MME 2001
MATERIALS SCIENCE 1
8.10.2015
outline
● overview of atomic structure
● periodic table: organization/groups/periods
ionization energy
electronegativity
● Classification of elements
● Metals
● Nonmetals
● Metalloids
● interatomic bonding
● primary vs secondary bonding
● covalent/ionic/metallic
● bonding energy-bonding force
Atomic Structure
HELIUM
ATOM
Shell
proton
nucleus
# electrons = # protons
++
N
-
Varying electrons: ions
electron
Varying neutrons: isotopes
N
neutron
ATOMIC MASS NUMBER = number
of protons + number of neutrons
ATOMIC NUMBER = number of
protons
rules
● Heisenberg’s uncertainty principle
we cannot precisely measure the momentum and
the position of an electron at the same time.
● Pauli exclusion principle
No 2 Electrons in an atom Can Have the same 4
Quantum Numbers; the same n, l, ml, and ms.
● The Aufbau principle
states that electrons are added to the lowest energy
orbitals first before moving to higher energy
orbitals.
● Hund’s rule,
we must fill each shell with one electron each
before starting to pair them up.
Atomic structure
The charge and mass number of an electron are:
a) charge = 0, Mass number = 1
b) charge = -1, Mass number = 0
c) charge = +1, Mass number = 1
d) charge = +1, Mass number = 0
The charge and mass number of a neutron are?
a) charge = +1, Mass number = 1
b) charge = 0, Mass number = 1
c) charge = +1, Mass number = 0
d) charge = -1, Mass number = 0
Atomic structure
The nucleus of the element having atomic
number 25 and atomic weight 55 will
contain?
a)
b)
c)
d)
25 protons and 30 neutrons
30 protons and 25 neutrons
55 protons
55 neutrons
Atomic structure
A beryllium atom has 4 protons, 5
neutrons, and 4 electrons. What is the
mass number of this atom?
a)
b)
c)
d)
e)
4
5
8
9
13
Atomic structure
Which has the least mass in an atom?
a) nucleus
b) proton
c) neutron
d) electron
If uranium-235, has 92 protons, how many
neutrons does the isotope uranium-238 have?
a) 92
b) 95
c) 143
d) 146
quantum numbers
every electron in an atom is characterized by four
quantum numbers.
There are three quantum numbers necessary to
describe an atomic orbital.
 The principal quantum number (n)
designates size
 The angular moment quantum number (l)
describes shape
 The magnetic quantum number (ml)
specifies orientation
Angular moment Quantum Number (l)
l signifies the subshell
l describes the shape of the orbital.
l values range from 0 to n – 1
Example: If n = 2, l can be 0 or 1.
n
l
subshell
energy state
1
0
2
0,1
s
1
s,p
3
3
4
5
6
0,1,2 0,1,2,3 0  4 0  5
s,p,d s,p,d,f s,p g s,p h
5
7
9
11
Quantum Numbers
To summarize quantum numbers:
principal (n) – size
angular (l) – shape
magnetic (ml) – orientation
Required to
describe an
atomic orbital
principal (n = 2)
2px
related to the magnetic
quantum number (ml )
angular momentum (l = 1)
electron spin (ms) direction of spin
Required to describe an electron in an atomic orbital
Electron Spin Quantum Number-ms
used to specify an electron’s spin.
There are two possible
directions of spin.
Allowed values of ms
are +½ and −½.
Atomic structure
The lowest principal quantum number for
an electron is?
a)
b)
c)
d)
e)
0
1
2
3
4
Quantum numbers
An electron with
n = 2, ℓ = 1, ml = −1, and ms = +1/2
is found in the same atom as a second electron with
n = 2, ℓ = 1, ml = −1.
What is the spin quantum number for the second
electron?
Since the first three quantum numbers are identical for
these two electrons, we know that they are in the same
orbital. As a result, the spin quantum number for the second
electron cannot be the same as the spin quantum number
for the first electron. This means that the spin quantum
number for the second electron must be ms = −1/2.
Atomic structure
Maximum number of electrons in a subshell with
l = 3 and n = 4 is
a)
b)
c)
d)
e)
10
12
14
16
18
Principal
Quantum No:
4
Subshell
n
l
s/0
p/1
d/2
f/3
No. of energy States:
ml
1/0
3 / -1,0,+1
5 / -2,-1,0,+1,+2
7 / -3,-2,-1,0,+1,+2,+3
Number of
Electrons
Per Subshell
2
6
10
14
Atomic structure
If n=3, and l=2, then what are the possible
values of ml ?
Since ml must range from –l to +l,
then ml can be: -2, -1, 0, 1, or 2.
Quantum Numbers:
A Macroscale Analogy
n
- indicates which train (shell)
 l
- indicates which car (subshell)
 ml - indicates which row (orbital)
 ms - indicates which seat (spin)
No two people can have exactly the
same ticket (sit in the same seat).
Electron Configurations
Q. the full electronic configuration of an element is
1s22s22p5.
How many electrons does it have in its outer
shell?
A. # of outer shell-valence electrons: 7
Q: the full electronic configuration of an element.
1s22s22p5.
What is its atomic number?
A. Atomic number: 9
Electronic Configurations
Fe-atomic # = 26
4d
4p
Energy
3d
4s
3p
3s
1s2 2s2 2p6 3s2 3p6 3d
6
4s2
N-shell n = 4
6 electrons left to be located
total # e-s
20 electrons
M-shell n = 3; 18 electrons
2p
2s
L-shell n = 2; 10 electrons
1s
K-shell n = 1; 2 electrons
Electron Configurations
rules for electron configurations:
● Electrons will reside in the
lowest possible energy orbitals
● Each orbital can accommodate
a maximum of two electrons.
● Electrons will not pair in
degenerate orbitals if an
empty orbital is available.
● Orbitals will fill in the order
