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Group 2 of the periodic table contains the alkaline earth metals: beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), and barium (Ba). At A-Level, the focus is on magnesium to barium, since beryllium behaves anomalously due to its very small atomic radius and high charge density. These elements share a common electron configuration ending in s², meaning each atom has two electrons in its outermost shell. This accounts for the +2 oxidation state that dominates their chemistry.
As you descend Group 2, atomic radius increases. Each successive element has an additional electron shell, so the outermost electrons are further from the nucleus.
| Element | Atomic Radius / pm | Electron Configuration | 1st IE / kJ mol⁻¹ | 2nd IE / kJ mol⁻¹ |
|---|---|---|---|---|
| Mg | 160 | 1s² 2s² 2p⁶ 3s² | 738 | 1451 |
| Ca | 197 | [Ar] 4s² | 590 | 1145 |
| Sr | 215 | [Kr] 5s² | 550 | 1064 |
| Ba | 222 | [Xe] 6s² | 503 | 965 |
The increase in radius has important consequences for reactivity, ionisation energy, and the properties of compounds.
First ionisation energy decreases down Group 2. Although nuclear charge increases (more protons), the effect is outweighed by increased shielding from inner electron shells and greater distance of the outer electrons from the nucleus. The outermost electrons are easier to remove in barium than in magnesium.
The second ionisation energy also decreases down the group for the same reasons. Since Group 2 elements form M²⁺ ions by losing both s² electrons, the decreasing ionisation energies explain the increasing reactivity down the group.
The jump from the 2nd to the 3rd ionisation energy is enormous for every Group 2 element (for Mg, the 3rd IE is 7733 kJ mol⁻¹ compared to 1451 kJ mol⁻¹ for the 2nd). This is because the third electron would be removed from an inner shell, which is much closer to the nucleus and far more strongly held. This is why Group 2 elements never form M³⁺ ions.
Melting points do not follow a simple trend in Group 2. The overall tendency is for melting point to decrease as the metallic radius increases and the bond strength per unit volume decreases, but changes in crystal packing cause irregularities.
| Element | Melting Point / °C | Crystal Structure |
|---|---|---|
| Be | 1287 | HCP |
| Mg | 650 | HCP |
| Ca | 842 | FCC |
| Sr | 777 | FCC |
| Ba | 727 | BCC |
The irregularity at calcium (higher than magnesium) arises because calcium adopts a face-centred cubic structure rather than hexagonal close packed, which affects the efficiency of metallic bonding.
The Group 2 metals react with water to form a metal hydroxide and hydrogen gas:
M(s) + 2H₂O(l) → M(OH)₂(aq) + H₂(g)
Reactivity increases down the group because ionisation energies decrease, making it easier for atoms to lose their two outer electrons.
Common exam mistake: Students sometimes write Mg(OH)₂ as the product when magnesium reacts with steam. With steam, the product is MgO (not the hydroxide), because the high temperature decomposes any hydroxide formed.
When Group 2 metal compounds are heated in a flame, electrons are promoted to higher energy levels and then fall back, emitting light of characteristic wavelengths.
| Element | Flame Colour | Approximate Wavelength / nm |
|---|---|---|
| Mg | No distinctive colour (white/bright) | — |
| Ca | Brick red / orange-red | 622 (orange-red) |
| Sr | Crimson / deep red | 675 (deep red) |
| Ba | Pale green / apple green | 524 (green) |
These flame colours are used as a diagnostic test to identify Group 2 metal ions. It is important to remember that the colours arise from electron transitions, not from the ion itself being coloured. The energy gap between the excited state and the ground state determines the wavelength of light emitted.
flowchart TD
A["Unknown Group 2 compound"] --> B{"Flame test colour?"}
B -->|"Brick red / orange-red"| C["Ca²⁺"]
B -->|"Crimson / deep red"| D["Sr²⁺"]
B -->|"Pale green / apple green"| E["Ba²⁺"]
B -->|"No distinctive colour"| F["Mg²⁺ (use other tests)"]
Two important solubility trends must be learned for Group 2 compounds. These trends are based on the balance between lattice enthalpy and hydration enthalpy of the ions, and both decrease down the group, but at different rates.
