bond energy- energy required to

break a chemical bond

-We can measure bond energy to

determine strength of interaction

ionic compound- a metal reacts

with a nonmetal

• Ionic bonds form when an atom that loses

electrons easily reacts with an atom that has

a high affinity for electrons. The charged

ions are held together by their mutual

attraction (Coulombic attraction).

• Ionic bonds form because the ion pair has

lower energy than the separated ions. All

bonds form in order to reach a lower energy

level.

Bond length- the distance where the

energy is at a minimum. We have a

balance among proton-proton repulsion,

electron-electron repulsion, and proton-

electron attraction.

In H2, the two e− will usually be found

between the two H atoms because they are

spontaneously attracted to both protons.

Therefore, electrons are shared by both

nuclei. This is called covalent bonding.

Polar covalent bonds occur when

electrons are not shared equally. One

end of the molecule may have a partial

charge. This is called a dipole.

+

H F H H

+ -

O

-

+ −

H—F polar H—H nonpolar

has dipole moment

O

+ +

H H O S

O bent, polar O

− has dipole moment planar

no dipole moment

CH4 tetrahedral NH3 trigonal pyramidal

no dipole moment has dipole moment

Electron Configurations:

Stable compounds usually have atoms with noble

gas electron configurations.

Two nonmetals react to form a covalent bond by

sharing electrons to gain valence electron

configurations.

When a nonmetal and a group A metal react to

form a binary ionic compound, the ions form

so that the valence electron configuration of

the nonmetal is completed and the valence

orbitals of the metal are emptied to give both

noble gas configurations.

Ions form to get noble gas configurations.

-exceptions in Group A metals:

Sn2+ & Sn4+

Pb2+ &Pb4+

Bi3+ & Bi5+

Tl+ & Tl 3+

Metals with d electrons will lose their highest

numerical energy level electrons before losing

their inner d electrons.

Size of Ions Positive ions (cations) are smaller than their

parent atoms since they are losing electrons. (More protons than electrons=greater nuclear pull)

Negative ions (anions) are larger than their

parent atoms since they are gaining electrons.

(Fewer protons than electrons= lower nuclear pull)

Think: Monster Ants & miniature cats

Ion size increases going down a

group.

Isoelectronic ions

–ions containing the same number of

electrons

O2−, F−, Na+, Mg2+, Al3+ all have the

Ne configuration. They are

isoelectronic.

*** For an isoelectronic series, size

decreases as Z increases.

Lattice energy- the change in energy that

takes place when separated gaseous ions

are packed together to form an ionic solid.

Na+(g) + Cl−(g) NaCl(s)

If exothermic, the sign will be negative and the

ionic solid will be the stable form.

We can use a variety of steps to determine the

heat of formation of an ionic solid from its

elements. This is called the Born-Haber cycle.

See examples on pages 366 & 368.

Lattice energy can be calculated using the following:

where k is a proportionality constant that depends on the

structure of the solid and the electron configuration of

the ions. Q1 & Q2 are the charges on the ions. r is

the distance between the center of the cation and the

anion.

Since the ions will have opposite charges, lattice energy

will be negative (exothermic).

The attractive force between a pair of oppositely

charged ions increases with increased charge on the

ions or with decreased ionic sizes.

r

QQk

21energy Lattice

The Structure of

Lithium Fluoride

The Relationship Between the Ionic Character of a Covalent Bond

and the Electronegativity Difference of the Bonded Atoms

Three Possible Types of Bonds

Polyatomic ions are held together by covalent

bonds. We call Na2SO4 ionic even though it

has 4 covalent bonds and 2 ionic bonds.

Ionic compound- any solid that conducts an

electrical current when melted or dissolved in

water

Salt- an ionic compound

A chemical bond is a model “invented” by

scientists to explain stability of compounds. A

bond really represents an amount of energy.

The bonding model helps us understand and

describe molecular structure. It is supported

by much research data. However, some data

suggests that electrons are delocalized. That

is, they are not associated with a particular

atom in a molecule.

• Single bond- one pair of shared electrons

• Double bond- two pair of shared electrons

• Triple bond- three pair of shared electrons

These values may be slightly different from those in your text. Use the

textbook values for your homework.

Looking at the chart on the

previous slide, what is the

relationship between bond

length and bond energy?

Is there a relationship between

number of bonds and bond

energy?

Bond energies and bond lengths are given in

tables on page 374.

We can use bond energies to calculate heats of

reaction.

H = D(bonds broken)- D(bonds formed)

2H2 + O2 2H2O

Ex. H = [2(432) + 495] –[4(467)] = −509 kJ

2 H−H O=O 4H−O

exothermic

Bonding Models:

Molecular Orbital Model-

Electrons occupy orbitals in a molecule in

much the same way as they occupy

orbitals in atoms.

Electrons do not belong to any one atom.

-very complex model

Localized electron model-

• molecules are composed if atoms that are

bound together by sharing pairs of electrons

using the atomic orbitals of the bound atoms

• traditional model

lone pair- pair of electrons

localized on an atom

(nonbonding)

shared pair or bonding pair-

electrons found in the space

between atoms

Lewis structure -shows how the

valence electrons are arranged

among the atoms in the molecule

The most important requirement for the

formation of a stable compound is that the atoms

achieve noble gas configurations

ionic [ Na ]+ [Cl]−

only valence electrons are included

molecular H2O

H – O - H

duet rule- hydrogen forms stable molecules

when it shares two electrons

H:H

-filled valence shell

Why does He not form bonds?

