Group 2

In a nutshell

The Group 2 metals (magnesium to barium) all have two electrons in an outer s sub-shell. In their reactions they lose both of them to form 2+ ions, so every reaction of a Group 2 metal is a redox reaction in which the metal is oxidised.

Reactivity increases down the group, because those two outer electrons get easier to remove. This subtopic also covers what Group 2 oxides do in water, and how Group 2 compounds are used as bases.

Assumed knowledge: Redox: oxidation numbers and oxidised vs reduced; Acids: bases, alkalis and neutralisation; Periodicity: first ionisation energy.

Core content

Group 2 metals: the s² outer shell and 2+ ions

Every Group 2 atom has an outer shell s² electron configuration: two electrons in the outermost s sub-shell.

When a Group 2 metal reacts, each atom loses its two outer s electrons to form a 2+ ion:

M → M²⁺ + 2e^-    e.g. Ca → Ca²⁺ + 2e^-

What this means for every reaction of the metal:

• The metal's oxidation number goes from 0 to +2, so the metal is oxidised (loss of electrons).
• Something else in the reaction gains those electrons and is reduced.
• The metal gives electrons away, so Group 2 metals act as reducing agents.
• In all their compounds, Group 2 elements have oxidation number +2.

Group 2 metals: redox reactions with oxygen

Group 2 metals burn in oxygen to form a metal oxide, MO:

2M(s) + O₂(g) → 2MO(s)

For example:

2Mg(s) + O₂(g) → 2MgO(s)

2Ca(s) + O₂(g) → 2CaO(s)

Track the oxidation numbers:

• Metal: 0 → +2, so the metal is oxidised.
• Oxygen: 0 → -2, so oxygen is reduced.

What you see:

• Magnesium burns with a bright white light and leaves a white solid (magnesium oxide).
• The Group 2 oxides are white ionic solids.
• The reaction becomes more vigorous down the group.

Group 2 metals: redox reactions with water

Group 2 metals react with water to form a metal hydroxide and hydrogen:

M(s) + 2H₂O(l) → M(OH)₂(aq) + H₂(g)

For example:

Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g)

Ba(s) + 2H₂O(l) → Ba(OH)₂(aq) + H₂(g)

Track the oxidation numbers:

• Metal: 0 → +2, so the metal is oxidised.
• Hydrogen: +1 → 0, but only the H atoms that end up in H₂. They are reduced.
• The H atoms that end up in OH^- stay at +1, and oxygen stays at -2.

Magnesium is the odd one out, because it reacts so slowly with cold water:

Note the product switch: water gives the hydroxide, but steam gives the oxide (MgO).

Observations for Ca, Sr and Ba with water, each more vigorous down the group:

• effervescence (fizzing, as hydrogen is given off)
• the metal dissolves
• the solution warms up
• the solution formed is alkaline, because it contains OH^- ions
• with calcium, the solution goes cloudy white, because calcium hydroxide is only slightly soluble and some of it stays as a solid

Group 2 metals: redox reactions with dilute acids

Group 2 metals react with dilute acids to form a salt and hydrogen:

metal + acid → salt + hydrogen

For example:

Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)

Ba(s) + 2HCl(aq) → BaCl₂(aq) + H₂(g)

Mg(s) + H₂SO₄(aq) → MgSO₄(aq) + H₂(g)

Track the oxidation numbers:

• Metal: 0 → +2, so the metal is oxidised.
• Hydrogen: +1 → 0, so hydrogen is reduced.
• Cl stays at -1 (and S stays at +6, O at -2): no change.

What you see: effervescence and the metal dissolving, both more vigorous down the group.

For these equations, use hydrochloric acid or dilute sulfuric acid. OCR limits this point to reactions that give a salt and hydrogen.

Group 2 metals: why reactivity increases down the group

Reactivity increases down the group because the first and second ionisation energies decrease, so the two outer electrons are lost more easily.

To form a 2+ ion, an atom must lose two electrons, so both ionisation energies matter. Using calcium as the example:

First ionisation energy: Ca(g) → Ca⁺(g) + e^-

Second ionisation energy: Ca⁺(g) → Ca²⁺(g) + e^-

Both equations need (g) on every species. The second one starts from the 1+ ion, not the atom.

The explanation is a chain. Each link is a separate point, so give all of them, in this order:

1. Atomic radius increases down the group: there are more shells, so the outer electrons are further from the nucleus.
2. There is more shielding by the extra inner shells of electrons.
3. So there is less nuclear attraction on the outer electrons. This outweighs the increase in nuclear charge (more protons).
4. So less energy is needed to remove the two outer electrons: the first and second ionisation energies decrease.
5. So the metal loses electrons more easily: it is oxidised more easily, so it is more reactive (a stronger reducing agent).

