Unit 1.4 - Bonding

The forces holding atoms together in compounds are known as chemical bonds

Bonds are broken and formed in chemical reactions

Bonding involves electrons in the outer shells of atoms

We will be studying ionic, covalent and metallic bonding

The octet rule

The first important thing to know about bonding is that atoms will generally react and form bonds in

such a way that they end up with an outer shell full of electrons. This principle is expressed in the

octet rule:

“When elements react they tend to do so in such a way that results in an _________ shell

containing __________ electrons”

Dot-cross diagrams

Dot-cross diagrams are an easy way of keeping track of what happens to ____________

during chemical reactions

Only the ___________ _________ electrons are shown as others are not involved in reacting

We will look at some examples drawing dot-cross diagrams for the formation of ionic bonds:

Sodium + chlorine

Lithium + fluorine

Magnesium + chlorine

Magnesium + oxygen

Ionic bonding

Ions are charged particles formed when atoms __________ or __________ electrons

Cations are ____________ ions

Anions are _____________ ions

Metals generally form ____________

Non-metals generally form ___________

Anions and cations are held together by

strong electrostatic forces of ____________

otherwise known as ionic bonds

Ionic radii

An atomic or ionic radius, measured in nanometers (nm, 10-9 m), is basically the size of an

atom or ion

The following table shows atomic and ionic radii for some common atoms and ions:

Cations Anions

Na Na+ F F-

Mg Mg2+ O O2-

K K+ Cl Cl2-

Ca Ca2+ N N3-

Trend

Cations are ______________ than the

corresponding atoms

Taking electrons away means that the

remaining electrons are held more tightly by

the __________________ force from the

______________ charged nucleus

Trend

Anions are _______________ than the

corresponding atoms

Adding more electrons means that the

__________________ force from the

_______________ is shared by more

electrons so they are held less tightly

Isoelectronic ions

The electronic configurations of common atoms and ions are shown below in pairs:

Remember… Think purrsitive!

Na+ 1s22s22p6

Ne 1s22s22p6Cl- 1s22s22p63s23p6

Ar 1s22s22p63s23p6

Because these pairs have the same electron configurations we say that they are ________________

Now try this…

1) Use a periodic table (p69) to determine the electron configuration of the following ions:

N3- 1s 2s 2p

O2- 1s 2s 2p

F- 1s 2s 2p

Na+ 1s 2s 2p

Mg2+ 1s 2s 2p

2) What do you notice about these different ions?

3) The circles in the diagram below represent these ions. Label the circles to show the order

of ionic radii:

0.072 nm

0.171 nm 0.140 nm 0.133 nm 0.102 nm

4) Can you explain why you chose to label the ions in this order?

Ionic compounds

Ionic compounds, when solid, form crystals with characteristic shapes and structures

We can use X-ray diffraction to study the structures of crystals

Hint: Think about the electrons and what holds them in place around the nucleus…

Lattice structures

We call the precise crystal structure of an ionic compound a ______________ structure

The lattice structure of a compound is the most _____________ way of arranging the ions

The forces of ______________ between anions and cations make lattices very strong and

rigid

unit cell . unit cell

Sodium chloride (NaCl) Caesium chloride (CsCl)

Face-centred cubic structure

a.k.a. rock salt structure

Body-centred cubic structure

Coordination number = 6 Coordination number = 8

What does all this mean?

Coordination number -

Unit cell -

Why are they different?

Hint: Think about the ionic radii (Na = 102 nm, Cs = 167 nm, Cl = 181 nm)

Properties of ionic compounds

The physical properties of all ionic compounds are very similar

Consider sodium chloride (NaCl):

Property Observation/ measurement

Explanation

Melting temperature801 C

Boiling temperature1413 C

Thermal conductivityPoor

Electrical conductivity

Solid -

Molten -

Solution -

Hardness

Malleability

NB Other ionic compounds behave similarly; you need to know their general properties and be able to explain how they arise

Evidence for the existence of ions

Ions are obviously far too small for us to see, so how do we know that they definitely exist?

