Chapter 12 1

Chemical BondingChemical Bonding

Chapter 12Chapter 12

Chapter 12 2

IntroductionIntroduction

The properties of many materials can be understood in terms of their microscopic properties. Microscopic properties of molecules include:

•the connectivity between atoms and •the 3D shape of the molecule.

Chapter 12 3

We consider three bonds “within molecules” (intramolecular forces):

• ionic bond (electrostatic forces which hold ions together, e.g. NaCl)

• covalent bond (results from sharing electrons between atoms, e.g. Cl2)

• metallic bonding (refers to metal nuclei floating in a sea of electrons, e.g. Na).

Chapter 12 4

In all chemical bonds, electrons are shared and transferred between atoms.

In bonding, electrons are involvedIn bonding, electrons are involved

Chapter 12 5

•The electrons involved in bonding are called valence electrons

•Valence electrons are found in the incomplete, outermost orbital of an atom.

Chapter 12 6

Ionic BondingIonic Bonding

Consider the reaction between sodium and chlorine:Na(s) + ½Cl2(g) NaCl(s)

Chapter 12 7

Ionic BondingIonic Bonding

Na(s) + ½Cl2(g) NaCl(s) H°f = -410.9 kJThe reaction is violently exothermic.We infer that the NaCl is more stable than its constituent elements. Why?Na has lost an electron to become Na+ and chlorine has gained the electron to become Cl. Note: Na+ has an Ne electron configuration and Cl has an Ar configuration.That is, both Na+ and Cl have an octet of electrons surrounding the central ion.

Chapter 12 8

Ionic BondingIonic Bonding

NaCl forms a very regular structure in which each Na+

ion is surrounded by 6 Cl ions.

Similarly, each Cl ion is surrounded by six Na+ ions.

There is a regular arrangement of Na+ and Cl in 3D.

Note that the ions are packed as closely as possible.

Note that it is not easy to find a molecular formula to

describe the ionic lattice.

Chapter 12 9

Ionic BondingIonic Bonding

Chapter 12 10

Covalent BondingCovalent Bonding

•In ionic bonding one atom completely loses an electron while the other gains the electron. •When two similar atoms bond, none of them wants to lose or gain an electron to form an octet.•When similar atoms bond, they share pairs of electrons to each obtain an octet.•Each pair of shared electrons constitutes one chemical bond.•Example: H + H H2 has electrons on a line connecting the two H nuclei.

Chapter 12 11

Covalent BondingCovalent Bonding

Chapter 12 12

Covalent BondingCovalent Bonding

Multiple BondsMultiple BondsIt is possible for more than one pair of electrons to be shared between two atoms (multiple bonds):

One shared pair of electrons = single bond (e.g. H2);Two shared pairs of electrons = double bond (e.g. O2);Three shared pairs of electrons = triple bond (e.g. N2).

Generally, bond distances decrease as we move from single through double to triple bonds.

H H O O N N

Chapter 12 13

•In a covalent bond, electrons are shared.

•Sharing of electrons to form a covalent bond does not imply equal sharing of those electrons.

•There are some covalent bonds in which the electrons are located closer to one atom than the other.

•Unequal sharing of electrons results in polar bonds.

Chapter 12 14

ElectronegativityElectronegativityThe preference one atom in a chemical bond for electrons is called electronegativity.Electronegativity is a scale from 0.7 (Cs) to 4.0 (F).

Electronegativity increases •across a period •up a group.

Chapter 12 15

ElectronegativityElectronegativity

Chapter 12 16

Difference in electronegativity is a gauge of bond polarity:

•electronegativity differences of 0-0.2 result in non-polar covalent bonds (equal or almost equal sharing of electrons);•electronegativity differences around 0.3-1.6 result in polar covalent bonds (unequal sharing of electrons);•electronegativity differences between 1.7-4.0 result in ionic bonds (transfer of electrons).

There is no sharp distinction between bonding types.

Chapter 12 17

The positive end (or pole) in a polar bond is represented + and the negative pole -.

QuickTime™ and aTIFF (Uncompressed) decompressor

are needed to see this picture.

Chapter 12 18

Lewis Symbols• Also known as electron dot

diagrams• A way of keeping track of

valence electrons.

How to write them1) Write the symbol.2) Put one dot for each valence

electron3) Don’t pair up until they have to

X

Chapter 12 19

We generally place the electrons one four sides of a square around the element symbol.

Octet rule: we know that s2p6 is a noble gas configuration. We assume that an atom is stable when surrounded by 8 electrons (4 electron pairs).

Chapter 12 20

Lewis symbol for nitrogen• Nitrogen has 5 valence

electrons.• First we write the symbol.• Then add 1 electron at a

time to each side.• Until they are forced to pair up.

