Chemical periodicity of other elements - Chemistry (9701) - Cambridge A-Level

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💪 Your A-Level Chemistry Study Guide: Chemical Periodicity of Other Elements

Hello future Chemists! This topic is where we really put your AS knowledge of periodicity to the test. You already know the rules for trends across Period 3, but now we apply those rules to the elements in different groups (vertically) – specifically Group 2 (Alkaline Earth Metals), Group 17 (Halogens), and some tricky elements like Nitrogen and Sulfur.

Don't worry if some of the concepts (like thermal stability or solubility) seem complex—we will break them down using simple analogies and focus on the fundamental energy factors that explain everything! Let’s dive in and see how the Periodic Table works its magic.

1. Group 2: The Alkaline Earth Metals (Mg to Ba)

1.1 General Reactivity and Physical Trends

The elements in Group 2 are Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), and Barium (Ba). They all form ions with a +2 charge ($M^{2+}$) by losing their two outer s-electrons.

Trend Down the Group: Reactivity Increases

Atomic Radius: Increases (due to more electron shells).
First/Second Ionisation Energy (IE): Decreases. The outer electrons are further from the nucleus and are better shielded, making them easier to remove.
Electronegativity: Decreases (less power to attract electrons).

Key Takeaway: Because IE decreases down the group, it takes less energy to form the $M^{2+}$ ion. This means the metals become more reactive as you move from Mg to Ba.

1.2 Reactions of Group 2 Elements (Mg to Ba)

The reactivity trend is clearly seen in their reactions:

1. Reaction with Water:

Magnesium (Mg): Reacts slowly with cold water, but quickly with steam.
$Mg(s) + 2H_2O(l) \rightarrow Mg(OH)_2(s) + H_2(g)$ (Slow, cold water)
$Mg(s) + H_2O(g) \rightarrow MgO(s) + H_2(g)$ (Fast, steam)
Calcium (Ca) onwards (Sr, Ba): React increasingly readily with cold water.
$Ca(s) + 2H_2O(l) \rightarrow Ca(OH)_2(s) + H_2(g)$

2. Reaction with Oxygen:

All burn in oxygen to form the metal oxide ($MO$).
$2Mg(s) + O_2(g) \rightarrow 2MgO(s)$

3. Reaction with Dilute Acids (HCl or H₂SO₄):

All react readily to form the salt and hydrogen gas.
$Mg(s) + 2HCl(aq) \rightarrow MgCl_2(aq) + H_2(g)$
Exception to note: $Ca, Sr,$ and $Ba$ react less readily with sulfuric acid because their sulfates are insoluble and form a protective layer on the metal surface, stopping the reaction.

1.3 Trends in Solubility of Group 2 Compounds

The solubility of Group 2 compounds shows opposite trends for hydroxides and sulfates.

Hydroxides ($M(OH)_2$): Solubility increases down the group.
(e.g., $Mg(OH)_2$ is almost insoluble; $Ba(OH)_2$ is soluble).
Sulfates ($MSO_4$): Solubility decreases down the group.
(e.g., $MgSO_4$ is soluble; $BaSO_4$ is highly insoluble—used in medicine as a "Barium meal").

💡 Memory Aid: For Group 2, the letter S (for Sulfates) means Solubility Sinks (decreases). The Hydroxides do the opposite!

1.4 Explaining Solubility Trends (A-Level Core Concept)

Solubility trends are explained by comparing two crucial energy changes:

1. Lattice Energy ($\Delta H_{latt}$): Energy released when gaseous ions form 1 mole of solid ionic lattice (exothermic).
2. Enthalpy of Hydration ($\Delta H_{hyd}$): Energy released when 1 mole of gaseous ions dissolve in water (exothermic).

The Enthalpy Change of Solution ($\Delta H_{sol}$) is related by an energy cycle (Syllabus 27.1, 27.2):
$$\Delta H_{sol} = \Delta H_{latt} + \sum \Delta H_{hyd} \quad \text{ (where } \Delta H_{latt} \text{ is represented as the energy to break the lattice)}$$

A compound is generally soluble if $\Delta H_{sol}$ is negative (exothermic) or slightly positive (endothermic). The key is how the magnitude of $\Delta H_{latt}$ and $\Delta H_{hyd}$ changes down the group:

Both $\Delta H_{latt}$ and $\Delta H_{hyd}$ become less negative (weaker) down the group because the ionic radius increases.
The smaller the ion, the more drastically its hydration energy changes with radius.