Its valence orbitals are already filled.

octet rule – most elements need 8 electrons to

complete their valence shell

Cl-Cl

Rules for writing Lewis structures

1. Add up the number of valence electrons from

all atoms.

2. Use 2 electrons to form a bond between each

pair of bound atoms. A dash represents a pair of

shared electrons.

3. Arrange the remaining electrons to satisfy the

duet rule for H and the octet rule for most others.

Ex. H2S

# of valence electrons: 1 + 1 + 6 = 8

H − S − H

Ex. CO2

# of valence electrons = 4 + 6 + 6 = 16

O – C – O This uses 20 electrons!

O = C = O

NH3 has 8 valence electrons

H

N− H

H

HCN

HCN has 10 valence electrons.

H−C≡N

NO+

NO+ has 5 + 6 −1 = 10 electrons

N≡O

CO32−

Carbonate has 4 + 18 + 2 = 24 valence electrons.

O 2−

C O

O

Exceptions:

Boron and beryllium tend to form compounds

where the B or Be atom have fewer than 8

electrons around them.

BF3 = 24 valence electrons

F

B F

F Common AP equation:

NH3 + BF3 H3NBF3

C, N, O, F always obey the

octet rule.

Some elements in Period 3 and beyond exceed

the octet rule.

Ex. SF6 S has 12 electrons around it

48 valence electrons

F

F F

F S F

F

d orbitals are used to accommodate the

extra electrons.

Elements in the 1st or 2nd period of the

table can’t exceed the octet rule

because there is no d sublevel.

If the octet rule can be exceeded, the

extra electrons are placed on the

central atom.

Examples of exceptions

Ex. I3−, ClF3, RnCl2

I - I - I F

F - Cl - F

Cl - Rn - Cl

Resonance-

-occurs when more than one valid Lewis

structure can be written for a particular molecule

actual structure is an average of all resonance

structures

-this concept is needed to fit the localized

electron model (electrons are really delocalized)

Ex. Benzene, C6H6

All bond lengths and angles are the same.

Ex. SO3

All 3 structures are equivalent. The bonds

can be thought of as 1 1/3 bonds.

Formal Charge

-used to determine the most accurate

Lewis structure

-is the difference between the # of

valence electrons on the free atom and

the # of valence electrons assigned to

the atom in the molecule

-atoms try to achieve formal

charges as close to zero as possible

-any negative formal charges are

expected to reside on the most

electronegative atoms

-Sum of the formal charges must

equal the overall charge on the

molecule (zero) or ion.

Ex. SO42−

O 2− O 2−

O S O O S O

O O

Formal charge only needs to be considered on the AP test if it is specifically asked for.

VSEPR-Valence Shell Electron Pair

Repulsion

-allows us to use electron dot structures to

determine molecular shapes

-the structure around a given atom is

determined primarily by minimizing

electron repulsions

-bonding and nonbonding pairs of electrons

around an atom position themselves as far

apart as possible

Steps:

1. Draw Lewis structure

2. Count effective electron pairs on central atom

(double and triple bonds count as one)

3. Arrange the electron pairs as far apart as

possible

Shapes

AX2 (A represents central atom, X represents

attached atom, E represents unshared electron

pair)

X – A – X linear 180o bond angle

O=C=O Cl – Be – Cl

AX3 Shape is trigonal planar

X X

A 120o bond angle

F F

X BF3 B

F

Any resonance SO3

structure can

be used to O− S = O

determine shape. O

AX2E Shape is bent

Bond angle is < 120o

X X

A

E

Ex. SnCl2 Cl Cl

Sn

AX4 Shape is tetrahedral

Bond angle is 109.5o

X Ex. CH4 H

X A X H C H

X H

Figure

8.14 The

Molecular

Structure

of

Methane

AX3E Shape is trigonal pyramidal

Bond angle is < 109.5o

Ex. NH3

H - N- H

H

Figure 8.15 The Molecular Structure of NH3

AX2E2 Shape is bent Bond angle is < 109.5o

Unshared electron pairs repel more than shared pair.

Lone pairs require more space than share pairs.

E Ex. H2O

X A X

E

H – O − H

Figure 8.16 The Molecular Structure of H2O

Figure 8.17 The Bond Angles in the CH4, NH3, and

H2O Molecules

AX5 Shape is trigonal bipyramidal

Bond angles are 120o(equatorial) and 90o(axial)

X

X A X

X

X

Ex. PCl5

Cl

Cl P Cl

Cl Cl

AX4E Shape is see-saw

Bond angles are <90o and <120o

X

E A X

X

X

Ex. SF4 34 electrons

F

S F

F

F

Figure 8.20 Three Possible Arrangements of the

Electron Pairs in the I3− Ion

AX3E2 Shape is T-shaped

Bond angle is <90o

X

E A X

E

X

Ex. ClF3

F

Cl F

F

AX2E3 shape is linear

bond angle is 180o

X

E A E

E

X

Ex. XeF2

F

Xe

F

Figure 8.19 Possible

Electron Pair

Arrangements

for XeF4

AX6 shape is octahedral

bond angle is 90o

X

X X

A

X X

X

Ex. SF6

F

F F

S

F F

F

AX5E Shape is square pyramidal

Bond angle is <90o

X

X X

A

X X

E

Ex. BrF5

F

F F

Br

F F

AX4E2 Shape is square planar.

Bond angle is 90o.

E

X X

A

X X

E

Youtube VSEPR annimation

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