Group 2 Reactivity

Spec Mapping — OCR H432 Module 3.1.2 — Group 2, covering the electron configurations of Group 2 elements, the trend in reactivity down the group (Mg → Ba) with water, oxygen and dilute acids, the explanation of the trend in terms of first and second ionisation energies, balanced equations for the relevant reactions, and the role of Group 2 metals as reducing agents (refer to the official OCR H432 specification document for exact wording).

Group 2 (the alkaline-earth metals) is OCR's standard example of a "down-the-group reactivity trend driven by ionisation energy": as you descend from beryllium (an inert metal that does not even react with water) through magnesium, calcium and strontium to barium (which behaves almost like a Group 1 alkali metal), reactivity with water, oxygen and dilute acids increases steadily. The driver is straightforward — every Group 2 reaction is a redox reaction in which the metal loses its two outer s electrons to form M²⁺, and losing those electrons becomes progressively easier as the outer-shell radius increases and shielding rises down the group. The first and second ionisation energies — quoted side-by-side in the previous lesson — fall down the group, and their sum (IE₁ + IE₂) is the thermodynamic quantity that ultimately determines reactivity. This lesson sets up the structural understanding of Group 2 reactivity and writes out every balanced equation OCR expects you to know.

Key Definition: the reactivity of a Group 2 metal is the rate and vigour with which it loses two electrons to form M²⁺(aq) ions; in OCR Module 3.1.2 it is operationally measured by the speed and vigour of reaction with water, dilute acids, and oxygen.

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Outer electron configurations

All Group 2 (alkaline-earth) metals have two electrons in their outer s sub-shell (ns²):

Element	Z	Full configuration	Outer config	Period
Be	4	1s² 2s²	2s²	2
Mg	12	[Ne] 3s²	3s²	3
Ca	20	[Ar] 4s²	4s²	4
Sr	38	[Kr] 5s²	5s²	5
Ba	56	[Xe] 6s²	6s²	6
(Ra)	88	[Rn] 7s²	7s²	7

In all their reactions, Group 2 elements lose their two outer s electrons to form 2+ ions with a noble-gas configuration. Hence the entire Group 2 chemistry is dominated by the half-reaction:

$\text{M(s)} \rightarrow \text{M}^{2+}\text{(aq)} + 2\text{e}^- \qquad \text{(oxidation)}$M(s)→M2+(aq)+2e−(oxidation)

Because they donate electrons, Group 2 metals function as reducing agents — they reduce another species (water, oxygen, an acid's H⁺) by transferring electrons to it. The strength of a Group 2 metal as a reducing agent depends on how readily it can shed those two electrons, which is measured by IE₁ + IE₂ in the gas phase and ultimately by the standard electrode potential $E^\circ$E∘ of M²⁺/M in aqueous solution.

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Trend in reactivity down the group

Reactivity increases down Group 2. A qualitative summary:

Metal	Behaviour with cold water	Behaviour with steam
Be	No reaction (even at high temperature; passivated by BeO surface)	Very slow
Mg	Very slow (passivated by surface Mg(OH)₂ coating)	Rapid; brilliant white flame
Ca	Moderate — fizzing as cold tap water is added	Rapid
Sr	Faster than Ca — vigorous effervescence	Very rapid
Ba	Very rapid; resembles a Group 1 alkali metal	Violent

Why reactivity increases down the group

Down Group 2, three quantities change in the same direction:

1. Atomic radius increases — extra shells (Be 2s², Mg 3s², Ca 4s², …) place the outer electrons further from the nucleus.
2. Shielding increases — each row adds a complete inner-shell to shield the outer electron from the nuclear charge.
3. Nuclear charge increases — Be has Z = 4; Ba has Z = 56. (This opposes the first two effects.)

