The Periodic Table

Ionisation Energies

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- 

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Factors Affecting Ionisation Energies

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
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Trends of Ionisation Energies

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.
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Exceptions of Trends

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
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Structure, Bonding and Properties - Graphite

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.
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Structure, Bonding and Properties - Graphene

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.
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Structure, Bonding and Properties - Diamond and Si

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.
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Simple Molecular Structure

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.
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Giant Covalent Structure

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.
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Melting and Boiling Points Across a Period

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.
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Group 2 - The Alkaline Earth Metals

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.
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Group 2 - Reactivity

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.
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Group 2 - Reactions

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.
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Group 2 - Oxides and Hydroxides

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.)

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Group 7 - The Halogens

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.
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Group 7 - Reactions

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.
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Group 7 - Displacement Reactions

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- 

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Group 7 - Disproportionation, Bleach and Halides

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.

 

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