..3p6/4s2/3d10/4p6/5s2/4d10/
5p6/6s2/4f14/5d10/6p6/7s2
Electron Configurations
F has a total of 9 electrons
Energy
2p
2s
1s
2p
2p
When there are one or more unpaired
electrons, as in the case of oxygen and
fluorine, the atom is called paramagnetic.
The ground state electron
configuration of F
1s22s22p5
Electron Configurations
Ne has a total of 10 electrons
Energy
2p
2s
1s
2p
2p
When all of the electrons in an atom are
paired, as in neon, it is called diamagnetic.
The ground state electron
configuration of Ne
1s 2s 2p
2
2
6
learning check
Write the electron configuration and give the orbital
diagram of a calcium (Ca) atom (Z = 20).
Z = 20, Ca has 20 electrons.
Each s subshell can contain a maximum of
two electrons, whereas each p subshell can
contain a maximum of six electrons.
Solution
Ca
1s22s22p63s23p64s2
1s2 2s2
2p6
3s2
3p6
Remember that the 4s orbital fills
before the 3d orbitals.
4s2
learning check
electron configuration for an arsenic atom (Z = 33) in the
ground state.
Z = 18 for Ar.
The order of filling beyond the noble gas
core is 4s, 3d, and 4p. Fifteen electrons go
into these subshells because there are 33 –
18 = 15 electrons in As beyond its noble gas
core.
Solution
As
[Ar]4s23d104p3
Arsenic is a p-block element; therefore,
we should expect its outermost electrons
to reside in a p subshell.
2
2
6
2
6
2
3
10
Valence electrons
They occupy the outermost shell.
They participate in the bonding between atoms
They dictate the physical and chemical properties
if the outermost or valence electron shell are
completely filled: stable electron configurations
occupation of the s and p states for the outermost
shell by a total of eight electrons, in neon (Ne),
argon (Ar), and krypton (Kr); inert, or noble, gases,
which are virtually unreactive chemically.
electron configuration?
Which one of the following is a proper orbital
configuration?
Learning check
The electrons with principle energy level n = 2 of a stable atom of
boron (atomic number = 5) would have an electron arrangement of
(a) ( ↑ ↓) ( ↑ ) (
)(
)
(b) ( ↑ ) ( ↑ ) ( ↑ ) (
)
(c) (
)(↑ )(↑ )(↑ )
(d) (
)(↑↓)(↑ )(
)
(e) ( ↑ ↓) ( ↑ ↓ ) ( ↑ ) ( ↑ )
Which of the following electron arrangements does not represent
an atom in its ground state?
(1s)
(2s)
(2p)
(3s)
(a) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ )
(b) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ ↓ )
(c) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ ) ( ↑ )
(d) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ ↓ ) ( ↑ ↓ ) (
)
beyond the d-orbitals
‘s’-groups
group
‘p’-groups
period
d-transition elements
1s2
2s2/2p6
3s2/3p6/
4s2/3d10/4p6
5s2/4d10/5p6/
6s2/4f14/5d10/6p6
lanthanides
actinides
f-transition elements
you will have a quiz
next week!
The Periodic Law
Mendeleev realized that:
When arranged by increasing atomic
number, the chemical elements display a
regular and repeating pattern of chemical
and physical properties.
what are these properties?
 Metallic vs nonmetallic character
 Atomic radius
 Ionization energies (energy necessary to remove
the outermost electron from the atom)
 Electron affinities (energy change when an
electron is added to a neutral atom)
 Reactivity
 Electronegativity
Organisation of the periodic table
The vertical columns: groups from 1 to 18.
Elements in the same group have similar valence
electron structures
and hence similar
chemical and
physical
properties.
groups
Organisation of the periodic table
elements are situated, with increasing atomic
number, in seven horizontal rows
called periods.
Each contains
elements with
electrons in the
Same outer shell.
periods
Periodic table
Periodic table
The first period contains elements with electrons in the first electron shell
only. Hydrogen and helium thus have behaviours very different to the lower
periods and are not easily classified into the groups used to describe the rest of
the table.
Periodic table
electrons fill the second electron shell which has both s and p orbitals. Elements
in this and the next period typically follow the so called octet rule, whereby
stable compounds are formed with 8 electrons in the outer shell.
Periodic table
elements with valence electrons in the 3s and 3p orbitals.
Elements in this period typically follow the so called octet rule, whereby stable
compounds are formed with 8 electrons in the outer shell.
Periodic table
elements with valence electrons in the 4s, 4p and 3d orbitals.
The octet rule is no longer applicable due to the introduction of the d subshell.
Periodic table
The fifth period contains elements with valence electrons in
the 5s, 5p and 4d orbitals.
Periodic table
32 elements with valence electrons in the 6s, 6p, 5d, and - including the
lanthanides - the 4f orbital. This period contains the last stable element, lead,
with all later elements being radioactive.
Periodic table
The seventh period contains 32 elements with valence electrons in the
7s,7p, 6d, and - including the actinides - the 5f orbital.
Periodic table
The alkali metals are soft, highly reactive metals with one electron in
their outermost s subshell. The reactivity of these elements increases
down the group
Periodic table
The alkaline earth metals are all reactive metals with two electrons in their