Although the nuclear charge rises, the first two effects — distance and shielding — dominate, so the outer 2s/3s/4s/5s/6s electrons are progressively easier to remove. The first and second ionisation energies fall accordingly:

Element	IE₁ / kJ mol⁻¹	IE₂ / kJ mol⁻¹	Sum IE₁ + IE₂
Be	899	1757	2656
Mg	738	1451	2189
Ca	590	1145	1735
Sr	549	1064	1613
Ba	503	965	1468

The sum IE₁ + IE₂ is what governs reactivity, because both electrons must be removed to form M²⁺. Notice the sum falls smoothly from 2656 kJ mol⁻¹ for Be to 1468 kJ mol⁻¹ for Ba — a ~45 % decrease over five elements. Hence the reactivity rises smoothly down the group.

This is the mirror-image of the halogen trend you will meet later in the course: halogens gain electrons to become reduced, and their reactivity decreases down the group because the incoming electron is held less tightly by the larger atoms.

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Reaction with water

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

$\text{M(s)} + 2\text{H}_2\text{O(l)} \rightarrow \text{M(OH)}_2\text{(aq)} + \text{H}_2\text{(g)}$M(s)+2H2​O(l)→M(OH)2​(aq)+H2​(g)

Example: calcium

$\text{Ca(s)} + 2\text{H}_2\text{O(l)} \rightarrow \text{Ca(OH)}_2\text{(aq)} + \text{H}_2\text{(g)}$Ca(s)+2H2​O(l)→Ca(OH)2​(aq)+H2​(g)

Observations:

• Metal sinks (briefly) then rises as hydrogen bubbles lift it.
• Vigorous effervescence (H₂ gas).
• Metal gradually dissolves; surface becomes whitish briefly as Ca(OH)₂ forms.
• A colourless solution forms; a faint white suspension may appear if Ca(OH)₂ exceeds its (low) saturation solubility.
• Solution becomes alkaline (pH ~12, because Ca(OH)₂ is a strong base in solution). Phenolphthalein indicator turns pink; universal indicator turns purple.
• Reaction is exothermic — the test tube becomes noticeably warmer.

Magnesium exception (cold water vs steam)

Magnesium reacts very slowly with cold water:

$\text{Mg(s)} + 2\text{H}_2\text{O(l)} \rightarrow \text{Mg(OH)}_2\text{(s)} + \text{H}_2\text{(g)} \quad \text{[very slow]}$Mg(s)+2H2​O(l)→Mg(OH)2​(s)+H2​(g)[very slow]

A coat of insoluble Mg(OH)₂ quickly forms on the metal's surface and passivates it, blocking further attack by water. The reaction is virtually undetectable over minutes; over weeks, magnesium ribbon in water does corrode slightly.

With steam at elevated temperature, however, Mg reacts rapidly to form magnesium oxide (not hydroxide) and hydrogen:

$\text{Mg(s)} + \text{H}_2\text{O(g)} \rightarrow \text{MgO(s)} + \text{H}_2\text{(g)}$Mg(s)+H2​O(g)→MgO(s)+H2​(g)

This is the brilliant white flame demonstration: a Mg ribbon held above a beaker of boiling water in a Bunsen flame burns with an intensely bright white flame, producing white MgO powder. The difference (MgO vs Mg(OH)₂ as product) arises because at steam temperature the hydroxide decomposes thermally: $\text{Mg(OH)}_2\text{(s)} \xrightarrow{\Delta} \text{MgO(s)} + \text{H}_2\text{O(g)}$Mg(OH)2​(s)Δ​MgO(s)+H2​O(g) so the final solid product is the oxide.

