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4-Inorganic Chemistry and the Periodic Table
The Periodic Table: Group 2
Trends in Group 2
Ionisation Energy
Atomic Radius
Reactivity
Melting Points
Reactions of Group 2 Metals
Reaction with Oxygen
Reaction with Chlorine
Reaction with Water
Group 2 Oxides and Hydroxides
Reaction of Oxides with Water
Reactions with Acids
Reaction of Limewater with Carbon Dioxide
Solubility of Group 2 Hydroxides and Sulphates
Thermal Stability of Group 1 and Group 2 Nitrates and Carbonates
Nitrates
Thermal Stability of Group 1 Nitrates
Thermal Stability of Group 2 Nitrates
Carbonates
Thermal Stability of Group 1 Carbonates
Thermal Stability of Group 2 Carbonates
Practical: Thermal Decomposition
Tests for Group 1 and Group 2 Cations
The Periodic Table: Group 7
Physical Properties of Halogens
Trends in Group 7
First Ionisation Energy:
First Electron Affinity:
Electronegativity:
Reactivity:
Bonding
Ionic Bonding
Covalent Bonding
Reactions of the Halogens
Reactions with Solutions of other Halides
Reactions with Metals
Reaction with Water
Reaction with Alkali
Cold Alkali
Hot Alkali
Test for Halogens
Hydrogen Halides
Physical Properties
Reaction with Water
4-Inorganic Chemistry and the Periodic Table 1
, Reaction with Ammonia
Halide Salts
Solubility
Test for Halides
Chemical Reactions of Halides
Halide Ions as Reducing Agents
Reaction of Halide Salts with Concentrated Sulfuric Acid
Analysis of Inorganic Compounds
Cations
Anions
The Periodic Table: Group 2
S-Block Elements - Group 1 and 2
Trends in Group 2
Ionisation Energy
The energy change for removing both s electrons -
the sum of the first and second ionisation energies.
First ionisation energy decreases down the group:
More shells - more shielding from electrons.
Outer electron further from nucleus - easier to
remove them.
Despite more protons - nuclear charge
increases.
In the same period, the first ionisation energy of
the group 2 metal is larger than the group 1 metal
- group 2 metal has a smaller atomic radius.
Atomic Radius
Atomic radius increases down the group:
More shells
More shielding
Despite more protons
Reactivity
Reactivity increases down the group:
Outer electron is further from nucleus - more
shielding - more easily lost.
Group 2 - forms 2+ ions - need to consider 1st
and 2nd ionisation energy - 2 electrons involved.
Melting Points
Melting points decrease down the group:
Metallic bonding
Ionic radius increases - weaker electrostatic
attraction - more shielding.
4-Inorganic Chemistry and the Periodic Table 2
, Magnesium (hexagonal lattice) increases to
calcium (cubic face centred lattice).
Magnesium - lower melting point - hexagonal
lattice has weaker bonds.
Reactions of Group 2 Metals
Group 2 metals - 2+ oxidation state.
Apart from beryllium - all metals and react to form positive ions of charge 2+.
Reactivity increases down the group - sum of the first and second ionisation energy decreases down the group - easier to
remove the two outer electrons.
Reaction with Oxygen
All group 2 metals burn in air - react with oxygen (redox) and form ionic oxides of formula - MO.
2Ca (s) + O2 (g) → 2CaO (s)
Mg can react slowly without a flame - forms a layer of magnesium oxide.
Reaction with Chlorine
All group 2 element react when heated in chlorine:
Ca + Cl2 → CaCl2
Redox reaction
Reaction with Water
Magnesium - reacts slowly with cold water to produce an alkaline suspension of magnesium hydroxide and hydrogen and
gas:
Mg (s) + 2H2 O (l) → Mg(OH)2 (aq) + H2 (g)
Calcium, strontium and barium react rapidly with cold water to produce alkaline solutions of the metal hydroxide and
bubbles of hydrogen gas:
Ca (s) + 2H2 O (l) → Ca(OH)2 (aq) + H2 (g)
Observations - bubbles, metal disappears and temp increases.