outermost s subshell. In general they are harder, denser, and have higher
melting points than their alkali metal analogues
Periodic table
A group of transition metal elements, the lightest two of which are exceptions
from the Aufbau principle (to determine the structure of the atom), showing
valence configurations of d5s1.
Periodic table
The chalcogens, or oxygen family is formed of non metals (oxygen and sulfur)
and metalloids and its elements are characterised by having 6 electrons in
their outer shell.
Periodic table
The halogens are a group of highly reactive elements with 7 electrons in their
outer shell. This is the only group which contains elements in all three states of
matter at room temperature and pressure.
Periodic table
The noble gases are typically relatively unreactive and are
characterised by a full outer electron shell.
Periodic table
The s block consists of elements with their valence electrons in s orbitals. Elements
within the s-block all behave fairly similarly, being soft, reactive metals. The s sub-shell
can contain a maximum of two electrons, and hence the block is two columns wide
Periodic table
The p block consists of elements with their valence electrons in p orbitals. The
characteristics of elements within the p-block are fairly varied, including metals and
non-metals and so called 'metalloids'. The p sub-shell can hold six electrons, in three
distinct orbitals known as px, py and pz
Periodic table
The d-block, also known as the ‘transition metals’ contains only metals,
typically capable of existing in at least two stable oxidation states. The d subshell can hold up to 10 electrons in 5 distinct orbitals.
Periodic table
The f block consists of the lanthanides and actinides which are all soft metals
many of which are not found in nature. The f sub-shell can contain up to 14
electrons in seven distinct orbitals
Periodic table
● The elements of the rightmost group, are the inert
gases, with filled electron shells and stable
electron configurations.
● Group VIIA and VIA elements are one and two
electrons deficient from having stable structures.
Periodic table
● The Group VIIA elements (F, Cl, Br, I, and At) are
sometimes termed the halogens.
● The alkali and the alkaline earth metals (Li, Na, K,
Be, Mg, Ca, etc.) are labeled as Groups IA and IIA,
with one and two electrons in excess of stable
structures.
Periodic table
● The elements in the three long periods, Groups
IIIB through IIB, are termed the transition
metals, which have partially filled d electron
states and in some cases one or two electrons in
the next higher energy shell.
● Groups IIIA, IVA, and VA (B, Si, Ge, As, etc.)
display characteristics that are intermediate
between the metals and nonmetals by virtue of
their valence electron structures.
Learning check
The elements in each vertical column on the periodic
table usually have similar properties and are called
a(n)
a) period
b) group
c) element
d) Property
Elements on the periodic table are arranged in order of
a) increasing density.
b) decreasing density.
c) increasing atomic number.
d) decreasing atomic number.
Learning check
An element has the electronic structure 2,8,4. Which
group is it in?
a) Group 3
b) Group 4
c) Group 5
d) Group 6
Which of these electronic structures belongs to a noble
gas?
a) 2
b) 2,2
c) 2,8,2
d) 2,8,4
Learning check
Two elements have these electronic structures: 2,1
and 2,8,1. What can you say about the elements?
a)
b)
c)
d)
They are both
They are both
They are both
They are both
in
in
in
in
group 1
group 2
period 1
period 2
nonmetallic character
İonization energy
Negative electron affinity
İonization energy
nonmetallic character
metallic character
Atomic radii
metallic character
Atomic radii
Negative electron affinity
Periodic Table
Ionization Energy
IE = energy required to remove an
electron from an atom in the gas phase.
Mg (g) + 738 kJ ---> Mg+ (g) + e-
Atomic Size
 Size goes UP on going down a group.
 Because electrons are added farther from
the nucleus, there is less attraction.
 Size goes DOWN on going across a period.
Atom size
increases!
Atom size
decreases!
atomic
radii
(picometer//
10-12m)
Electronegativity
● the tendency of an atom to attract electrons
towards itself.
● An atom's electronegativity is affected by both
its atomic number and the distance at which
its valence electrons reside from the charged
nucleus.
● The higher the associated electronegativity
number, the more an element or compound
attracts electrons towards it.
Electronegativity
electronegativity increases in moving from
left to right and from bottom up.
Atoms are more likely to accept electrons
if their outer shells are almost full, and if
they are less “shielded” from (i.e., closer
to) the nucleus.
electronegativity
increases!
Electronegativity
Electronegativity
Electronegativity
Learning check
Which elements are the most electronegative
element of those shown in the diagrams below?
Learning check
Fluorine has a lower electronegativity than
a) Oxygen
b) Chlorine
c) Lithium
d) None of the above
Metals, Nonmetals & Metalloids
1
Nonmetals
2
3
4
5
Metals
6
7
Metalloids
Metals
88 elements are metals or metal like element
Physical properties:






