flask -- and you have a "triple point" with fixed temperature and pressure.
Questions:
1. In what states does matter exist?
2. Explain the terms: boiling point, equilibrium vapour pressure, sublimation.
3. Draw and desribe a phase diagram of a single-component system.
4. What does mean the term triple point?
5. Write the Gibbs phase rule.
References:
1. El Sair R.R.: Fundamentals of Chemistry. 2012, Romain Elsair & Ventus Publishing ApS, ISBN:
978-87-403-0105-2.
2. Nelson P.G.: Introduction to Inorganic Chemistry. 2011, Peter G. Nelson & Ventus Publishing
ApS, ISBN: 978-87-7681-732-9.
3. Beier S.P., Hede P.D.: Essentials of Chemistry. 2010, BookBoon, ISBN: 978-87-403-0322-3.
31
4 CHEMICAL REACTION
Chapter mission
This chapter presents an overview of knowledge about the classification of chemical reactions
based on phases of reactants and mechanism of the changes.
Special attention is paid to
protolytic and oxidation-reduction reactions. For a complete characterization of chemical processes
an overview of knowledge from the kinetics and thermodynamics are provided, too.
Chapter objectives
After studying this chapter you should know to:

classify the chemical reaction according to reaction phases and mechanism of changes,

characterize protolytic and oxidation-reduction reactions,

describe relations that characterize the chemical reaction rate and chemical equilibrium.
A chemical reaction is a process that is usually characterized by a chemical change in which the
starting materials (reactants) are different from the products. Chemical reactions tend to involve the
motion of electrons, leading to the formation and breaking of chemical bonds. There are several
different types of chemical reactions and more than one way of classifying them. Here are some
common reaction types.
4.1 Clasification of chemical reactions
4.1.1 Chemical reactions according to number of reaction phases
Homogeneous - that occur in a single phase (gaseous, liquid, or solid),
2 NO (g) + O2 (g)  2 NO2 (g)
(4.1)
Heterogenous that occur on the interface of phases
CuSO4 (aq) + Fe (s)  FeSO4 (aq) + Cu (s)
(4.2)
4.1.2 Chemical reactions according to type of changes
Synthesis reaction - two or more simple substances combine to form a more complex substance
NH3 (g) + HCl (g)  NH4Cl (s)
(4.3)
Decomposition reaction - a more complex substance breaks down into its more simple parts.
CaCO3 (s)  CaO (s) + CO2 (g)
Single replacement reaction - a single uncombined element replaces another in a compound.
32
(4.4)
Zn (s) + H2SO4 (l)  ZnSO4 (s) + H2 (g)
(4.5)
Double replacement reaction - parts of two compounds switch places to form two new compounds.
Two reactants yield two products.
2 NaCl (s) + H2SO4 (l)  Na2SO4 (s) + 2 HCl (g)
(4.6)
4.2 Reaction mechanism
In chemistry, a reaction mechanism is the step by step sequence of elementary reactions by which
overall chemical change occurs.
Classification of chemical reactions regarding to reaction mechanism
1. protolytical (acid-base) reaction,
2. oxidation-reduction,
3. precipitation reaction,
4.2.1 Protolytical reakction
An acid-base reaction is type of double displacement reaction that occurs between an acid and a
+
-
base. The H ion in the acid reacts with the OH ion in the base to form water and an ionic salt:
HA + BOH → H2O + BA
(4.7)
The Arrhenius Theory of acids and bases
The theory

Acids are substances which produce hydrogen ions in solution.

Bases are substances which produce hydroxide ions in solution.
Neutralisation happens because hydrogen ions and hydroxide ions react to produce water.
+
-
H (aq) + OH (aq) →H2O(l)
(4.9)
The Bronsted-Lowry Theory of acids and bases
The theory

An acid is a proton (hydrogen ion) donor.

A base is a proton (hydrogen ion) acceptor.
The relationship between the Bronsted-Lowry theory and the Arrhenius theory

The Bronsted-Lowry theory doesn't go against the Arrhenius theory in any way - it just adds to
it.

Hydroxide ions are still bases because they accept hydrogen ions from acids and form water.

An acid produces hydrogen ions in solution because it reacts with the water molecules by
giving a proton to them.
When hydrogen chloride gas dissolves in water to produce hydrochloric acid, the hydrogen chloride
molecule gives a proton (a hydrogen ion) to a water molecule. A co-ordinate (dative covalent) bond is
formed between one of the lone pairs on the oxygen and the hydrogen from the HCl. Hydroxonium
+
ions, H3O , are produced.
33
+
(4.10)
-
H2O + HCl → H3O + Cl
When an acid in solution reacts with a base, what is actually functioning as the acid is the
hydroxonium ion. For example, a proton is transferred from a hydroxonium ion to a hydroxide ion to
make water.
+
(4.11)
-
H3O (aq) + OH (aq) → 2 H2O(l)
Showing the electrons, but leaving out the inner ones:
It is important to realise that whenever you talk about hydrogen ions in solution, H
+
(aq),
what you are
actually talking about are hydroxonium ions.
The hydrogen chloride / ammonia problem
This is no longer a problem using the Bronsted-Lowry theory. Whether you are talking about the
reaction in solution or in the gas state, ammonia is a base because it accepts a proton (a hydrogen
ion). The hydrogen becomes attached to the lone pair on the nitrogen of the ammonia via a coordinate bond.
If it is in solution, the ammonia accepts a proton from a hydroxonium ion:
+
+
NH3(aq) + H3O →NH4 (aq) + H2O(l)
(4.12)
If the reaction is happening in the gas state, the ammonia accepts a proton directly from the hydrogen
chloride:
+
-
NH3(g) + HCl(g) → NH4 (s) + Cl (s)
Either way, the ammonia acts as a base by accepting a hydrogen ion from an acid.
34
(4.13)
Conjugate pairs
When hydrogen chloride dissolves in water, almost 100% of it reacts with the water to produce
hydroxonium ions and chloride ions. Hydrogen chloride is a strong acid, and we tend to write this as a
one-way reaction:
+
(4.14)
-
H2O + HCl → H3O + Cl
In fact, the reaction between HCl and water is reversible, but only to a very minor extent. In order to
generalise, consider an acid HA, and think of the reaction as being reversible.
+
(4.15)
-
HA + H2O ↔H3O + A
Thinking about the forward reaction:

The HA is an acid because it is donating a proton (hydrogen ion) to the water.

The water is a base because it is accepting a proton from the HA.
-
But there is also a back reaction between the hydroxonium ion and the A ion:

The H3O is an acid because it is donating a proton (hydrogen ion) to the A ion.

The A ion is a base because it is accepting a proton from the H3O .
+
-
-
+
The reversible reaction contains two acids and two bases. We think of them in pairs, called conjugate
pairs.
-
-
When the acid, HA, loses a proton it forms a base, A . When the base A accepts a proton back again
it obviously refoms the acid HA. These two are a conjugate pair. Members of a conjugate pair differ
from each other by the presence or absence of the transferable hydrogen ion.
-
-
If you are thinking about HA as the acid, then A is its conjugate base. If you are thinking about A as
the base, then HA is its conjugate acid. The water and the hydroxonium ion are also a conjugate pair.
Thinking of the water as a base, the hydroxonium ion is its conjugate acid because it has the extra
hydrogen ion which it can give away again.
Thinking about the hydroxonium ion as an acid, then water is its conjugate base. The water can
accept a hydrogen ion back again to reform the hydroxonium ion.
A second example of conjugate pairs
This is the reaction between ammonia and water that we looked at earlier:
35
Think first about the forward reaction. Ammonia is a base because it is accepting hydrogen ions from
the water. The ammonium ion is its conjugate acid - it can release that hydrogen ion again to reform
the ammonia. The water is acting as an acid, and its conjugate base is the hydroxide ion. The
hydroxide ion can accept a hydrogen ion to reform the water. Looking at it from the other side, the
ammonium ion is an acid, and ammonia is its conjugate base. The hydroxide ion is a base and water
is its conjugate acid.
4.2.2 Oxidation-reduction reakction
Redox reactions, or oxidation-reduction reactions, have a number of similarities to acid-base
reactions. Fundamentally, redox reactions are a family of reactions that are concerned with the
transfer of electrons between species. Like acid-base reactions, redox reactions are a matched set -you don't have an oxidation reaction without a reduction reaction happening at the same time.
Oxidation refers to the loss of electrons, while reduction refers to the gain of electrons. Each reaction
by itself is called a "half-reaction", simply because we need two (2) half-reactions to form a whole
reaction. In notating redox reactions, chemists typically write out the electrons explicitly:
Cu (s) → Cu
2+
+2e
-
(4.16)
This half-reaction says that we have solid copper (with no charge) being oxidized (losing electrons) to
form a copper ion with a plus 2 charge. Notice that, like the stoichiometry notation, we have a
"balance" between both sides of the reaction. We have one (1) copper atom on both sides, and the
-
charges balance as well. The symbol "e " represents a free electron with a negative charge that can
now go out and reduce some other species, such as in the half-reaction:
+
-
2 Ag (aq) + 2 e → 2 Ag (s)
(4.17)
Here, two silver ions (silver with a positive charge) are being reduced through the addition of two (2)
electrons to form solid silver. The abbreviations "aq" and "s" mean aqueous and solid, respectively.
We can now combine the two (2) half-reactions to form a redox equation:
2+
-
(4.18)
2Ag (aq) + 2e →2Ag(s)
(4.19)
Cu(s) → Cu (aq) + 2e
+
-
-
2+
Cu(s) + 2Ag+ + 2e → Cu (aq) + 2Ag(s) + 2e
-
(4.20)
2+
or Cu(s) + 2Ag+ → Cu (aq) + 2Ag(s)
(4.21)
We can also discuss the individual components of these reactions as follows. If a chemical causes
+
another substance to be oxidized, we call it the oxidizing agent. In the equation above, Ag is the
oxidizing agent, because it causes Cu(s) to lose electrons. Oxidants get reduced in the process by a
+
reducing agent. Cu(s) is, naturally, the reducing agent in this case, as it causes Ag to gain electrons.
36
4.2.3 Precipitation reaction
A precipitation reaction is a reaction in which soluble ions in separate solutions are mixed together to
form an insoluble compound that settles out of solution as a solid.
That insoluble compound is called a precipitate.
example:
AgNO3 + KCl  AgCl + KNO3
(4.22)
4.3 Chemical kinetics
Chemical kinetics is the study of chemical reactions with respect to reaction rates, effect of various
variables, re-arrangement of atoms, formation of intermediates etc. At the macroscopic level, we are
interested in amounts reacted, formed, and the rates of their formation. At the molecular or
microscopic level, the following considerations must also be made in the discusion of chemical
reaction mechanism.
Molecules or atoms of reactants must collide with each other in chemical reactions. The molecules
must have sufficient energy (discussed in terms of activation energy) to initiate the reaction. In some
cases, the orientation of the molecules during the collision must also be considered.
The activation energy of a reaction is the amount of energy needed to start the reaction. It represents
the minimum energy needed to form an activated complex during a collision between reactants
(Fig.4.1)
Fig. 4.1 Activation energy
(http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch22/activate.html)
4.3.1 Reaction rate
Chemical reaction rates are the rates of change in concentrations or amounts of either reactants or
products. The main factors that influence the reaction rate include: the physical state of the reactants,
the concentrations of the reactants, the temperature at which the reaction occurs, and whether or not
any catalysts are present in the reaction.
Factors influence reaction rate:
1. Nature of reactants

Acid-base reactions, formation of salts, and exchange of ions are fast reactions.
37

Reactions in which large molecules are formed or break apart are usually slow.

Reactions breaking strong covalent bonds are also slow.
2. Concentration effect
The rate of a chemical reaction is the change in concentration over the change in time and is a metric
of the "speed" at which a chemical reactions occurs and can be defined in terms of two observables:
In the simple reaction
A+B→C+D
(4.23)
dc A
dc
 B
dt
dt
(4.24)
The reaction rate can be defined thusly:
vc  
(Note this is negative because it measures the rate of disappearance of the reactants. )
vc 
dcC dc D

dt
dt
(4.25)
(This is the rate at which the products are formed)
They both are linked via the balanced chemical reactions and can both be used to measure the
reaction rate.
The reaction rate is often found to have the form:
 dA
 kA
dt
(4.26)
The Arrhenius equation gives the quantitative basis of the relationship between the activation energy
and the reaction rate at which a reaction proceeds. The rate constant is then given by

k  Ae
Ea
RT
(4.27)
where
k - rate constant,
R - the ideal gas constant [8.314 J/(mol-° K)],
T- the temperature in degrees Kelvin,
Ea - the activation energy in joules per mol,
A - constant called the frequency factor; which is related to the fraction of collisions between reactants
having the proper orientation to form an activated complex.
Catalysis is the increase in the rate of a chemical reaction due to the participation of an additional
substance called a catalyst. With a catalyst, reactions occur faster and with less energy. Because
38
catalysts are not consumed, they are recycled. Often only tiny amounts are required. Katalysts
substances used to facilitate reactions by activation energy decreasing (Fig. 4.2)
Fig. 4.2 Decreasing activation energy by katalysts
(http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch22/activate.html)
In heterogeneous reactions, reactants are present in more than one phase and the rates of chemical
reactions is affected by surface areas.
Chemical equilibrium
In a chemical reaction, chemical equilibrium is the state in which both reactants and products are
present in concentrations which have no further tendency to change with time. Usually, this state
results when the forward reaction proceeds at the same rate as the reverse reaction. The reaction
rates of the forward and backward reactions are generally not zero, but equal. Thus, there are no net
changes in the concentrations of the reactant(s) and product(s). Such a state is known as dynamic
equilibrium.
In the simple reaction
A + B  C + D
(4.28)
concentrations of reactants and products remain constant.
The rate of forward reaction v1 = k1 A B is equal to the rate of reverse reaction v2 = k2 C D
If v1 = v2,
than
k 1 CD

 Kc
k 2 AB
(4.29)
Kc - equilibrium constant (varies with temperature.)
Thermochemistry
Thermochemistry is the study of the energy and heat associated with chemical reactions and/or
physical transformations. A reaction may release or absorb energy, and a phase change may do the
39
same, such as in melting and boiling. Thermochemistry focuses on these energy changes, particularly
on the system's energy exchange with its surroundings. Thermochemistry is useful in predicting
reactant and product quantities throughout the course of a given reaction. In combination with entropy
determinations, it is also used to predict whether a reaction is spontaneous or non-spontaneous,
favorable or unfavorable.
Endothermic reactions absorb heat. Exothermic reactions release heat. Thermochemistry coalesces
the concepts of thermodynamics with the concept of energy in the form of chemical bonds. The
subject commonly includes calculations of such quantities as enthalpy, entropy, and free energy.
Enthalpy
Enthalpy, H, is a statefunction used to describe the heat changes that occur in a reaction
underconstant pressure. When a reaction is allowed to take place in an open container a quantity of
heatproportional to the quantity of matterpresent, will be released or absorbed. This flow of heat is the
enthalpy change, ∆H. The units for ∆H are kJ (or kJ/mol).