When heated in steam - magnesium burns, producing magnesium oxide and hydrogen:
Mg (s) + H2 O (g) → MgO (s) + H2 (g)
Bright light/glow, white powder produced
Group 2 Oxides and Hydroxides
Group 2 oxides and hydroxides are bases.
General formula - MO and M(OH)2 .
Reaction of Oxides with Water
Magnesium oxide reacts slowly to form a slightly alkaline suspension of magnesium hydroxide:
MgO (s) + H2 O (l) → Mg(OH)2 (s)
Calcium oxide (quicklime) - reacts exothermically with water to form an alkaline suspension of calcium hydroxide:
CaO (s) + H2 O (l) → Ca(OH)2 ⇌ Ca2+ (aq) + 2OH− (aq)
Calcium hydroxide solution - limewater.
Strontium and barium oxides react with water to form alkaline solutions of the hydroxide:
BaO (s) + H2 O (l) → Ba2+ (aq) + 2OH− (aq)
4-Inorganic Chemistry and the Periodic Table 3
, Reactions with Acids
Oxides and hydroxides of the group 2 metals are bases - react with acids to form salts and water:
MO (s) + 2H+ (aq) → M2+ (aq) + H2 O (l)
M(OH)2 (s) + 2H+ (aq) → M2+ (aq) + 2H2 O (l)
Magnesium oxide reacts with dilute sulfuric acid to form magnesium sulfate and water:
MgO (s) + H2 SO4 (aq) → MgSO4 (aq) + H2 O (l)
Calcium hydroxide reacts with nitric acid solution to form calcium nitrate and water:
Ca(OH)2 (s) + 2HNO3 (aq) → Ca(NO3 )2 (aq) + 2H2 O (l)
Beryllium oxide is amphoteric - acts as a base and an acid:
BeO (s) + 2H+ (aq) → Be2+ (aq) + H2 O (l)
BeO (s) + 2OH− (aq) + H2 O (l) → Be(OH)4 2− (aq)
Group 1 hydroxides react with acids in the same way to form salts - are water soluble and form alkaline solutions:
2NaOH (aq) + H2 SO4 (aq) → Na2 SO4 (aq) + 2H2 O (l)
Reaction of Limewater with Carbon Dioxide
Carbon dioxide - acidic oxide.
Reacts with the base calcium hydroxide to form the salt, calcium carbonate and water.
Calcium carbonate - insoluble and appears as a milky precipitate.
Test for carbon dioxide:
Ca(OH)2 (aq) + CO2 (g) → CaCO3 (s) + H2 O (l)
If carbon dioxide is passed into a suspension of calcium carbonate, the milkiness slowly clears as the soluble acid salt,
calcium hydrogen carbonate is formed:
CaCO3 (s) + H2 O (l) + CO2 (g) → Ca(HCO3 )2 (aq)
Calcium hydrogen carbonate is soluble in water - calcium carbonate is insoluble.
To test between a carbonate and a hydrogen carbonate of a group 1 metal.
A solution of the test substance is added to a solution of calcium chloride:
If carbonate - a white precipitate of calcium carbonate is seen.
If hydrogen carbonate - no precipitate forms until the mixture is heated.
Hydrogen carbonate ions decompose forming carbonate ions, carbon dioxide and water:
2HCO3 2− (aq) → CO3 2− (aq) + CO2 (g) + H2 O (l)
Carbonate ions then react with the calcium ions to form a precipitate of insoluble calcium carbonate:
CO3 2− (aq) + Ca2+ (aq) → CaCO3 (s)
Solubility of Group 2 Hydroxides and Sulphates
Solubility of group 2 hydroxides increases down the group.
Solubility of group 2 sulphates decreases down the group.
Sodium sulphate + solution of calcium, strontium or barium compounds → white precipitate of an insoluble sulphate
produced.