1. Electron density maps

Sodium chloride (NaCl) 4-methoxybenzoic acid Electron density maps are produced using X-ray diffraction

X-rays are fired through crystals and the interference patterns produced tell us something

about where the ___________________ are

The big difference between electron density maps for ionic and covalent compounds is…

This shows that…

2. Electrolysis Watch the electrolysis of copper(II) chromate(VI), draw a diagram and use

it to explain how this experiment suggests the existence of ions

Lattice energy “The enthalpy of formation of one mole of an ionic compound from

gaseous ions under standard conditions”

Making bonds is an _____thermic process, in other words: energy is _______________ as

things become more stable

Forming an ionic lattice is a very _____thermic process as the ions take up their most stable

possible configuration

The energy released when forming an ionic lattice is known as the _______________ energy

Born-Häber cycles

We can determine lattice energies experimentally by measuring a number of other enthalpies and

combining them in a diagram called a Born-Häber cycle to give the lattice enthalpy:

Born-Häber cycle for the formation of NaCl:

energy

Definitions of terms:

∆Hθf = Standard enthalpy change of formation, measured using a bomb

calorimeter

∆Hθat = Standard enthalpy change of atomisation - changing a substance into a

monatomic gas

∆Hθi1 = First ionisation energy - energy required to remove 1 electron and form a cation,

from spectroscopic measurements

∆Hθe = First elecron affinity - energy required to add 1 electron and form and anion, also

from spectroscopic measurements

∆Hθlat = The lattice enthalpy of the substance - the enthalpy of formation of one mole of an

ionic compound from gaseous ions under standard conditions, calculated from the B-

H cycle

Steps involved in drawing a B-H cycle:

1) Start with the reactants in their standard forms

2) Draw in the enthalpy of formation (downwards)

3) Atomise the metal (upwards)

4) Ionise the metal (upwards)

5) Atomise the non-metal (upwards)

6) Add an electron to the non-metal (downwards)

7) Use the gap to calculate the lattice enthalpy:

Prep - Use the values in the table below to construct Born-Häber cycles for the formation of potassium iodide and potassium chloride and hence calculate the lattice energies:

∆Hθf[KI(s)] = -162.8 ∆Hθ

at[K(s)] = 89.0 ∆Hθat[1/2 Cls(g)] =

122.0∆Hθ

e[Cl(g)] = -349.0

∆Hθf[KCl(s)] = -436.7 ∆Hθ

i1[K(g)] = 418.8 ∆Hθat[1/2 Is(g)] =

107.0∆Hθ

e[I(g)] = -295.2

All values given in kJ mol-1

What factors affect lattice energy?

Compound

Radius of cation (nm)

Radius of anion (nm)

Lattice energy (kJ mol-1)

Compound Radius of cation (nm)

Radius of anion (nm)

Lattice energy (kJ mol-1)

NaF 0.102 0.133 -918 NaCl 0.102 0.180 -780

NaCl 0.102 0.180 -780 KCl 0.138 0.180 -711

NaBr 0.102 0.195 -742 MgF2 0.072 0.133 -2957

NaI 0.102 0.215 -705 CaF2 0.100 0.133 -2630

Trend in lattice enthalpies of sodium halides Trend in lattice enthalpies of group 1 & 2 halides

a) Size of cation…

b) Size of anion…

c) Charge on cation/anion…

This can all be explained using Coulomb’s law:

In other words:

Smaller ions can get __________ together so the force of

attraction between them is _____________

The greater the ____________ on an ion the greater its power

to ________________ another oppositely charged ion

Predicting stability - Why not NaCl 2?

Q ~ What are the formulae of the following compounds:

i. Sodium chloride

ii. Magnesium chloride

iii. Potassium chloride

iv. Calcium chloride

Why do these compounds always have the same formulae? Why is sodium chloride never NaCl2 for

example or magnesium fluoride never MgF3?