| Compound | Solubility at 25 °C / mol dm⁻³ | Description |
|---|---|---|
| Mg(OH)₂ | 2.0 × 10⁻⁴ | Almost insoluble; forms a suspension |
| Ca(OH)₂ | 1.5 × 10⁻² | Slightly soluble; limewater is a saturated solution |
| Sr(OH)₂ | 3.4 × 10⁻² | Moderately soluble |
| Ba(OH)₂ | 1.5 × 10⁻¹ | Most soluble; strongly alkaline |
This trend means that as you go down Group 2, the hydroxides dissolve more readily and produce more alkaline solutions. Ba(OH)₂ solution is strongly alkaline (pH ≈ 13–14 depending on concentration).
| Compound | Solubility at 25 °C / mol dm⁻³ | Description |
|---|---|---|
| MgSO₄ | 2.1 | Very soluble (Epsom salts) |
| CaSO₄ | 4.7 × 10⁻² | Slightly soluble |
| SrSO₄ | 5.3 × 10⁻⁴ | Sparingly soluble |
| BaSO₄ | 9.4 × 10⁻⁶ | Insoluble |
The insolubility of BaSO₄ is exploited in the test for sulfate ions: adding BaCl₂ solution to a sample acidified with HCl produces a white precipitate of BaSO₄ if sulfate ions are present.
For both hydroxides and sulfates, lattice enthalpy decreases down the group (as the cation gets bigger, the ions are further apart and the lattice is less strongly held). Hydration enthalpy also decreases (less energy is released when larger ions are hydrated). The key is that for hydroxides, the lattice enthalpy decreases faster than hydration enthalpy, so dissolving becomes more favourable. For sulfates, the hydration enthalpy decreases faster than lattice enthalpy, so dissolving becomes less favourable.
flowchart LR
subgraph "Hydroxide Solubility"
A1["Mg(OH)₂ – Insoluble"] --> B1["Ca(OH)₂ – Slight"] --> C1["Sr(OH)₂ – Moderate"] --> D1["Ba(OH)₂ – Soluble"]
end
subgraph "Sulfate Solubility"
A2["MgSO₄ – Soluble"] --> B2["CaSO₄ – Slight"] --> C2["SrSO₄ – Sparingly"] --> D2["BaSO₄ – Insoluble"]
end
A student dissolves a white solid in water. The resulting solution is strongly alkaline (pH 13). They add dilute H₂SO₄ to a portion and observe a white precipitate.
Analysis:
Answer: The compound is barium hydroxide, Ba(OH)₂.
If the solution had been only weakly alkaline (pH ≈ 10–11) and no precipitate formed with dilute H₂SO₄, the compound could be Ca(OH)₂, since CaSO₄ is slightly soluble and may not produce a clear precipitate.
Group 2 shows clear periodic trends driven by increasing atomic radius and decreasing ionisation energy. Reactivity with water increases down the group, hydroxide solubility increases, and sulfate solubility decreases. Flame tests provide a quick way to identify individual Group 2 elements. Always include actual data values (ionisation energies, solubilities, decomposition temperatures) in your exam answers to demonstrate thorough knowledge.
Edexcel 9CH0 specification Topic 4 — Inorganic Chemistry and the Periodic Table, sub-topic 4.1 covers the reactions of Group 2 elements (Mg, Ca, Sr, Ba) with oxygen, water and dilute acids, the trend in reactivity descending the group, and the link between reactivity, atomic radius and first/second ionisation energies (refer to the official specification document for exact wording). This material is examined principally in Paper 1 (9CH0/01: Advanced Inorganic and Physical Chemistry) and reappears synoptically in Paper 3 (9CH0/03: General and Practical Principles) where qualitative analysis pulls in CP3 observations. Group 2 also intersects Topic 13 (Energetics II) for lattice energies of MO and M(OH)₂, and Topic 1 (Atomic Structure) for the periodic justification of trends.