For Hydroxides ($M^{2+}$ and small $OH^-$):

As we move down (e.g., from $Mg^{2+}$ to $Ba^{2+}$), the cation radius increases significantly compared to the small, constant $OH^-$ anion. This means the *change* in the cation's $\Delta H_{hyd}$ (which is high for small ions) is larger than the *change* in $\Delta H_{latt}$. The energy required to break the lattice becomes less dominant, making the total $\Delta H_{sol}$ more favorable (more soluble).

For Sulfates ($M^{2+}$ and large $SO_4^{2-}$):

The $SO_4^{2-}$ anion is very large. As we move down the group, the cation radius increases, becoming more similar to the anion radius. This causes both $\Delta H_{latt}$ and $\Delta H_{hyd}$ to decrease at roughly the same rate. However, the Lattice Energy dominates the trend, becoming too large to overcome, resulting in decreasing solubility.

1.5 Trends in Thermal Stability

This specifically refers to the stability of nitrates ($M(NO_3)_2$) and carbonates ($MCO_3$) when heated.

Trend: Stability Increases Down the Group

When heated, they decompose:

Carbonates: $MCO_3(s) \rightarrow MO(s) + CO_2(g)$
Nitrates: $2M(NO_3)_2(s) \rightarrow 2MO(s) + 4NO_2(g) + O_2(g)$

Explanation (A-Level Core Concept: Polarisation):

Thermal decomposition occurs when the heat causes the large anion ($CO_3^{2-}$ or $NO_3^-$) to distort and break down. This distortion is called polarisation, and it is caused by the small, positive metal ion ($M^{2+}$).

High Polarising Power: Small, highly charged cations (like $Mg^{2+}$) have a high charge density. They effectively 'pull' the electron cloud of the large, floppy anion towards them, distorting and weakening the bonds within the anion.
Low Polarising Power: Large cations (like $Ba^{2+}$) have a lower charge density and cause less distortion.

Since $Mg^{2+}$ has high polarising power, $MgCO_3$ is easily polarised, weak bonds break easily, and it decomposes at a lower temperature (low thermal stability).

Since $Ba^{2+}$ has low polarising power, $BaCO_3$ is harder to break, and it decomposes at a higher temperature (high thermal stability).

Key Takeaway G2: Reactivity increases down the group. Solubility of hydroxides increases, but solubility of sulfates decreases. Thermal stability increases down the group, explained by decreasing polarising power of the cation.

2. Group 17: The Halogens (Cl, Br, I)

Group 17 elements (F, Cl, Br, I) are powerful oxidising agents and form diatomic molecules ($X_2$).

2.1 Physical Properties of Halogens

Colour Trend (Syllabus 11.1):
- Chlorine ($Cl_2$): Pale green gas
- Bromine ($Br_2$): Reddish-brown volatile liquid
- Iodine ($I_2$): Grey/black volatile solid (forms purple vapour)
Volatility/Boiling Point Trend: Increases down the group (Gas $\rightarrow$ Liquid $\rightarrow$ Solid).
Explanation: Down the group, the atoms become larger and have more electrons. This leads to stronger instantaneous dipole-induced dipole forces (London dispersion forces), requiring more energy to separate the molecules.
Bond Strength (Syllabus 11.1): The trend is not smooth. $Cl_2$ has a stronger bond than $F_2$.
Trend: $Cl_2 > Br_2 > F_2 > I_2$
Explanation: In the small $F_2$ molecule, the lone pairs of electrons on the two adjacent F atoms are very close together. This causes strong inter-electron repulsion, weakening the F-F bond significantly.

2.2 Chemical Properties: Oxidising Power

Halogens are oxidising agents (they accept electrons: $X_2 + 2e^- \rightarrow 2X^-$).

Trend: Oxidising Power Decreases Down the Group

This is because as you go down the group, the atoms get larger and the outer shell is further away from the nucleus, meaning the incoming electron is less attracted, making reduction harder.

Displacement Reactions: A more reactive (stronger oxidising) halogen will displace a less reactive halogen from its halide salt in solution.

$Cl_2$ displaces $Br^-$ and $I^-$:
$Cl_2(aq) + 2Br^-(aq) \rightarrow 2Cl^-(aq) + Br_2(aq)$
$Cl_2(aq) + 2I^-(aq) \rightarrow 2Cl^-(aq) + I_2(aq)$
$Br_2$ displaces only $I^-$:
$Br_2(aq) + 2I^-(aq) \rightarrow 2Br^-(aq) + I_2(aq)$

Did you know? You can use an organic solvent like tetrachloromethane to help see the colours of the halogens formed in these solutions.

2.3 Halide Ions ($X^-$) as Reducing Agents

Halide ions ($Cl^-$, $Br^-$, $I^-$) can donate electrons, acting as reducing agents.