Test for sulphate ions:
BaCl2 (aq) + Na2 SO4 (aq) → 2NaCl (aq) + BaSO4 (s)
Ba2+ (aq) + SO4 2− (aq) → BaSO4 (s)
Sodium Hydroxide Sodium Sulphate
4-Inorganic Chemistry and the Periodic Table 4
4-Inorganic Chemistry and the Periodic Table
The Periodic Table: Group 2
Trends in Group 2
Ionisation Energy
Atomic Radius
Reactivity
Melting Points
Reactions of Group 2 Metals
Reaction with Oxygen
Reaction with Chlorine
Reaction with Water
Group 2 Oxides and Hydroxides
Reaction of Oxides with Water
Reactions with Acids
Reaction of Limewater with Carbon Dioxide
Solubility of Group 2 Hydroxides and Sulphates
Thermal Stability of Group 1 and Group 2 Nitrates and Carbonates
Nitrates
Thermal Stability of Group 1 Nitrates
Thermal Stability of Group 2 Nitrates
Carbonates
Thermal Stability of Group 1 Carbonates
Thermal Stability of Group 2 Carbonates
Practical: Thermal Decomposition
Tests for Group 1 and Group 2 Cations
The Periodic Table: Group 7
Physical Properties of Halogens
Trends in Group 7
First Ionisation Energy:
First Electron Affinity:
Electronegativity:
Reactivity:
Bonding
Ionic Bonding
Covalent Bonding
Reactions of the Halogens
Reactions with Solutions of other Halides
Reactions with Metals
Reaction with Water
Reaction with Alkali
Cold Alkali
Hot Alkali
Test for Halogens
Hydrogen Halides
Physical Properties
Reaction with Water
4-Inorganic Chemistry and the Periodic Table 1
, Reaction with Ammonia
Halide Salts
Solubility
Test for Halides
Chemical Reactions of Halides
Halide Ions as Reducing Agents
Reaction of Halide Salts with Concentrated Sulfuric Acid
Analysis of Inorganic Compounds
Cations
Anions
The Periodic Table: Group 2
S-Block Elements - Group 1 and 2
Trends in Group 2
Ionisation Energy
The energy change for removing both s electrons -
the sum of the first and second ionisation energies.
First ionisation energy decreases down the group:
More shells - more shielding from electrons.
Outer electron further from nucleus - easier to
remove them.
Despite more protons - nuclear charge
increases.
In the same period, the first ionisation energy of
the group 2 metal is larger than the group 1 metal
- group 2 metal has a smaller atomic radius.
Atomic Radius
Atomic radius increases down the group:
More shells
More shielding
Despite more protons
Reactivity
Reactivity increases down the group:
Outer electron is further from nucleus - more
shielding - more easily lost.
Group 2 - forms 2+ ions - need to consider 1st
and 2nd ionisation energy - 2 electrons involved.
Melting Points
Melting points decrease down the group:
Metallic bonding
Ionic radius increases - weaker electrostatic
attraction - more shielding.
4-Inorganic Chemistry and the Periodic Table 2
, Magnesium (hexagonal lattice) increases to
calcium (cubic face centred lattice).
Magnesium - lower melting point - hexagonal
lattice has weaker bonds.
Reactions of Group 2 Metals
Group 2 metals - 2+ oxidation state.
Apart from beryllium - all metals and react to form positive ions of charge 2+.
Reactivity increases down the group - sum of the first and second ionisation energy decreases down the group - easier to
remove the two outer electrons.
Reaction with Oxygen
All group 2 metals burn in air - react with oxygen (redox) and form ionic oxides of formula - MO.
2Ca (s) + O2 (g) → 2CaO (s)
Mg can react slowly without a flame - forms a layer of magnesium oxide.
Reaction with Chlorine
All group 2 element react when heated in chlorine:
Ca + Cl2 → CaCl2
Redox reaction
Reaction with Water
Magnesium - reacts slowly with cold water to produce an alkaline suspension of magnesium hydroxide and hydrogen and
gas:
Mg (s) + 2H2 O (l) → Mg(OH)2 (aq) + H2 (g)
Calcium, strontium and barium react rapidly with cold water to produce alkaline solutions of the metal hydroxide and
bubbles of hydrogen gas:
Ca (s) + 2H2 O (l) → Ca(OH)2 (aq) + H2 (g)
Observations - bubbles, metal disappears and temp increases.