Still don't get it? · why more protons does not mean a stronger hold

Imagine a fridge magnet holding a sheet of paper. Now swap it for a stronger magnet, but put five thick books between the magnet and the paper. The stronger magnet loses its grip, because the paper is now much further away and there is a lot of stuff in between.

Rebuild it for the atom, one step at a time:

1. The magnet is the nucleus. Going down Group 2 it gets "stronger": barium has 56 protons, magnesium only 12.
2. The paper is the two outer electrons, the ones that must be removed to make the 2+ ion.
3. The books are the inner shells of electrons. Each new period adds a whole extra shell, so barium has far more of them.
4. Two things weaken the grip on the outer electrons: they are further away (bigger atomic radius), and the inner shells shield them (inner electrons repel the outer ones).
5. Those two effects win over the extra protons. So the outer electrons are held less strongly, and it takes less energy to pull them off.
6. A metal reacts by losing those electrons, so the easier they come off, the more reactive the metal.

Back to the exam wording: down the group, atomic radius increases and there is more shielding, so there is less nuclear attraction on the outer electrons (this outweighs the increased nuclear charge). The first and second ionisation energies decrease, so the metal loses electrons more easily and is more reactive.

Group 2 compounds: oxides with water

Group 2 oxides react with water to form the metal hydroxide:

MO(s) + H₂O(l) → M(OH)₂(aq)

e.g. CaO(s) + H₂O(l) → Ca(OH)₂(aq)

The hydroxide that dissolves releases OH^- ions, so the solution is alkaline:

Ca(OH)₂(aq) → Ca²⁺(aq) + 2OH^-(aq)

This is not a redox reaction: no oxidation numbers change (the metal stays +2, O stays -2, H stays +1).

The alkalinity increases down the group: the pH of the solution goes up from Mg to Ba.

The reason: the hydroxides become more soluble down the group, so more OH^- ions are released into solution, giving a higher pH.

The exact pH depends on how much solid is added and how much dissolves. What matters is that every solution is alkaline (pH above 7) and the pH rises down the group.

Still don't get it? · why the pH goes up down the group

Imagine stirring the same spoonful of four different sugary powders into four glasses of water. Some powders dissolve completely; one barely dissolves and mostly sits at the bottom. Only the part that dissolves sweetens the drink, so the glass with the least dissolved powder tastes least sweet.

Rebuild it for Group 2, one step at a time:

1. Every Group 2 oxide reacts with water to make the same kind of product: a hydroxide, M(OH)₂.
2. Only the hydroxide that dissolves can release OH^- ions into the water. Solid sitting at the bottom does nothing to the pH.
3. Magnesium hydroxide barely dissolves, so very few OH^- ions get into the water: the solution is only weakly alkaline.
4. Going down the group, each hydroxide dissolves more than the one above it, so more OH^- ions are released.
5. More OH^- ions in solution means a more alkaline solution, which means a higher pH.

Back to the exam wording: down the group the hydroxides are more soluble, so more OH^- ions are released, so the solutions are more alkaline and have a higher pH. Do not explain it by saying "the metal is more reactive": that is a different trend.

Group 2 compounds: uses as bases

Several Group 2 compounds are used as bases to neutralise acids. These are neutralisation reactions, not redox: no oxidation numbers change.

Calcium hydroxide and acid soil. Soil acidity is H⁺(aq) ions:

Ca(OH)₂(s) + 2H⁺(aq) → Ca²⁺(aq) + 2H₂O(l)

As a full equation, for example with nitric acid:

Ca(OH)₂(s) + 2HNO₃(aq) → Ca(NO₃)₂(aq) + 2H₂O(l)

Antacids. Stomach acid is hydrochloric acid, so antacid equations use HCl:

Mg(OH)₂(s) + 2HCl(aq) → MgCl₂(aq) + 2H₂O(l)

CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)

Two points that carry marks:

• A hydroxide + acid gives salt + water only.
• A carbonate + acid gives salt + water + carbon dioxide. Do not forget the CO₂.

The spec's examples are "for example (but not limited to)", so you can be given another Group 2 base and asked for its equation. The patterns above always apply.

Worked examples

1. Ionisation energy equations and the energy to form a 2+ ion

Using the chart above, write the equation for the second ionisation energy of barium, and calculate the energy needed to turn 1 mol of gaseous barium atoms into gaseous Ba²⁺ ions. Compare it with magnesium.