Reactions that release heat are termed exothermic, they have negative values of ∆H.

Reactions that absorb heat are termed endohermic, they have positive values of ∆H.
The endothermic reaction shown below indicates that 92.2 kJ are absorbed when 2 moles of NH 3
decompose to form 1 mole of N2 and 3 moles of H2.
2 NH3 (g) → N2 (g) + 3 H2 (g)
∆H = +92.2 kJ
(4.30)
Standard Heat of Formation
A substance in its most stable form at one atmosphere is said to be in its standard state (the
temperature need not be 25 °C, but this temperature is most commonly used to tabulate data). The
enthalpy change which occurs when one mole of a compound is produced from its elements in their
o
standard states is called the standard heat of formation, ∆H f . The equation for the standard heat of
formation for ammonia, NH3, is given below:
o
1/2 N2 (g) + 3/2 H2 →1 NH3 (g)
By definition all elements in their standard states have a ∆H
∆H f = -46.1 kJ
o
f
(4.31)
equal to zero.
Entropy
Entropy, S, is a quantitative measure of the disorder or randomness of a system. The greater the
number of ways that matter can be arranged, the greater entropy. Listed below are some criteria for
predicting changes in entropy.

Entropy increases with temperature.

Solids have the lowest entropy, while gases have the most. Liquids and aqueous
solutions have intermediate values.

In any reaction where the number of moles gas increases the entropy will
increase.

When pure substances form mixtures entropy will increase.
40
Reactions like, 2HI (g) → H2 (g) + I2 (g), where the number of moles of gas are constant, but the
number of different types of gas increases, proceed with an increase in entropy.
Gibbs free energy
The Gibbs free energy, G, is a measure of the energy available to the system to do useful work. The
difference in free energy between the products and the reactants is called the f r e e energy change,
∆G. ∆G is expressed in kJ (or kJ/mol). When the free energy of the reactants is greater than that of the
products a reaction will be spontaneous. It is the sign of ∆G which may be used to determine the
spontaneity of a reaction.

If ∆G is negative, the reaction is spontaneous.

If ∆G is positive, the reaction is nonspontaneous.

If ∆G is zero, the reaction is at equilibrium.
Although it is the change in free energy, ∆G, which determines the direction of a reaction, it is the
activation energy which determines its rate.
The value of ∆G may be obtained by the formula below if ∆H, T, and ∆S is known:
∆G = ∆H - T∆S
(4.31)
When using this formula the temperature must be in Kelvin, and the units of energy for ∆H and ∆S
must be the same.
Questions:
1. How are clasified chemical reactions according to number of reaction phases?
2. Define the protolytical reactions.
3. Characterize oxidation-reduction reactions.
4. What is the matter of chemical kinetics?
5. Name the factors which influence reaction rate.
6. Write the formula of the chemical equilibrium.
7. Characterize Entropy, Enthalpy and Gibbs free energy.
References:
1. Walker, D.: Chemical reactions. 2007, Evans Brothers, ISBN: 0237530015
2. Houston, P. L.: Chemical Kinetics and Reaction Dynamics. 2006, Dover Books on Chemistry,
ISBN: 0486453340
3. Koretsky, M. D.: Engineering and Chemical Thermodynamics. 2010, John Wiley & Sons Inc,
ISBN: 0471385867
41
5 WATER
Chapter mission
This chapter is aimed on the influence of structure of water molecule on its physical – chemical
properties. Water is characterized as a polar solvent which is connected with solubility also
negligible impurities cause hardnes of water. Chapter presents methods of water softening. At the
end is explained pH scale.
Chapter objectives
After studying this chapter you should know to:
characterize water structure,
describe properties of water solutions and methods of softening.
define pH scale of solutions.
Water covers 70.9% of the Earth's surface, and is vital for all known forms of life. On Earth, 96.5% of
the planet's water is found mostly in oceans; 1.7% in groundwater; 1.7% in glaciers and the ice caps of
Antarctica and Greenland. Water is also an essential requirement for all forms of life. Human beings,
for example, consist of about two-thirds water. Water is a liquid at ambient conditions, but it often coexists on Earth with its solid state, ice, and gaseous state (water vapour or steam). Water is a
tasteless, odorless liquid at ambient temperature and pressure, and appears colorless in small
quantities, although it has its own intrinsic very light blue hue. Ice also appears colorless, and water
vapor is essentially invisible as a gas.
5.1 Water structure
Water is the chemical substance with chemical formula H2O, one molecule of water has two hydrogen
atoms covalently bonded to a single oxygen atom. Water is primarily a liquid under standard
conditions, which is not predicted from its relationship to other analogous hydrides of the oxygen
family in the periodic table, which are gases such as hydrogen sulfide. The elements surrounding
oxygen in the periodic table, nitrogen, fluorine, phosphorus, sulfur and chlorine, all combine with
hydrogen to produce gases under standard conditions. The reason that water forms a liquid is that
oxygen is more electronegative than all of these elements with the exception of fluorine. Oxygen
attracts electrons much more strongly than hydrogen, resulting in a net positive charge on the
hydrogen atoms, and a net negative charge on the oxygen atom. The presence of a charge on each of
these atoms gives each water molecule a net dipole moment. Electrical attraction between water
molecules due to this dipole pulls individual molecules closer together, making it more difficult to
separate the molecules and therefore raising the boiling point.
42
In H2O, only two of the six outer-shell electrons of oxygen are used for this purpose, leaving four
electrons which are organized into two non-bonding pairs. The four electron pairs surrounding the
oxygen tend to arrange themselves as far from each other as possible in order to minimize repulsions
between these clouds of negative charge. The two non-bonding pairs remain closer to the oxygen
atom, these exert a stronger repulsion against the two covalent bonding pairs, effectively pushing the
two hydrogen atoms closer together. The result is a distorted tetrahedral arrangement in which the
H—O—H angle is 104.5° (instead 109.5° in tetrahedral geometry) (Fig. 5.1).
Fig. 5.1 Water structure
5.1.1 Hydrogen bonds in water
In a polar covalent bond, the electrons are closer to the oxygen nucleus than the hydrogen nucleus.
This is because of the geometry of the molecule and the great electronegativity difference between the
Hydrogen atom and the oxygen atom.
The result of this pattern of unequal electron association is a charge separation in the molecule, where
one part of the molecule, the oxygen, has a partial negative charge and the hydrogen have a partial
positive charge.
A hydrogen bond is the electrostatic attraction between polar molecules that occurs when a hydrogen
(H) atom bound to a highly electronegative atom such as oxygen (O) is attracted to some other nearby
highly electronegative atom (Fig. 5.2). These hydrogen-bond attractions can occur between molecules
(intermolecular) or within different parts of a single molecule (intramolecular). The hydrogen bond (5 to
30 kJ/mole) is stronger than a van der Waals interaction, but weaker than covalent or ionic bonds.
Intermolecular hydrogen bonding is responsible for the high boiling point of water (100 °C) compared
to the other group 16 hydrides that have no hydrogen bonds.
Fig. 5.2 Hydrogen bonds in water
The solid form of most substances is denser than the liquid phase; thus, a block of most solids will sink
in the liquid. However, a block of ice floats in liquid water because ice is less dense. Upon freezing,
43
the density of water decreases by about 9%.This is due to the 'cooling' of intermolecular vibrations
allowing the molecules to form steady hydrogen bonds with their neighbors and thereby gradually
locking into positions reminiscent of the hexagonal packing achieved upon freezing to ice (Fig.5.3).
Fig. 5.3 Crystal structure of hexagonal ice - gray dashed lines indicate hydrogen bonds
5.2 Water solutions
In chemistry, a solution is a homogenous mixture composed of only one phase. In such a mixture, a
solute is dissolved in another substance, known as a solvent. Water is a good solvent due to its
polarity.
When an ionic or polar compound enters water, it is surrounded by water molecules. The partially
negative dipoles of the water are attracted to positively charged components of the solute, and vice
versa for the positive dipoles (Fig. 5.4).
Fig. 5.4 Dissolving of NaCl
5.2.1 Solubility
The ability of one compound to dissolve in another compound is called solubility (Fig. 5.5). When a
liquid is able to completely unfavorable dissolve in another liquid the two liquids are miscible. Two
substances that can never mix to form a solution are called immiscible (Fig.5.6).
44
Fig. 5.5 An example of solubility
Fig.5.6 An example of immiscibility
The interactions between different molecules or ions may be energetically favored or not. If
interactions are unfavorable, then the free energy decreases with increasing solute concentration. At
some point the energy loss outweighs the entropy gain, and no more solute particles can be dissolved;
the solution is said to be saturated (Fig. 5.7, Fig. 5.8).
Fig. 5.7 Saturated solution of lead(II)
Fig. 5.8 Saturated solution of NaCl
chromate
However, the point at which a solution can become saturated can change significantly with different
environmental factors, such as temperature, pressure and contamination. Usually, the greater the
temperature of the solvent, the more of a given solid solute it can dissolve (Fig. 5.9).
Fig. 5.9 Dependence solubility on temperature 1 - CuSO4.5H2O, 2 - NaCl, 3 - Na2SO4,
45
5.2.2 Water hardness
Water is a good solvent and picks up impurities easily. As water moves through soil and rock, it
dissolves very small amounts of minerals and holds them in solution. Calcium and magnesium
dissolved in water are the two most common minerals that make water "hard."
The degree of hardness becomes greater as the calcium and magnesium content increases and is
related to the concentration of multivalent cations dissolved in the water.
There are two types of hardness, temporary and permanent, which depend on the type of salts
present.
1. Temporary hardness is usually result of dissolved calcium and magnesium hydrocarbonates
Ca(HCO3)2, Mg(HCO3)2 ,
Water softening: hydrocarbonates can be converted to carbonates by heating according to the
equation:
Ca(HCO3)2 (aq)  CaCO3 + CO2 + H2O
(5.1)
2. Permanent hardness contains the sulphates of calcium and magnesium, which remain soluble
even when heated.
Water softening by lime and sodium carbonate is based on the precipitate production (e.g.
limestone, gypsum)
Ca(OH)2 + Ca(HCO3)2  2 CaCO3 + 2 H2O
Ca(OH)2 + MgSO4 + 2 H2O  Mg(OH)2 + CaSO4.2 H2O
CaSO4 + Na2CO3  CaCO3 + Na2SO4
(5.2)
(5.3)
(5.4)
Adding of sodium phosphate resulted in precipitate of calcium phosphate
Na3PO4 + 3 Ca(HCO3)2  Ca3(PO4)2 + 6 NaHCO3
2 Na3PO4 + 3 CaSO4  Ca3(PO4)2 + 3 Na2SO4
(5.5)
(5.6)
Water softening by ion Exchange
The presence of calcium (Ca) and/or magnesium (Mg) in water results in water being considered
“hard.” Calcium and magnesium ions in water react with heat, metallic plumbing, and chemical agents
such as detergents to decrease the effectiveness of nearly any cleaning task. Hard water can be
softened using an ion exchange softening process.
Ion exchange involves removing the hardness ions calcium and magnesium and replacing them with
non-hardness ions, typically sodium supplied by dissolved sodium chloride salt, or brine. The softener
contains a microporous exchange resin, usually sulfonated polystyrene beads that are supersaturated
with sodium to cover the bead surfaces. As water passes through this resin bed, calcium and
magnesium ions attach to the resin beads and the loosely held sodium is released from the resin into
the water. The softening process is illustrated in Fig. 5.10 and Fig. 5.11.
After softening a large quantity of hard water the beads become saturated with calcium and
magnesium ions. When this occurs, the exchange resin must be regenerated, or recharged. To
regenerate, the ion exchange resin is flushed with a salt solution . The sodium ions in the salt brine
solution are exchanged with the calcium and magnesium ions on the resin and excess calcium and
magnesium is flushed out with wastewater.
46
Fig. 5.10 The principle of ion exchange
Fig. 5.11 Ion exchange in praxis
5.3 Autoionization of water and pH scale
The molecules in pure water autodisociate into hydronium and hydroxide ions in the following
equilibrium:
H2O + H2O  H3O
Kc 
+
+
OH
-
(5.7)
H O . OH 


