CIE AS/A Level Chemistry 10.1 Similarities and trends in the properties of the Group 2 metals, magnesium to barium, and their compounds Study Notes- 2025-2027 Syllabus
CIE AS/A Level Chemistry 10.1 Similarities and trends in the properties of the Group 2 metals, magnesium to barium, and their compounds Study Notes – New Syllabus
CIE AS/A Level Chemistry 10.1 Similarities and trends in the properties of the Group 2 metals, magnesium to barium, and their compounds Study Notes at IITian Academy focus on specific topic and type of questions asked in actual exam. Study Notes focus on AS/A Level Chemistry latest syllabus with Candidates should be able to:
describe reactions with oxygen, water and dilute HCl and H₂SO₄
describe reactions of oxides, hydroxides and carbonates
describe thermal decomposition of nitrates and carbonates and trends in stability
predict trends in physical and chemical properties
state trends in solubility of hydroxides and sulfates
Reactions of Elements with Oxygen, Water and Dilute Acids
Metals react with oxygen, water and dilute acids in ways that reflect their position in the reactivity series. More reactive metals react more readily and vigorously, while less reactive metals show limited or no reaction.
Higher reactivity → faster and more vigorous reactions
1. Reactions with Oxygen
Metals react with oxygen to form metal oxides. The reaction becomes less vigorous as reactivity decreases.
- Very reactive metals react rapidly.
- Less reactive metals require heating.
- Some metals form protective oxide layers.
Examples:
\( \mathrm{4Na + O_2 \rightarrow 2Na_2O} \)
\( \mathrm{2Mg + O_2 \rightarrow 2MgO} \)
\( \mathrm{4Al + 3O_2 \rightarrow 2Al_2O_3} \)
2. Reactions with Water
The reaction of metals with water depends strongly on their reactivity.
- Very reactive metals react with cold water.
- Moderately reactive metals react with steam.
- Unreactive metals do not react.
Examples:
Cold water (sodium):
\( \mathrm{2Na + 2H_2O \rightarrow 2NaOH + H_2} \)
Steam (magnesium):
\( \mathrm{Mg + H_2O \rightarrow MgO + H_2} \)
3. Reactions with Dilute Hydrochloric Acid
Reactive metals react with dilute hydrochloric acid to form metal chlorides and hydrogen gas.
General equation:
\( \mathrm{Metal + 2HCl \rightarrow Metal\ chloride + H_2} \)
Examples:
\( \mathrm{Mg + 2HCl \rightarrow MgCl_2 + H_2} \)
\( \mathrm{Zn + 2HCl \rightarrow ZnCl_2 + H_2} \)
4. Reactions with Dilute Sulfuric Acid
Reactive metals also react with dilute sulfuric acid to form metal sulfates and hydrogen gas.
General equation:
\( \mathrm{Metal + H_2SO_4 \rightarrow Metal\ sulfate + H_2} \)
Examples:
\( \mathrm{Mg + H_2SO_4 \rightarrow MgSO_4 + H_2} \)
\( \mathrm{Fe + H_2SO_4 \rightarrow FeSO_4 + H_2} \)
Important Notes
- Copper and metals below hydrogen do not react with dilute acids.
- Aluminium may not appear to react due to an oxide coating.
- Hydrogen gas is produced in all metal–acid reactions.
Example
Write the equation for the reaction of magnesium with dilute hydrochloric acid.
▶️ Answer / Explanation
\( \mathrm{Mg + 2HCl \rightarrow MgCl_2 + H_2} \)
Example
Explain why magnesium reacts with steam but only very slowly with cold water.
▶️ Answer / Explanation
Magnesium is moderately reactive.
Cold water does not provide enough energy to overcome the activation energy, but steam does.
Example
Describe how the reactions of sodium, magnesium and aluminium differ with oxygen, water and dilute acids.
▶️ Answer / Explanation
Sodium reacts vigorously with oxygen and cold water, forming sodium oxide and sodium hydroxide.
Magnesium reacts readily with oxygen, slowly with cold water, and more rapidly with steam and acids.
Aluminium reacts with oxygen to form a protective oxide layer and reacts with acids only when this layer is removed.
These differences are due to decreasing reactivity and increasing oxide stability.
Reactions of Oxides, Hydroxides and Carbonates with Water and Dilute Acids
Metal oxides, hydroxides and carbonates show characteristic reactions with water and dilute acids. These reactions depend on whether the compounds are basic, amphoteric or acidic in nature.
Basic compounds neutralise acids → salt + water (and carbon dioxide for carbonates)
1. Reactions of Oxides
(a) With Water
- Basic metal oxides react with water to form hydroxides.
- Some oxides do not react due to strong bonding.
\( \mathrm{Na_2O + H_2O \rightarrow 2NaOH} \)
\( \mathrm{MgO + H_2O \rightarrow Mg(OH)_2} \)
\( \mathrm{Al_2O_3} \): no reaction with water
(b) With Dilute Acids
- Basic oxides neutralise acids.