good conductors of heat and electricity
shiny
ductile (can be stretched into thin wires)
malleable (can be pounded into thin sheets)
High density (heavy for their size)
High melting point
Metals
Metals
chemical properties:
 Easily lose electrons
 Form positive (+) ions
 Corrode easily
Non-metals
Non-metals are on the right of the stairstep line.
Their characteristics are opposite to those of metals.
Physical Properties of Nonmetals:
 No luster (dull appearance)
 Poor conductor of heat and electricity
 Brittle (breaks easily)
 Not ductile
 Not malleable
 Low density
 Low melting point
 Many non-metals are gases.
Non-metals
Chemical Properties of Non-metals:
 Tend to gain electrons
 metals that tend to lose electrons but nonmetals
that tend to gain electrons, to form compounds
with each other.
 These compounds are called
ionic compounds.
 When two or more
nonmetals bond with
each other, they form
a covalent compound.
Metalloids
 Metalloids (metal-like) have properties of both
metals and non-metals.
 They are solids
 can be shiny or dull
 They conduct heat
and electricity better
than non-metals but
not as well as metals
 They are ductile and
malleable
Non-metals
Learning check
Which of the following is a property of alkali metals?
a) They are so hard they cannot be cut.
b) They are very reactive.
c) They are stored in water.
d) They have few uses.
Most of the elements in the periodic table are
a) metals.
b) metalloids.
c) gases.
d) nonmetals.
Learning check
The elements to the right of the zigzag line on the
periodic table are called
a) metalloids.
b) conductors.
c) metals.
d) nonmetals.
Transition metals are
a) good conductors of thermal energy.
b) more reactive than alkali metals.
c) not good conductors of electric current.
d) used to make aluminum.
Learning check
What do the elements on the far right of the
table (He, Ne, Ar, and Kr) have in common?
a)They are liquid in normal conditions
b)They are metals that rust easily
c) They are very reactive gases
d)They do not generally react with other elements
Learning check
Which of the following statements describes most
metals?
a) They are easily shattered.
b) They are gases at room temperature.
c) They are dull.
d) They are good conductors of electric current.
Elements lying along the zigzag line on a periodic table
are
a) metals
b) nonmetals
c) metalloids
d) noble gases
Learning check
Elements in a period have …….
a) a wide range of chemical properties
b) the same atomic radius
c) similar chemical properties
d) the same number of protons
The elements in Group 1 of the periodic table are
commonly called the…..
a) alkali metals
b) transition metals
c) alkaline earth metals
d) rare earth metals
Learning check
Elements in a group have
a) a wide range of chemical properties
b) the same atomic radius
c) similar chemical properties
d) the same number of protons
What do the elements on the far right of the table
(He, Ne, Ar, and Kr) have in common?
a) They are liquid in normal conditions
b) They are metals that rust easily
c) They are very reactive gases
d) They do not generally react with other elements
bonding and properties
● Some general behaviors of the various material
types (i.e., metals, ceramics, polymers) may be
explained by bonding type.
● For example, metals are good conductors of both
electricity and heat, as a consequence of their
free electrons.
● By way of contrast, ionically and covalently
bonded materials are typically electrical and
thermal insulators because of the absence of
large numbers of free electrons.
interatomic bonding
● the bonding involves the valence electrons
● the nature of the bond depends on the electron
structures of the constituent atoms.
● There are three types of bonding: each bonding
type arises from the tendency of the atoms to
assume stable electron structures.
● Secondary or physical forces and energies are
weaker than the primary ones, but nonetheless
influence the physical properties of some
materials.
interatomic bonding
 Ionic
 Metal (cation) with non-metal (anion)
Transfer of electron(s)
 Strong bond  high melting point
 Covalent
 Non-metal with non-metal
 Sharing of electron(s)
 Non-polar (equal distribution of electrons)
 Polar (uneven electron distribution)
 Weak bonds…low melting points
 Metallic (nuclei in a “sea” of shared electrons)