The second ionisation energy starts from the 1+ ion:

Ba⁺(g) → Ba²⁺(g) + e^-

Forming Ba²⁺(g) from Ba(g) needs the first and second ionisation energies:

$\text{energy} = \text{1st IE} + \text{2nd IE}$ $\text{Ba}: \; 503 + 965 = +1468 \ \text{kJ mol}^{-1}$ $\text{Mg}: \; 738 + 1451 = +2189 \ \text{kJ mol}^{-1}$ $\text{difference} = 2189 - 1468 = 721 \ \text{kJ mol}^{-1}$

Barium needs 721 kJ mol^-1 less energy to form its 2+ ion, which is why it is the more reactive metal.

Where marks go:

• Ionisation energies are always positive (energy must be put in), so the total is +1468 kJ mol^-1, never negative.
• Writing Ba(g) → Ba²⁺(g) + 2e^- for the "second ionisation energy" is wrong: that equation is the sum of the first and second.

2. Model answer: explain the trend in reactivity

"Explain why barium is more reactive than magnesium." A full answer makes these points, in this order:

1. Barium atoms have more shells, so a larger atomic radius: the outer electrons are further from the nucleus.
2. Barium has more shielding from inner shells.
3. So there is less nuclear attraction on barium's outer electrons, even though barium has a greater nuclear charge.
4. So barium has lower first and second ionisation energies: less energy is needed to remove its two outer electrons.
5. So barium loses electrons more easily (is oxidised more easily), so it is more reactive.

Each point must be comparative ("larger", "more", "less", "lower"). "Barium has a big atom and lots of shielding" does not compare it with anything.

3. Synoptic calculation: hydrogen from a Group 2 metal and water

Calculate the volume of hydrogen, in dm³ at RTP, produced when 0.250 g of calcium reacts completely with water. (A_r of Ca = 40.1; molar gas volume = 24.0 dm³ mol^-1 at RTP, both given on the data sheet.)

Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g)

$n(\text{Ca}) = \frac{m}{M} = \frac{0.250 \ \text{g}}{40.1 \ \text{g mol}^{-1}} = 6.234 \times 10^{-3} \ \text{mol}$ $\text{ratio Ca} : \text{H}_2 = 1 : 1$ $n(\text{H}_2) = 6.234 \times 10^{-3} \ \text{mol}$ $V(\text{H}_2) = n \times 24.0 \ \text{dm}^3 \ \text{mol}^{-1} = 6.234 \times 10^{-3} \times 24.0 = 0.1496 \ \text{dm}^3$ $V(\text{H}_2) = 0.150 \ \text{dm}^3 \ \text{(3 s.f.)}$

Where marks go:

• The ratio comes from the balanced equation: 1 Ca gives 1 H₂. Students who use the 2 in "2H₂O" halve or double the answer.
• Keep the unrounded moles until the last line, then give 3 s.f., matching 0.250 g and 24.0.
• The unit is dm³. If the question asks for cm³, multiply by 1000: 150 cm³.

4. Synoptic calculation: which antacid neutralises more acid per gram?

Show which neutralises more hydrochloric acid per gram: Mg(OH)₂ or CaCO₃. (A_r: H = 1.0, C = 12.0, O = 16.0, Mg = 24.3, Ca = 40.1.)

Both react with HCl in a 1 : 2 ratio:

Mg(OH)₂(s) + 2HCl(aq) → MgCl₂(aq) + 2H₂O(l)

CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)

Magnesium hydroxide:

$M_r[\text{Mg(OH)}_2] = 24.3 + 2(16.0 + 1.0) = 58.3$ $n[\text{Mg(OH)}_2] = \frac{1.00 \ \text{g}}{58.3 \ \text{g mol}^{-1}} = 1.715 \times 10^{-2} \ \text{mol}$ $n(\text{HCl}) = 2 \times 1.715 \times 10^{-2} = 3.43 \times 10^{-2} \ \text{mol per gram}$

Calcium carbonate:

$M_r(\text{CaCO}_3) = 40.1 + 12.0 + 3(16.0) = 100.1$ $n(\text{CaCO}_3) = \frac{1.00 \ \text{g}}{100.1 \ \text{g mol}^{-1}} = 9.990 \times 10^{-3} \ \text{mol}$ $n(\text{HCl}) = 2 \times 9.990 \times 10^{-3} = 2.00 \times 10^{-2} \ \text{mol per gram}$

Magnesium hydroxide neutralises more acid per gram (3.43 × 10^-2 mol vs 2.00 × 10^-2 mol), because its M_r is smaller and both bases react with 2 mol of HCl.

Where marks go: forgetting the × 2 from the equation, or writing Mg(OH)₂ as "MgOH₂" and getting M_r = 42.3 instead of 58.3.

