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Chapter 3

Chemical Bonding

Ref: Ralph H. Petrucci, F. Geoffrey Herring, Jeffry D. Madura, Carey Bissonnette. General Chemistry:

Principles and Modern Applications. Prentice Hall 2010, English 10th ed.

Theodore E. Brown, H. Eugene H LeMay, Bruce E. Bursten, Catherine Murphy, Patrick Woodward.

Chemistry: The Central Science. Prentice Hall 2011, English 12th ed.

Raymond Chang and Jason Overby. General Chemistry: The Essential Concepts. McGraw-Hill

Science/Engineering/Math 2010, 6th ed.

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Lewis Theory1. Electrons, especially those of the outermost (valence) electronic shell, play a

fundamental role in chemical bonding.

2. In some cases, electrons are transferred from one atom to another. Positive

and negative ions are formed and attract each other through electrostatic forces

called ionic bonds.

3. In other cases, one or more pairs of electrons are shared between atoms. A

bond formed by the sharing of electrons between atoms is called a covalent

bond.

4. Electrons are transferred or shared in such a way that each atom acquires an

especially stable electron configuration. Usually this is a noble gas configuration,

one with eight outer-shell electrons, or an octet.

Lewis Theory

A Lewis symbol consists of a chemical symbol to represent the nucleus and core

(inner-shell) electrons of an atom, together with dots placed around the symbol to

represent the valence (outer-shell) electrons.

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Lewis Structure

Write the Lewis Structure for nitrogen trifluoride (NF3), nitric acid (HNO3),

carbonate (CO32-)

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Some Exceptions to the Octet Rule

1. The Incomplete Octet

Fewer than eight :

Group 2A (stable with 4 e-), BeH2,

Group 3A (stable with 6 e-), BF3

2. Odd-Electron Molecules: NO, NO2

3. The Expanded Octate

Atoms of the second-period elements cannot have more than eight

valence electrons around the central atom, but atoms of elements in

and beyond the third period of the periodic table form some

compounds in which more than eight electrons surround the central

atom. SF6, PCl5, XeF4

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Formal Charge and Lewis Structure

Formal Charge is the electrical charge difference between the valence

electrons in an isolated atom and the number of electrons assigned to the atom

in a Lewis Structure.

FC = [(number valence e- in free atom) – (number lone-pair e-) - ½

(number bond-pair e-)]

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Formal Charge and Lewis Structure

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Calculate the formal charge of the following molecules:

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Resonance

The two resonance structure are equivalent.14

Resonance

Which resonance form of azide anion, N3- are the most stable.

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Covalent Bonding

Bonding between a hydrogen atom and a chlorine atom involves the sharing of

electrons, which leads to a covalent bond.

The sharing of a single pair of electrons between bonded atoms produces a

single covalent bond.

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Coordinate Covalent Bonds

The Lewis theory of bonding describes a covalent bond as the sharing of a pair of

electrons, but this does not necessarily mean that each atom contributes an

electron to the bond. A covalent bond in which a single atom contributes both of

the electrons to a shared pair is called a coordinate covalent bond.

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Multiple Covalent Bonds

The Covalent Bond

A bond in which two electrons are shared by two atoms

Low ∆EN, H2, F2,

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Types of Covalent Bonds

Single Bond – two atoms are held together by one electron

pair.

Multiple bonds – two atoms share two or more pairs of

electrons.

Double bond – two atoms share two pairs of electrons

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Types of covalent bonds

Multiple bonds – two atoms share two or more pairs of electrons.

Triple bond – two atoms share three pairs of electrons

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Electronegativity

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Polar Covalent Bond

Even though, covalent bond is formed from the two atoms sharing

their electrons, if the atoms have different EN the electrons

will spend more time in the vicinity of one atom that has higher

EN than the other.

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Bond Enthalpy

A measurement of the stability of a molecule is its “bond enthalpy”

which is the enthalpy change required to break a particular bond

in 1 mole of gaseous molecules.

Bond enthalpies in solids and liquids are affected by neighboring

molecules.

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Bond Enthalpy

Measuring the strength of covalent bonds in polyatomic molecules

is more complicated.

In each case, and O-H bond is broken, but the first step is more

endothermic than the second. The difference between the two ∆Ho values

suggests that the second O-H bond itself has undergone change, because of

the changes in the chemical environment. Thus, for polyatomic molecules

the bond enthalpy is the average of a particular bond.