Bonding forces
● physical properties of materials = f (interatomic
forces that bind the atoms together)
● two isolated atoms interact as they are brought
close together from an infinite separation.
● At large distances, interactions are negligible,
because the atoms are too far apart to have an
influence on each other; however, at small
separation distances, each atom exerts forces
on the other.
Bonding forces
● The origin of an attractive force FA depends on
the particular type of bonding that exists
between the two atoms.
● Repulsive forces (FR) arise from interactions
between the negatively charged electron clouds
for the two atoms
● they are important only at small values of r as
the outer electron shells of the two atoms begin
to overlap.
Bonding force and bonding energy
The minimum energy
corresponds to the
equilibrium spacing, r0.
the bonding energy for
these two atoms, E0,
corresponds to the
Equilibrium
energy at this
interatomic spacing
minimum point;
it represents the energy
required to move these
two atoms to an infinite
separation.
Bonding forces
The net force FN between the two atoms is just the
sum of both attractive and repulsive components
FN = FA+ FR
When FA and FR balance, or become equal, there is
no net force; implying a state of equilibrium
FA + FR = 0
The centers of the two atoms will remain separated
by the equilibrium spacing r0
ionic bonding
● Forms between metallic and nonmetallic
elements; elements at the horizontal extremities
of the periodic table.
● a metallic atom easily gives up its valence
electrons to the nonmetallic atoms.
● In the process all the atoms acquire stable
configurations and become ions.
● Ionic bonding is non-directional (magnitude of the
bond is equal in all directions around the ion)
● Ceramic materials exhibit ionic bonding
Ionic Bonding
• Occurs between + and - ions.
• Requires electron transfer.
• Large difference in electronegativity required.
Na (metal)
Unstable
11 electrons
electron
Cl (nonmetal)
Unstable
17 electrons
Na (cation) +
Cl (anion)
stable
stable
Coulombic
Attraction
positive and negative ions, by virtue of their net
electrical charge, attract one another
ionic bonding
Schematic
representation of
ionic bonding in
sodium chloride
(NaCl).
ionic bond - Electronegativity
ionic bond : metal
+
donates
electrons
nonmetal
accepts
electrons
Dissimilar electronegativities
ex: MgO
Mg 1s2 2s2 2p6 3s2
Mg2+ 1s2 2s2 2p6
O 1s2 2s2 2p4
O2- 1s2 2s2 2p6
Ionic Bonding - examples
• Predominant bonding in Ceramics
NaCl
MgO
CaF 2
CsCl
Give up electrons
Acquire electrons
Li Be
Na Mg
K Ca Sc
Rb Sr
Y
Cs Ba
accept 2e
accept 1e
give up 3e
H
give up 2e
give up 1e
Columns: Similar Valence Structure
inert gases
ionic bonding - Periodic Table
He
O
F
S
Cl Ar
Se
Br Kr
Te
Po
I
Ne
Xe
At Rn
Fr Ra
Electropositive elements:
Readily give up electrons
to become + ions.
Electronegative elements:
Readily acquire electrons
to become - ions.
Electronegativity
• Ranges from 0.7 to 4.0,
• Large values: tendency to acquire electrons.
Smaller electronegativity
Larger electronegativity
Ion Sizes
+
Li,152 pm
3e and 3p
Li + , 78 pm
 e  2e and 3 p
Forming
a cation.
 CATIONS are SMALLER than the atoms from
which they come.
 The electron/proton attraction has gone UP
and so size DECREASES.
Ion Sizes
F, 71 pm
9e and 9p
F- , 133 pm
+ e  10 e and 9 p
Forming
an anion.
 ANIONS are LARGER than the atoms from
which they come.
 The electron/proton attraction has gone DOWN
and so size INCREASES.
 Trends in ion sizes are the same as atom sizes.
learning check
Which atom or ion has the smallest radius?
a)
b)
c)
d)
O2+
O+
O
O2–
Learning check
How do the size of a negative ion compare to the
size of the atom that formed it?
a) it's smaller
b) it's larger
c) it's the same size
d) it varies
ionic bonding
● for ionic materials to be stable, all positive ions
must have as nearest neighbors negatively
charged ions in a three dimensional scheme.
● The predominant bonding in ceramic materials
is ionic.
● Ionic materials are characteristically hard and
brittle and, electrically and thermally
insulative.
● These properties are directly related to
electron configurations and/or the nature of the
ionic bond.
ionic bonding
For two isolated ions, the attractive energy EA is a
function of the interatomic distance
EA = - A/r
An analogous equation for the repulsive energy is
ER = B/rn
A, B, and n are constants whose values depend on
the particular ionic system.
The value of n is approximately 8.
Ionic Bonding
 Energy – minimum energy most stable
 Energy balance of attractive and repulsive terms
EN = EA + ER =