Trend: Reducing Power Increases Down the Group

As the ion size increases down the group, the outer electrons are further from the nucleus and less tightly held, making them easier to lose (i.e., making them stronger reducing agents).

Reaction of Halide Ions with Concentrated Sulfuric Acid ($H_2SO_4$)

Concentrated $H_2SO_4$ is a strong acid and a moderate oxidising agent. The reaction depends entirely on the reducing power of the halide ion.

Halide ($X^-$)	Product ($H_2SO_4$ Reduction)	Reducing Power	Reaction Type
$Cl^-$ (Weakest)	No reduction of $H_2SO_4$	Too weak	Only acid-base reaction: $NaCl + H_2SO_4 \rightarrow NaHSO_4 + HCl$
$Br^-$ (Medium)	Sulfur Dioxide ($SO_2$)	Strong enough to reduce $H_2SO_4$ to $SO_2$	Redox reaction: $2HBr + H_2SO_4 \rightarrow Br_2 + SO_2 + 2H_2O$
$I^-$ (Strongest)	$SO_2, S,$ and Hydrogen Sulfide ($H_2S$)	Strong enough to reduce $H_2SO_4$ to $SO_2$, and then further reduce $SO_2$ to $S$ and $H_2S$	Redox reaction: $8HI + H_2SO_4 \rightarrow 4I_2 + H_2S + 4H_2O$

Important Note: This means you can distinguish $Cl^-$, $Br^-$, and $I^-$ by the gases produced when reacted with concentrated $H_2SO_4$.

Reaction of Halide Ions with Aqueous Silver Nitrate ($AgNO_3$)

This is the standard test for halide ions, which form insoluble silver halides. The solubility of these precipitates in aqueous ammonia is key to distinguishing them.

$Ag^+(aq) + Cl^-(aq) \rightarrow AgCl(s)$ (White ppt, soluble in dilute $\text{NH}_3$)
$Ag^+(aq) + Br^-(aq) \rightarrow AgBr(s)$ (Cream ppt, soluble in concentrated $\text{NH}_3$, insoluble in dilute $\text{NH}_3$)
$Ag^+(aq) + I^-(aq) \rightarrow AgI(s)$ (Pale yellow ppt, insoluble in concentrated $\text{NH}_3$)

2.4 The Chemistry of Chlorine (Disproportionation)

A disproportionation reaction is a redox reaction where the same element is simultaneously oxidised and reduced.

Chlorine with Aqueous Sodium Hydroxide (NaOH)

Chlorine reacts differently depending on the temperature of the aqueous NaOH:

1. Cold Aqueous NaOH:

Chlorine is oxidised (to +1 in $ClO^-$) and reduced (to -1 in $Cl^-$).

$$\text{Oxidation States:} \quad \stackrel{0}{Cl}_2 \rightarrow \stackrel{-1}{Cl}^- \quad \text{(Reduction)} \quad \text{and} \quad \stackrel{0}{Cl}_2 \rightarrow \stackrel{+1}{ClO}^- \quad \text{(Oxidation)}$$
$$\text{Equation:} \quad Cl_2(aq) + 2NaOH(aq) \rightarrow NaCl(aq) + NaClO(aq) + H_2O(l)$$

The solution formed contains sodium hypochlorite ($NaClO$), which is the active ingredient in bleach.

2. Hot Aqueous NaOH:

Chlorine disproportionates further (to +5 in $ClO_3^-$) and is reduced (to -1 in $Cl^-$).

$$\text{Equation:} \quad 3Cl_2(aq) + 6NaOH(aq) \rightarrow 5NaCl(aq) + NaClO_3(aq) + 3H_2O(l)$$

Chlorine in Water Purification

Chlorine gas is used to kill bacteria in water supplies (Syllabus 11.4).

When chlorine dissolves in water, it also undergoes a disproportionation reaction:

$$Cl_2(aq) + H_2O(l) \rightleftharpoons HOCl(aq) + HCl(aq)$$

The active species that kill the bacteria are hypochlorous acid (HOCl) and the hypochlorite ion ($ClO^-$). Both are powerful oxidising agents.

Key Takeaway G17: Halogen oxidising power decreases down the group. Halide reducing power increases down the group, visible in the progressive reduction of concentrated $H_2SO_4$ by $Br^-$ and $I^-$.

3. Nitrogen and Sulfur

These two non-metals have unique chemistry, often dominating environmental discussions (Syllabus 12.1).

3.1 Nitrogen: Lack of Reactivity

Nitrogen gas ($N_2$) makes up about 78% of the atmosphere but is remarkably unreactive.