With hydrochloric acid:
\( \mathrm{Na_2O + 2HCl \rightarrow 2NaCl + H_2O} \)
With sulfuric acid:
\( \mathrm{MgO + H_2SO_4 \rightarrow MgSO_4 + H_2O} \)
2. Reactions of Hydroxides
(a) With Water
- Group 1 hydroxides dissolve readily in water.
- Group 2 hydroxides are less soluble.
\( \mathrm{NaOH(s) \rightarrow Na^+(aq) + OH^-(aq)} \)
(b) With Dilute Acids
- Hydroxides neutralise acids.
With hydrochloric acid:
\( \mathrm{NaOH + HCl \rightarrow NaCl + H_2O} \)
With sulfuric acid:
\( \mathrm{Mg(OH)_2 + H_2SO_4 \rightarrow MgSO_4 + 2H_2O} \)
3. Reactions of Carbonates
(a) With Water
- Most metal carbonates are insoluble in water.
- Group 1 carbonates are soluble.
Example:
\( \mathrm{Na_2CO_3(s) \rightarrow 2Na^+(aq) + CO_3^{2-}(aq)} \)
(b) With Dilute Acids
- Carbonates react with acids to produce carbon dioxide.
- Effervescence is observed.
With hydrochloric acid:
\( \mathrm{Na_2CO_3 + 2HCl \rightarrow 2NaCl + H_2O + CO_2} \)
With sulfuric acid:
\( \mathrm{MgCO_3 + H_2SO_4 \rightarrow MgSO_4 + H_2O + CO_2} \)
Example
Write the equation for the reaction of sodium hydroxide with dilute hydrochloric acid.
▶️ Answer / Explanation
\( \mathrm{NaOH + HCl \rightarrow NaCl + H_2O} \)
Example
Describe the reaction between magnesium oxide and dilute sulfuric acid.
▶️ Answer / Explanation
Magnesium oxide neutralises sulfuric acid to form magnesium sulfate and water.
\( \mathrm{MgO + H_2SO_4 \rightarrow MgSO_4 + H_2O} \)
Example
Describe how you could distinguish between a metal oxide and a metal carbonate using dilute hydrochloric acid.
▶️ Answer / Explanation
Both will neutralise the acid, but a carbonate will produce carbon dioxide, causing effervescence.
The oxide will form only salt and water with no gas produced.
Thermal Decomposition of Nitrates and Carbonates
When heated, many nitrates and carbonates decompose. The temperature required for decomposition varies systematically down a group. This variation in thermal stability can be explained by the size and charge density of the metal ion.
Lower thermal stability → decomposes easily on heating
Higher thermal stability → requires stronger heating
1. Thermal Decomposition of Carbonates
Most metal carbonates decompose on heating to form a metal oxide and carbon dioxide.
General equation:
\( \mathrm{MCO_3(s) \rightarrow MO(s) + CO_2(g)} \)
Examples:
\( \mathrm{MgCO_3 \rightarrow MgO + CO_2} \)
\( \mathrm{CaCO_3 \rightarrow CaO + CO_2} \)
Exceptions:
- Group 1 carbonates (except lithium carbonate) are thermally stable and do not decompose easily.
- \( \mathrm{Li_2CO_3} \) decomposes on heating.
\( \mathrm{Li_2CO_3 \rightarrow Li_2O + CO_2} \)
2. Thermal Decomposition of Nitrates
Nitrates decompose in different ways depending on the metal.
(a) Group 1 Nitrates
- All Group 1 nitrates (except lithium) form nitrites and oxygen.
\( \mathrm{2NaNO_3 \rightarrow 2NaNO_2 + O_2} \)
Lithium nitrate:
\( \mathrm{4LiNO_3 \rightarrow 2Li_2O + 4NO_2 + O_2} \)
(b) Group 2 and Transition Metal Nitrates
- Decompose to the metal oxide, nitrogen dioxide and oxygen.
\( \mathrm{2Mg(NO_3)_2 \rightarrow 2MgO + 4NO_2 + O_2} \)
\( \mathrm{2Cu(NO_3)_2 \rightarrow 2CuO + 4NO_2 + O_2} \)
Trend in Thermal Stability
Down a group, both carbonates and nitrates become more thermally stable.
- Metal ions become larger.
- Charge density decreases.
- Less polarisation of the carbonate or nitrate ion.
Explanation:
Small, highly charged metal ions strongly polarise the carbonate or nitrate ion, weakening the \( \mathrm{C{-}O} \) or \( \mathrm{N{-}O} \) bonds and making decomposition easier. Larger ions polarise less, increasing thermal stability.
Example
Write the equation for the thermal decomposition of magnesium carbonate.
▶️ Answer / Explanation
\( \mathrm{MgCO_3 \rightarrow MgO + CO_2} \)
Example
Describe the products formed when copper(II) nitrate is heated.
▶️ Answer / Explanation
Copper(II) nitrate decomposes to copper(II) oxide, nitrogen dioxide and oxygen.
\( \mathrm{2Cu(NO_3)_2 \rightarrow 2CuO + 4NO_2 + O_2} \)
Example
Explain why barium carbonate is more thermally stable than magnesium carbonate.