3
H 2O2
(5.8)
The concentration of water in a water solution is constant and this expression simplifies to:
+
-
Kw= [H ] [OH ]
+
or
(5.9)
-
Kw = c(H3O ) . c(OH )
(5.10)
where Kw is called the dissociation constant of water and equals 1.00x10
+
-
o
-14
at room temperature.
-7
The molar concentrations of H3O and OH at 25 C are equal and found to be 1.10 mol.dm
+
-
-7
c(H3O ) = c(OH ) = 10 mol.dm
47
-3
-3
(5.11)
5.3.1 pH scale
In a solution pH is equal to the negative logarithm (base 10) of the molar concentration of dissolved
hydronium ions;
+
pH = - log c(H3O )
(5.12)
A low pH indicates a high concentration of hydronium ions, while a high pH indicates a low
concentration. This negative of the logarithm matches the number of places behind the decimal point,
so, for example, 0.1 molar hydrochliric acid should be near pH 1 and 0.0001 molar HCl should be near
pH 4.
In pure water, there is an equal number of hydroxide and hydronium ions, so it has a neutral pH of 7.
A pH value less than 7 indicates an acidic solution, and a pH value more than 7 indicates a basic
solution (Table 5.1).
Table 5.1 Acidity and alkality of solutions
Solution
Concentration
pH
neutral
c(H3O ) = c (OH )
7
acidic
c(H3O )  c (OH )
7
basic
c(H3O )  c (OH )
7
+
-
+
-
+
-
Questions:
1. Describe the water structure.
2. How can hydrogen bonds influence the physical properties of water?
3. How are defined chemical solutions and which are their main components?
4. What does the term water hardness mean and how is water hardness devided?
5. What methods of water hardness removal do you know?
6. Define pH scale, acidity and alkality of solutions.
References:
1. Eisenberg, D., Kauzmann, W.: The structure and properties of water. 2005, Oxford University
Press, ISBN: 0198570260.
2. Benjamin, M. M.,
Lawler, D. F.: Water quality engineering: physical / chemical treatment
processes. 2013, Wiley, ISBN: 978-1-118-16965-0
3. McEachern, R., Wist, W., Lehr, J. H.: Water Softening with Potassium Chloride: Process, Health,
and Environmental Benefits. 2009, Wiley, ISBN: 978-0-470-08713-8.
4. Kohlmann, F. J.: What is pH, and how is it measured? 2003. GLI International.
48
6 HETEROGENOUS SYSTEMS
Chapter mission
This chapter is focused on characterization of disperse systems and basic physical – chemical
processes on the interfaces solid/gas or solid/liquid phases.
This knowledge is used for
determination of physical parameters of building materials e.g. specific surface determination,
wetting of surfaces and the pollutants removal from solutions by sorption.
Chapter objectives
After studying this chapter you should know to:
characterize disperse systems,
explain the main differences between physical and chemical adsorptions,
describe determination of specific surface of powder materials,
explain the wetting of solid surfaces,
explain principle of adsorption on the interface of solid/liquid phases.
6.1 Disperse systems
Disperse system is defined as a system in which one substance (The Dispersed Phase) is distributed, in
discrete units, throughout a second substance (the continuous Phase or dispersion medium). Each
phase can exist in solid, liquid, or gaseous state. Forms of disperse system are presented in Table 63.1.
Table 6.1 Forms of disperse system
continuos phase
disperse phase
gaseous
gaseous
complete solubility
between gases, hence no
multi phase systems
possible
liquid
foam
solid
expanded material,
thermal isolation,
activated coal,
silica gel
liquid
solid
aerosols, clouds
fog
dust, smoke
dispersion
emulsion
suspension
organic tissue
heterogenous
alloy,
ceramic,
natural minerals
Many materials used today are disperse systems where one substance, often a particulate, is
dispersed in another phase. Examples are common in areas as diverse as adhesives, agrochemicals,
49
cement, ceramics, colloids, cosmetics and personal care formulations, food and drink, mining and
mineral slurries, paints, inks and surface coatings, pharmaceuticals and polymer systems.
The physical properties of the dispersed particles – average particle size, size distribution, charge on
the particles, and particle shape (Fig. 6.2).
Fig. 6.2 Particle shape
Based on the average particle size dispersions can be classified as:

Molecular dispersions,

Colloids,

Coarse dispersions (suspensions).
6.1.1 Molecular dispersion
Molecular dispersion is a true solution of a solute phase in a solvent. The dispersed phase (solute) is
in form of separate molecules homogeneously distributed throughout the dispersion medium (solvent).
The molecule size is less than 1 nm.
The examples of molecular dispersions: air (a molecular mixture of Oxygen, Nitrogen and some other
gases), electrolytes (aqueous solutions of salts (Fig.6.3)), metal alloys, solid solutions.
Fig. 6.3 Aqueous solution of Cu (II)
6.1.2 Colloids
Colloids are micro-heterogeneous dispersed systems, in which the size of
the dispersed phase
particles is within the range 1 - 1000 nm. The colloids phases cannot be separated under gravity,
50
centrifugal or other forces. Dispersed phase of colloids may be separated from the dispersion medium
by micro-filtration.
The examples of colloids: milk (Fig. 6.4) (emulsion of fat and some other substances in water (Fig.
6.5)), fog (aerosol of water micro-droplets in air), opal (colloidal silica).
Fig. 6.4 Colloid - Milk
Fig. 6.5 Oil – water emulsion
6.1.3 Coarse dispersions
Coarse dispersions are heterogeneous dispersed systems, in which the dispersed phase particles are
larger than 1000 nm.
Coarse dispersions are characterized by relatively fast sedimentation of the dispersed phase (Fig. 6.6)
caused by gravity or other forces. Dispersed phase of coarse dispersions may be easily separated
from the continuous phase by filtration (Fig. 6.7).
Fig. 6.6 Coarse dispersion
Fig. 6.7 Filtration of coarse disperse system
6.2 Adsorption
Adsorption is the process by which a molecule or atom adsorb onto a surface of substrate.
Surface onto which adsorption can occur (example: catalyst surface, activated carbon, alumina) is
called substrate or adsorbent:
51
Adsorbates are molecules or atoms that adsorb onto the substrate (example: nitrogen, hydrogen,
carbon monoxide, water)
Depending on the nature of attractive forces existing between the adsorbate and adsorbent,
adsorption can be classified as:
a) Physical adsorption – bonding between molecules and surface is by weak van der Waals
forces (Fig. 6.8).
b) Chemical adsorption – chemical bonds is formed between molecules and surface (Fig. 6.9)
Fig. 6.8 Physical adsorption
Fig. 6.9 Chemical adsorption
Differences between physical and chemical adsorption are follows:
Physical adsorption:

Van der Waals attraction between adsorbate and adsorbent,

The attraction is not fixed to a specific site and the adsorbate is relatively free to move on the
surface,

This is relatively weak, reversible, adsorption capable of multilayer adsorption.
Chemical adsorption:

Some degree of chemical bonding between adsorbate and adsorbent characterized by strong
bonds,

Adsorbed molecules are not free to move on the surface,

There is a high degree of specificity and typically a monolayer is formed,

The process is seldom reversible.
6.2.1 Adsorption at the gas/solid interface
Adsorption equilibria
If the adsorbent and adsorbate are contacted long enough an equilibrium will be established between
the amount of adsorbate adsorbed and the amount of adsorbate in solution.
The equilibrium relationship is described by isotherms:
1. Freundlich Adsorptiom Isoterm
In 1909, Freundlich gave an empirical expression representing the isothermal variation of Adsorption
of a quantity of gas adsorbed by unit mass of solid adsorbent with pressure (1).
52
1
x
 kp n
m
(6.1)
Where x is the mass of the gas adsorbed on mass m of the adsorbent at pressure p and k, n are
constants whose values depend upon adsorbent and gas at particular temperature.
It is valid only at low pressure when the extent of adsorption is directly proportional to pressure. An
example of Froundlich adsorption isoterms at variout temperatures is in Fig. 6.10
Fig. 6. 10 Freundlich adsorption isoterms at variout temperatures (adsorption of carbon monooxide on
active carbon)
2. Langmuir isotherm
relates the coverage or adsorption of molecules on a solid surface to gas presure or concentration of a
medium above the solid surface at a fixed temperature. The equation is stated as:

a
bp

am 1  bp
(6.2)
θ or theta is the fractional coverage of the surface, p is the gas pressure or concentration, b is a
constant, am is the maximum loading corresponding to complete coverage of the surface by adsorbate
The constant b is the Langmuir adsorption constant and increases with an increase in the binding
energy of adsorption and with a decrease in temperature. An example of Langmuir adsorption
isoterms is in Fig. 6.11.
Fig. 6.11 Langmuir izoterm of hydrogen adsorption on copper at temperature 25 °C
53
6.2.2 Interaction at the liquid /solid interface
6.2.2.1 Surface Energy
Existence of Surface energy is an explanation of why material accumulates at the interface is based
on the excess energy associated with particles at interfaces (Fig. 6.12). For example, in the case of
pure water and air, the water molecules at the air-water interface have higher energy than water
molecules in the interior of the water phase.The reason that these surface molecules have higher
energy is that, unlike the interior molecules, they have an unbalanced force component (on the airside
of the molecule). These surface molecules have additional energy to balance the forces. This excess
energy is called surface tension. Since it takes energy to create interfacial surfaces, the system will
try to minimize the total interfacial surface area. Hence, we see spherical droplet.
Fig. 6.12 Surface tension
6.2.2.2 Wetting
The contact angle  is the angle at which a liquid/vapour interface meets a solid surface (Fig. 6.13).
The contact angle is specific for any given system and is determined by the interactions across the
three interfaces.
Most often the concept is illustrated with a small liquid droplet resting on a flat horizontal solid surface.
Fig. 6.13 The contact angle  at the liquid/solid interface
If the liquid is very strongly attracted to the solid surface (for example water on a strongly hydrophilic
solid) the droplet will completely spread out on the solid surface and the contact angle will be close to
0° (Fig. 6.14).
Less strongly hydrophilic solids will have a contact angle up to 90°. On many highly hydrophilic
surfaces, water droplets will exhibit contact angles of 0° to 30°.
If the solid surface is hydrophobic, the contact angle will be larger than 90° (Fig. 6.15)
54
Fig. 6.14 A strongly hydrophilic surface
Fig. 6.15 Hydrophobic surface
6.2.2.3 Thermodynamics of surface adsorption
In solutions certain particles tend to concentrate at the surface.These particles are those that have low
affinity for the solvent. These are hydrophobic molecules. Because they have low affinity for the
solvent they can get to the surface easily since they have low bond energy in the bulk phase.
The water system prefers to have these molecules at the surface because the placement at the
surface requires less energy than a water molecule -- hydrophobic molecules decrease surface
energy (surface tension) relative to a pure water system.
On the other hand if a particle has a high affinity for the solvent phase (hydrophilic) it will tend to
remain in the bulk solution because of its strong bond with water. In fact, this bonding makes the
water bonding stronger and, therefore, there is a larger energy required to get water molecules to the
surface-- therefore, hydrophilic molecules increase surface tension, e.g. salts such as NaCl.
Hydrophilic substances are those that can enter into a charged interaction with water molecules
6.2.2.4 Isoterm models for solutions
There are four common models for isotherms:
1. Linear model (Fig. 6.16)
Fig. 6.16 Linear model for solutions
qe= mass of material adsorbed (at equilibrium) per mass of adsorbent;Ce= equilibrium concentration
in solution when amount adsorbed equals qe
2. Langmuir Isotherm
This model assumes monolayer coverage and constant binding energy between surface and
adsorbate (Fig. 6.17).
55
KQ0a C e
qe 
1  KCe
(6.3)
Fig. 6.17 Langmuir isotherm for solutions
qe= mass of material adsorbed (at equilibrium) per mass of adsorbent;
o
Q a: represents the maximum adsorption capacity (monolayer coverage)(g solute/g adsorbent);
Ce= equilibrium concentration in solution when amount adsorbed equals q e;
K : Langmuir equilibrium constant
3. BET (Brunauer, Emmett and Teller) isotherm
This is a more general, multi-layer model (Fig. 6.18). It assumes that a Langmuir isotherm applies to
each layer and that no transmigration occurs between layers. It also assumes that there is equal
energy of adsorption for each layer except for the first layer.
K BCe Q0a
qe 
(6.4)
Cs  Ce 1  K B  1Ce / Cs 
CS =saturation (solubility limit) concentration of
the solute. (mg/liter)
KB = a parameter related to the binding intensity
for all layers.
Note: when Ce << CS and KB >> 1 and K = KB/Cs
BET isotherm approaches Langmuir isotherm.
Fig. 6.18 BET isotherm for solutions
1.
Freundlich Isotherm
For the special case of heterogeneous surface energies (particularly good for mixed wastes) in
which the energy term, “KF”, varies as a function of surface coverage we use the Freundlich model
(Fig. 6.19).
56
q e  K FC
1
n
e
(6.5)
n and KF are system specific constants.
Fig. 6.19 Freundlich isotherm for solutions
Questions:
1. Define disperse systems and give some examples.
2. How are clasified dispersions according to the average particle size.
3. Define the term Adsorption.
4. Characterize physical and chemical adsorption.
5. Describe adsorption isoterms valid for interaction at gas/solid interface.
6. Characterize wetting on the solid surfaces.
7. Describe isoterm models for chemical solutions.
References:
1. Tadros, T. F.: Formulation of Disperse Systems: Science and Technology. 2014, Wiley, ISBN:
978-3-527-67830-3.
2. Toth, J.:Adsorption: Theory, Modeling, and Analysis. 2002, Marcel Dekker, ISBN:
0824707478.
3. Masel, R. I.: Principles of Adsorption and Reaction on Solid Surfaces. 1996, Wiley, ISBN:
0471303925
57
7 BUILDING MATERIALS
Chapter mission
This chapter gives basic information about the composition and properties of compounds used as
the building materials, physical - chemical processes of setting and hardening of inorganic binders,
which are essential for understanding and the right selection of building materials for various
purposes in building industry.
Chapter objectives
After studying this chapter you should know to:

describe the composition and properties compounds used as building materials,

characterize kind of binders used in building industry,

write the basic chemical equations of the binder preparation,

explain processes of setting and hardening of inorganic binders,
7.1 The important compounds used as building materials
7.1.1 Silicon dioxide
Silicon dioxide, or silica, is the most abundant of all minerals and is a component of the group of
silicates, among which are the clays. It is present in cements in a state of combination, and is derived
from the clay or shale used as a raw material. It is the main component of pozzolanas and is used in
the form of sand as a constituent of mortar.
Pure crystalline silica occurs in nature as quartz, forming hexagonal crystals belonging to the trigonal
system. The specific gravity is 2.654 and hardness 7. Modification changes of quartz, otherwise
known as low-quartz or α-quartz, in dependence on temperature are illustrated in schema
870C
-Quarz 
  573C
-Quarz
1470C
-Tridymite

