Properties of Ionic Compounds

Ionic bonding is the net electrostatic force of attraction between ions. This is a strong force of attraction and explains why ionic compounds:

 Are hard crystalline substances

 Have high melting and boiling temperatures

 Are often soluble in water and other polar solvents, but insoluble in non polar solvents

 Do not conduct electricity when solid because their ions cannot move away from their fixed positions in the giant lattice

 Conduct electricity when molten or dissolved in water because the charged ions are free to move

Recap Electrolysis

When molten sodium chloride conducts electricity positive sodium ions move towards the negative terminal (cathode) while the negative ions move to the positive electrode (anode)

At the cathode

+ Na (l) + 2e- Na(l) At the anode

2Cl-(l) Cl2(g) + 2e-

Read p82-3 Summarise the HSW Evidence for existence of ions

Ionic Radii

Atomic radii varies across and down the periodic table. Xray diffraction techniques can measure the spacing between ions in a crystal. From the diffraction patterns it is possible to measure the radii of individual ions.

The radius of a positive ion is smaller than its atomic radius because it has lost an electron and the remaining electrons are being more tightly attracted to the positive nucleus. Similarly the radius of the negative ion is larger than its atomic radius because it has gained electrons. The additional negative charge means the electrons are being held less tightly to the nucleus.

This is further explained by the isoelectronic ions shown above they all have 10 e- but their ionic radius is different because of the attraction from the nucleus.

N3- has 7 protons attracting 10 electrons they are held more loosely therefore the radius is bigger. Mg2+ has lost 2 electrons and therefore its 12 protons are holding on to the remaining 10 electrons more tightly hence the radius is smaller.

Complete Q 10-12 and Activity

Types of Lattice structure

Electron density maps show the arrangement of ions in a lattice. The arrangement varies depending on the size of the ions present.

Sodium chloride has the face centred cubic arrangement

Each ion has 6 neighbours and therefore a coordination number of 6. The Na+ ions are much smaller than the Cl- ions and they fit in the spaces between the Cl- ions. The packing of ions in sodium chloride is known as the rock salt structure. Caesium chloride is different the Cs+ ion is larger than the Na+ ion and more Cl- can fit around the Cs+. Each Cs+ has 8 Cl- ions around it so the Cs+ ion has a coordination number 8. The structure is known as a body centred cubic structure.

We can use the crystal structure to work out the formula of caesium chloride. Each unit cell contains one Cs ion. At each corner of the cell is a Cl- ion which is shared by 7 other unit cells. Meaning that 1/8 th of each Cl- ion belongs to the unit cell.

1xCs+ ions= 1Cs-

Joins with

8x1/8th Cl- ion=1Cl- meaning the formula of a unit cell is CsCl. Energy Changes

The diagram above shows the formation of sodium chloride from its elements but it is a very simplistic illustration. The diagram ignores the fact that when this reaction takes place

 Sodium is actually part of a giant metallic bonded structure  Energy is needed to separate the sodiums from each other (Enthalpy change of atomization)

 Energy is required to remove electrons (the first ionization energy)to form sodium ions

 Chlorine is actually made up of diatomic molecules

 Energy is required to separate the chlorine atoms from each other ( Enthalpy change of atomization)

 Energy is required to add electrons to each chlorine atom to form a chloride ion (the electron affinity)

Where does all this energy come from to turn sodium metal and chlorine molecules into sodium chloride?

Ionic crystals are actually stable because of the large release of energy when oppositely charged ions ɵ come together to form a crystal lattice. This energy is known as the Lattice Energy, ΔH lattice

Two German scientists Max Born(1882-1970) and Fritz Haber(1868-1934) analysed the energy changes when ionic compounds were formed. Their work resulted in Born-Haber cycles which are thermochemical cycles that allow us to investigate the stability and the bonding of ionic compounds.

The strength of ionic bonds are measured as lattice energies in k Jmol-1.

The energy released is greater and the force of attraction between the ions stronger if:

 The charges on the ions are large

 The ions are small allowing them to get close to each other