Question (8 marks):
(a) Calcium reacts vigorously with cold water; magnesium reacts only very slowly with cold water but rapidly with steam. Explain this difference, referring to ionisation energies and the products formed. (5)
(b) Write balanced equations for the reaction of (i) calcium with cold water and (ii) magnesium with steam. State the role of the metal in each. (3)
Solution with mark scheme:
(a) Step 1 — identify the trend.
Reactivity with water increases down Group 2 because the first and second ionisation energies decrease as atomic radius increases and the outer 4s electrons (Ca) are further from the nucleus and more shielded than the 3s electrons (Mg).
M1 — explicit statement that ionisation energies decrease down the group, supported by atomic radius and shielding.
A1 — linkage of lower ionisation energy to easier loss of two electrons, hence faster reaction with water.
Step 2 — explain the cold-water observation.
Calcium loses its two outer electrons more readily, so the reaction Ca + 2H₂O → Ca(OH)₂ + H₂ proceeds at room temperature. Magnesium's higher ionisation energies mean the activation energy for electron loss to water is too high; with steam, the higher temperature supplies the activation energy, and the product is the oxide MgO + H₂ rather than the hydroxide.
M1 — different products noted (Ca(OH)₂ vs MgO).
A1 — connection of higher temperature (steam) to overcoming activation energy.
Step 3 — link redox identity.
The metals are oxidised (M → M²⁺ + 2e⁻); water is reduced (2H₂O + 2e⁻ → H₂ + 2OH⁻ or → H₂ + O²⁻ in steam). Increased ease of oxidation down the group is the redox restatement of the ionisation-energy trend.
A1 — explicit redox framing.
(b) (i) Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g). B1
(ii) Mg(s) + H₂O(g) → MgO(s) + H₂(g). B1
In each, the metal is the reducing agent (it is oxidised; it reduces water). B1 — role correctly identified for both.
Total: 8 marks (M2 A3 B3).
Question (6 marks): A sample of strontium is added to a beaker of cold water.
(a) State two observations and write the balanced equation. (3)
(b) Predict, with reasoning, whether the reaction would be faster or slower than the analogous reaction of barium under the same conditions. (3)
Mark scheme decomposition by AO:
| Mark | AO | Awarded for |
|---|---|---|
| (a) B1 | AO1 | Effervescence / fizzing observed (H₂ evolved) |
| (a) B1 | AO1 | Metal sinks/dissolves; solution becomes alkaline / cloudy with Sr(OH)₂ (limited solubility relative to Ba(OH)₂) |
| (a) B1 | AO2 | Sr(s) + 2H₂O(l) → Sr(OH)₂(aq) + H₂(g) — balanced, state symbols |
| (b) M1 | AO2 | Reaction with Ba would be faster |
| (b) M1 | AO2 | Ba has lower 1st + 2nd ionisation energies (larger atomic radius, more shielding) |
| (b) A1 | AO3 | Therefore lower activation energy for M → M²⁺ + 2e⁻; oxidation easier, faster rate |
Total: 6 marks split AO1 = 2, AO2 = 3, AO3 = 1. Group 2 reactivity questions on Paper 1 lean heavily on AO2 (linking observations to explanation) — pure recall is rarely worth more than a third of the marks.
| AO | Typical share on Group 2 questions | Earned by |
|---|---|---|
| AO1 (knowledge) | 30–40% | Stating products of M + H₂O / O₂ / HCl, naming the trend in reactivity |
| AO2 (application) | 40–50% | Linking ionisation energy / atomic radius to rate observations; predicting unfamiliar reactions |
| AO3 (analysis) | 15–25% | Multi-step reasoning across redox + ionisation + thermodynamics |
Examiner-rewarded phrasing: "1st and 2nd ionisation energies decrease down the group because…"; "the outermost s-electrons are further from the nucleus and more shielded"; "the metal is oxidised (M → M²⁺ + 2e⁻); water is reduced". Phrases that lose marks: bare "more reactive" without mechanism; "loses electrons more easily" without invoking shielding or radius; confusing first ionisation energy alone with the two-electron process needed for M²⁺.
A Paper 1 pattern to watch: a question stating "explain in terms of ionisation energy" forbids alternative explanations (electronegativity, electron affinity); answers that drift into other quantities lose marks even when factually correct.