-Cristobalite
  1720C
melting
α quartz is transformed at 573° into high-quartz or β-quartz with a 2 per cent increase in volume. At
870° tridymite becomes the stable form and at 1470° this is further transformed to cristobalite which
melts at 1720°. These transformations to tridymite and cristobalite are very sluggish and, unless
58
heated very slowly, quartz simply passes into a viscous glass at about 1500-1600°; the true meltingpoint of quartz is not known accurately, but is probably below 1470°. There are also a number of other
minor modifications of silica which need not be discussed here and numerous variations in the order in
which the transformations can occur. The question whether tridymite is a stable phase or whether
alkali or water are needed to stabilize the structure remains controversial. Hydrothermal studies have
suggested that tridymite is stable but other studies indicate that impurities are needed. The tridymite
and cristobalite structures are closely related.
Quartz is chemically a very inert substance at ordinary temperatures, but when strongly heated it
reacts vigorously with bases. Sand behaves in mortars as an indifferent material, being bound
together mechanically by the particles of lime or cement, but not entering into chemical reaction with
them. At high temperatures silica behaves as an acid oxide and is capable of combining with bases to
form silicates, and of expelling other acids from their compounds.
Silicon dioxide is an extremely stable solid. Silicon dioxide consists of a tetrahedral network of
covalently bonded silicon and oxygen atoms in which no individual SiO 2 units or molecules can be
distinguished. This structure is analogous to that of diamond.
Silicon dioxide is present in quartz, sand, sandstone, flint, jasper, amethyst and agate. Silicon dioxide
dissolves in water very slowly, and only to a slight extent to form orthosilicic acid, H 4SiO4, or
metasilicic acid, H2SiO3. Orthosilicic acid is hydrated metasilicic acid, H2SiO3(H2O). When either of
these acids is completely dehydrated, silica forms. This dehydration reaction is the source of the silica
in petrified wood and in the skeletons of small one-celled animals called diatoms.
Pure quartz crystals and some other crystalline solids generate an electric potential if they are
mechanically deformed in a certain direction. Substances having this property are piezoelectric.
Since silicon dioxide is an ingredient of glass, hydrofluoric acid also reacts with glass. This reaction is
used to etch glass. The surface of a piece of glass is covered with a layer of wax because wax does
not react with hydrofluoric acid. The wax is then removed from certain areas.
7.1.2 Silicates and Silicones; Polymeric Network Structures
Silica (SiO2) reacts with sodium carbonate at about 1300°C to form sodium silicate, Na 2SiO3;
Na2CO3(s) + SiO2(s)  Na2SiO3(s) + CO2(g)
(7.1)
Many metal silicates exist in nature. Silicate ions have polymeric network structures in which each
silicon atom is surrounded by four oxygen atoms in a tetrahedral arrangement as shown in Figure 1.
The SiO4 tetrahedral share CaSiO3. When molten glass is poured on a smooth layer of pure molten
tin, it hardens to plate glass. The surface of the molten tin is perfectly smooth and so is the glass
floating on it. Plate glass produced in this way does not need to be polished after hardening.
59
Figure 1 Structure of silicate chain
When silica is dissolved in an aqueous solution of NaOH, mixture called water glass is obtained:
SiO2(s) + 2NaOH(aq)  “Na2SiO3(aq)I” + H2O (l)
(7.2)
The quotation marks indicate that water glass is not a pure compound but a mixture in which the
sodium and silicon content are variable. Water glass 60 fit60d as a fire retardant on fabrics, as an
adhesive in the manufacture of cardboard cartons, and as an egg preservative that acts by sealing the
pores in egg shells. When an aqueous solution of a metal silicate is treated with strong aqueous acid,
a semisolid substance called silica gel is formed. When silica gel is dried, a porous solid is obtained
that is widely used as a drying agent and as an absorbent for certain gases.
Hundreds of silicate minerals also contain aluminum. These minerals are called aluminosilicates.
They include such minerals as feldspars and clays. There are dozens of other silicate minerals of
general interest, 60 fit60d mica and various gems, including zircon, topaz, emerald, aquamarine, and
tourmaline. Micas are silicates in which the molecules form layers of SiO 4 tetrahedra, so mica has
a layer structure.
Clay contains a considerable amount of water. When the water is driven off by strong heating,
a glasslike material called a ceramic is produced. Ceramics have many applications, from fine
porcelain vases and china to the tiles that cover space shuttles to protect them from the fierce heat
they generate when reentering the earth’s atmosphere.
Another important class of oxygen-containing compounds of silicon are the silicones. These are
synthetic polymeric substances containing alternating silicon and oxygen atoms in – O – Si – O –
chains. Hydrocarbon groups, 60 fit60d methyl groups (-CH3), are attached to the two remaining
bonding positions of each silicon atom. Silicones are chemically inert, water repellant, heat resistant,
and good electric insulators. These properties make them useful as protective coatings, lubricants,
insulators, sealants, and adhesives. Silicones are biologically inert. They are therefore used
cosmetically for certain surgical implants.
7.1.3 Carbonates and Hydrogen Carbonates
Carbon forms two common oxoanions; one is carbonate ion and the other is hydrogen carbonate ion,
commonly named bicarbonate ion. Examples of naturally occurring metal carbonates include CaCO 3,
Na2CO3, MgCO3 (magnesite), and FeCO3 (siderite).
Sodium carbonate, or soda ash, is a constituent of some soaps and soap powders. Sodium
carbonate decahydrate, Na2CO3.10 H2O, is known as washing soda. Like soda ash, 60 fit used in
60
laundering. Sodium hydrogen carbonate, NaHCO3, is baking soda, a principal component of baking
powders. Magnesium carbonate 61 fit61d in toothpaste, in talcum powder, and for glass manufacture.
Calcium carbonate is found principally as limestone, but 61 fit also a component in other minerals,
including marble, calcite, pearls, coral, and chalk.
Limestone is not soluble in water, but 61 fit is treated with water and carbon dioxide, it slowly converts
to a solution of calcium hydrogen carbonate:
CaCO3(s) + H2O (l) + CO2(g)  Ca (aq) + 2HCO3 (aq)
2+
-
(7.3)
This reaction is mainly responsible for the formation of limestone caves. When water percolates
through rock containing calcium carbonate in the ground above a cavern, some of the slat dissolves,
and the solution of calcium hydrogen carbonate may start dripping from the ceiling of the cavern. As
water evaporates from this solution in an open cavern, the solution of calcium hydrogen carbonate
converts back to solid calcium carbonate:
Ca (aq) + 2HCO3 I(aq)  CaCO3(s) + H2O(g) + CO2(g)
2+
-
(7.4)
7.1.4 Calcium oxide - Lime
The burning of lime (Calcium oxide - CaO) consists essentially in the calcinations o one of the
naturally occurring forms of calcium carbonate at a sufficiently high temperature to decompose the
calcium carbonate. Calcium carbonate occurs naturally in forms of varying purity as marble, chalk, and
limestones.
The reaction
CaCO3  CaO + CO2
(7.5)
attains a dissociation pressure equal to atmospheric pressure at 894°C. The decomposition is carried
out in shaft and rotary kilns, or by more primitive forms of burning, to produce the various classes of
limes. The composition and properties of the resulting product depend both on the composition of the
chalk or limestone used and on the efficiency with which the burning process is carried out.
It is a white, amorphous and friable material with a specific gravity varying from 3.08 to 3.30, the value
being greater the higher the temperature at which the lime is burnt. The chemical reactivity of lime, like
its specific gravity, depends on the temperature at which it is prepared. Lime which has not been
heated above 1000° ‘slakes’ instantly on the addition of water, whilst a more strongly ignited product
requires a considerable time for hydration. A similar inactivity is observed towards steam or carbon
dioxide.
Fat (high-calcium) limes are obtained from the calcinations of limestones of a high degree of purity
and contains 95 per cent and upwards of calcium oxide. On the addition of water they slake rapidly
with the evolution of much heat, the lumps breaking down to form a lime putty.
61
Hydraulic limes are obtained from the burning of limestones that contain a proportion of clay. At the
burning temperature of 1000-1200° the reaction products include β 2CaO . SiO2, and less basic
silicates, 2 CaO.Al2O.SiO2, 4CaO.Al2O3.Fe2O3 and calcium aluminates. The hydraulic properties are to
be attributed to the dicalcium silicate and in some degree to the aluminates. There exists an almost
continuous series of limes varying from fat limes with a content of alumina and silica below 1-2 per
cent to eminently hydraulic limes with up to 50 per cent of these constituents.
Fat limes when used as mortars harden only by absorption of carbon dioxide from the air with
formation of calcium carbonate, a process which is very slow to penetrate beyond the surface. They
will not set under water. Hydraulic limes, by virtue of the alumina and silica compounds they contain,
harden slowly without such adventitious aid and can be used under water
7.1.5 Calcium hydroxide
The product of the hydration of lime by water is calcium hydroxide
CaO + H2O  Ca(OH)2 + 67 kJmol
-1
(7.6)
This reaction 62 tis62d for lime milk preparation. Calcium hydroxide decomposes into lime and water
vapour at about 400° in a free atmosphere.
62 tis therefore not found in cement clinker immediately after burning. 62 tis usually produced to
a small extent during the grinding of cements by the action of water released from the added gypsum.
62 tis obtained in the slaking of lime as a white, amorphous powder. If crystallised slowly, as in the
hydration of Portland cement, it forms large, well-developed crystals. Hexagonal plate crystals of
calcium hydroxide can be observed in set Portland cements, and from old cement briquettes very
good crystals have been obtained. Calcium hydroxide reacts with carbon dioxide to form calcium
carbonate:
Ca(OH)2 + CO2 = CaCO3 + H2O
(7.7)
and this reaction is the cause of the hardening of high-calcium lime mortars. It only takes place in the
presence of moisture.
7.1.6 Magnesium oxide
Magnesium oxide or magnesia, MgO, is only present in small quantities in Portland cement, being
derived from magnesium carbonate present in the original limestone in the form of dolomite,
CaCO3.MgCO3, or to a small extent from the clay or shale.Thermal decomposition
of dolomite
proceeds in two steps as follows
600C
CaCO3.MgCO3 
CaCO3 + MgO + CO2
(7.8)
900C
CaCO3 + MgO  CaO + MgO + CO2
62
(7.9)
Pure magnesium oxide, if not heated to too high a temperature, possesses distinct hydraulic
properties. Light, porous magnesia combines with water to form the hydroxide Mg(OH) 2, but without
setting, whilst dense magnesia prepared at full red heat sets, yielding a coherent mass, although of
low tensile strength. Magnesite, MgCO3, when dead burnt at 1400-1500° yields a product which reacts
only very slowly with water; 63 tis used as a refractory.
Magnesia is soluble to some extent in lime at high temperatures and lime is rather less soluble in
magnesia. These solid solutions decompose at lower temperatures and for all practical purposes it
may be assumed that the compounds crystallize independently from their fused mixtures. If the double
,
carbonate, dolomite is ignited at 650-750° a mixture of magnesium oxide and calcium carbonate is