F = k q1q2

r2

Where q1 and q2 are the charges on the ions, r is the distance between them and F is the force of attraction

Dot-cross diagram for NaCl Dot-cross diagram for NaCl2

Sodium loses _____ electron Sodium loses _____ electrons

Use the data-books (p20) to find

out the following information:

Sodium chloride

First ionisation energy of sodium

Second ionisation energy of sodium

Theoretical lattice energy of NaCl -2526 kJ mol-1

NaCl2 and MgCl3 will not form

because…

Magnesium chloride

First ionisation energy of magnesium

Second ionisation energy of magnesium

Third ionisation energy of magnesium

Theoretical lattice energy of MgCl3 -4500 kJ mol-1

Polarisation of bonds

Born-Häber cycles allow us to calculate lattice energies from ________________ results

Coulomb’s law allows us to calculate ________________ lattice energies

The table below compares experimental and theoretical lattice energies:

Q ~ What do you notice about the numbers in the table?

CompoundLattice energy / kJ mol-1

DifferenceBorn-Häber Theoretical

NaF -918 -912

NaCl -780 -770

NaBr -742 -735

NaI -705 -687

AgF -958 -920

AgCl -905 -833

AgBr -891 -816

AgI -889 -778

Hmmm….I think I see a pattern developing!

A polar bear!

But why is there a difference? We need to look closely at the bonding involved, specifically whether

there is polarisation:

(a) Totally ionic bond (b) Polarised ionic bond (c) Covalent bond

Ionic bonds can be distorted by the attraction of the positive ______________ for the outer

electrons of the negative _____________

This means that some ionic bonds show a degree of ________________ character

These polarised bonds are ______________ than totally ionic bonds and give rise to more

negative lattice energies than predicted theoretically; they are ____________ stable than

expected

What causes ionic bonds to become polarised?

Cation

Ionic radius - Small cations have

____________ charge density and so are

more polarising

Charge - Cations with a large positive charge

are __________ polarising because they

have stronger attraction for the outer

electrons

Anion

Ionic radius - Large anions hold their outer

electrons _________ tightly, so the larger the

anion, the more __________ it is polarised

Effect on physical properties

1. Because theoretical lattice enthalpies are calculated by assuming that the ions are spherical

and separate they cannot account for polarisation and so are generally lower than lattice

enthalpies calculated from experimental data using Born-Häber cycles.

2. The melting temperatures of silver halides are about 20% lower than those of the sodium

halides, which is in agreement with the idea that the silver compounds have a greater covalent

character to their bonding.

Covalent bonding

Covalent bonding, unlike ionic bonding which involves the _______________ of electrons,

involves the ______________ of electrons between atoms

Atoms share electrons in order to gain ________ ________ _________

Ionic or covalent?

Whether a compound involves ionic or covalent bonding is related to the potential lattice energy for

that compound:

Ionisation

energy

Lattice energy Ionisation

energy

Lattice energy

Ionic Covalent

Dot-cross diagrams

You’ve had lots of practice so you’ll find it easy to draw dot-cross diagrams for these:

Chlorine Water

Methane Oxygen

Electron density maps

Atoms consist of a nucleus containing __________

and ___________ which is surrounded by orbiting

_____________

In covalent bonds electrons are most likely to be

found in the region between the _____________

charged nuclei

The area of electron density between the nuclei

has a _____________ charge, which ____________ the nuclei inwards and keeps them

_______________ bonded together

How strong and long are covalent bonds?

The distance between two atoms involved in covalent bonding depends on two factors:

1. The attraction between the bonding _____________ and the nuclei

2. The repulsion between the like-charged _______________

How the two forces balance out can be shown in the following diagram:

Covalent bonds are of such a

length that attraction and repulsion

are balanced

Bond enthalpy is a measure of

bond strength

Bond Bond length / nm Bond enthalpy / kJmol-1

H-H 0.74 436

C-C 0.154 348

C=C 0.134 612

C-H 0.109 412

Dative covalent bonds

Sometimes both of the ______________ that make up a covalent bond come from the same atom,

this is called a dative covalent bond:

Ammonia

(NH4+)

Aluminium

chloride dimer

(Al2Cl6)

En

erg

y

Separation / nm

We represent dative bonds with an arrow:

Ammonia Aluminium chloride

Dative bonds are the same length and strength as normal covalent bonds

Another example of a compound with dative bonding is carbon monoxide (CO)

Shapes of covalent molecules

Different molecules have different shapes, remember methane, which is tetrahedral

Draw dot-cross diagrams for the following compounds:

Compoun

dDot-cross diagram Electron density diagram Shape

BeCl2

CH4

BF3

H2O

There is a strong force of _______________ between electron pairs

As a result of this the bonding and non-bonding pairs of electrons will spread out to be as far

apart from each other as possible

This gives rise to the molecular geometries shown above

Giant atomic structures

Many covalently bonded compounds exist as molecules e.g.