Repulsive energy ER
Interatomic separation r
Net energy EN
Attractive energy EA
A
r

B
rn
Bonding Forces and Energies
calculate the force of attraction between a K+ and
an O2- ion separated by r0 =1.5 nm.
The attractive force between two ions FA is the
derivative with respect to the interatomic separation
of the attractive energy expression,
dEA
FA= 
dr
A
)


(
d
2
A
r
=  = (  )
r
dr
A
1
=  (
4 0

Z1e )( Z2e )
0 is the permittivity of vacuum (8.85x10-12 F/m).
Z1 & Z2 are the valences of the two ion types
e is the electronic charge (1.602x10-19 C).
Bonding Forces and Energies
Since the valences of the K+ and O2- ions
(Z1 and Z2) are +1 and -2, respectively,
Z1 = 1 and Z2 = 2, then
FA
(Z1e) (Z 2 e)
=
40r 2
(1)(2)(1.602  1019 C) 2
=
(4)() (8.85  1012 F/m) (1.5  109 m) 2

=2.05  10-10 N
covalent bonding
● stable electron configurations are assumed by the
sharing of electrons between adjacent atoms.
● Two atoms that are covalently bonded will each
contribute at least one electron to the bond, and
the shared electrons may be considered to belong
to both atoms.
● The covalent bond is directional; it is between
specific atoms and may exist only in the direction
between one atom and another that participates
in the electron sharing.
Covalent Bonding
similar electronegativity  share electrons
bonds are determined by valence –
s & p orbitals dominate bonding
shared electrons
Example: CH4
H
of carbon atom
CH 4
H
C: has 4 valence e-,
∙∙
needs 4 more
H: C : H
H
H
C
∙∙
H: has 1 valence e-,
H
needs 1 more
H
shared electrons
Electronegativities
of hydrogen atom
are comparable.
●
●
●
●
Covalent Bonding
●
●
●
●
The bonds between oxygen and hydrogen in a
water molecule are covalent bonds.
There are two covalent bonds in a water
molecule, between the oxygen and each of the
hydrogen atoms.
Each bond represents one
electron.
In a covalent bond,
electrons are shared
between atoms,
not transferred.
Covalent Bonding
Fluorine has seven valence electrons
 A second atom also has seven
 By sharing electrons
 Both end with full orbitals

F F
8 Valence
electrons
Writing Lewis Structures
The Lewis structure contains the element symbol with dots
representing electrons.
The only electrons
shown are those on
the outer energy
level or valence
electrons.
The electrons are
placed around the
element symbol,
one at a time,
clockwise or
counterclockwise,
and then grouped in pairs as more electrons are added.
Covalent bonding in H2 molecule
Covalent bonding in H2O molecule
covalent bonding
● Covalent bonds may be very strong, as in
diamond, which is very hard and has a very high
melting temperature, 3550 C, or they may be
very weak, as with bismuth, which melts at about
270 C.
● Polymeric materials typify this bond, the basic
molecular structure often being a long chain of
carbon atoms that are covalently bonded together
with two of their available four bonds per atom.
● The remaining two bonds normally are shared
with other atoms, which also covalently bond.
covalent bonding
● interatomic bonds may be partially ionic and
partially covalent.
● very few compounds exhibit pure ionic or covalent
bonding.
● the degree of either bond type depends on the
relative positions of the components in the periodic
table or the difference in their electronegativities.
● The wider the separation (the greater the difference
in electronegativity), the more ionic the bond.
● the closer they are (the smaller the difference in
electronegativity), the greater the degree of
covalency.
interatomic bonding
No electronegativity difference between two
atoms leads to a purely non-polar covalent bond.
A
B
A small electronegativity difference leads to a
polar covalent bond.
A
B
A large electronegativity difference leads to an
ionic bond.
ionic-covalent mixed bonding
 Ionic-Covalent Mixed Bonding
ionic character =

