CHAPTER 8 AND 9

Bonding

TYPES OF BONDSIonic Bonds: Formed between ions. Transfer of electrons occurring.

Covalent Bonds: Molecules form between atoms that share electrons.

Metallic Bonds: Form multiple bonds using the free flowing electrons of a metal.

LEWIS SYMBOLS AND OCTET RULE

Atoms tend to gain and lose electrons until they have 8 valence electrons similar to noble gases. This is a stable electron configuration as shown experimentally by the high ionization energy and low electron affinity.

IONIC BONDING An atom with a low ionization energy reacts with an atom with high electron affinity.

The electron moves. Opposite charges hold the atoms together.

COULOMB’S LAW

Q is the charge. d is the distance between the

centers. If charges are opposite, E is

negativeexothermic

Same charge, positive E, requires energy to bring them together.

FORMING IONIC COMPOUNDS Lattice energy - the energy

associated with making a solid ionic compound from its gaseous ions.

M+(g) + X-(g) ® MX(s) This is the energy that “pays” for

making ionic compounds. Energy is a state function so we

can get from reactants to products in a round about way.

NA(S) + ½F2(G) ® NAF(S) First sublime Na Na(s) ® Na(g)

DH = 109 kJ/mol Ionize Na(g) Na(g) ® Na+(g) + e-

DH = 495 kJ/mol Break F-F Bond ½F2(g) ® F(g)

DH = 77 kJ/mol Add electron to F F(g) + e- ® F-(g)

DH = -328 kJ/mol Formation of solid Na+ + F- ® NaF (s)

DH = -575 kJ/mol Lattice energy Na(s) + ½F2(g) ® NaF(s)

DH = -928 kJ/mol

COVALENT BONDING Electrons are shared by atoms. Ionic and Covalent are extremes. In between are polar covalent bonds. The electrons are not shared evenly. One end is partially positive, the other

negative. Indicated using small delta .d

H - Fd+ d-

H - F

d+

d-

H - Fd+

d-

H -

F

d+

d-

H - F d+

d-

H - Fd+d-

H - F

d+d-

H - F d+

d-

+-

H - Fd+ d-

H - Fd+ d-

H - Fd+ d- H - F

d+ d-

H - Fd+ d-

H - Fd+ d-

H - Fd+ d-

H - Fd+ d-

- +

Electronegativity difference

Bond Type

Zero

Intermediate

Large

Covalent

Polar Covalent

Ionic

Covale

nt C

hara

cte

r d

ecre

ases

Ion

ic C

hara

cte

r incre

ases

COVALENT BONDS

The electrons in each atom are attracted to the nucleus of the other.

The electrons repel each other, The nuclei repel each other. They reach a distance with the lowest

possible energy. The distance between is the bond length.

0

En

erg

y

Internuclear Distance

Bond Length

0

En

erg

y

Internuclear Distance

Bond Energy

ELECTRONEGATIVITY The ability of an electron to attract

shared electrons to itself. Tends to increase left to right. Decreases as you go down a group. Noble gases aren’t discussed. Difference in electronegativity between

atoms tells us how polar.

DIPOLE MOMENTS A molecule with a center of negative

charge and a center of positive charge is dipolar (two poles),

We say that is has a dipole moment. Center of charge doesn’t have to be on

an atom. Will line up in the presence of an electric

field.

H - Fd+ d-

% I

on

ic C

hara

cte

r

Electronegativity difference

25%

50%

75%

LiBr

LiF

HCl

THE COVALENT BOND The forces that cause a group of

atoms to behave as a unit. Why? Due to the tendency of atoms to

achieve the lowest energy state. It takes 1652 kJ to dissociate a mole of

CH4 into its ions Since each hydrogen is hooked to the

carbon, we get the average energy = 413 kJ/mol

COVALENT BOND ENERGIES We made some simplifications in describing

the bond energy of CH4 Each C-H bond has a different energy.

CH4 ® CH3 + H DH = 435 kJ/mol

CH3 ® CH2 + H DH = 453 kJ/mol

CH2 ® CH + H DH = 425 kJ/mol

CH® C + H DH = 339 kJ/mol

Each bond is sensitive to its environment.

AVERAGES single bond one pair of electrons is shared. double bond two pair of electrons are shared. triple bond three pair of electrons are shared. More bonds, shorter bond length.

USING BOND ENERGIES We can find DH for a reaction. It takes energy to break bonds, and end

up with atoms (+). We get energy when we use atoms to form

bonds (-). If we add up the energy it took to break

the bonds, and subtract the energy we get from forming the bonds we get the DH.

Energy and Enthalpy are state functions.

FIND THE ENERGY FOR THIS

2 CH2 = CHCH3 +2NH3 O2+ ® 2 CH2 = CHC º N+ 6 H2O

C-H 413 kJ/mol

C=C 614kJ/mol

N-H 391 kJ/mol

O-H 467 kJ/mol

O=O 495 kJ/mol

CºN 891 kJ/mol

C-C 347 kJ/mol

LOCALIZED ELECTRON MODEL

Simple model, easily applied. A molecule is composed of atoms

that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.

Three Parts1) Valence electrons using Lewis

structures2) Prediction of geometry using VSEPR3) Description of the types of orbitals

LEWIS STRUCTURE Shows how the valence electrons are

arranged. One dot for each valence electron. A stable compound has all its atoms with

a noble gas configuration. Hydrogen follows the duet rule. The rest follow the octet rule. Bonding pair is the one between the

symbols.