Equations for the other Group 2 metals

$\text{Ca(s)} + 2\text{H}_2\text{O(l)} \rightarrow \text{Ca(OH)}_2\text{(aq)} + \text{H}_2\text{(g)}$Ca(s)+2H2​O(l)→Ca(OH)2​(aq)+H2​(g)

$\text{Sr(s)} + 2\text{H}_2\text{O(l)} \rightarrow \text{Sr(OH)}_2\text{(aq)} + \text{H}_2\text{(g)}$Sr(s)+2H2​O(l)→Sr(OH)2​(aq)+H2​(g)

$\text{Ba(s)} + 2\text{H}_2\text{O(l)} \rightarrow \text{Ba(OH)}_2\text{(aq)} + \text{H}_2\text{(g)}$Ba(s)+2H2​O(l)→Ba(OH)2​(aq)+H2​(g)

State symbols are mandatory — OCR awards a separate mark for them in mark schemes.

flowchart LR
  A[Group 2 metal M] --> B[Add to cold water]
  B --> C{Position in group}
  C -->|Be| D[No reaction passivated by BeO]
  C -->|Mg| E[Slow; surface Mg OH 2 layer]
  C -->|Ca, Sr, Ba| F[Effervescence increasing down group]
  F --> G[M + 2H2O to M OH 2 aq + H2 g]
  G --> H[Solution alkaline; phenolphthalein turns pink]
  E --> I[Mg + steam to MgO + H2 brilliant white flame]

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Reaction with oxygen

Group 2 metals burn in oxygen to form ionic oxides (M²⁺ O²⁻):

$2\text{M(s)} + \text{O}_2\text{(g)} \rightarrow 2\text{MO(s)}$2M(s)+O2​(g)→2MO(s)

Examples:

$2\text{Mg(s)} + \text{O}_2\text{(g)} \rightarrow 2\text{MgO(s)} \quad \text{[brilliant white flame, white MgO powder]}$2Mg(s)+O2​(g)→2MgO(s)[brilliant white flame, white MgO powder]

$2\text{Ca(s)} + \text{O}_2\text{(g)} \rightarrow 2\text{CaO(s)} \quad \text{[red-orange flame]}$2Ca(s)+O2​(g)→2CaO(s)[red-orange flame]

$2\text{Ba(s)} + \text{O}_2\text{(g)} \rightarrow 2\text{BaO(s)}$2Ba(s)+O2​(g)→2BaO(s)

Reactivity increases down the group: Ba is so reactive with oxygen and water that it must be stored under oil. With excess oxygen, Sr and Ba can form peroxides (MO₂) and superoxides (MO₂ with extra electrons), but at A-Level you only need the simple oxides.

The flame colours in these combustions (Mg white, Ca brick-red, Sr crimson, Ba green) form the basis of the flame test for cations — OCR does not require recall of flame test colours (see the qualitative-analysis lesson at the end of this course), but they are useful diagnostic cues in practical work.

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Reaction with dilute acids

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

$\text{M(s)} + 2\text{HCl(aq)} \rightarrow \text{MCl}_2\text{(aq)} + \text{H}_2\text{(g)}$M(s)+2HCl(aq)→MCl2​(aq)+H2​(g)

Examples with dilute hydrochloric acid:

$\text{Mg(s)} + 2\text{HCl(aq)} \rightarrow \text{MgCl}_2\text{(aq)} + \text{H}_2\text{(g)}$Mg(s)+2HCl(aq)→MgCl2​(aq)+H2​(g)

$\text{Ca(s)} + 2\text{HCl(aq)} \rightarrow \text{CaCl}_2\text{(aq)} + \text{H}_2\text{(g)}$Ca(s)+2HCl(aq)→CaCl2​(aq)+H2​(g)

With dilute sulfuric acid:

$\text{Mg(s)} + \text{H}_2\text{SO}_4\text{(aq)} \rightarrow \text{MgSO}_4\text{(aq)} + \text{H}_2\text{(g)}$Mg(s)+H2​SO4​(aq)→MgSO4​(aq)+H2​(g)

$\text{Ca(s)} + \text{H}_2\text{SO}_4\text{(aq)} \rightarrow \text{CaSO}_4\text{(s)} + \text{H}_2\text{(g)}$Ca(s)+H2​SO4​(aq)→CaSO4​(s)+H2​(g)

The Ca + H₂SO₄ reaction slows to a halt because CaSO₄ is insoluble — it coats the metal surface and stops further attack. This is a good link to the solubility trend in the next lesson (Group 2 sulfate solubility decreases down the group, so MgSO₄ stays in solution but CaSO₄ does not).