(X A X B )2 
 

4
1
e






x (100 %)
where XA & XB are Pauling electronegativities
Ex: MgO

XMg = 1.3
XO = 3.5

(3.5 1.3 )2


4
% ionic character  1  e




 x (100%)  70.2% ionic


metallic bonding
● found in metals and their alloys.
● Metallic materials have one, two, or at most,
three valence electrons.
● these valence electrons are more or less free to
drift throughout the entire metal. They may be
thought of as forming a “sea of electrons” or an
“electron cloud”.
● The remaining nonvalence electrons and atomic
nuclei form what are called ion cores, which
possess a net positive charge equal in magnitude
to the total valence electron charge per atom.
metallic bonding
● the metallic bond is nondirectional in character. In
addition, these free electrons act as a “glue” to
hold the ion cores together.
● Bonding may be weak or strong; energies range
from 68 kJ/mol (0.7 eV/atom) for mercury to 849
kJ/mol (8.8 eV/atom) for tungsten.
Their respective melting temperatures are 39 and
3410 C.
● Metallic bonding is found in the periodic table for
Group IA and IIA elements and, in fact, for all
elemental metals.
metallic bonding
Metallic bonding
Primary Bonding-summary
• Covalent bonds can be strong
e.g, Diamond melting point >3550°C
or, covalent bonds can be weak
e.g, Bismuth melting point: 270°C
• Polymers: Covalent bonds
• Partially ionic + partially covalent: possible
• Wider separation in the periodic table: Ionic
• Closer together in the periodic table: Covalent
interatomic bonding
Bonding energies correlate with melting points.
Secondary-van der waals-bonding
Secondary, van der Waals, or physical bonds are
weak in comparison to the primary or chemical
ones; bonding energies are typically on the order
of only 10 kJ/mol (0.1 eV/atom).
Secondary bonding exists between virtually all
atoms or molecules, but its presence may be
obscured if any of the three primary bonding types
is present.
Secondary bonding is evidenced for the inert gases,
which have stable electron structures, and, in
addition, between molecules in molecular
structures that are covalently bonded.
Secondary-van der waals- bonding
Secondary bonding forces arise from atomic or
molecular dipoles. In essence, an electric dipole
exists whenever there is some separation of positive
and negative portions of an atom or molecule.
The bonding results from the coulombic attraction
between the positive end of one dipole and the
negative region of an adjacent one
Hydrogen bonding, a special type of secondary bonding, is found to exist
between some molecules that have hydrogen as one of the constituents.
Schematic illustration of van der Waals
bonding between two dipoles
Secondary bonding
Molecular dipoles occur due to the unequal sharing of
electrons between atoms in a molecule.
More electronegative atoms pull the bonded electrons closer
to themselves.
This results in a molecular dipole in which one side of the
molecule possesses a partially negative charge and the other
side a partially positive charge.
Arises from interaction between dipoles
ex: liquid H2
asymmetric electron
H2
clouds
+
-
secondary
bonding
+
-
H H
H2
H H
secondary
bonding
Properties linked with Bonding:Tm
● Melting Temperature, Tm
Energy
ro
r
smaller Tm
higher Tm
Tm is larger if Eo is larger.
Properties From Bonding : 
• Coefficient of thermal expansion, 
length, L o
unheated, T 1
coeff. thermal expansion
DL
DL
=  ( T 2 -T 1)
Lo
heated, T 2
Energy
• a ~ symmetry at ro
Eo
Eo
unstretched length
ro
r
larger 
smaller 
 is larger if Eo is smaller.
Properties From Bonding: E
Elastic modulus
F
L
=E
Ao
Lo
E similar to spring constant
Energy
unstretched length
ro
r
E is larger if curvature is larger.
smaller Elastic Modulus
larger Elastic Modulus
Summary
● Polymers, with weak, secondary, intermolecular
bonds (low melting points) have very high
expansion coefficients.
● Ceramics which are strongly bonded (i.e., ionic
or network covalent) have low thermal
expansion coefficients.
● Metals with high melting points (strong bonding)
have low thermal expansion coefficients. Low
melting point metals have high thermal
expansion coefficients.
Summary
The Periodic Table
● Elements in each of the columns (or
groups) of the periodic table have
distinctive electron configurations.
● For example, Group 0 elements (the inert
gases) have filled electron shells, and
● Group IA elements (the alkali metals) have
one electron other than a filled electron
shell.
Summary
● There is a sharing of valence electrons between
adjacent atoms when bonding is covalent.
● Polymers and some ceramic materials bond
covalently.
● The percent ionic character (%IC) of a bond
between two elements (A and B) depends on
their electronegativities (X’s).
● Relatively weak van der Waals bonds result from
attractive forces between electric dipoles, which
may be induced or permanent.
Summary: Bonding
Type
Bond Energy
Comments
Ionic
Large!
Nondirectional (ceramics)
Covalent
Variable
large-Diamond
small-Bismuth
Directional
(semiconductors, ceramics
polymer chains)
Metallic