Others, like diamond (diagram right) exist as atomic

crystals or molecular crystals

These giant atomic structures form crystal lattices

held together by ___________ bonds

Other examples of giant atomic structures are silicon

carbide (SiC) and silicon(IV) oxide (SiO2)

Compounds with giant atomic structures all have

similar physical properties:

Metallic bonding

Properties of metals

How does the theory of metallic bonding explain the following properties of metals? (p92)

Electrical conductivity:

Most simple model of bonding in metals

Metal crystal containing metal cations

surrounded by a delocalised ‘sea’ of the

electrons that they have lost

High thermal conductivity:

High melting/boiling temperatures:

Malleability and ductility:

Predicting the properties of metals

The theory of _______________ bonding can be used to predict the properties of metals

The strength of metallic bonding depends on the charge density of the cloud of delocalised

_________________ and the metal ________

Higher charge density means _______________ electrostatic attraction

Lower charge density means ____________ electrostatic attraction

Q ~ Use the theory of metallic bonding to explain the differences shown in the following table:

a) Melting temperature…

Metal Melting temp / ºC Thermal conductivity / W/cm/K

Sodium 98 1.35

Magnesium 649 1.5

Copper 1083 3.85

b) Thermal conductivity…

Summary of chapter 1.4 - Bonding

Ionic bonding

Covalent bonding

Practice questions:

Ionic bonding1. Magnesium reacts spontaneously and extremely vigorously with fluorine gas

a) Write an equation, including state symbols, for the reaction of magnesium with fluorine.

b) Draw a dot-cross diagram to show the bonding in magnesium fluoride.

2. Compare and contrast the structures of sodium chloride and caesium chloride. Why do they differ?

3. Sodium chloride is, under standard conditions, an ionic solid and is used as a flavour enhancera) Under what conditions does sodium chloride conduct electricity? Explain your answer.

b) For which compound would you expect the lattice enthalpy to be more negative: sodium chloride or magnesium chloride? Explain your choice.

Metallic bonding

c) Why does sodium chloride always exist as NaCl and never NaCl2?

4. Why is the theoretical lattice enthalpy for silver iodide so different to the experimentally determined value?

Covalent bonding1.Define the term covalent bond.

2. Nitrogen is the most abundant gas in the Earth’s atmosphere. It occurs as a covalently bonded molecule, N2

a) Draw an electron density diagram for nitrogen

c) Draw a dot-cross diagram to show the bonding in nitrogen

3. Beryllium chloride, BeCl2, is a chemical compound that can be said to be ‘electron-deficient’. It forms pairs of linked molecules known as dimers when solid.

a) Draw a dot-cross diagram for a single BeCl2 molecule

b) Draw a dot-cross diagram for a Be2Cl4 dimer

c) What type of bonding is present in Be2Cl4?

4. Covalent bonds vary in length depending on what atoms are involved. Describe the factors which influence the length of a covalent bond; you can include a diagram as part of your answer.

Metallic bonding1. Sodium is shiny, grey metal with a relatively low melting temperature. It is also relatively soft and can be cut with a sharp knife at room temperature.

a) Describe the type of bonding that is present in solid sodium metal

b) Explain why iron is much harder than sodium

c) Explain why copper conducts thermal energy more readily than sodium

2. Most metals are said to be malleable. What does the word malleable mean and how does the theory of metallic bonding account for this property?

Now go away and REVISE! There is a test on this chapter just around the corner and that’s going to be closely followed by the school exam, which will also include questions on… guess what… BONDING!