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Redox summary

Every Group 2 reaction is a redox reaction in which the metal is oxidised (loses electrons) and hydrogen (or oxygen) is reduced:

flowchart LR
  A[M oxidation state 0] -->|loses 2e-| B[M2+ oxidation state +2]
  C[2H+ ox state +1] -->|gains 2e-| D[H2 ox state 0]
  E[O2 ox state 0] -->|gains 4e- per O2| F[2 O2- ox state -2]

Half-equations for the calcium + water reaction:

$\text{Ca(s)} \rightarrow \text{Ca}^{2+}\text{(aq)} + 2\text{e}^- \quad \text{(oxidation)}$Ca(s)→Ca2+(aq)+2e−(oxidation)

$2\text{H}_2\text{O(l)} + 2\text{e}^- \rightarrow \text{H}_2\text{(g)} + 2\text{OH}^-\text{(aq)} \quad \text{(reduction)}$2H2​O(l)+2e−→H2​(g)+2OH−(aq)(reduction)

Sum to give the overall ionic equation: $\text{Ca(s)} + 2\text{H}_2\text{O(l)} \rightarrow \text{Ca}^{2+}\text{(aq)} + 2\text{OH}^-\text{(aq)} + \text{H}_2\text{(g)}$Ca(s)+2H2​O(l)→Ca2+(aq)+2OH−(aq)+H2​(g)

The metal is the reducing agent in every Group 2 reaction; its reducing power increases down the group as IE₁ + IE₂ falls.

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Worked examples

Example 1 — Sr with water

Q: Predict and explain whether strontium will react faster or slower than calcium with water. Write a balanced equation (with state symbols) and state the observations.

Answer: Strontium reacts faster than calcium. Sr is below Ca in Group 2; its outer (5s²) electrons are further from the nucleus and more shielded than Ca's 4s² electrons. The sum IE₁ + IE₂ for Sr (~1613 kJ mol⁻¹) is lower than for Ca (~1735 kJ mol⁻¹), so Sr loses its two outer electrons more readily and is therefore the stronger reducing agent / more reactive.

Equation: $\text{Sr(s)} + 2\text{H}_2\text{O(l)} \rightarrow \text{Sr(OH)}_2\text{(aq)} + \text{H}_2\text{(g)}$Sr(s)+2H2​O(l)→Sr(OH)2​(aq)+H2​(g).

Observations: Sr sinks then rises; vigorous effervescence (faster than with Ca); metal disappears quickly; colourless alkaline solution forms (turns phenolphthalein pink); reaction is noticeably exothermic.

Example 2 — Mg with steam

Q: Write a balanced equation, with state symbols, for the reaction of magnesium with steam, and explain why the product is different from that of magnesium with cold water.

Answer: With steam: $\text{Mg(s)} + \text{H}_2\text{O(g)} \rightarrow \text{MgO(s)} + \text{H}_2\text{(g)}$Mg(s)+H2​O(g)→MgO(s)+H2​(g)

The product is MgO (oxide), not Mg(OH)₂ (hydroxide), because at the elevated temperature of steam the hydroxide is thermally unstable and decomposes: $\text{Mg(OH)}_2\text{(s)} \xrightarrow{\Delta} \text{MgO(s)} + \text{H}_2\text{O(g)}$Mg(OH)2​(s)Δ​MgO(s)+H2​O(g)

The kinetics with steam are dramatically faster than with cold water — at room temperature, Mg(OH)₂ rapidly forms a passivating layer on the metal surface that halts further reaction; at steam temperature, the layer decomposes to MgO (which is less coherent on the metal surface) and the underlying Mg continues to react.

Example 3 — Mg with dilute HCl

Q: A student adds a small piece of magnesium ribbon to 25 cm³ of 1.0 mol dm⁻³ HCl. Describe the observations and write a balanced ionic equation.