General Chemistry I
Dr. PHAN TẠI HUÂN
Faculty of Food Science and Technology
Nong Lam University
Module 2:
Chemical bonds and molecular configurations
• Chemical bonds (ionic, covalent, metallic, hydrogen…).
• Valence Bond Theory, VSEPR model, three-dimensional
molecular structures.
• Covalent bond, Hybridization.
• Molecular orbital theory: molecular orbitals, Sigma and Pi
orbitals, Homonuclear and Hetero-nuclear diatomic
molecules.
• Examples upon configurations of organic and inorganic
compounds.
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Chemical bond
What are chemical bonds?
• Chemical bonds are strong attractive forces that enable
atoms or groups of atoms to hold together.
• The two major categories of chemical bonds are ionic and
covalent bonds.
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Lewis Electron-Dot Formulas
• Lewis
electron-dot
formulas
are
diagrammatic
representations of the atoms involved and their valence
electrons.
• The valence electrons (the electrons in the outermost
occupied s and p orbitals) are usually represented as dots
around the elemental symbol.
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Resonance
• Resonance is an important concept in chemistry. Though
we represent definite Lewis structures of molecules, in
actuality the electrons are not localized. They are shared
and delocalized by atoms in such a way as to be in the
most stable electron distribution. This is called resonance.
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3
Ionic Bond
• The major force behind the formation of an ionic bond is
the electrostatic attractive force that exists between
negative and positive ions.
• It is formed by the transfer of one or more electrons from
one atom to another.
• The atom that donates the electrons becomes positive
(cation), and the counterpart atom that receives those
electrons becomes negative (anion).
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Ionic Bond
• Many metals are easily oxidized—that is, they lose
electrons to form cations.
• Many nonmetals are readily reduced—that is, they gain
electrons to form anions.
• When the electronegativity difference between two
elements is large, as between a metal and a nonmetal, the
elements are likely to form a compound by ionic bonding.
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4
Group IA metals and Group VIIA nonmetals
mp 801°C
(melting point)
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Group IA metals and Group VIA nonmetals
mp > 1700°C
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Group IIA metals and Group VIA nonmetals
mp 2580°C
Ca
Ca2+ + 2e-
and
O + 2e-
O2-
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Simple Binary Ionic Compounds
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The attractive energy in an ionic bond
• The energy of interaction between a pair of ions can be
calculated by using Coulomb´s law.
Where V has units of joules, r is the distance between the ion
centers in nanometers, Q1 and Q2 are the numerical ion
charges, and εo is the permittivity of the vacuum.
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Exercise
• Sodium cloride (r = 276 pm).
• How much is the ionic energy per pair of ions?
• How much is the energy of interaction for a mole of pairs
of Na+ and Cl- ions?
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Crystal lattice energy
This energy refers to a mole of Na+..Cl- ion pairs in the gas phase
where a given pair is far from any other pair.
In solid NaCl, which contains a large array of closely packed Na+
and Cl- ions, where a given ion is close to many oppositely
charged ions, the energy associated with ionic bonding is much
greater than 504 kJ/mol because of the larger numbers of
interaction ions. The crystal lattice energy of NaCl is 789 kJ/mol.
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Energy relationships in ionic bonding
• There is a general tendency in nature to achieve stability.
One way to do this is by lowering potential energy; lower
energies generally represent more stable arrangements.
• Ionic bonding occurs between elements with low ionization
energies and those with high electronegativities.
Example:
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Covalent bond – Hydrogen model
• When electrons are in the region between two nuclei,
attractive electrical forces exceed repulsive electrical
forces, leading to the stable arrangement of a chemical
bond.
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Covalent bond formation
• The separation distance where the molecule is most stable
is known as the bond length.
• The amount of stability at this separation distance is the
bond energy, or strength of the bond.
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Covalent bond formation
H2
F2
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Covalent bond
• A covalent bond is formed as a result of the sharing of a
pair of electrons between atoms.
•
Covalent bonds result when the difference in
electronegativities between the bonding atoms is very
small.
• Though the intramolecular bonds of covalent compounds
are significant, the intermolecular forces are relatively
weak. Because of this, covalent compounds have relatively
lower boiling and melting points when compared to ionic
compounds.
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Octet rule
• The tendency to achieve an octet of electrons in
the outermost shell is called the octet rule.
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Octet rule and Lewis formulas
• How to place the electrons around the bonded atoms?
– How many of the available valence electrons are
bonding electrons (shared)?
– How many are unshared electrons (associated with only
one atom)?
• A pair of unshared electrons in the same orbital is called a
lone pair.
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Octet rule and Lewis formulas
The total number (S) of electrons shared in the molecule or
polyatomic ion:
S=N- A
• N is the total number of valence shell electrons needed by
all the atoms in the molecule or ion to achieve noble gas
configurations (N= 8 x number of atoms that are not H,
plus 2 x number of H atoms).
• A is the number of electrons available in the valence shells
of all of the atoms.