Question: Explain why barium reacts more vigorously with water than calcium.
Grade C response (~120 words):
Barium is below calcium in Group 2. Going down the group, the atoms get bigger, so the outer electrons are further from the nucleus and easier to lose. Barium therefore loses its two outer electrons more easily than calcium, so it reacts faster with water producing Ba(OH)₂ and H₂.
Examiner commentary: 2/3. The candidate identifies size and ease of electron loss but does not name ionisation energy explicitly or mention shielding — both expected at A-Level. They lose the AO2 mark for not connecting ease of oxidation to reaction rate via activation energy.
Grade A response (~150 words):*
Descending Group 2 from Ca to Ba, atomic radius increases and inner-shell shielding of the outer 6s electrons (Ba) is greater than for the 4s electrons (Ca). Both the first and second ionisation energies therefore fall (Ca: 590 + 1145 = 1735 kJ mol⁻¹; Ba: 503 + 965 = 1468 kJ mol⁻¹). Since the rate-determining step in M + H₂O is the loss of two electrons (M → M²⁺ + 2e⁻), the lower combined ionisation energy of Ba lowers the activation energy and the reaction proceeds faster. The redox restatement: Ba is the stronger reducing agent (E°(Ba²⁺/Ba) = −2.91 V vs −2.87 V for Ca).
Examiner commentary: 3/3. Quantitative ionisation energies, explicit two-electron loss, redox potentials — the answer reads as a chemist's full justification rather than a recall sentence.
Question: Compare the reactions of magnesium and calcium with (i) cold water and (ii) dilute hydrochloric acid, giving products and observations. Explain the difference using ionisation energies.
Grade A response (~280 words):*
(i) Cold water: Mg reacts very slowly, producing a few bubbles of H₂ over hours and a dilute Mg(OH)₂ solution that turns universal indicator pale blue (Mg(OH)₂ has low solubility, ~1.4 × 10⁻⁴ mol dm⁻³). Ca reacts steadily, generating visible effervescence within seconds and a cloudy suspension of Ca(OH)₂ (limewater, ~0.02 mol dm⁻³) that is more strongly alkaline (pH ≈ 12). Equations: Mg(s) + 2H₂O(l) → Mg(OH)₂(aq) + H₂(g); Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g).
(ii) Dilute HCl: Both metals react vigorously, but Ca is faster. Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g); Ca(s) + 2HCl(aq) → CaCl₂(aq) + H₂(g). The solutions remain clear because both chlorides are soluble. The difference in rate is less dramatic than with water because the high [H⁺] makes both reactions thermodynamically very favourable.
Explanation: Ca's first and second ionisation energies (590 + 1145 = 1735 kJ mol⁻¹) are lower than Mg's (738 + 1451 = 2189 kJ mol⁻¹) because Ca's outer electrons are in the 4s subshell, further from the nucleus and shielded by [Ne]3s²3p⁶, whereas Mg's outer 3s electrons experience shielding only from [Ne]. Lower ionisation energy means lower activation energy for M → M²⁺ + 2e⁻, hence faster oxidation and faster reaction.
Examiner commentary: 6/6. Quantitative IEs, both reactions covered with correct equations, observations linked to product solubility, and the underpinning electronic-structure argument made explicit.
Question: Discuss how the reactions of Group 2 elements with oxygen, water and dilute hydrochloric acid illustrate the trend in reducing power down the group. Refer to ionisation energies, electrode potentials, and at least two specific element comparisons.
Grade A response (~360 words):*
Group 2 elements are oxidised in all three reactions: with O₂ to MO, with H₂O to M(OH)₂ (or MO + H₂ for Mg with steam), and with HCl to MCl₂ + H₂. In each, the metal loses two electrons (M → M²⁺ + 2e⁻), so any trend in reaction rate or vigour reflects the trend in ease of electron loss — the reducing power.