obtained.
The melting-point of MgO is 2852°. 63 tis isometric, shows perfect cubic cleavage, and has
a refractive index of 1.737. It occurs in nature as the mineral periclase.
When magnesia and a solution of magnesium chloride react together
is obtained the product –
magnesium oxychloride cement. Magnesite is calcined so as to give a lightly burnt reactive product
and the ground material mixed as required with a strong solution (about 20 per cent anhydrous salt) of
magnesium chloride. Combination of the magnesia and magnesium chloride takes place with the
development of heat and results in the formation of magnesium oxychloride 3MgO.MgCl 2.11H2O. In all
probability the magnesium oxide and chloride first react in solution and a supersaturated solution of
the oxychloride is formed from which the solid separates out.
7.1.7 Sulphate
Gypsum, CaSO4.2H2O is one of important compounds of sulphate typ. During the partial or complete
dehydration of gypsum are produced the cementing materials-plasters. When finely ground gypsum is
heated to about 150° in open the product obtained is the hemihydrate CaSO 4.½H2O.
CaSO4. 2 H2O  CaSO4. 0,5 H2O + 1,5 H2O
(7.10)
Hardening process of hemihydrate CaSO4.½H2O is connected with its hydratation
2(CaSO4. 0,5 H2O) + 3 H2O  2(CaSO4. 2 H2O) H = -38,6 kJmol
-1
(7.11)
If calcination is carried out at a higher temperature the remaining water of crystallization is lost and first
soluble and then insoluble anhydrite is formed. Soluble anhydrite CaSO4, which forms the main basis
of the anhydrous gypsum plasters is predominant when the temperature is allowed to rise to 190-200°.
63cta very hydroscopic and absorbs water vapour very rapidly to form the hemihydrate. When soluble
anhydrite is heated to higher temperatures its reactivity steadily diminishes until at about 600° the
product is relatively inert and is known as insoluble anhydrite. In requires the addition of a suitable
catalyst, an accelerator, to render it reactive. When gypsum is calcined at 1100-1200° some
63
dissociation into sulphur trioxide and lime occurs, leaving free lime dispersed in the product to 64cta s
an accelerator of set. The quality of the gypsum coming from a flue gas desulphurization plant
Ca(OH)2 + SO2  Ca SO3 + H2O
(7.12)
SO2 + Ca(OH)2 + l/2 H2O  Ca SO4 + H2O
(7.13)
is comparable to that of natural gypsum. It can be used as a building material especially as: building
plaster; plaster board; plaster wall plates; plaster slabs; and mortar for the mining industry. In addition
it can be used as an additive in the cement industry (solidification control) or it can be further
processed to a cement-like binding agent in the concrete and building material industry.
If flue gas desulphurization gypsum is used in the plaster industry a larger amount of reagents or setup agents is necessary because of the fineness and grain form of the crystals in contrast to natural
gypsum. The flue gas desulphurization system has to be designed for forced oxidation in order to
convert practically all of the dissolved calcium sulphite into calcium sulphate, which crystallizes as
gypsum (CaSO4 . 2H20).
7.1.8 Aluminium oxide
Aluminium oxide or alumina, Al2O3, occurs in nature as corundum (α Al2O3). Alumina in a combined
state is an important constituent of cements, in which it behaves as an acid. It occurs combined with
silica in all clays and in these it may be regarded as a base. It also occurs free in a hydrated form,
mixed with a proportion of ferric oxide and smaller amounts of titania and silica, in bauxite.
7.1.9 Ferric oxide
The oxides of iron only occur to a relatively small extent in Portland cements, being derived from the
clay or shale, but are an important constituent of high-alumina cements where they are derived from
the bauxite. Ferrous oxide (FeO) does not usually occur in more than small amounts, ranging from a
trace to about 0.4 per cent, in Portland cement, but it is present in high-alumina cement in amounts
from 2 to 3 per cent upwards.
Ferric oxide (Fe2O3) is a constituent of both Portland and high-alumina cements. It resembles alumina
in character and acts as an acid radical in cements, being combined with the bases (mainly CaO)
present.
7.2 Portland
cement
–
Manufacture,
types
properties,
and
specifications
Portland cement has become one of the most important construction materials during the last 150
years or so, primarily because concretes can be used advantageously for so many different purposes.
Cement is produced from an appropriate combination of a lime-containing material, such as limestone,
and clayey materials by burning this mixture, then grinding the resulting clinker along with a small
amount of gypsum. A typical tiny Portland cement grain consists of numerous microscopic crystals
64
called clinker minerals. About three-fourths of these minerals are calcium silicates (alite and belite),
the rest being calcium aluminate, iron compounds, and minor constituents. The amounts of these
minerals can be measured directly, or they can be calculated from oxide analysis data.
7.2.1 Manufacture
The raw materials of Portland cement consist principally of limestone or some other lime-containing
material such as marl, chalk, or shells, and clay or shale or some other argillaceous (clayey) material
such as ashes or slag. Calcium oxide, silica, alumina, and ferric oxide must be present within narrowly
defined limits, and other constituents, such as magnesia and alkalies, must not exceed specified
limits.
All raw materials must be ground to an impalpable powder and intimately mixed before burning. In the
dry process, all grinding and blending operations are done with dry materials and the final mixing is
accomplished chiefly in the grinding mills. In the wet process, the final grinding and blending are
brought about in a water slurry, and mixing is accomplished both in the grinding mills and by stirring in
large vats.
The maximum temperature in the kiln is between 1250 and 1900°C. During the passage of the mixture
down the length of the kiln, several reactions take place at various temperature levels. These are
summarized in Table 1.The applied temperature is such that only a minor proportion of the reacting
mixture is in the molten condition. The liquid that forms during the burning process causes the charge
to agglomerate into nodules of various sizes, usually in (5-25 mm) in diameter, and characteristically
black, glistening hard. This material is known as Portland cement clinker. The charge drops from the
end of the kiln into some form of cooler and is then ground usually with 3-5 % of gypsum (CaSO4 .
2H2O), to a fine powder.
7.2.2 Composition of portland cement clinker
The composition of a Portland cement depends on the composition and proportions of the raw
materials as well as on details of the manufacturing process, such as burning temperature, cooling
rate.
Table 1 Reaction in the Kiln
Temperature (°C)
Process
Thermal change
100 and below
500 and above
900 and above
Evaporation of free water
Dehydroxylation of clay minerals
Crystallization of products of clay
Mineral dehydroxylation
Decomposition of CaCO3
Reaction between CaCO3 or CaO and
aluminosilicates
Beginning of liquid formation
Further liquid formation and completion
of formation of cement compounds
Endothermic
Endothermic
900-1200
1250-1280
Above 1280
65
Exothermic
Endothermic
Exothermic
Endothermic
Probably endothermic
on balance
7.2.2.1 Oxide Composition
When producing cement, the chemical composition of cement is controlled by the content of silica
(SiO2), lime (CaO), alumina (Al2O3) and iron (Fe2O3). These oxides become characteristic clinker
minerals which during the addition of gypsum will be ground to cement. About 95% of Portland cement
clinker is made of certain compounds of four oxides. Other so-called minor constituents or impurities
include, among others, magnesia, sodium, and potassium oxides (the alkalis); titania; sulfur;
phosphorous, and manganese oxides.
The chemical composition of a Portland cement is evaluated.
by Silica modulus =
SiO2
Al 2 O3  Fe 2 O3
(7.14)
where the oxides represent the amounts determined by chemical analysis of the cement expressed in
percent by weight. Values of the silica modulus may vary approximately within 1.7 to 2.5, the usual
value being within 2.0 to 2.5.
Iron modulus (or alumina ratio) =
Al 2 O3
Fe 2 O3
(7.15)
It usual value in standard Portland cements is between 1 and 3.
A third ratio that can be used is
Lime saturation factor =
CaO  0.7 SO3
2.8SiO2  1.2 Al 2 O3  0.65Fe 2 O3
(7.16)
This factor represents the ratio of the quantity of lime actually available in the cement to the maximum
permissible lime content. The lime saturation factor of commercial Portland cements is around 0.850.90.
Portland cement is made up of four main clinker compounds: Tricalcium silicate 3 CaO.SiO2 (allite)
(C3S), Dicalcium silicate 2 CaO.SiO2 (belite)(C2S), Tricalcium aluminate 3 CaO.Al2O3 (C3A)
and Tetra-calcium aluminoferrite (C4AF). Where C stands for calcium oxide, S for silica, A for
alumina, and F for iron oxide. The equilibrium is reached at the clinkering temperature and maintained
during cooling.
Small amounts of uncombined lime, magnesia, alkalis and minor amounts of other elements (titanium,
manganese etc.) are also present. The composition of Portland cements falls within the range of 60 to
67 percent lime, 17 to 25 percent silica, 2 to 8 percent alumina, and 0 to 6 percent iron oxide together
with 1 to 7 percent sulphur trioxide, derived mainly from the added gypsum, 0.1 to 5 percent
magnesia, and 0.1 to 1.5 percent alkalis.
The two calcium silicates C3S and C2S form about 70-80% of a Portland cement. The internal structure
of these silicates can be described as SiO 4 tetrahedrons separated and connected by calcium ions. In
reality the clinker minerals in a Portland cement are not in the form of pure compounds. The calcium
silicates, for instance, contain small amounts of alumina, magnesia, and possibly some other oxides.
Each of the clinker minerals has important individual characteristics. Alite paste attains the greater
66
part of its strength in a week and little increase occurs at longer ages. The heat development
produced by reactions between alite and water (“heat of hydration”) is also quite intensive. Belite
produces little strength until after several weeks, but gains steadily in strength at later ages until it
approaches equality with alite. Also, the development of its heat of hydration is much slower.
Tricalcium aluminate alone attains very little strength, but in mixes with calcium silicates shows a
favorable effect on early strength development. It is not clear yet how tricalcium aluminate contributes
to the hardening and strength. Due to the presence of the ferrite phase in Portland cement is derived
its characteristic gray color. Portland cement without iron is white.
The term “minor” refers to the quantity of these constituents in Portland cement rather than to their
importance. They may have significant effects on the quality of cement.
To obtain high early strengths, relatively high alite and C3A (up to 15% potential) contents are
desirable. A Portland cement of low heat of hydration, such as would be used in a massive dam
construction, requires a relatively high belite content (at least 40%) at the expense of the alite
(maximum 35%) and C3A (maximum 7%) contents.
The different standard specifications for Portland cement contain different requirements of chemical
and physical properties: MgO, SO3, alkalis as Na2O, loss in ignition, insoluble residue, Bogue
composition, fineness (Blaine), soundness, autoclave expansion, compressive strength and initial and
final setting. Besides these requirements a standard specification can contain more specific
requirements.
C3S