Explanation:

Nitrogen exists as a diatomic molecule held together by a triple covalent bond ($N \equiv N$).
This bond has an extremely high bond energy, requiring huge amounts of energy to break, thus inhibiting most reactions at ambient temperatures.
The $N_2$ molecule is non-polar, meaning it has little attraction for other molecules, contributing to its inertness.

3.2 Ammonia ($NH_3$) and Basicity

Ammonia is a crucial base.

Brønsted-Lowry Definition: A base is a proton ($H^+$) acceptor.
Ammonia acts as a base by accepting a proton from water (or an acid) using the lone pair of electrons on the N atom, forming the ammonium ion ($NH_4^+$).
The bonding in the ammonium ion involves three normal covalent bonds and one coordinate (dative covalent) bond formed when $\text{NH}_3$ donates its lone pair to $H^+$.
$$\text{Basicity:} \quad NH_3(aq) + H_2O(l) \rightleftharpoons NH_4^+(aq) + OH^-(aq)$$

Displacement from Ammonium Salts:

Ammonium salts can react with an acid or base (usually a strong base like NaOH or KOH) to displace ammonia gas (used in the laboratory test for $NH_4^+$).

$$NH_4Cl(s) + NaOH(aq) \xrightarrow{heat} NaCl(aq) + H_2O(l) + NH_3(g)$$

3.3 Oxides of Nitrogen (NOx) and the Environment

Oxides of nitrogen ($NO_x$, mainly $NO$ and $NO_2$) are significant pollutants (Syllabus 12.1.3).

1. Formation and Sources:

In nature, $NO_x$ is formed during lightning. Man-made sources are primarily from internal combustion engines (cars) and power plants. High temperatures and pressures inside engines cause nitrogen and oxygen from the air to react:

$$N_2(g) + O_2(g) \xrightarrow{high \ temperature} 2NO(g)$$

This $NO$ can then react further in the atmosphere:

$$2NO(g) + O_2(g) \rightarrow 2NO_2(g)$$

2. Catalytic Removal (Catalytic Converters):

To reduce pollution, car exhaust systems contain catalytic converters (using precious metals like Palladium, Platinum, and Rhodium). These catalysts convert harmful gases into less harmful ones:

$$2NO(g) + 2CO(g) \xrightarrow{Pt/Rh \ catalyst} N_2(g) + 2CO_2(g)$$

3. Photochemical Smog:

Smog is a cocktail of pollutants, one key component of which is Peroxyacetyl nitrate (PAN).

Atmospheric $NO$ and $NO_2$ react with unburned hydrocarbons (fuel) in the presence of sunlight to form PAN. PAN is an eye irritant and respiratory toxin.

3.4 Sulfur and Acid Rain

Sulfur dioxide ($SO_2$) is the primary contributor to industrial acid rain, formed by burning sulfur-containing fuels (like coal).

The Role of Nitrogen Oxides in Acid Rain:

$NO$ and $NO_2$ play a catalytic role in the oxidation of atmospheric sulfur dioxide ($SO_2$) to sulfur trioxide ($SO_3$), which then forms sulfuric acid ($H_2SO_4$).

Step 1: Nitrogen dioxide oxidises sulfur dioxide:
$$SO_2(g) + NO_2(g) \rightarrow SO_3(g) + NO(g)$$ Step 2: Nitrogen monoxide is regenerated (catalytic step) and reacts with oxygen:
$$2NO(g) + O_2(g) \rightarrow 2NO_2(g)$$
The sulfur trioxide then reacts with rainwater to form acid rain:
$$SO_3(g) + H_2O(l) \rightarrow H_2SO_4(aq)$$

Note: $NO_2$ itself also contributes directly to acid rain by reacting with water to form nitric acid ($HNO_3$).

Key Takeaway N & S: Nitrogen's inertness is due to the strong triple bond. $NO_x$ pollutants are catalytic in forming acid rain and are key precursors for photochemical smog.

📌 Quick Review Checklist: Chemical Periodicity

Can you explain why Group 2 reactivity increases down the group? (IE decreases)
Do you know the solubility trends for Group 2 hydroxides (increases) and sulfates (decreases)?
Can you link thermal stability of carbonates/nitrates to polarisation? (Smaller cation = less stable).
Can you explain the bond strength anomaly in $F_2$? (Lone pair repulsion).
Do you know the products of halide ions ($Cl^-$, $Br^-$, $I^-$) with concentrated $H_2SO_4$? (Only $I^-$ gives $H_2S$).
Can you write the equation for the disproportionation of $Cl_2$ in cold NaOH?
Can you explain the role of $NO_x$ in acid rain formation? (Catalytic oxidation of $SO_2$).

