Lecture 3: Structure and Bonding (Chemistry in Context 8.1–8.5 & 10.1–10.8 & ChemFactsheet 5)

The Chemical Bond

-classification of bond type

The Ionic Bond

-formation of positive and negative ions

-ionic bonding & lattices

The Covalent Bond

-dot and Cross Diagrams

-simple & Giant covalent structures

-dative Covalent Bonding

The Metallic Bond

-metallic lattices and electron delocalisation

-alloys

Intermediate Bond Types

-electronegativity and polarized bonds

-polarised ionic bonds and Fajan’s Rules

Periodicity (4.2-4.4)

-repeating chemical properties of the

chemical elements that occurs when

layed out in order of atomic number (Z)

is called periodicity

-elements in same valence electron

configuration similar chemistry

-blocks of periodic

table are named

after the orbitals

housing valence

electrons:

s-block f-block d-block p-block

Cl

Br

Sr

Mg

Ca

I

s2 s2p5

Cation and Anion formation

first

ionisation

second

ionisation

M M+ M2+

X first

electron

affinity

second

electron

affinity

X- X2-

(endothermic)

(exothermic)

(endothermic)

(endothermic)

Ionic Bonding (8.2)

when K atom and Cl atom

brought together it eventually

becomes favourable for electron

to jump from K to Cl:

(now both K+ and Cl- have stable

octet (noble gas configuration)

(Lewis formula of KCl)

K Cl K Cl+_

+DH

heat evolved

(lattice energy)

Ionic Bonding (8.2)

-both metal and

non-metal seek stable

noble gas electron

configuration

+

Cl

Cl -

[Ar]4s1

[Ar]

[Ne]2s22p5

[Ar]

chloride ion

chlorine

Sylvite

(KCl)

Ionic Bonding (8.2)

-both metal and

non-metal seek stable

noble gas electron

configuration

-MX ionic salts form

favourably when

M has a small ionisation

energy and X has a

large, exothermic

electron affinity

+

Cl

Cl -

[Ar]4s1 [Ne]2s22p5

[Ar]

chloride ion

chlorine

[Ar]

[Ar]

Dot-Cross Diagrams (8.2)

dot/cross diagrams show:

-compound’s electronic structure

-where the electrons originated

e.g. CaCl2:

[Ar]

[Ar]

Dot-Cross Diagrams (8.2)

dot/cross diagrams show:

-compound’s electronic structure

-where the electrons originated

e.g. CaCl2:

What salt is this?

[Ar]

[Ar]

Dot-Cross Diagrams (8.2)

dot/cross diagrams show:

-compound’s electronic structure

-where the electrons originated

e.g. CaCl2:

e.g. MgO:

[Ne] [Ne]

-salts are crystalline:

ions are arranged

regularly giving solids

with flat faces and

straight edges

Ionic Lattices (10.7)

small NaCl

crystals

-ionic bonds are very polar

dissolve well in polar solvents

solid salt (fixed ions)

electrical insulators

aqueous salt solution or molten salt

(ions free to move in liquid)

conduct electricity

-ionic bonds are strong!

high melting points

(750 °C for KCl, 3000 °C for MgO)

Physical Properties of Salts (10.8)

giant CaSO4 crystals, Mexico

What is a chemical bond? (8.1)

-an attractive force holding together atom(s) or ions(s) that makes

them function as a unit.

-bond forms if it makes the system is lower in energy than when

atoms are apart.

-energy input required to break bond = bond strength or bond energy

e.g. the ClCl bond has a strength of 243 kJ/mol

the O=O bond has a strength of 499 kJ/mol

the lattice energy of KCl is 715 kJ/mol

the lattice energy of MgO is 3930 kJ/mol

The way atoms are bonded together shapes physical and chemical

properties of a compound.

e.g. graphite is grey, soft, conductor of electricity

diamond is a transparent insulator, the hardest substance known

e.g. C and Si have similar chemistry but:

SiO2 is a brittle unreactive crystalline solid

CO2 is a gas

Types of Chemical Bond

-atoms can bond to each other in three main ways:

Ionic bonding (e.g. NaCl, CaCO3)

Covalent bonding (e.g. H2O, PCl5)

Metallic bonding (e.g. Fe, K, Hg)

Why Study Bonding?