The head clinker component in cement, typical more than 50 %

Quick development of strength - C3S reacts more quickly than C2S

High contribution to the final strength

Resistant to sulphur attack

25 weight % water bind under hydration of C3S

Heat development: 500 kJ/kg

C3S = 4.071*CaO-(7.600*SiO2+6.718*Al2O3+1.430*Fe2O3+2.852*SO3) according to
Bogue's methods
C2S
C3A

Second clinker component in cement, between 10 - 60 %

Slow development of strength - C2S reacts more slow than C3S

High contribution to the final strength

Resistant to sulphur attack

20 weight % water bind under hydration of C2S

Heat development: 250 kJ/kg

C2S = 2.87*SiO2-0.754*(Ca3SiO5) according to Bogue's methods

Range in the cement between 3 - 10 %

High contribution to the early strength

Low contribution to the final strength
67
C4AF

Not resistant to sulphur attack

40 - 210 weight % water bind under hydration of C3A

Fast and high heat development: 900 kJ/kg

C3A = 2.65*Al2O3-1.69*Fe2O3 according to Bogue's methods

Range in the cement between 5 - 10 %

Small contribution to the development of strength

37 - 70 weight % water bind under hydration of C4AF

Moderate to low heat development: 300 kJ/kg

Hydration of C4AF + 13H = C4AFH13

C4AF = 3.04*Fe2O3 according to Bogue's methods
The chemical composition of cement influences the characteristics of cement. The clinker mineral C 3S
is the head clinker component in cement. The strength developed by Portland cement depends on its
composition of C3S, C2S, C3A, C4AF and the fineness (Blaine) to which it is ground.
7.2.3 Technically important properties of portland cement
Portland cement must have a number of qualities to be acceptable. In addition to chemical
composition, a Portland cement, to be acceptable, should comply with the requirements of fineness,
time of setting, soundness, air content, and strengths, to which other requirements, heat of hydration
for instance, may be added. In order to fulfill all these requirements, pertinent standards usually
contain specifications for the chemical composition and for the physical properties of
cements. Most of the chemical requirements for Portland cement
Portland
is in connection with the
composition of Portland cement clinker. From the standpoint of the user, physical properties provide
more direct information concerning the acceptability of a cement. The physical properties of Portland
cement are fineness time of setting, soundness, air content, and strength. There are also tentative
requirements concerning heat of hydration and false set.
Influence of Portland cement on concrete properties
Effects of cement on the most important concrete properties are presented in Table 3. Cement
composition and fineness play a major role in controlling concrete properties. Fineness of cement
affects the placeability, workability, and water content of a concrete mixture much like the amount of
cement used in concrete does.
Cement composition affects the permeability of concrete by controlling the rate of hydration. However,
the ultimate porosity and permeability are unaffected. The coarse cement tends to produce pastes
with higher porosity than that produced by finer cement. Cement composition has only a minor effect
on freeze-thaw resistance. Corrosion of embedded steel has been related to C 3A content. The higher
the C3A, the more chloride can be tied into chloroaluminate complexes—and thereby be unavailable
for catalysis of the corrosion process.
68
Table 3 Effects of cements on concrete properties.
Cement Property
Cement Effects
Placeability
Cement amount, fineness, setting characteristics
Strength
Cement composition (C3S, C2S and C3A), loss on ignition,
fineness
Drying Shrinkage
SO3 content, cement composition
Permeability
Cement composition, fineness
Resistance to sulfate
C3A content
Alkali Silica Reactivity
Alkali content
Corrosion of embedded steel
Cement Composition (esp. C3A content)
7.2.4 Hydration of portland cement
The term hydraulic cement refers to a powdery material that reacts with water and as a result,
produces a strong as well as water-insoluble solid. Portland cement and high-alumina cement, among
others, are hydraulic cements. Gypsum in the amount of 3-6% should be added to the clinker to slow
down this reaction and control the time of setting as well as the strength development.
Both the setting and the hardening of a Portland cement paste are the result of a series of
simultaneous and consecutive reactions between the water and the constituents of the cement. These
reactions all together are covered by the term hydration of portland cement. The two most important
chemical reactions during the early part of the hydration are (1) the reaction between the C 3A as well
as the gypsum of the cement and water; and (2) the hydration of alite in the cement and water.
C3A reacts very rapidly with water, giving two hydrated products:
2C3A + 21H = C4AH13 + C2AH8
(7.17)
These forms platelets within the cement, and convert to C3AH6, which forms very quickly, and is
responsible for the initial formation of a crystalline network. In the presence of free lime in the cement,
the formation of C4AH13 is favoured. This slows the formation of C3AH6, but even so the formation of
C4AH13 can causing the cement to set too quickly. To avoid speed setting is gypsum added the cement
and the mineral ettringite is formed on hydration:
C3A + 3CaSO42H2O + 25-26H2O = Ca6Al2O6(SO4)31-32H2O
(7.18)
Ca3SiO5 + (y+z) H2O = z Ca(OH)2 + Ca(3-z)SiO(5-z).y H2O
(7.19)
Hydration of C3S
are to some extent dependent on the presence of C3A and gypsum. Both C3A and gypsum stimulate
the hydration of C3S. Also Alkalis have some influence at the hydration.
On hydration, C2S shows similar behaviour to C 3S, but is slower to react. It does however continue to
hydrate late in the setting period, and may then contribute to the strength of the cement.
69
When the Portland cement is insufficiently retarded, the reactions are too fast, causing an undesirable
“quick set”. The predominant reactions at later ages are the hydrations of the calcium silicates, which
continue for many months at a diminishing rate.
In the course of hydration every Portland cement grain breaks up into a billion particles, forming
mostly a poorly crystallized, porous solid, so-called CSH (calcium silicate hydrate) gel. The
mechanism of gel formation is as follows: On contact with the still unhydrated part of a cement grain,
water dissolves a portion of it; this solution diffuses out from the grain surface toward larger spaces
through the very small pores of the solid shell of previously created hydration products around the
cement grains; then the new hydration products precipitate from the solution. The very fine texture and
the resulting high specific surface are the most significant characteristics of the CSH gel. One cubic
3
centimeter solid volume of cement can develop at least 2 cm hydration products. Therefore, as a
consequence of hydration, the volume of solids within the boundaries of paste specimen increases,
producing interlocking laths and a reduction in the overall porosity of the paste. These are the primary
sources of the stiffening and strength development of the cement paste.
The complex reactions between water and cement, thus the chemical composition as well as the
structure of the hardened paste, vary to a certain extent with time, temperature, water-cement ratio,
and several other factors. Regardless of the conditions, however, the hardened paste always contains
a considerable amount of pores of different sizes, namely, gel pores, capillary pores, and air voids.
The hydration of cement compounds is exothermic. The heat developed is called the “heat of
hydration”. The measurement of heat evolution is particularly suitable for the investigation of the early
stages of hydration. During a short period beginning when Portland cement and water are first brought
into contact at room temperature, and during the time of mixing, relatively rapid chemical reactions
occur primarily between the C3A of the cement and water. The water rapidly by hydrolysis of alite as
well as of calcium aluminate, by CaSO4 . 2H2O from the gypsum, and by other compounds. Eventually
crystalline Ca(OH)2 appears. Then the gypsum reacts with calcium aluminates to form solid calcium
aluminate sulfate hydrates, for instance, ettringite, which act as a protective coating on the surfaces of
the cement particles. This period, which has been called the dormant period, normally lasts 40-120
min., depending on the characteristics of cement. When the portland cement is insufficiently retarded,
the time of initial setting is considerably less than 1 hr.
It appears likely that the chemical reactions that are predominant during the regular setting period are
(1) the reaction between C3A and CaSO4 to form a calcium aluminate sulfate hydrate at a slow rate as
the liquid part in the paste is saturated with calcium hydroxide and gypsum; and (2) the hydration of
calcium silicates, mainly of alite present in the cement, producing a poorly crystallized, porous but
stable gel, called CSH (calcium silicate hydrate) gel, or sometimes tobermorite gel. In this gel are
embedded several more or less well-crystallized hydrates, mainly calcium hydroxide, and unhydrated
cement particles. Water is also present in the system in various states. The produced gel particles are
so small that they are invisible under an ordinary microscope; they are visible, however, under an
electron microscope.
After final set, chemical reactions continue at a diminishing rate until one or more of the conditions
necessary to reaction are lacking. This stage of hydration is called the hardening process during
70
which the predominant reaction is the continuing hydration of the calcium silicates. The decrease in
rate is the result of two effects: (1) the surface area of unhydrated cement particles decreases as the
smaller particles become completely hydrated and the larger particles become smaller; and (2) a layer
of CSH gel forms on the surfaces of the cement particles. slowing down further reaction by forming a
protective coating.
2C3S + 6H = C3S2H3 + 3Ca(OH)2
(7.20)
where C3S2H3 represents a CSH gel; or
C3A + H10 + CaSO4 . H2O = C3A . CaSO4 . H12
(7.21)
where the right side of the equation is calcium aluminate monosulfate hydrate.
The main difference between the hydration of tricalcium silicate and that of dicalcium silicate is that the
former develops more Ca(OH)2. Also x-ray analysis indicates that the CSH gel developed from
dicalcium silicate has lower CaO/SiO2 ratio than that developed from tricalcium silicate during the first
half of the hydration process, but this difference becomes negligible as the process approaches
complete hydration. In the early stages of hydration of Portland cement, C 3A usually produces
ettringite, that is, calcium aluminate trisulfate hydrates. The compositions of the other hydrated calcium
aluminates and ferrites are quite complex.
Questions:
1. Characterize the properties of silicon dioxide.
2. Describe production of lime.
3. Describe the process of production and hardening of plasters.
4. Describe manufacture of Portland cement.
5. Characterize the main clinker compounds.
6. Characterize hydration of Portland cement.
7. Name the basic components of concrete.
References:
1. Varghese, P. C.: Buildings material. 2005, PHI Learning Pvt. Ltd., ISBN: 8120328485
2. UK Limes Team: Hydraulic Lime Mortar: for Stone, Brick and Block Masonry. 2003, Donhead,
ISBN: 978187339464
3. Day, K. W.: Concrete Mix Design: Quality Control & Specification. 2013, Design Media, ISBN:
9780415504997.
71
8 CONCRETE DETERIORATION
Chapter mission
This chapter presents an overview of factors that have negative impact on concrete deterioration.
Special attention is paid to influence of acid, carbonatation and sulphate on concrete leaching and
destruction. Also corrosion of reinforcing steel in concrete is described. The end of the chapter is
focused on deterioration of concrete from freeze thaw actions.
Chapter objectives
After studying this chapter you should know to:

characterize types of chemical reaction in concrete affecting durability,

describe concrete corrosion by leaching, carbonatation andsulfate attack,

explain corrosion of concrete reinforcement.
8.1 Concrete durability
Durability is dedined as the quality of materials or structures of continuing to be useful after an
extended period of time and usage. Chemical reaction in concrete affecting durability can be devided
as follows:

Reactions with larger products than the reactants, causing expansion and as a result cracking.

(Transformation) reactions with products having a lower strength (adhesion).

Dissolution (leaching) or other chemical reactions, leading to higher porosity and therefore lower
strength.
8.2 Concrete corrosion
8.2.1 Leaching (dissolution of concrete)
Cement hydration is a dissolution-precipitation reaction: Clinker partly dissolves until the water
becomes (over) saturated with respect to the CSH and Ca(OH) 2
phases, and these phases
precipitate. Fresh tap or distilled water is undersatured with respect to CSH and Ca(OH) 2 . Therefore,
these phases will dissolve (go into solution) in tap water. This process is called leaching of concrete,
and will result in a higher porosity and therefore lower strenght. Among important factors belongs:

Initial porosity/permeability of concrete (water needs to go in);

Properties of water (Ca and CO3 concentrations, pH);

Type and rate of water flow.
2+
2-
Order of decreasing dissolution rate: Ca(OH) 2 > monosulfate > ettringite > CSH
72
8.2.2 Acid attack
Acids (pH < 7) have a leaching action in principle similar to that of water, but a much stronger one.
Acids dissolve Ca(OH)2 according to the following reaction:
2 HX + Ca(OH) 2 -> CaX2 + 2H2O
(8.1)
(X is the negative ion of the acid)
Aggresive acids are: HCl, HNO3, H2SO4 (sulphuric acid), Sulphuric acid is very damaging to concrete
as it causes both acid and sulphate attack.
Examples of other corrosive solutions: ammonium sulphate, sodium sulfate, magnesium sulfate,
sodium nitrate, CaCl2, MgCl2, and NaCl.
Acid attack can be diagnosed by two main features:
•
Absence/depletion of calcium hydroxide in the cement paste,
•
Surface dissolution of cement paste exposing aggregates.
8.2.3 Carbonation of concrete
Carbon dioxide (CO2) from the environment diffuses into concrete and dissolves in the pore solution of
2-
2+
cement paste, producing CO3 ions, which react with Ca to produce calcite (CaCO3):
Ca(OH) 2 + CO2 → CaCO3 + H2O
(8.2)
If the concrete is too dry (RH <40%) CO 2 cannot dissolve and no carbonation occurs. If on the other
hand it is too wet (RH >90%) CO2 cannot enter the concrete and the concrete will not carbonate.
Optimal conditions for carbonation occur at a RH of 50% (range 40-90%).
Carbonation can have positive or negative effects on concrete durability:
•
the compressive strength of carbonated concrete may be higher than that of uncarbonated
concrete
•
In carbonated zones in concrete the pH falls to 8 or lower, which is insufficient to maintain the
protective layer on the reinforcement steel → Fe-corrosion.
8.2.4 Corrosion of concrete reinforcement
In order to understand the mechanisms behind corrosion of reinforcing steel in concrete, one has to
examine the chemical reactions involved. In concrete, the presence of abundant amount of calcium
hydroxide and relatively small amounts of alkali elements, such as sodium and potassium, gives
concrete a very high alkalinity-with pH of 12 to 13. It is widely accepted that, at the early age of the
concrete, this high alkalinity results in the transformation of a surface layer of the embedded steel to a
tightly adhering film, that is comprised of an inner dense spinel phase in epitaxial orientation to the
steel substrate and an outer layer of ferric hydroxide. As long as this film is not disturbed, it will keep
the steel passive and protected from corrosion.
When a concrete structure is often exposed to deicing salts, salt splashes, salt spray, or seawater,
chloride ions from these will slowly penetrate into the concrete, mostly through the pores in the
hydrated cement paste. The chloride ions will eventually reach the steel and then accumulate to
73
beyond a certain concentration level, at which the protective film is destroyed and the steel begins to
corrode, when oxygen and moisture are present in the steel-concrete interface.
Once corrosion sets in on the reinforcing steel bars, it proceeds in electrochemical cells formed on the
surface of the metal and the electrolyte or solution surrounding the metal. Each cell is consists of a
pair of electrodes (the anode and its counterpoint, the cathode) on the surface of the metal, a return
circuit, and an electrolyte. Basically, on a relatively anodic spot on the metal, the metal undergoes
oxidation (ionization), which is accompanied by production of electrons, and subsequent dissolution.
These electrons move through a return circuit, which is a path in the metal itself to reach a relatively
cathodic spot on the metal, where these electrons are consumed through reactions involving
substances found in the electrolyte. In a reinforced concrete, the anode and the cathode are located
on the steel bars, which also serve as the return circuits, with the surrounding concrete acting as the
electrolyte.
Fig. 8.1 Corrosion of concrete reinforcement
Corrosion can also occur even in the absence of chloride ions. For example, when the concrete
comes into contact with carbonic acid resulting from carbon dioxide in the atmosphere, the ensuing
carbonation of the calcium hydroxide in the hydrated cement paste leads to reduction of the alkalinity,
to pH as low as 8.5, thereby permitting corrosion of the embedded steel.
8.2.5 Sulphate attack
Sulfate attack is a chemical breakdown mechanism where sulfate ions attack components of the
cement paste. Sulphate attack can be 'external' or 'internal':
External: due to penetration of sulphates in solution, in groundwater for example, into the concrete
from outside. Sulphates are common in areas of mining operations, paper industries. It may also be
found in soils and waters (ground water, waste water). Common sulphates found in ground water are
calcium, sodium, potassium and magnesium
Sulphates (in solution with water) permeate the concrete and react chemically with:

The cement paste’s hydrated lime Ca(OH)2
 Calcium aluminate C3AnH
74
Formation of gypsum and ettringite results in expansion, stresses, cracking, scaling, loss of bond
aggregate-paste.
Fig. 8.2 Sulphate corrosion
Internal: due to a soluble sulphate source being incorporated into the concrete at the time of mixing,
gypsum (CaSO4.2H2O) in the aggregate, for example. In some cases the gypsum added to cement (to
control setting) can cause sulphate attack.
Calcium hydroxide and alumina-bearing phases of hydrated portland cement are most vulnerable to
attack by sulfate ions.
Expansive reaction (stress-formation):
C3A + CaSO4.2H2O → C4ASH12 (monosulfate) → C6AS3H32 (ettringite)
+
(8.3)
2+
Depending on the cation type associated with the sulphate solution (Na or Mg ), both Ca(OH)2 and
C-S-H present in the hydrated portland cement paste may be converted to gypsum by sulphate attack:
Na2SO4 +Ca(OH) 2 + 2H2O → CaSO4.2H2O + 2NaOH
(8.4)
or:
MgSO4 +Ca(OH) 2 + 2H2O → CaSO4.2H2O + Mg(OH)2
3MgSO4 + 3C-S-H + 8H2O → 3CaSO4.2H2O + 3Mg(OH) 2 + 2SiO2.H2O
(8.5)
(8.6)
Fig. 8.3 Magnesium based corrosion
Thaumasite formation
Thaumasite (CaCO3·CaSO4·CaSiO3·15H2O) is formed during sulphate attack in the presence of
carbonates at low temperatures. Unlike Ettringite the quantity of thaumasite that can form is not limited
by the Al2O3 content, but only by CaO and SiO2. Can cause severe damage, but is not believed to be
the principal cause of concrete deterioration by sulphate attack
75
Factors influencing sulphate attack:

amount and nature of the sulphate present,

level of the water table and its seasonal variation,

flow of groundwater and soil porosity,

form of construction, quality (porosity) of concrete.