Ionisation energies. Combined first and second ionisation energies decrease down the group: Mg 2189, Ca 1735, Sr 1614, Ba 1468 kJ mol⁻¹. The cause is increasing atomic radius (the outer s-electrons occupy progressively higher principal quantum shells) and increasing shielding by filled inner shells. The effective nuclear charge experienced by the outer electrons is approximately constant at +2, but the larger r in the Coulombic term reduces attraction.
Electrode potentials. E°(M²⁺/M) becomes progressively more negative: Mg²⁺/Mg = −2.37 V, Ca²⁺/Ca = −2.87 V, Sr²⁺/Sr = −2.89 V, Ba²⁺/Ba = −2.91 V. More negative E° means the metal is a stronger reducing agent — exactly the redox restatement of the ionisation-energy trend.
Comparison 1 (Mg vs Ca with water): Mg + cold water is barely perceptible; Ca + cold water produces steady effervescence. Mg requires steam (Mg + H₂O(g) → MgO + H₂) because activation energy is too high at 25 °C. The IE difference (2189 − 1735 = 454 kJ mol⁻¹) maps directly to a higher activation energy for Mg.
Comparison 2 (Ca vs Ba with O₂): Ca burns with a brick-red flame to give white CaO; Ba burns with a pale-green flame to give BaO with traces of BaO₂ (the peroxide is stable for Ba because the larger Ba²⁺ stabilises the larger O₂²⁻ — a lattice-energy argument). The fact that Ba can form a peroxide while Mg and Ca do not is a direct cation-size effect.
The unified picture: descending Group 2, decreasing IE → more negative E° → faster, more vigorous redox reactions across all three reagents.
Examiner commentary: 9/9. Quantitative throughout, two distinct comparisons, redox + thermodynamic framing, and synthesis to a unified mechanism.
Oxbridge interview prompt: "Strontium is a Group 2 metal yet ⁹⁰Sr is one of the most dangerous fission products. Why does its chemical similarity to calcium make it biologically hazardous, and what does this say about the limitations of group-trend reasoning in real-world systems?"
This lesson maps directly to Core Practical 3 (CP3): Investigating the rates of reaction of Group 2 elements with water and acids. CP3 asks candidates to time the production of hydrogen (gas-syringe or inverted measuring cylinder over water) when measured masses of Mg ribbon, Ca turnings and Sr (where available) react with cold water or dilute acid. The data illustrate the order Mg ≪ Ca < Sr < Ba in initial rate, and the resulting graphs (volume vs time) feed into kinetics questions on Paper 1 + 3. Quantitative outcomes — e.g. moles of H₂ evolved compared with stoichiometric expectation — also link to CP16 (Qualitative analysis of inorganic ions) when the alkaline product is tested with universal indicator or BaCl₂ to confirm Ba²⁺. CP3 reinforces the Group 2 narrative: descending the group, ionisation energy falls, activation energy falls, rate rises. Students should record temperature, volume of acid/water and surface area of the metal — uncontrolled variables that examiners frequently target in evaluative AO3 questions.
This content is aligned with the Pearson Edexcel GCE A Level Chemistry (9CH0) specification, Paper 1 — Advanced Inorganic and Physical Chemistry, Topic 4: Inorganic Chemistry and the Periodic Table (sub-topic 4.1, Group 2). For the most accurate and up-to-date information, please refer to the official Pearson Edexcel specification document.
graph TD
A["Group 2 metal<br/>Mg, Ca, Sr, Ba"] --> B{"Reagent?"}
B -->|"O2"| C["MO solid<br/>2M + O2 to 2MO"]
B -->|"H2O cold"| D["M(OH)2 + H2<br/>slow for Mg<br/>fast for Ca/Sr/Ba"]
B -->|"H2O steam"| E["MO + H2<br/>(only Mg in practice)"]
B -->|"dilute HCl"| F["MCl2(aq) + H2<br/>vigorous for all"]
C --> G["Down group:<br/>radius increases<br/>shielding increases"]
D --> G
E --> G
F --> G
G --> H["IE1 + IE2 decrease<br/>E°(M2+/M) more negative"]
H --> I["Reducing power increases<br/>rate increases"]
style D fill:#27ae60,color:#fff
style I fill:#3498db,color:#fff