Covalent Bonding

-a covalent bond forms when partially filled orbitals overlap:

-electrons are shared so that each contributing atom can experience a full outer shell of electrons (noble gas configuration)

-each H obtains full valence shell (n = 1) and the same electron

configuration as He H2 is stable

-other homonuclear diatomic molecules form by making similar

covalent bond

e.g. F has seven valence electrons so shares one electron:

dot-cross

structure of F2

each F obtains octet of electrons and electron configuration [Ne]

note each F has three non-bonded or lone pairs of electrons in the

valence shell (n = 2)

Octet rule: atoms proceed as far as possible toward completing

their octets by sharing electron pairs in covalent bonds

F

+ +

+

+

++

++ +

F

+ +

+ +

++

++

F +F

F is much more

electronegative than H

HF is a polar covalent bond

-in some diatomic molecules electrons are not distributed

evenly e.g. HF

H F

electron-poor electron-rich

dot-cross diagram

-covalent HF bond allows both

atoms to have full outer shell

(H obtains full n = 1 shell,

F obtains full n = 2 shell)

Electronegativity: the tendency of an atom to attract

electron density towards itself in a chemical bond

increasing electronegativity

What is Electronegativity? (8.2)

F is the most electronegative, Cs the least

The greater the difference in electronegativity between

two elements, the greater the polarity of the bond

If electronegativity difference >1.8

compound is mostly ionic:

P2O5 (difference = 1.4) completely covalent

LiI (difference = 1.5) partly ionic, partly covalent

Al2O3 (difference = 2.0) mostly ionic, slightly covalent

CsF (difference = 3.3) completely ionic

Polar Covalent Bonds

non-polar covalent: homonuclear diatomic

molecules e.g. H2, Cl2, both atoms identical

so electron density is arranged symmetrically

ionic bond – other extreme – very different

atoms so electrons completely transferred

very different electronegativity

H F

polar covalent: electrons

not transferred between

atoms but is unequally

shared due to slightly

different electronegativity

-polar bonds can give polarity to the overall molecule e.g.

-importantly for life, water is

a polar molecule:

polar molecules spontaneously align

themselves in an electric or magnetic field:

Cl

CClCl

Cl

H

CClCl

Cl

non polar

CCl4

polar

CHCl3

non-polar liquids unpeturbed

e.g. Hg, Br2, hydrocarbons

polar liquids deflected

e.g. H2O, alcohols, acetone

statically-charged

plastic rod

Dot-cross diagrams of polyatomic molecules (8.4)

-octet rule is obeyed for all atoms, even in larger molecules e.g.

ammonia (NH3)

N (1s22s22p3) is in group 5

must share 3 electrons to gain full octet

structure features three bonding pairs and one lone (non-bonded) pair

carbon dioxide (CO2)

O (1s22s22p4) is in group 6

must share 2 electrons to gain full octet

structure features four bonding pairs

(two double bonds) and

four lone (non-bonded) pairs

CO2 Lewis structure

Dative Bonding (8.5)

-sometimes both electron pairs from a covalent bond

are provided by one atom

e.g. formation of

ammonium ion:

-boron trifluoride is electron-deficient compound

(incomplete octet; only 6 electrons in outer shell)

forms dative bond with electron donors e.g. H2O, NH3

BF3

Metallic Bonding (10.5)

- close packing of metals has important

consequences:

-atoms’ outer (valence) electrons are

delocalised

(can move freely throughout lattice)

-metal structure is array of positive ions

immersed in a sea of electrons

-delocalised electrons:

-bind atoms strongly - strong high melting points (e.g. Fe 1530 ºC, W 3500 ºC)

-allow conduction of heat and electricity throughout solid (especially Ag, Cu)

delocalised ‘sea’

of electrons

Metallic Bonding (10.5)

-bonding between atoms is not

directional

so metals are

bendy and:

i) malleable - can be beaten into

thin sheets without fracturing

ii) ductile - can be drawn into

fine wires without fracturing

-close-packing of atoms explains metals’

high density:

8 gcm-3 for Cu and Fe

as high as: 22 gcm-3 for Os, Au & Ir

Alloys (10.5)