• For ionic charges. We add electrons to account for
negative charges and subtract electrons to account for
positive charges.
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Exercise
Write the Lewis formula for F2
Write the Lewis formula for HF
Write the Lewis formula for H2O
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Exercise
Write the Lewis formula for CO2
Write the Lewis formula for NH4+
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Coordinate covalent bond
• Coordinate covalent bond is formed when the two
electrons that are shared in the formation of the bond are
donated by one group or atom involved in the bond.
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Polar and nonpolar covalent bonds
• Nonpolar covalent compound: when two atoms combine,
just like in the formation of a hydrogen molecule (H2), the
atoms are one and the same and they have the same
electronegativity.
• Polar covalent compound: if the two atoms that are
combined via covalent bond are different (ex: HF), then
there is unequal sharing of electrons due to the
electronegativity difference between those atoms.
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Dipole moment
• The delta+ (δ+) indicates the “ partial positive charge“ of
the H atom, and delta-(δ-) indicates the “ partial negative
charge “ of the F atom.
Electronegativity
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Dipole moment
• The dipole moment, μ, is defined as the product of
the distance, d, separating charges of equal magnitude
and opposite sign, and the magnitude of the charge, q.
μ= d x q
1 D (debye) = 3.336 x 10-30 C.m
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The continuous range of bonding types
• The degree of electron sharing or transfer depends on the
electronegativity difference between the bonding atoms.
• Nonpolar covalent bonding (involving equal sharing of
electron pairs) is one extreme, occurring when the atoms
are identical ((EN) is zero).
• Ionic bonding (involving complete transfer of electrons)
represents the other extreme, and occurs when two
elements with very different electronegativities interact
((EN) is large).
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Dipole-Dipole bonding
• Permanent dipole–dipole interactions occur between polar
covalent molecules because of the attraction of the atoms of
one molecule to the atoms of another molecule.
• Average dipole–dipole energies are approximately 4 kJ per
mole of bonds.
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Hydrogen Bonding
• A hydrogen bond (H bond) forms when a hydrogen
atom is shared by two electronegative atoms. The atom
to which the hydrogen is covalently bonded is referred to
as the hydrogen bond donor, and the other atom is
referred to as the hydrogen bond acceptor.
• Typical hydrogen-bond energies are in the range 15 to
20 kJ/mol.
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Van der Waals bond
• The distribution of electrons around an atom is not fixed;
rather, the character of the so-called electron cloud
fluctuates with time.
• Through these fluctuations, a transient asymmetry of
electron distribution, or dipole moment, can be established.
• When atoms are close enough together, this asymmetry on
one atom can influence the electronic distribution of
neighboring atoms.
• The result is a similar redistribution of electron density in
the neighbors, hence an attractive force between the atoms
is developed. This attractive force, referred to as a Van der
Waals bond.
• Typically a van der Waals bond is worth only about 1
kcal/mol.
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Valence shell electron pair repulsion
(VSEPR) theory
• The sets of valence shell electrons on the central atom
repel one another. They are arranged about the central
atom so that repulsions among them are as small as
possible.
=> This results in maximum separation of the regions of
high electron density about the central atom.
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VSEPR Theory
The number of regions of high electron density around
the central atom (steric number):
1. Each bonded atom (ligand) is counted as one region of
high electron density, whether the bonding is single,
double, or triple.
2. Each unshared pair of valence electrons (lone pair) on the
central atom is counted as one region of high electron
density.
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Examples
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VSEPR Theory
• According to VSEPR theory, the structure is most stable
when the regions of high electron density on the central
atom are as far apart as possible.
• The arrangement of these regions of high electron density
around the central atom is referred to as the electronic
geometry of the central atom.
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Bond angles
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VSEPR Theory
• After we know the electronic geometry, we consider how
many of these regions of high electron density connect
(bond) the central atom to other atoms.
• This lets us deduce the arrangement of atoms around the
central atom, called the molecular geometry.
• Name of the molecular structure is always
based on the positions of the atoms
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Example of NH3
• The molecular structure of NH3 is a trigonal
pyramid, rather than a tetrahedron.
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More about bond angles
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Example SF4
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Polar molecules:
The influence of molecular geometry
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Examples
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Valence bond (VB) Theory
• VSEPR theory to describe the orientations of the regions
of high electron density.
• VB theory (Localized electron model) to describe the
atomic orbitals that overlap to produce the bonding with
that geometry.
=> Thus, the two theories work together to give a fuller
description of the bonding.
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Valence bond (VB) Theory
• When other atoms are nearby as in a molecule or ion, an
atom can combine its valence shell orbitals to form a new
set of orbitals that is at a lower total energy in the presence
of the other atoms than the pure atomic orbitals would be.
• This process is called hybridization, and the new orbitals
that are formed are called hybrid orbitals.
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sp3 hybridization
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General Features of Hybridization
1. The number of valence orbitals generated by the
hybridization process equals the number of valence
atomic orbitals participating in hybridization.
2. The steric number (region of high electron density) of an
inner atom uniquely determines the number and type of
hybrid orbitals.