▶️ Answer / Explanation
Barium ions are larger than magnesium ions and have a lower charge density.
This results in less polarisation of the carbonate ion.
The carbonate ion is therefore more stable and decomposes at a higher temperature.
Trends in Physical and Chemical Properties and Making Predictions
The physical and chemical behaviour of elements and their compounds follows predictable trends. By understanding these trends, it is possible to describe observed reactions and to predict the behaviour of unfamiliar elements or compounds.
Trend → explanation → prediction
Elements Involved in (Reactions with Oxygen, Water and Acids) 
The elements involved are mainly metals. Their behaviour depends on their position in the reactivity series.
- More reactive metals react more vigorously with oxygen, water and dilute acids.
- Less reactive metals react slowly or not at all.
- Formation of hydrogen gas indicates metal–acid reactions.
Prediction:
If a metal is placed higher in the reactivity series, it will react faster with water and acids and form more stable oxides.
Compounds in (Oxides, Hydroxides and Carbonates + Acids/Water)
These compounds are mainly basic or amphoteric.
- Metal oxides and hydroxides neutralise acids.
- Carbonates produce carbon dioxide when reacting with acids.
- Solubility decreases down Group 2.
Prediction:
A carbonate of a reactive metal will effervesce with dilute acid, producing \( \mathrm{CO_2} \).
Compounds in (Thermal Decomposition of Carbonates and Nitrates)
Thermal stability depends on the size and charge density of the metal ion.
- Thermal stability increases down a group.
- Small, highly charged ions destabilise anions.
- Carbonates and nitrates decompose more easily for smaller ions.
Prediction:
A nitrate of a metal higher in Group 2 will require stronger heating to decompose.
Compounds in (Acid–Base Behaviour of Oxides, Hydroxides and Carbonates)
Acid–base behaviour changes with bonding type.
- Ionic oxides and hydroxides are basic.
- Aluminium compounds are amphoteric.
- Carbonates act as bases.
Prediction:
An oxide formed by a metal with high charge density is likely to show amphoteric behaviour.
Linking Physical Properties to Chemical Behaviour
- High melting point → ionic or giant lattice → strong bonds.
- Low melting point → simple molecular → weak intermolecular forces.
- High solubility → strong hydration of ions.
Key Idea: Physical properties often give direct clues about chemical reactivity.
Example
A metal reacts rapidly with cold water. Predict its position relative to magnesium.
▶️ Answer / Explanation
The metal is more reactive than magnesium and is higher in the reactivity series.
Example
A carbonate requires strong heating to decompose. Predict the size of the metal ion and explain.
▶️ Answer / Explanation
The metal ion is large.
Large ions have low charge density and polarise the carbonate ion less, increasing thermal stability.
Example
An unknown oxide reacts with both dilute hydrochloric acid and aqueous sodium hydroxide. Predict its nature and the type of metal involved.
▶️ Answer / Explanation
The oxide is amphoteric.
This suggests it is formed by a small, highly charged metal ion such as aluminium.
Variation in the Solubilities of Hydroxides and Sulfates
The solubilities of hydroxides and sulfates show clear and predictable trends down Group 2 of the Periodic Table. These trends are important for explaining reactions in aqueous solution.
Down Group 2 → regular change in solubility
1. Solubility of Group 2 Hydroxides
The solubility of Group 2 hydroxides increases down the group.
- \( \mathrm{Mg(OH)_2} \) is only very slightly soluble.
- \( \mathrm{Ca(OH)_2} \) is sparingly soluble.
- \( \mathrm{Ba(OH)_2} \) is much more soluble.
Overall trend:
Solubility of hydroxides increases down Group 2
2. Solubility of Group 2 Sulfates
The solubility of Group 2 sulfates decreases down the group.
- \( \mathrm{MgSO_4} \) is soluble in water.
- \( \mathrm{CaSO_4} \) is sparingly soluble.
- \( \mathrm{BaSO_4} \) is very insoluble.
Overall trend:
Solubility of sulfates decreases down Group 2
Key Points to Remember
- Hydroxides become more soluble down the group.
- Sulfates become less soluble down the group.
- These are empirical trends you are expected to state at A-level.
Example
Which is more soluble in water: \( \mathrm{Mg(OH)_2} \) or \( \mathrm{Ba(OH)_2} \)?
▶️ Answer / Explanation
\( \mathrm{Ba(OH)_2} \) is more soluble.
Solubility of Group 2 hydroxides increases down the group.
Example
State and explain the trend in solubility of Group 2 sulfates.
▶️ Answer / Explanation
The solubility of Group 2 sulfates decreases down the group.
This is observed from soluble \( \mathrm{MgSO_4} \) to insoluble \( \mathrm{BaSO_4} \).
Example
A white solid forms when aqueous \( \mathrm{Ba^{2+}} \) ions are mixed with aqueous sulfate ions. Use solubility trends to explain this observation.
▶️ Answer / Explanation
Solubility of Group 2 sulfates decreases down the group.
\( \mathrm{BaSO_4} \) is very insoluble, so it precipitates as a white solid.