-man-made mixture of metals

-made by mixing molten metals

in desired proportions

-metallic bonding but altered properties

e.g. bronze (10% Sn in Cu) – stronger than Cu

Cu-Be alloys are

strong and spark

resistant (used

on oil rigs)

Cr

Cr C C

weak (layers slide)

iron stainless steel

strengthens and

makes corrosion

resistant

solder (Pb-Sn)

very low melt point)

bronze

statue

Network Solids (Giant Covalent Lattices) (10.6)

-bonding is covalent and in an infinite lattice – ‘one giant molecule’

-e.g. diamond, quartz, ruby, sapphire – brittle but very hard

Quartz (SiO2) – a three-dimensional infinite

lattice of tetrahedral Si atoms linked by O atoms

-sand is mostly made of silica (10.10)

amethyst

Network Solids (Giant Covalent Lattices)

-bonding is covalent and in an infinite lattice – ‘one giant molecule’

-e.g. diamond, quartz, ruby, sapphire – brittle but very hard

diamond – a three-dimensional lattice of tetrahedral carbon atoms

-lattice very rigid: the hardest substance known

used in drill-bits and (powdered) as an abrasive

good thermal conductor (rigidity transfers atomic vibrations)

but electrical insulator (no delocalised electrons)

graphite - a two-dimensional array of trigonal carbon atoms

-hexagonal rings - strong bonding within the layers but weak

bonding between them

graphite is soft (used in pencils and as lubricant).

non-bonded electrons are delocalised throughout plane

graphite conducts electricity

examples bonding and

structure

conduct-

ivity

m.p./

b.p.

strength

metals Cu, Al,

Os, Hg,

brass,

solder

metallic lattice -

positive metal

ions held

together by

delocalised

electrons

good high malleable,

ductile,

bendy

ionic solids NaCl,

CaCO3

ionic lattice very low

(except in

solution)

high hard, rigid,

brittle

network

solids

graphite

diamond,

SiO2, BN

covalent lattice very low

(except

graphite)

very

high

very hard,

rigid,

brittle

molecular

solids

ice, sugar,

wax, I2,

PCl5, S8.

covalent

molecules held

by lattice of

weak forces

very low low soft, brittle

Summary: Bonding in Solids

-bonding in solids can be grouped into four categories :

metallic, ionic, giant covalent and simple covalent

Al CaCO3

S

C

What kind of bonding is

holding together the following?

candle

wax

SA GA

limestone

cliff

brass

ice

Hg

Classifying Bonding

‘bonding’ is general term referring to forces that hold together any

type of chemical species: molecules, groups of molecules, atoms or ions

Intramolecular bonding holds together atoms e.g.

-ionic bonding

-metallic bonding

-covalent bonding

-strong (150-500 kJ/mol)

Intermolecular bonding acts between molecules e.g.

-van der Waals’ forces

-dipole-dipole forces

-ion dipole forces

-hydrogen bonding

-weak (2-25 kJ/mol) but often responsible

for bulk, physical properties of matter

molecule’s functional groups

Functional Groups and Physical Properties

determine

strength of intermolecular forces present

determines compound’s

physical properties

e.g.

H3CO

CH3O

HH

H3CCH3

CH3H3C

H3CCl

CH3Cl

INTERMOLECULAR FORCES SUMMARY

(non-covalent interactions)

van der Waals’ force

(1-5 kJ/mol)

dipole-dipole interaction

(5-10 kJ/mol)

hydrogen bonding interaction

(10-30 kJ/mol)

d- d+

d+ d-

d+ d-

d- d+

weakest force

(lowest boiling

point)

stronger force

(higher boiling

point)

-what determines if one compound is soluble in another?

Functional Groups and Intermolecular Forces: Solubility

-polarity of a molecule’s bonds determine in which solvents

it is soluble

i.e. compound likely to be soluble in particular solvent if the

intermolecular forces are similar in compound itself and the solvent:

e.g.

water and

short-chain

alcohols are

miscible

methanol and

acetone are

miscible

“like dissolves like”

Functional Groups and Intermolecular Forces: Solubility

H3CO

H

OC

H3C

CH3

HO

H

HO

C2H5

H

OH5C2