cement content: high sulphate , high alkali, high MgO, cement fineness, high C 3A,
high C3S rete deterioration by sulphate attack
8.2.6 Freeze-thaw damage
Deterioration of concrete from freeze thaw actions may occur when the concrete is critically saturated,
which is when approximately 91% of its pores are filled with water. When water freezes to ice it
occupies 9% more volume than that of water. If there is no space for this volume expansion in a
porous, water containing material like concrete, freezing may cause distress in the concrete.
Distress to critically saturated concrete from freezing and thawing will commence with the first freezethaw cycle and will continue throughout successive winter seasons resulting in repeated loss of
concrete surface.
To protect concrete from freeze/thaw damage, it should be air-entrained by adding a surface active
agent to the concrete mixture.
This creates a large number of closely spaced, small air bubbles in the hardened concrete. The air
bubbles relieve the pressure build-up caused by ice formation by acting as expansion chambers.
About 4% air by volume is needed and the air-bubbles should be well distributed and have a distance
between each other of less than 0.25 mm in the cement paste.
8.2.7 Fire-resistance of concrete
Concrete probably has the best fire resistance of all common building material. When exposed to fire it
does not emit toxic fumes, smoke or partly melt.
Concrete has a relatively low thermal conductivity and high specific heat and is a good protection for
steel against fire.
Fire can result in 'spalling' of concrete (i.e., breaking of layers or pieces from the concrete surface),
due to (i) aggregate expansion, (ii) gas pressure build-up, or (ii) phase transitions.
Fire may causes extreme initial drying and shrinkage and thereby cause shrinkage cracking by selfrestraint (i.e., shrinkage gradients) or substrate-restraint.
Questions:
1. Characterize types of chemical reaction in concrete affecting durability
2. Describe concrete corrosion by leaching.
3. Describe concrete corrosion by carbonatation
4. Describe external concrete corrosion by sulfate atack.
5. Describe internal concrete corrosion by sulfate atack.
76
References:
1. Dhir, R. K., Newlands, M. D.: Controlling concrete degradation. 1999, Thomas Telford Ltd,
ISBN:9780727728197.
2.
Bertolini, L., Elsener,B., Pedeferri, P., Redaelli, E., Polder, R. B.: Corrosion of steel in
concrete: prevention, diagnosis, repair. 2013, Wiley, ISBN: 978-3-527-33146-8.
77
9 METALS
Chapter mission
This chapter is focused on the characterisation of the metal properties and their chemical basis.
Special attention is paid on the ways of metal production.
Chapter objectives
After studying this chapter you should know:

to describe the properties of metals,

to explain the oxidation/reduction properties of metals,

to describe the principle of metals production.
The majority of the chemical elements in periodic table are classified as metal (Fig. 9.1). Metals are
defined as luster, good thermal and electrical conductivity, and the capability of being permanently
shaped or deformed at room temperature.
Figure 9.1 Localisation of metals in periodic table
Metals are located on the left side and the middle of the periodic table. Group IA and Group IIA (the
alkali metals) are the most active metals. The transition elements, groups IB to VIIIB, are also
considered metals. The basic metals are the element to the right of the transition metals. The bottom
78
two rows of elements beneath the body of the periodic table are the lanthanides and actinides, which
are also metals.
Figure 9.2 Sodium
9.1 Properties of metals
Metals are shiny solids are room temperature (except mercury, which is a shiny liquid element), with
characteristic high melting points and densities. Many of the properties of metals, including large
atomic radius, low ionization energy, and low electronegativity, are due to the fact that the electrons in
the valence shell of a metal atoms can be removed easily. One characteristic of metals is their ability
to be deformed without breaking. Malleability is the ability of a metal to be hammered into shapes.
Ductility is the ability of a metal to be drawn into wire. Because the valence electrons can move freely,
metals are good heat conductors and electrical conductors.
Metals combine with other metals and some non-metallic elements to form a vast number of alloys
that enhance the properties of metals in specific applications, e.g., the combination of iron, nickel and
chromium provides a series of stainless steel alloys that are in common use. Metals such as nickel,
vanadium, molybdenum, cobalt, rare earths and the platinum group metals enable the catalytic
reactions for the synthesis of many organic chemicals from petroleum. A wide variety of metal
compounds and salts impart beneficial properties to products like plastics in terms of colour,
brightness, flame resistance and resistance to degradation. Photography has been made possible by
the effect of light on metal salts.
The properties of strength and ductility enable the extensive use of metals in structures and
machinery. Metals and alloys exhibit ductility, malleability and the ability to be deformed plastically
(that is, without breaking), making them easy to shape into beams (steel beams for construction),
extrusions (aluminum frames for doors and windows), coins, metal cans and a variety of fasteners
(nails and paper-clips). The strength of metals under pressure (compression), stretching (tensile) and
sheer forces makes them ideal for structural purposes in buildings, automobiles, aircraft frames, gas
pipelines, bridges, cables, and some sports equipment.
Metals are excellent conductors of both heat and electricity. In general, conductivity increases with
decreasing temperature, so that, at absolute zero (-273°C), conductivity is infinite; in other words,
metals become superconductors. Thermal conductivity is harnessed in automobile radiators and
cooking utensils. Electrical conductivity provides society with the ability to transmit electricity over long
distances to provide lights and power in cities remote from electrical generating stations. The circuitry
in household appliances, television sets and computers relies on electrical conductivity.
79
Figure 9.3 Copper conductor wire
Fatigue resistance - the ability to resist breaking after repeated deformation such as bending - enables
the use of metals in springs, levers and gears. Temperature resistance makes metals suitable for jet
engines and filaments in light-bulbs. Optical Characteristics: Metals are uniformly lustrous and, except
for copper and gold, are silvery or greyish. This is because all metals absorb light at all frequencies
and immediately radiate it. Metals impart mirrors with their reflective surface. The lustre of metals
gives them the attractive appearance that is so important in jewellery and coins. (Interestingly, metals
also provide the intangible, distinctive "metallic ring" that is associated with coins.)
Ferromagnetism is exhibited by iron and several other metals. In addition, other metals and alloys can
be magnetized in an electrical field to exhibit paramagnetism. Magnetic properties are employed in
electric motors, generators, and speaker systems for audio equipment.
Metals emit electrons when exposed to radiation (e.g. light) of a short wavelength or when heated to
sufficiently high temperatures. These phenomena are exploited in television screens, using rare earth
oxides and in a variety of electronic devices and instruments. Conversely, the ability of metals such as
lead to absorb radiation is employed in shielding, for example in the apron provided by dentists during
an X-ray examination.
9.2 Standard potential of metals
In an electrochemical cell, an electric potential is created between two dissimilar metals. This potential
is a measure of the energy per unit charge which is available from the oxidation/reduction reactions to
drive the reaction. It is customary to visualize the cell reaction in terms of two half-reactions, an
oxidation half-reaction and a reduction half-reaction.
The cell potential has a contribution from the anode which is a measure of its ability to lose electrons it will be called its "oxidation potential". The cathode has a contribution based on its ability to gain
electeons, its "reduction potential". The cell potential can then be written
Ecell = oxidation potential + reduction potential
If we could tabulate the oxidation and reduction potentials of all available electrodes, then we could
predict the cell potentials of voltaic cells created from any pair of electrodes. Actually, tabulating one or
the other is sufficient, since the oxidation potential of a half-reaction is the negative of the reduction
80
potential for the reverse of that reaction. Two main hurdles must be overcome to establish such a
tabulation
1. The electrode potential cannot be determined in isolation, but in a reaction with some other
electrode.
2. The electrode potential depends upon the concentrations of the substances, the temperature,
and the pressure in the case of a gas electrode.
In practice, the first of these hurdles is overcome by measuring the potentials with respect to a
standard hydrogen electrode. It is the nature of electric potential that the zero of potential is arbitrary; it
is the difference in potential which has practical consequence. Tabulating all electrode potentials with
respect to the same standard electrode provides a practical working framework for a wide range of
calculations and predictions. The standard hydrogen electrode is assigned a potential of zero volts.
The second hurdle is overcome by choosing standard thermodynamic conditions for the measurement
of the potentials. The standard electrode potentials are customarily determined at solute
concentrations of 1 Molar, gas pressures of 1 atmosphere, and a standard temperature which is
usually 25°C. The standard cell potential is denoted by a degree sign as a superscript.
1. Measured against standard hydroden electrode.
2. Concentration 1 Molar
°
E Cell
3. Pressure 1 atmosphere
4. Temperature 25°C
A series in which metals are arranged in the decreasing order of reduction potential is called also as a
Galvanic Corrosion Chart. The chart contains the galvanic or electrochemical series ranks of metals
according to their potential, is generally measured with respect to the Standard Calomel Electrode
(S.C.E.).
This Galvanic Corrosion Chart says that the "anodic" or "less noble" metals at the negative end of the
series - at the right of this diagram, such as magnesium, zinc and aluminium - are more likely to be
attacked than those at the "cathodic" or "noble" end of the series such as gold and graphite.. The
Table 9.1 shows some of the extreme values for standard cell potentials.
Table 9.1 Galvanic Corrosion Chart
Cathode
Half-Reaction
+
-
+
-
(Reduction) Standard Potential
°
E (volts)
Li (aq) + e -> Li(s)
-3.04
K (aq) + e -> K(s)
2+
-2.92
-
Ca (aq) + 2e -> Ca(s)
+
-2.76
-
Na (aq) + e -> Na(s)
2+
-
2+
-
-2.71
Zn (aq) + 2e -> Zn(s)
-0.76
Cu (aq) + 2e -> Cu(s)
+
0.34
-
O3(g) + 2H (aq) + 2e -> O2(g) + H2O(l) 2.07
-
-
F2(g) + 2e -> 2F (aq)
2.87
81
The values for the table entries are reduction potentials, so lithium at the top of the list has the most
negative number, indicating that it is the strongest reducing agent. The strongest oxidizing agent is
fluorine with the largest positive number for standard electrode potential.
Figure 9.4 Oxidising and reducing properties of metals according to its reduction potentials
The most active (most strongly reducing) metals appear on top, and least active metals appear on the
bottom. A more active metal (such as Zn) will donate electrons to the cation of a less active metal
(Cu
2+,
for example.)
Notice the special role of hydrogen here; although H2 does not have the physical properties of a metal,
it is capable of being "displaced" (a rather archaic term seldom used in modern chemistry) from H2O or
+
H -containing (acidic) solutions. Note that the "active" metals are all "attacked by acids"; what this
+
really means is that they are capable of donating electrons to H .
Based on the value of the reduction potential the metals are devided to noble and non-noble metals.
Noble metals are any of several metallic chemical elements that have outstanding resistance to
oxidation, even at high temperatures; the grouping is not strictly defined but usually is considered to
include rhenium, ruthenium, rhodium, palladium, silver, osmium, iridium, platinum, and gold; i.e., the
metals of groups VIIb, VIII, and Ib of the second and third transition series of the periodic table.
Silver and gold, which with copper are often called the coinage metals, and platinum, iridium, and
palladium comprise the so-called precious metals, which are used in jewelry.
9.3 Metal production
The term metal production refers to all of the processes involved in the conversion of a raw material,
such as a metallic ore, to a final form in which the metal can be used for some commercial or industrial
purpose. In some instances, metal production involves relatively few steps since the metal already
occurs in an elemental form in nature. Such is the case with gold, silver, platinum, and other so-called
noble metals. These metals normally occur in nature uncombined with other elements and can
therefore be put to some commercial use with comparatively little additional treatment.
82
In the majority of cases, however, metals occur in nature as compounds, such as the oxide or the
sulfide, and must first be converted to their elemental state. They may then be treated in a wide variety
of ways in order to make them usable for specific practical applications.
The process of producing metals from ores is based on the reduction reactions. The reduction of
metals was originally understood to be the reactions used to obtain metals from their oxides by using
substances having greater affinity for oxygen than the metal. The simplest example is the production
of iron from its protoxide:
FeO + C = Fe + CO
(9.1)
This reaction takes place, in particular, in blast furnaces.
The possibility of reducing metals is determined by the free energy of the reaction
MeO + R = Me + RO
(9.2)
where MeO is the metal oxide and R is the reducing agent. If in this reaction (at constant temperature
and pressure) the total free energy for Me and RO is less than for MeO and R, the process proceeds
from left to right, with formation of metal. The process is facilitated if the final product, which is metal,
is present in the dissolved state (solid or liquid), since dissolution is accompanied by a decrease in
free energy. This explains why, in the reduction of metals, some particularly stable oxides yield the
corresponding alloys as end products. Thus, the reduction of metals requires the presence of a
definite thermodynamic stimulus. In addition, great importance also attaches to the kinetic conditions
of reduction, which are determined by crystallochemical changes (in the case of solid oxides), the
mechanism of the chemical reactions at the phase boundaries, and the mass-transfer conditions for
the reagents—for example, diffusion.
In a more general chemical sense, the reduction of metals consists of the addition of electrons to an
atom or group of atoms. Therefore, reduction of metals also includes processes in which metals are
obtained at a cathode by electrolysis of salt melts or solutions—for example, in the case of copper:
++
Cu + 2e = Cu
(9.3)
where e is an electron.
In nonferrous metallurgy, reduction of metals is carried out in the production of metals from sulfides,
chlorides, and other compounds. Since the electrons given off by the reducing agent are necessary for
reduction, reduction processes are inseparably connected with oxidizing processes.
Metals are often extracted from the Earth by means of mining, resulting in ores that are relatively rich
sources of the requisite elements. Ore is located by prospecting techniques, followed by the
exploration and examination of deposits. Mineral sources are generally divided into surface mines,
which are mined by excavation using heavy equipment, and subsurface mines.
Once the ore is mined, the metals must be extracted, usually by chemical or electrolytic reduction.
Pyrometallurgy uses high temperatures to convert ore into raw metals, while hydrometallurgy employs
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aqueous chemistry for the same purpose. The methods used depend on the metal and their
contaminants.
Metals always occur in their oxidized state in ores, often as the oxide or sulfide of the metal. In order
to convert an ore to its elemental state, therefore, it must be reduced. Reduction is a chemical reaction
that is the opposite of oxidation. Metals can be reduced in a variety of different ways.
When a metal ore is an ionic compound of that metal and a non-metal, the ore must usually be
smelted — heated with a reducing agent — to extract the pure metal. Many common metals, such as
iron, are smelted using carbon as a reducing agent. Some metals, such as aluminium and sodium,
have no commercially practical reducing agent, and are extracted using electrolysis instead. In
technology the most important examples of such processes are the production of aluminum by
electrolysis of alumina from a melt and the production of copper from aqueous solutions of CuSO 4.
Sulfide ores are not reduced directly to the metal but are roasted in air to convert them to oxides.
Thus the ways of metals production is always based on the chemical reduction process and include:
1. Reduction from the oxides and sulphides using carbon as reducing agent,
2. Reduction using aluminium as reducing agent,
3. Electrolysis.
9.4 Iron
Iron is the most used of all the metals, including 95 % of all the metal tonnage produced worldwide.
Thanks to the combination of low cost and high strength it is indispensable. Its applications go from
food containers to family cars, from scredrivers to washing machines, from cargo ships to paper
staples. Steel is the best known alloy of iron, and some of the forms that iron takes include: pig iron,
cast iron, carbon steel, wrought iron, alloy steels, iron oxides.
Iron is a lustrous, ductile, malleable, silver-gray metal (group VIII of the periodic table). It is known to
exist in four distinct crystalline forms. Iron rusts in damp air, but not in dry air. It dissolves readily in
dilute acids. Iron is chemically active and forms two major series of chemical compounds, the bivalent
iron (II), or ferrous, compounds and the trivalent iron (III), or ferric, compounds.
With ores of iron, reduction can be accomplished by reacting oxides of iron with carbon and carbon
monoxide. One of the common devices used for this purpose is the blast furnace. The blast furnace is
a tall cylindrical vessel into which is fed iron ore (consisting of oxides of iron), coke (nearly pure
carbon) and limestone. The temperature in the blast furnace is then raised to more than 1000°C (Fig.
10.5). At this temperature, carbon reacts with oxygen to form carbon monoxide, which in turn, reacts
with oxides of iron to form pure iron metal. The limestone in the original mixture added to the blast
furnace reacts with and removes silicon dioxide (sand), an impurity commonly found with iron ore.
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Figure 9.5 Blast furnnace
9.5 Non-ferrous metals
Non-ferrous metals include aluminum, brass, copper, nickel, tin, lead, and zinc, as well as precious
metals like gold and silver. While non-ferrous metals can provide strength, they are primarily used
where their differences from ferrous metals can provide an advantage.
For instance, non-ferrous metals are much more malleable than ferrous metals. Non-ferrous metals
are also much lighter, making them well-suited for use where strength is needed, but weight is a
factor, such as in the aircraft or canning industries. Because they contain no iron, non-ferrous metals
have a higher resistance to rust and corrosion, which is why you’ll find these materials in use for
gutters, water pipes, roofing, and road signs. Finally, they are also non-magnetic, which makes them
perfect for use in small electronics and as electrical wiring.
As far as recycling goes, aluminum is the third most recycled material in the world. However, many
other non-ferrous materials like copper, brass and lead are relatively scarce, and metallurgists rely
heavily on scrap material recycling to make new ones.
9.5.1 Copper
Copper provides a diverse range of properties: good thermal and electrical conductivity, corrosion
resistance, ease of forming, ease of joining, and color. However, copper and its alloys have relatively
low strength-to-weight ratios and low strengths at elevated temperatures. Some copper alloys are also
susceptible to stress-corrosion cracking unless they are stress relieved. Next to silver, copper is the
next best electrical conductor. It is a yellowish red metal that polishes to a bright metallic luster. It is
tough, ductile and malleable. Copper has a disagreeable taste and a peculiar smell. Copper is
resistant to corrosion in most atmospheres including marine and industrial environments. It is corroded
by oxidizing acids, halogens, sulphides and ammonia based solutions.
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Copper and its alloys -- the brasses and bronzes -- are available in rod, plate, strip, sheet, tube
shapes, forgings, wire, and castings.
Figure 9.6 Examples of the application of copper
9.5.2 Aluminum
Pure aluminum is a silvery-white metal with many desirable characteristics. It is light, nontoxic (as the
metal), nonmagnetic and nonsparking. It is easily formed, machined, and cast. Pure aluminum is soft
and lacks strength, but alloys with small amounts of copper, magnesium, silicon, manganese, and
other elements have very useful properties. Aluminum is an abundant element in the earth's crust, but
it is not found free in nature. The Bayer process is used to refine aluminum from bauxite, an aluminum
ore. Because of aluminum's mechanical and physical properties, it is an extremely convenient and
widely used metal. Some Common Uses in Building & Construction Industry:

door and window frames

wall cladding, roofing, awnings
Figure 9.7 Aluminium and an example of its application in civil industry
9.5.3 Zinc
Zinc is a silvery blue-grey metal with a relatively low melting point (419.5°C) and boiling point (907°C).
When unalloyed, its strength and hardness is greater than that of tin or lead, but appreciably less than
that of aluminium or copper. The pure metal cannot be used in stressed applications due to low creep-
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resistance. For these reasons most uses of zinc are after alloying with small amounts of other metals
or as a protective coating for steel.
One of the most useful characteristics of zinc is its resistance to atmospheric corrosion, and just over
half of its use is for the protection of steelwork. In addition to its metal and alloy forms, zinc also
extends the life of other materials such as steel (by hot dipping or electrogalvanizing), rubber and
plastics (as an aging inhibitor), and wood (in paints). Zinc is also used to make brass, bronze, and diecasting alloys in plate, strip, and coil; foundry alloys; superplastic zinc; and activators and stabilizers
for plastics.
Figure 9.8 Zinc and magnesium Magnesium
As the lightest structural metal available, magnesium has a high strength-to-weight ratio. With its low
modulus of elasticity combined with moderate strength, magnesium alloys can absorb energy
elastically, providing excellent dent resistance and high damping capacity. Magnesium has good
fatigue resistance and performs particularly well in applications involving a large number of cycles at
relatively low stress. The metal is sensitive to stress concentration, however, so notches, sharp
corners, and abrupt section changes should be avoided. Magnesium alloys are the easiest of the
structural metals to machine and they can be shaped and fabricated by most metalworking processes,
including welding.
Questions:
1. List some of the metals and its properties.
2. Define the standard reduction potential of metals E0.
3. Explain the standard reduction potential role in metal behaviour.
4. Characterize the Galvanic Corrosion Chart.
5. Characterize the iron and steel production.
6. Define the difference between the term “pig iron” and “steel”.
7. Characterise the non-ferrous metals and their application in building industry.
References:
1. Beier S.P., Hede P.D.: Essentials of Chemistry. 2010, BookBoon, ISBN: 978-87-403-0322-3.
2. El Sair R.R.: Fundamentals of Chemistry. 2012, Romain Elsair & Ventus Publishing ApS, ISBN:
978-87-403-0105-2.
3. Landolt D.: Corrosion and Surface Chemistry of Metals. 2007, CRC Press, ISBN: 1439807884
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10 CORROSION OF METALS
Chapter mission
This chapter is focused on the corrosion process of metals.
Chapter objectives
After studying this chapter you should know:

Describe what the principle of corrosion process is.

Classify the corrosion types.

Explain a dry corrosion.

Explain a wet corrosion.
10.1 Definition of corrosion
Corrosion is defined as the damage or deterioration of a material (usually a metal) due to a reaction or
interaction with the environment. Corrosion may be defined also as: The destruction of a metal or its
properties by chemical or electro-chemical reaction with its environment.
Fig. 10.1 Deterioration of materials by corrosion
The basic cause of corrosion is the instability of metal in its refined form. The process of corrosion is
the tendency of a metal to revert to its natural state. What dictates the level of corrosion is the
combination of the material type and the environment it is exposed to. The force that makes metals
corrode is a natural consequence of their temporary existence in the metallic form. To attain the
metallic state from their natural form as various chemical compounds, it is necessary to expend
varying amounts of energy from external sources, this process normally being carried out in a furnace.
The more reactive a metal is with its environment, the more energy is required to bring about the
change from its natural state to the metallic form. It is relatively high for metals such as magnesium,
aluminium and iron and relatively low for metals such as copper, silver and gold. All environments are
corrosive in some manner. Understanding the environment helps to determine what factors contribute
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to corrosion activity and what the appropriate control methods could be. Corrosion environments can
be placed into four major categories:

liquid,

underground,

atmospheric and

high temperature.
In most industrial applications, the process system is exposed to many, if not all of these
environments. A material that is inert in one environment may not be in another. It is for this reason
that material selection is important to ensure that adequate performance characteristics, especially life
span, are obtained. Cost and availability dictate the materials that are used in industrial processes.
This trade-off is what causes most corrosion problems.
In most cases, the combination of the metals found in equipment and structures, combined with the
wide range of possible environments, will result in more than one form of corrosion within a system.
10.2 Classification of corrosion
Corrosion is often classified as

Chemical Corrosion or Dry Corrosion,

Electro chemical Corrosion or Wet Corrosion.
Wet corrosion occurs when a liquid phase is present and dry corrosion occurs in the absence of a
liquid phase or above the dew point of the environment.
10.2 Dry and wet corrosion
10.2.1
Dry corrosion
Dry corrosion or oxidation occurs when a gas e.g. oxygen in the air reacts with metal without the
presence of a liquid. These are generally metal/gas or metal/vapour reactions involving non-metals
such as oxygen, halogens, hydrogen sulphide, sulphur vapour, etc. and oxidation, scaling and
tarnishing are the more important forms.
The basic reaction involved in dry corrosion is:
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(10.1)
where M is a metal element.
The metal loses electrons to form an ion and some free electrons. The ionic metal can then react with
oxygen to form a metal oxide. A characteristic of these reactions is that the initial oxidation of the
metal, reduction of the non-metal, and formation of compound must occur at one and the same place
at the metal/non-metal interface. Should the compound be volatile or discontinuous, further interaction
at the interface (or through a thin film of constant thickness – Fig. 10.3) is possible and in most cases
the reaction rate will tend to remain constant with time (linear law).
The mechanisms involved make different metals corrode at varying rates. When oxygen gas comes
into contact with a metal surface, this causes an oxide layer to form around the metal. The layer of
metal oxide limits the amount of oxygen able to reach the metal surface, thus the rate of corrosion is
reduced.
10.3 Thin film formation on the metal surface
If the film is continuous it will present a barrier to the reactants and further interaction will necessitate
passage of the reactants through the film by (a) diffusion of the non-metal or (a) diffusion and
migration of ions of the reactants.
10.2.2
Wet corrosion
With the exception of some forms of high-temperature corrosion, all forms of wet corrosion occur
through the action of the electrochemical cell. An electrochemical cell consists of an anode, a cathode,
a connection, and an electrolyte.
The anode is the metal that corrodes. It undergoes oxidation and therefore loses electrons.
The cathode can be a metal or any other conducting material. It undergoes reduction and therefore
gains electrons. The reaction that occurs at the cathode is not necessarily related to the material that it
is made from. The connection is necessary for the electrons to travel between the anode and cathode
and can be either physical direct contact or some form of wire. An electrolyte must also be present to
allow for migration of ions between the cathode and anode and participate in the formation of
corrosion products.
This cell contains what is known as an oxidation/reduction reaction. In this reaction, an exchange of
electrons (due to a difference in potential) occurs, where an anode is the site of oxidation and a
cathode is the site of reduction. The electrons given off at the anode travel through the metal to the
cathode, where they are consumed in a reduction reaction.
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10.4 Corrosion of iron
The basic process comprises a reaction between the metal, oxygen and water to produce a soluble
form of the metal or one of its oxides. Rust is an iron oxide, usually red oxide formed by
the redox reaction of iron and oxygen in the presence of water or air moisture. Several forms of rust
are distinguishable both visually and by spectroscopy, and form under different circumstances.[1] Rust
consists of hydrated iron(III) oxides Fe2O3·nH2O and iron(III) oxide-hydroxide(FeO(OH), Fe(OH)3).
In the corrosion process, two reactions take place. In one, the anodic reaction, metal atoms are
ionised and pass into solution leaving their electrons within the original metal surface. In the second,
the cathodic reaction, the free electrons within the metal are taken up by chemical species such as O 2
and H2O in reduction reactions. The overall reaction proceeds as follows:
Fe + 2H2O → Fe (OH)2 + H2
(10.2)
The overall reaction can be broken down into the oxidising anodic reaction
2+
Fe → Fe
-
+ 2e
(10.3)
and the reducing cathodic reaction
-
2H2O + 2e → H2 + 2(OH)
-
(10.4)
The reaction 10.3 and 10.4 are called ‘half cell’ reactions. Reaction 10.3 is the half of the process
which is responsible for the damage during corrosion. The speed at which this reaction proceeds is
directly related to the corrosion rate.
10.3 The factors affecting the rate of corrosion
The factors that affect the rate of corrosion are
1. Factors connected with the metal and its surface,
2. Factors connected with the atmosphere,
3. Factors connected with the corrosion product.
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Factors connected with the metal include:

Position of the metal in the E.M.F Series
The type of impurity present in it and its electropositive nature decides the corrosion of a
metal. For example when iron has impurities like copper, tin, etc. iron corrodes since iron is
more electropositive than metals like copper and tin. On the other hand when iron is coupled
with zinc, zinc corrodes since zinc is more electropositive than iron. Generally pure metal does
not corrode, as there is no cathode spot available to induce corrosion. A rough surface
corrodes readily as it collects more dirt and provides more cathode spot for corrosion. A
polished surface does not corrode easily.

Purity of the metal
Generally pure metal does not corrode, as there is no cathode spot available to induce
corrosion.

Surface of the metal
A rough surface corrodes readily as it collects more dirt and provides more cathode spot for
corrosion. A polished surface does not corrode easily.
Factors connected with the atmosphere include

Nature of the atmosphere,

Temperature and pH of the atmosphere,

Amount of moisture in the atmosphere,

Amount of oxygen in the atmosphere,

Amount of chemical fumes in the atmosphere.
Questions:
1. Define the corrosion of metals.
2. Write the principle of chemical and electrochemical corrosion.
3. Describe dry corrosion processes.
4. Describe wet corrosion processes.
5. Give an expamle of the factors affercting rate of corrosion.
References:
1. Landolt D.: Corrosion and Surface Chemistry of Metals. 2007, CRC Press, ISBN: 1439807884
2. Kaesche, H.: Corrosion of Metals. 2003, Springer Verlag Berlin, ISBN 978-3-642-96038-3
3. Nelson P.G.: Introduction to Inorganic Chemistry. 2011, Peter G. Nelson & Ventus Publishing
ApS, ISBN: 978-87-7681-732-9.
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