• The number of electrons available for bonding are indicated by unpaired dots

• Nitrogen would have 3 bonding electrons

N

Chapter 12 21

Drawing Lewis Structures of MoleculesDrawing Lewis Structures of Molecules•Add the valence electrons.•Identify the central atom (usually the one with the highest molecular mass and closest to the center of the periodic table).•Place the central atom in the center of the molecule and add all other atoms around it.•Place one bond (two electrons) between each pair of atoms.•Complete the octet for the central atom.•Complete the octets for all other atoms. Use double bonds if necessary.

Chapter 12 22

• To draw the electron dot structure of an ion, you must add the charges to the number of electrons (for a negative ion)

See p.342 in Heath

QuickTime™ and aTIFF (Uncompressed) decompressor

are needed to see this picture.

Chapter 12 23

Formal ChargeFormal Charge

It is possible to draw more than one Lewis structure with the octet rule obeyed for all the atoms.To determine which structure is most reasonable, we use formal charge.Formal charge is the charge on an atom that it would have if all the atoms had the same electronegativity.To calculate formal charge, electrons are assigned as follows:

•All nonbonding electrons are assigned to the atom on which they are found.•Half the bonding electrons are assigned to each atom in a bond.

Chapter 12 24

Drawing Lewis StructuresDrawing Lewis StructuresFormal ChargeFormal ChargeFormal charge is:

valence electrons - number of bonds - lone pair electronsConsider:

For C: There are 4 valence electrons (from periodic table).In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure.Formal charge: 4 - 5 = -1.

C N

Chapter 12 25

Drawing Lewis StructuresDrawing Lewis StructuresFormal ChargeFormal ChargeConsider:

For N:There are 5 valence electrons.In the Lewis structure there are 2 nonbonding electrons and 3 from the triple bond. There are 5 electrons from the Lewis structure.Formal charge = 5 - 5 = 0.

We write:

C N

C N

Chapter 12 26

Drawing Lewis StructuresDrawing Lewis StructuresFormal ChargeFormal ChargeThe most stable structure has:

•the smallest formal charge on each atom,•the most negative formal charge on the most electronegative atoms.

Resonance StructuresResonance StructuresSome molecules are not well described by Lewis Structures.Typically, structures with multiple bonds can have similar structures with the multiple bonds between different pairs of atoms.

Chapter 12 27

More on Resonance Structures

Check this out:Check this out:

Chapter 12 28

Drawing Lewis StructuresDrawing Lewis StructuresResonance StructuresResonance StructuresExample: experimentally, ozone has two identical bonds whereas the Lewis Structure requires one single (longer) and one double bond (shorter).

O

OO

Chapter 12 29

Drawing Lewis StructuresDrawing Lewis StructuresResonance StructuresResonance StructuresResonance structures are attempts to represent a real structure that is a mix between several extreme possibilities.

Chapter 12 30

Drawing Lewis StructuresDrawing Lewis StructuresResonance StructuresResonance StructuresExample: in ozone the extreme possibilities have one double and one single bond. The resonance structure has two identical bonds of intermediate character.

Common examples: O3, NO3-, SO4

2-, NO2, and benzene.O

OO

O

OO

Chapter 12 31

Drawing Lewis StructuresDrawing Lewis StructuresResonance in BenzeneResonance in BenzeneBenzene consists of 6 carbon atoms in a hexagon. Each C atom is attached to two other C atoms and one hydrogen atom.There are alternating double and single bonds between the C atoms.

Experimentally, the C-C bonds in benzene are all the same length.

Chapter 12 32

Drawing Lewis StructuresDrawing Lewis StructuresResonance in BenzeneResonance in BenzeneWe write resonance structures for benzene in which there are single bonds between each pair of C atoms and the 6 additional electrons are delocalized over the entire ring:

Benzene belongs to a category of organic molecules called aromatic compounds (due to their odor).

or

Chapter 12 33

Drawing Lewis StructuresDrawing Lewis StructuresResonance in BenzeneResonance in Benzene

Chapter 12 34

Exceptions to the Octet RuleExceptions to the Octet RuleThere are three classes of exceptions to the octet rule:

•Molecules with an odd number of electrons;•Molecules in which one atom has less than an octet;•Molecules in which one atom has more than an octet.

1) Odd Number of Electrons1) Odd Number of ElectronsFew examples. Generally molecules such as ClO2, NO, and NO2 have an odd number of electrons.

Chapter 12 35

NO has 5+6=11 valence electrons to place:

Nitrogen ends up with less than an octet.

Chapter 12 36

Exceptions to the Octet RuleExceptions to the Octet Rule

Less than an OctetLess than an OctetMolecules with less than an octet are typical for compounds of Groups 1A, 2A, and 3A.Most typical example is BF3.

BF3 has (7*3)+3=24 valence electrons. Boron ends up with less than an octet.

Chapter 12 37

More than an OctetMore than an OctetThis is the largest class of exceptions.Atoms from the 3rd period onwards can accommodatemore than an octet.Beyond the third period, the d-orbitals are low enough in energy to participate in bonding and accept the extra electron density.