Keep practising those comparisons and explanations! You've got this!

Halide (\(X^-\))	Product (\(H_2SO_4\) Reduction)	Reducing Power	Reaction Type
\(Cl^-\) (Weakest)	No reduction of \(H_2SO_4\)	Too weak	Only acid-base reaction: \(NaCl + H_2SO_4 \rightarrow NaHSO_4 + HCl\)
\(Br^-\) (Medium)	Sulfur Dioxide (\(SO_2\))	Strong enough to reduce \(H_2SO_4\) to \(SO_2\)	Redox reaction: \(2HBr + H_2SO_4 \rightarrow Br_2 + SO_2 + 2H_2O\)
\(I^-\) (Strongest)	\(SO_2, S,\) and Hydrogen Sulfide (\(H_2S\))	Strong enough to reduce \(H_2SO_4\) to \(SO_2\), and then further reduce \(SO_2\) to \(S\) and \(H_2S\)	Redox reaction: \(8HI + H_2SO_4 \rightarrow 4I_2 + H_2S + 4H_2O\)

💪 Your A-Level Chemistry Study Guide: Chemical Periodicity of Other Elements

1. Group 2: The Alkaline Earth Metals (Mg to Ba)

1.1 General Reactivity and Physical Trends

Trend Down the Group: Reactivity Increases

1.2 Reactions of Group 2 Elements (Mg to Ba)

1.3 Trends in Solubility of Group 2 Compounds

1.4 Explaining Solubility Trends (A-Level Core Concept)

1.5 Trends in Thermal Stability

Trend: Stability Increases Down the Group

2. Group 17: The Halogens (Cl, Br, I)

2.1 Physical Properties of Halogens

2.2 Chemical Properties: Oxidising Power

Trend: Oxidising Power Decreases Down the Group

2.3 Halide Ions (\(X^-\)) as Reducing Agents

Trend: Reducing Power Increases Down the Group

Reaction of Halide Ions with Concentrated Sulfuric Acid (\(H_2SO_4\))

Reaction of Halide Ions with Aqueous Silver Nitrate (\(AgNO_3\))

2.4 The Chemistry of Chlorine (Disproportionation)

Chlorine with Aqueous Sodium Hydroxide (NaOH)

Chlorine in Water Purification

3. Nitrogen and Sulfur

3.1 Nitrogen: Lack of Reactivity

3.2 Ammonia (\(NH_3\)) and Basicity

3.3 Oxides of Nitrogen (NOx) and the Environment

1. Formation and Sources:

2. Catalytic Removal (Catalytic Converters):

3. Photochemical Smog:

3.4 Sulfur and Acid Rain

The Role of Nitrogen Oxides in Acid Rain:

📌 Quick Review Checklist: Chemical Periodicity

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Chemistry (9701) | Chemical periodicity of other elements

💪 Your A-Level Chemistry Study Guide: Chemical Periodicity of Other Elements

1. Group 2: The Alkaline Earth Metals (Mg to Ba)

1.1 General Reactivity and Physical Trends

Trend Down the Group: Reactivity Increases

1.2 Reactions of Group 2 Elements (Mg to Ba)

1.3 Trends in Solubility of Group 2 Compounds

1.4 Explaining Solubility Trends (A-Level Core Concept)

1.5 Trends in Thermal Stability

Trend: Stability Increases Down the Group

2. Group 17: The Halogens (Cl, Br, I)

2.1 Physical Properties of Halogens

2.2 Chemical Properties: Oxidising Power

Trend: Oxidising Power Decreases Down the Group

2.3 Halide Ions (\(X^-\)) as Reducing Agents

Trend: Reducing Power Increases Down the Group

Reaction of Halide Ions with Concentrated Sulfuric Acid (\(H_2SO_4\))

Reaction of Halide Ions with Aqueous Silver Nitrate (\(AgNO_3\))

2.4 The Chemistry of Chlorine (Disproportionation)

Chlorine with Aqueous Sodium Hydroxide (NaOH)

Chlorine in Water Purification

3. Nitrogen and Sulfur

3.1 Nitrogen: Lack of Reactivity

3.2 Ammonia (\(NH_3\)) and Basicity

3.3 Oxides of Nitrogen (NOx) and the Environment

1. Formation and Sources:

2. Catalytic Removal (Catalytic Converters):

3. Photochemical Smog:

3.4 Sulfur and Acid Rain

The Role of Nitrogen Oxides in Acid Rain:

📌 Quick Review Checklist: Chemical Periodicity

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