Common exam mistakes

• Saying the nuclear charge decreases down the group. It increases (more protons). The point is that greater distance and more shielding outweigh it, so the nuclear attraction on the outer electrons decreases.
• Stopping the reactivity chain early. "The outer electrons are further away" alone is not enough. Link it all the way: further away and more shielded, less nuclear attraction, lower ionisation energies, electrons lost more easily.
• Saying ionisation energy increases down the group. Both the first and second ionisation energies decrease from Mg to Ba.
• Writing the second ionisation energy as M(g) → M²⁺(g) + 2e^-. The second ionisation energy starts from the 1+ ion: M⁺(g) → M²⁺(g) + e^-. And every species needs (g).
• Swapping the magnesium products. Magnesium with steam gives MgO + H₂. Magnesium with (cold) water gives Mg(OH)₂ + H₂, very slowly.
• Wrong formulae. The ion is M²⁺, so it is Ca(OH)₂ (not CaOH), MgCl₂ (not MgCl), Ba(OH)₂ (brackets needed). Hydrogen gas is H₂, never H.
• Writing "oxidising" for "oxidised". "Calcium is oxidising" means calcium is an oxidising agent, the opposite of what happens. Calcium is oxidised (0 to +2); it is a reducing agent.
• Calling every reaction in this topic redox. Metal + oxygen, water or acid is redox. Oxide + water, and the base + acid reactions (Ca(OH)₂ on soil, antacids), are not: no oxidation numbers change.
• Explaining the pH trend with reactivity. The pH rises down the group because the hydroxides are more soluble, so more OH^- ions are released. "Barium is more reactive" does not explain the pH.
• Reversing the pH trend. MgO in water gives the lowest pH (about 9 to 10, because Mg(OH)₂ is only slightly soluble); BaO gives the highest.
• Leaving CO₂ out of the carbonate equation. CaCO₃ + 2HCl gives CaCl₂ + H₂O + CO₂.
• Mixing up the uses. Ca(OH)₂ treats acid soil; Mg(OH)₂ and CaCO₃ are the antacids.

Key definitions

• First ionisation energy: the energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions (the removal of 1 mol of electrons from 1 mol of gaseous atoms).
• Second ionisation energy (equation only; a definition is not required by OCR here): M⁺(g) → M²⁺(g) + e^-, e.g. Ca⁺(g) → Ca²⁺(g) + e^-.
• Oxidation: loss of electrons, or an increase in oxidation number.
• Reduction: gain of electrons, or a decrease in oxidation number.
• Reducing agent: a species that gives electrons to another species, and is itself oxidised. Group 2 metals are reducing agents.
• Base: a substance that neutralises an acid to form a salt (e.g. metal oxides, metal hydroxides, carbonates).
• Alkali: a substance that releases OH^- ions in aqueous solution (a soluble base).

Specification

• I can state that Group 2 elements have an outer shell s² electron configuration, and explain that they lose these two electrons in redox reactions to form 2+ ions. (3.1.2 a)
• I can write equations for the reactions of Mg to Ba with oxygen, and use oxidation numbers to show what is oxidised and reduced. (3.1.2 b i)
• I can write equations for the reactions of Mg to Ba with water (including Mg with steam), describe what is seen, and use oxidation numbers to show they are redox. (3.1.2 b ii)
• I can write equations for the reactions of Mg to Ba with dilute acids to form a salt and hydrogen, and use oxidation numbers to show they are redox. (3.1.2 b iii)
• I can state that reactivity increases from Mg to Ba, as shown by these redox reactions. (3.1.2 b)
• I can explain the increase in reactivity down the group in terms of the first and second ionisation energies (atomic radius, shielding, nuclear attraction), and write the equation for a second ionisation energy. (3.1.2 c)
• I can write equations for the action of water on Group 2 oxides, give the approximate pH of the solutions, and explain the trend of increasing alkalinity down the group. (3.1.2 d)
• I can describe uses of Group 2 compounds as bases, with equations, including Ca(OH)₂ to neutralise acid soils and Mg(OH)₂ and CaCO₃ as antacids. (3.1.2 e)

Related notes

• Electron structure: s sub-shells and electron configurations
• Redox: oxidation numbers, and metals reacting with acids
• Acids: bases, alkalis and neutralisation equations
• Amount of substance: moles and gas volumes, used in the synoptic worked examples
• Periodicity: first ionisation energy and its trend down a group
• The halogens: the Group 7 non-metals, which act as oxidising agents
• Qualitative analysis: test-tube tests for ions, including PAG4 on Group 2 salts