When heated in steam - magnesium burns, producing magnesium oxide and hydrogen:
Mg (s) + H2 O (g) → MgO (s) + H2 (g)
Bright light/glow, white powder produced
Group 2 Oxides and Hydroxides
Group 2 oxides and hydroxides are bases.
General formula - MO and M(OH)2 .
Reaction of Oxides with Water
Magnesium oxide reacts slowly to form a slightly alkaline suspension of magnesium hydroxide:
MgO (s) + H2 O (l) → Mg(OH)2 (s)
Calcium oxide (quicklime) - reacts exothermically with water to form an alkaline suspension of calcium hydroxide:
CaO (s) + H2 O (l) → Ca(OH)2 ⇌ Ca2+ (aq) + 2OH− (aq)
Calcium hydroxide solution - limewater.
Strontium and barium oxides react with water to form alkaline solutions of the hydroxide:
BaO (s) + H2 O (l) → Ba2+ (aq) + 2OH− (aq)
4-Inorganic Chemistry and the Periodic Table 3
, Reactions with Acids
Oxides and hydroxides of the group 2 metals are bases - react with acids to form salts and water:
MO (s) + 2H+ (aq) → M2+ (aq) + H2 O (l)
M(OH)2 (s) + 2H+ (aq) → M2+ (aq) + 2H2 O (l)
Magnesium oxide reacts with dilute sulfuric acid to form magnesium sulfate and water:
MgO (s) + H2 SO4 (aq) → MgSO4 (aq) + H2 O (l)
Calcium hydroxide reacts with nitric acid solution to form calcium nitrate and water:
Ca(OH)2 (s) + 2HNO3 (aq) → Ca(NO3 )2 (aq) + 2H2 O (l)
Beryllium oxide is amphoteric - acts as a base and an acid:
BeO (s) + 2H+ (aq) → Be2+ (aq) + H2 O (l)
BeO (s) + 2OH− (aq) + H2 O (l) → Be(OH)4 2− (aq)
Group 1 hydroxides react with acids in the same way to form salts - are water soluble and form alkaline solutions:
2NaOH (aq) + H2 SO4 (aq) → Na2 SO4 (aq) + 2H2 O (l)
Reaction of Limewater with Carbon Dioxide
Carbon dioxide - acidic oxide.
Reacts with the base calcium hydroxide to form the salt, calcium carbonate and water.
Calcium carbonate - insoluble and appears as a milky precipitate.
Test for carbon dioxide:
Ca(OH)2 (aq) + CO2 (g) → CaCO3 (s) + H2 O (l)
If carbon dioxide is passed into a suspension of calcium carbonate, the milkiness slowly clears as the soluble acid salt,
calcium hydrogen carbonate is formed:
CaCO3 (s) + H2 O (l) + CO2 (g) → Ca(HCO3 )2 (aq)
Calcium hydrogen carbonate is soluble in water - calcium carbonate is insoluble.
To test between a carbonate and a hydrogen carbonate of a group 1 metal.
A solution of the test substance is added to a solution of calcium chloride:
If carbonate - a white precipitate of calcium carbonate is seen.
If hydrogen carbonate - no precipitate forms until the mixture is heated.
Hydrogen carbonate ions decompose forming carbonate ions, carbon dioxide and water:
2HCO3 2− (aq) → CO3 2− (aq) + CO2 (g) + H2 O (l)
Carbonate ions then react with the calcium ions to form a precipitate of insoluble calcium carbonate:
CO3 2− (aq) + Ca2+ (aq) → CaCO3 (s)
Solubility of Group 2 Hydroxides and Sulphates
Solubility of group 2 hydroxides increases down the group.
Solubility of group 2 sulphates decreases down the group.
Sodium sulphate + solution of calcium, strontium or barium compounds → white precipitate of an insoluble sulphate
produced.
Test for sulphate ions:
BaCl2 (aq) + Na2 SO4 (aq) → 2NaCl (aq) + BaSO4 (s)
Ba2+ (aq) + SO4 2− (aq) → BaSO4 (s)
Sodium Hydroxide Sodium Sulphate
4-Inorganic Chemistry and the Periodic Table 4