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Shapes of Molecules

Think about the molecules in 3D

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Valence shell electron-pair repulsion

(VSEPR)

The basis of the VSEPR approach is that electrons repel each

other because they are negatively charged.

Molecules adopt geometries such that valence electron pairs

position themselves as far from each other as possible to

minimize electron–electron repulsions.

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Valence shell electron-pair repulsion

(VSEPR)

A is the central atom, X stands for any atom or group of atoms

surrounding the central atom, and E represents a lone pair of electrons.

The steric number (SN) = m + n

CO2 --SN = 2

SO3 -- SN = 3

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Valence-Shell Electron-Pair Repulsion

(VSEPR) Theory

Central Atoms with Two Electron Groups

Central Atoms with Three Electron Groups

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Valence-Shell Electron-Pair Repulsion

(VSEPR) Theory

Central Atoms with Three Electron Groups

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Valence-Shell Electron-Pair Repulsion

(VSEPR) TheoryCentral Atoms with Four Electron Groups

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Lone-pair repulsion

The VSEPR model predicts that electron-pair repulsions involving

lone pairs ( lp ) are stronger than those involving bonding pairs ( bp ) in

the order

lp-lp repulsions > lp-bp repulsions > bp-bp repulsions

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Possible Structures of SF 4

Possible

Structures of ClF3

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Multiple bonds

The VSEPR model considers double and triple bonds to have

slightly greater repulsive effects than single bonds because of the

repulsive effect of ¶ electrons that increase the electron density

between the bonded atoms beyond that present in a σ bond.

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Electronegativity and atomic size effects

Electronegativity is a measure of an atom’s ability to attract

electrons from a neighboring atom to which it is bonded; it can

be viewed as the ability of an atom to win the competition to

attract shared electrons.

If the central atom remains the same, molecules that have a larger

difference in electronegativity values between their central and

outer atoms have smaller bond angles.

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Molecules Having Steric Number = 5

For main group atoms having a steric number of 5, it is

instructive to consider the relative bond lengths for axial and

equatorial positions.

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There is a tendency for less electronegative groups to occupy

equatorial positions, similar to lone pairs and multiply bonded

atoms.

Molecules Having Steric Number = 5

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MOLECULES HAVING STERIC NUMBER = 5

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Valence-Shell Electron-Pair Repulsion

(VSEPR) Theory

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Dipole Moment

There is a shift of electron density from H to F because the F

atom is more electronegative than the H atom.

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Hybrid Orbitals

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HYBRID ORBITALS

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Molecular Orbitals

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Ref: Ralph H. Petrucci, F. Geoffrey Herring, Jeffry D. Madura, Carey Bissonnette. General Chemistry:

Principles and Modern Applications. Prentice Hall 2010, English 10th ed.

Theodore E. Brown, H. Eugene H LeMay, Bruce E. Bursten, Catherine Murphy, Patrick Woodward.

Chemistry: The Central Science. Prentice Hall 2011, English 12th ed.

Raymond Chang and Jason Overby. General Chemistry: The Essential Concepts. McGraw-Hill

Science/Engineering/Math 2010, 6th ed.

Formation of Molecular Orbitals from Atomic Orbitals

Schrodinger equations can be written for electrons in molecules.

Approximate solutions to these molecular Schrodinger equations

can be constructed from linear combinations of atomic

orbitals (LCAO), the sums and differences of the atomic wave

functions.

are atomic wave functions for atoms a and b

are adjustable coefficients that quantify the contribution of

each atomic orbital to the molecular orbital.62

Sigma Bonding of s-Orbitals

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Sigma Bonding of p-Orbitals

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Pi Bonding of p-Orbitals

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Bonding from P-orbitals Interaction

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How does this interaction cause the bonding?

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Sigma Bonding

The σ (sigma) notation indicates orbitals that are symmetric to rotation about the line connecting the nuclei.

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Pi Bonding

The ¶ (pi) notation indicates a change in sign of the wave function

with C2 rotation about the bond axis

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Molecular Orbitals from d-orbitals

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Energy Match and Molecular Orbital Formation

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Homonuclear Diatomic Molecules

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Orbital

mixing

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