The Periodic Table

Ionisation

Removal of an electron.

First ionisation energies

The energy required to remove 1 mole of an electron from 1 mole of gaseous atoms - endothermic process.

X (g) --> X+ (g) + e- 

Factors affecting ionisation energies

• Atomic radius
  • Greater radius = smaller IE

• Nuclear charge
  • More protons = greater nuclear charge = large IE

• Electron shielding
  • Inner chells of electrons repel outer shell electrons
  • More shielding = smaller IE

Trends of first ionisation energies down a group

• DECREASES.
• Number of shells increases.
• More shielding.
• Atomic radius increases.

Trends of first ionisation energies across a period

• INCREASES.
• Nuclear charge increases as more protons are added.
• More electrons are added to the same shell.
• No shielding difference.
• Atomic radius decreases.

Beryllium to Boron / Magnesium to Aluminium

• The outer electron is in a p orbital - slightly higher energy level than s orbital
• Further away from the nucleus
• Shielding from s orbital
• DECREASE in IE

Nitrogen to Oxygen / Phosphorus to Sulphur

• Two electrons in the outer orbital
• Repulsion of electron makes it easier to remove
• DECREASE in IE

1) Graphite

• Arranged in sheets covalently bonded - 3 bonds each with one delocalised electron.
• Sheets bonded with weak induced dipole-dipole forces.

Properties

• Dry lubricant - weak forces between layers allow sheets to slip over each other.
• Can conduct electricity - delocalised electron is free to move.
• Less dense - layers held far apart.
• High melting point - strong covalent bonds.
• Insoluble - covalent bonds are too strong to be broken.

2) Graphene

• Sheet of carbon atoms joined together in hexagons.
• One atom thick - 2D.

Properties

• Can conduct electricity - delocalised electron can carry charge.
• Strong - delocalised electron strengthens the covalent bond.
• Transparent and light - touchscreens.
• High strength and low mass - high speed electronics, aircraft technology.

3) Diamond

• Made up of carbon atoms - each C atom is covalently bonded to 4 others.
• Tetrahedral shape.
• Crystal lattice structure.

Properties

• High melting point - strong covalent bonds.
• Hard - diamond tipped drills.
• Thermal conductor - vibrations travel easily through.
• Cannot conduct electricity - outer electron held in localised bond.
• Will not dissolve in any solvents.

4) Sillicon

• Similar to diamond - 4 Si covalently bonded.

Simple molecular structure

• Made up of small simple molecules e.g. NH2, H2, O2.
• 3D - bonded by weak van der waals forces.
• Covalent bonds between atoms.
• Melting and boiling point depends on the strength of the induced dipole-dipole forces.

Properties

• Low melting and boiling points - weak van der waals forces.
• No electrical conductivity - no charged particles can move.
• Soluble in non-polar solvents - van der waals forces form.

Giant covalent structure

• 3D structure of atoms bonded by strong covalent bonds.

Properties

• High melting and boiling points - many strong covalent bonds.
• Non conductors of electricity - no free charged particles can move.
• Solubility - insoluble in both polar and non-polar solvents as covalent bonds are too strong to break.

Melting and boiling points across a period

Metals (Li, Be, Na, Mg, Al)

• Melting point increases.
• Metallic bonding is stronger - as atomic radius increases, number of delocalised electrongs increases.

Giant covalent ( Si, C)

• Strong covalent bonds.

Simple molecular structure 

• Melting point decreases.
• Weak intermolecular forces to overcome.

Noble gases 

• Lowest melting point- monoatomic and have weak induced dipole-dipole forces.

Group 2 - The alkaline earth metals

• Form 2+ ions - have 2 electrons in the outer shell.
• Have a high m.p and b.p.
• Light metals with low densities and colourless compounds.
• First ionisation energies DECREASE down the group.

Melting point and boiling point decrease for...

• Mg - decrease in metallic strength.
• Ca - increase in atomic radius, less attraction for delocalised electrons.

Reactivity of elements in group 2

Strong reducing agents

• Loses 2 electrons from outer shell.
• M --> M2+ + 2e- (general eq: M is group 2 element)
• Oxidised as the oxidaition number changes from 0 to 2+ - therefore stronger reducing agent.

Reactivity increases down the group

• Due to increasing ease of losing electrons.

1) With water.

• Forms hydroxides, H2 gas formed. 
• M (s) + 2H2O (l) --> M(OH)2 (aq) + H2 (g)
• M: 0 to +2 so is oxidisd. H: +1 to 0 so is reduced.
• Reactivity INCREASES down the group.

2) With oxygen.

• Forms oxides (MO), white solid.
• 2M (s) + O2 (g) --> 2MO (s)
• M: 0 to +2 so is oxidised. O2: 0 to -2 so is reduced.

3) Dilute acid.

• Forms a salt and H2.
• M (s) + 2HCl (aq) --> MCl (aq) + H2 (g)
• M: 0 to +2 so is oxidised. H: +1 to 0 so is reduced.
• Reactivity INCREASES down the group.

Are bases.

Neutralised by acids to form salt and water.

Oxides

• React with water to form a slution of the metal hydroxide.
• Gets more soluble down the group.

Hydroxides

• Dissolve in water to form alkaline solutions.
• Solubility increases down the group - more alkaline solutions formed.

(Carbonates decompose at a higher temperature.)

The Halogens

• Highly reactive non-metals.
• Low m.p and b.p
• Exist as diatomic molecules.
• Number of electrons increase - induced dipole-dipole increases between molecules.

• F2 - pale yellow, gas at RTP
• Cl2 -  green, gas at RTP
• Br2 - red-brown, liquid at RTP.
• I2 - grey, solid at RTP.

1) As oxidising agents

• Strong
• Remve electrons.
• Reactivity decreases down the group - atomic radius increases, shielding increases, ability to gain electron decreases.
• X (g) + e- --> X- (g)

2) Redox reactions

• Often displacement reactions with an organic solution.

• Cl2: with water = pale green > with cyclohexane = pale green.
• Br2: with water = orange > with cyclohexane = orange.
• I2: with water= brown > with cyclohexane = violet.

Displacement Reactions

Chlorine displaced both Br- and I-

• Cl2 + 2Br- --> 2Cl- + Br2
• Cl2 + 2I- --> 2Cl- + I2

Bromine displaces I- only

• Br2 + 2I- --> 2Br- + I2

Iodine doesn't displace Cl- or Br- 

Disproportionation

The reaction of something being both oxidised and reduced at the same time. 

• Halogens undergo this with alkalis - cold and dilute.
• X + 2NaOH --> NaXo + NaX + H2O
• X2 + 2OH- --> XO- + X- + H20 
• X goes from 0 to +1 (NaXO) and -1 (NaX) so is oxidised and reduced.

Chlorine

• Cl2 + H2O <--> HClO + HCl
• 2NaOH +Cl2 --> NaCl + NaClO + H2O
• HClO + 2H2O <--> CLO- H3O+ (Bleach - chlorate ions kill bacteria in swimming pools but may react with organic matter linked with causes of cancer)

Halides

• NaCl - common calt and NaF - in toothpaste to prevent decay.