3. Hybrid orbitals form localized bonds by overlap with
atomic orbitals or with other hybrid orbitals.
4. There is no need to hybridize orbitals on outer atoms.
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sp2 hybridization
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sp hybridization
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sp3d hybridization
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sp3d2 hybridization
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Describing a molecule with VB theory
In summary:
• Draw the Lewis structure(s).
• Determine the arrangement of electron pairs, using the
VSEPR theory.
• Specify the hydrid orbitals needed to accomodate the
electron pairs.
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General Features of Hybridization
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Compounds containing double bonds
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Compounds containing double bonds
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Sigma (σ) and Pi (π) bond
• A sigma (σ) bond is a bond resulting from head-on overlap
of atomic orbitals. The region of electron sharing is along
and cylindrically around an imaginary line connecting the
bonded atoms.
• A pi (π) bond is a bond resulting from side-on overlap of
atomic orbitals. The regions of electron sharing are on
opposite sides of an imaginary line connecting the bonded
atoms and parallel to this line.
(A pi bond can form only if there is also a sigma bond
between the same two atoms.)
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Compounds containing triple bonds
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Molecular Orbital Theory
Valence bond theory has some problems:
• It incorectly assumes that electrons are localized, the concept
of resonance must be added.
• It is not easily to deal with moleculars containing unpaired
electrons.
• It gives no direct information about bond energies.
⇒ Molecular orbital theory:
Molecules have a set of molecular orbitals made from the
combination of atomic orbitals.
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Molecular Orbital Theory
• Some molecular orbitals are called “bonding orbitals“,
some are “nonbonding“ and still others are called
“antibonding orbitals“.
• Recall that the orbital shape can be described using a wave
function.
• When waves meet they can add to each other or, if they are
out of phase, cancel each other.
• Bonding orbitals are additive, antibonding orbitals are
those in which the waves cancel each other.
• When there is no overlap, the waves cannot interat and the
orbitals are nonbonding.
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Molecular orbital of hydrogen
When two hydrogen orbitals interact, they generate two
molecular orbitals:
• One bonding σ
MO1 = 1sA + 1sB
• One antibonding σ*
MO2 = 1sA - 1sB
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Evidence for Antibonding Orbitals
• The Lewis structure and valence bond theory predict that
all the electrons in O2 are paired and that O2 should be
diamagnetic.
• The experiment shows that liquid oxygen adheres to the
poles of a magnet. Attraction to a magnetic field shows that
molecular oxygen is paramagnetic.
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Homonuclear diatomic molecule
• When two identical atoms form a molecule, we call it a
homonuclear diatomic molecule.
• The energy diagram here (for O2) can be used to decribe the
bonding in p-block molecules of the type X2.
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Bond order
• Bond order is the quantitative indicator of molecular
stability (bond strength) of a diatomic molecule.
• Bond order = ½ (number of bonding electron – number of
antibonding electrons).
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Homonuclear diatomic molecule
For second-row diatomic molecules and ions:
• σp is lower in energy than π when Zaverage > 7.
• π is lower in energy than σp when Zaverage ≤ 7.
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Metallic bonding: Band theory
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Conductors
Conductors:
• Have a partially filled valence band.
• Require a small amount of energy to
jump to unfilled energy levels.
• Conduct electricity through emotion
of electrons in signly occupied levels.
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Insulators and Semiconductors
Insulators:
• Have a filled valence band.
• Have an empty conduction band.
• Have a large energy gap (≥ 500 kJ/mol)
between valence and conduction bands,
allowing no flow of electrons through filled
levels.
Semiconductors:
• Have moderate band energy gaps (50 – 300
kJ/mol).
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Summary
After you have studied this module, you should be able to
• Write Lewis dot representations of atoms.
• Predict whether bonding between specified elements will
be primarily ionic, covalent, or polar covalent.
• Compare and contrast characteristics of ionic and covalent
compounds.
• Describe how the elements bond by electron transfer (ionic
bonding).
• Describe energy relationships in ionic compounds.
• Write Lewis dot and dash formulas for molecules and
polyatomic ions.
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Summary
• Write formal charges for atoms in covalent structures.
• Describe resonance, and know when to write resonance
structures and how to do so.
• Relate the nature of bonding to electronegativity
differences.
• Describe the basic ideas of the valence shell electron pair
repulsion (VSEPR) theory.
• Use the VSEPR theory to predict the electronic geometry.
• Use the VSEPR theory to predict the molecular geometry.
• Describe how elements bond by sharing electrons
(covalent bonding).
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Summary
• Describe the relationships between molecular shapes and
molecular polarities.
• Predict whether a molecule is polar or nonpolar.
• Describe the basic ideas of the valence bond (VB) theory.
• Analyze the hybrid orbitals used in bonding in polyatomic
molecules and ions.
• Use hybrid orbitals to describe the bonding in double and
triple bonds.
• Describe the basic ideas of the Molecular orbital.
• Describe the basic ideas of the Metallic band theory.
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