RULES1. Sum the valence electrons.2. Use a pair to form a bond between each pair

of atoms.3. Arrange the rest to fulfill the octet rule (except

for H and the duet).4. Place left over electrons on the central atom.5. If there are not enough electrons to give the

central atom an octet, try multiple bonds. H2O CO2

CN-

FORMAL CHARGE

The difference between the number of valence electrons on the free atom and that assigned in the molecule.

We count half the electrons in each bond as “belonging” to the atom.

Molecules try to achieve as low a formal charge as possible.

Negative formal charges should be on electronegative elements.

SO42-

RESONANCE Sometimes there is more than one

valid structure for an molecule or ion.

Use double arrows to indicate it is the “average” of the structures.

It doesn’t switch between them. Localized electron model is based

on pairs of electrons, doesn’t deal with odd numbers.

NO3-

NO2-

EXAMPLES

XeO3

NO4-3

SO2Cl2

EXCEPTIONS TO THE OCTET1. Odd numbers of electrons

NO2. Atoms with fewer than octet

Be and B often do not achieve octetBF3

3. Atoms with more than octet Third row and larger elements can exceed the

octet SF6

Use 3d orbitals?

EXCEPTIONS TO THE OCTET When we must exceed the octet, extra

electrons go on central atom.

ClF3

XeO3

ICl4-

BeCl2

VSEPR Lewis structures tell us how the atoms

are connected to each other. They don’t tell us anything about shape. The shape of a molecule can greatly

affect its properties. Valence Shell Electron Pair Repulsion

Theory allows us to predict geometry Molecules take a shape that puts

electron pairs as far away from each other as possible.

VSEPR Have to draw the Lewis structure to

determine electron domains. bonding nonbonding lone pair Lone pair takes more space. Multiple bonds count as one domain.

The number of domains determinesbond anglesunderlying structure

The number of atoms determines actual shape

VSEPRElectronDomains

BondAngles

UnderlyingShape

2 180° Linear

3 120° Trigonal Planar4 109.5° Tetrahedral

5 90° &120°

Trigonal Bipyramidal

6 90° Octagonal

ELECTRON DOMAIN GEOMETRY

ELECTRON DOMAIN VS. MOLECULAR SHAPEElectron Domain Just count the number of electron

domains and the geometry corresponds to that number.

Molecular Shape Electron domain geometry is many

times not the shape of the molecule. Here we are worried about just the atoms.

MOLECULAR GEOMETRY Multiple molecular geometries can result

from the same electron geometry.

NON-BONDING ELECTRONS Non-bonding pairs take up more space

than bonding pairs and therefore change the angles within the atom.

MULTIPLE BONDS Multiple bonds create a high electron

density on one side of the molecule which also changes angles.

DIFFERENT MOLECULAR GEOMETRIES Linear Trigonal Planar Bent Tetrahedral Trigonal pyramidal Trigonal bipyramidal Seesaw T-shaped Octahedral Square pyramidal Square planar

POLARITY Individual bonds can have bond dipoles

but the overall molecules polarity may be different.

POLARITY Individual bond dipoles (vectors) can be

added to give the entire molecular dipole.

HYBRIDIZATIONBeryllium: Has no single occupied orbitals therefore should not bond.

If it absorbs energy to excite an electron to 2p it has 2 electrons that can form a bond.

HYBRID ORBITALSThe s and p orbitals can combine to make a hybrid orbital called the sp hybridized orbital.

-sp orbitals have a linear position and is consistent with molecular geometries.-it has 2 degenerate orbitals.

HYBRID ORBITALS Other hybrid orbitals that can be made

include sp2 and sp3. This occurs when there is more than one p orbital overlapping with an s orbital.

SIGMA AND PI BONDS Sigma (σ) bonds overlap head to head

and electron density is in line with the internuclear axis.

Pi (π) bonds overlap side to side and electron density is above and below internuclear axis.

TYPES OF BONDS Single bonds are all sigma bonds. Multiple bonds have one sigma bond

and the rest are pi.

DELOCALIZED PI BONDING Pi electrons are

not localized over one N-O bond but delocalized over all 3 bonds.

This is related to resonance.

MOLECULAR ORBITAL (MO) THEORY In MO theory, we invoke the wave nature of

electrons. If waves interact constructively, the resulting

orbital is lower in energy: a bonding molecular orbital.

MO THEORY In hydrogen the

electrons go into the bonding orbital because it is lower in energy.

The bond order is one half the difference between the number of bonding and antibonding electrons.

For hydrogen, with two electrons in the bonding MO and none in the antibonding MO, the bond order is 2-0/2 = 1

MO THEORY

In the case of He2, the bond order would be

12

(2 − 2) = 0

• Therefore, He2 does not exist.

MO THEORY

For atoms with both s and p orbitals, there are two types of interactions:The s and the p orbitals

that face each other overlap in fashion.

The other two sets of p orbitals overlap in fashion.

MO THEORY

The resulting MO diagram looks like this (Fig. 9.41).

There are both s and p bonding molecular orbitals and s* and * antibonding molecular orbitals.

SECOND-ROW MO DIAGRAMS