Variable
large-Tungsten
small-Mercury
Nondirectional (metals)
Secondary
smallest
Directional
inter-chain (polymer)
inter-molecular
Summary: Primary Bonds
Large bond energy
Ceramics
large Tm
(Ionic & covalent bonding):
large E
small a
Variable bond energy
Metals
moderate Tm
(Metallic bonding):
moderate E
moderate a
Polymers
(Covalent & Secondary):
Directional Properties
Secondary bonding dominates
small Tm
small E
large a
Learning check
Explain why covalently bonded materials are
generally less dense than ionically or
metallically bonded ones.
because covalent bonds are directional in
nature whereas metallic and ionic bonds are
not; when bonds are directional, the atoms
cannot pack together in as dense a manner,
yielding a lower mass density.
Learning check
If the difference in electronegativities between two atoms
is zero, the bonds are
a) Non polar covalent
b) Polar covalent
c) Mostly ionic
d) Slightly ionic
The bond between O and H in OH
a) Nonpolar covalent
b) Very polar
c) Slightly polar covalent
d) Mostly ionic
Learning check
Electronegativity refers to
a) the degree of negative charge on an electron
b) the energy required to remove an electron from a gaseous
atom in the ground state
c) the ability of an atom to attract the electrons in a
covalent bond toward itself
d) the energy change that occurs when an electron is
accepted by a gaseous atom to form an anion
In general, electronegativity _______ going left to
right across a row in the periodic table
a) decreases
b) İncreases
c) Does not change
d) None of the above
Learning check
The most electronegative elements are
a) found in the upper right corner of the periodic table
b) the alkali metals
c) The alkaline earth metals
d) The transition elements
The level of attraction of one atom for electrons when
bonding with another atom is called
a) ionization energy
b) An ionic bond
c) A nonpolar covalent bond
d) electronegativity
Learning check
When sodium and chlorine react, chlorine removes
sodium's valence electron and __________ forms
between them
a) a covalent bond
b) a nonpolar covalent bond
c) an ionic bond
d) A polar covalent bond
When an electron pair is shared between two atoms of
equal electronegativity,
a) a nonpolar covalent bond is formed
b) An ionic bond is formed
c) a polar covalent bond is formed
d) electron transfer occurs
Learning check
A polar covalent bond results from
a) a transfer of electrons to the atom of least
electronegativity
b) an equal sharing of an electron pair between two atoms
c) the formation of oppositely charged ions
d) None of the above
The type of bond that involves a cation and an anion is
_____.
a) nonpolar covalent
b) Metallic
c) Polar covalent
d) ionic
Learning check
Of the following, the most likely pair to form an ionic
bond is ….
a) an alkali metal and an alkaline earth metal
b) a halogen and an alkaline earth metal
c) a halogen and a metalloid
d) an alkaline earth metal and a transition element
Which of the following is in an ionic bond?
a) F2
b) MgCl2
c) NO
d) H2O
Learning check
In a crystalline compound, each anion is surrounded
by…...
a) Negative ions
b) Molecules
c) Positive ions
d) Dipoles
When a metal forms an ionic bond with a non-metal, the
nonmetal atom will …………...
a) gain an electron and become a positive ion
b) lose an electron and become a positive ion
c) lose and electron and become a negative ion
d) gain an electron and become a negative ion
Learning check
The less the electronegativity differences between two
bonded atoms, the greater the ……………...
a) polar character
b) ionic character
c) Metallic character
d) Covalent character
What type of bonding does NaCl have?
a) polar covalent
b) Metallic
c) Nonpolar covalent
d) ionic
Learning check
How are bond length and bond energies related?
a) the higher the bond energy, the shorter the bond
length
b) the lower the bond energy, the shorter the bond length
c) they are not related
d) the higher the bond energy, the longer the bond length
What determines bond length?
a) the distance at which potential energy is at a minimum
b) the distance at which the two atoms are as close as
possible
c) the distance at whch potential energy is at a maximum
d) the point at which the attraction forces outweighs the
repulsion forces
Learning check
How are thermal expansion coefficient () and bond
energies related?
a) the higher the bond energy, the smaller 
b) the lower the bond energy, the smaller 
c) they are not related
d) the higher the bond energy, the larger 
Bonding between nonmetals and nonmetals primarily
involves?
a. interactions between protons, electrons, and neutrons
b. interactions between protons
c. interactions between protons and electrons
d. transfer of electrons
e. sharing of electrons
Learning check
A bond in which electrons are shared equally is known
as ……………..
a) polar covalent
b) Metallic
c) İonic
d) Non-polar covalent
Mobile electrons within bonding networks best
describes which type of bond?
a) Metallic
b) İonic
c) Polar covalent
d) Non-polar covalent
see you next week!