Chemistry (9701) | Group 2

👋 Welcome to Group 2: The Alkaline Earth Metals!

Hey there, future chemist! This chapter dives deep into Group 2 of the Periodic Table, commonly known as the Alkaline Earth Metals. You’ve already met these elements (like Magnesium in your Biology classes!), but now we’re going to explore the fascinating chemical trends that make them so predictable and useful. Understanding Group 2 is essential because it sets the stage for mastering all the periodic trends in inorganic chemistry!

We will cover everything from how they react with water to why some of their compounds refuse to dissolve. Don't worry if the physical chemistry explanations get a little complex; we'll break down the concepts of lattice energy and hydration step-by-step!

---

Section 1: The Group 2 Elements – Fundamentals and Trends

1.1 Electronic Structure and General Properties

Group 2 elements include Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), and Barium (Ba). (Radium is sometimes included but rarely assessed due to its radioactivity).

Key Feature:

• Every element in Group 2 has two electrons in its outermost (valence) shell.
• To achieve a stable, full outer shell (like a Noble Gas), they lose these two electrons easily, forming ions with a +2 charge (e.g., \({\text{Mg}}^{2+}\), \({\text{Ca}}^{2+}\)).
• They are highly reactive metals, strong reducing agents (they cause reduction by losing electrons themselves), and form ionic compounds (since they are metals reacting with non-metals).

1.2 Trends Down the Group (Mg to Ba)

As you move down Group 2, the number of electron shells increases. This increase governs all the major trends:

Atomic and Ionic Radius

Trend: Atomic and ionic radii increase down the group.

• Explanation: More electron shells are added going from Mg down to Ba, making the atoms and ions physically larger.

First Ionisation Energy (IE)

Trend: First Ionisation Energy (IE) decreases down the group.

• Reminder: IE is the energy needed to remove the first electron from a gaseous atom.
• Explanation: As the atom gets larger, the outermost electron is further away from the nucleus (less nuclear attraction). Crucially, there is also more shielding from the increased number of inner electron shells. This makes the outer electron easier to remove.

---

Section 2: Reactions of Group 2 Metals

Since reactivity increases down the group, you will see Magnesium reacting reluctantly, while Barium reacts vigorously.

2.1 Reaction with Oxygen (O₂)

Group 2 metals burn in oxygen to form white ionic oxides (\(\text{MO}\)).

\[2\text{M}(\text{s}) + \text{O}_{2}(\text{g}) \rightarrow 2\text{MO}(\text{s})\]

Example: Magnesium burns brilliantly in air (a characteristic bright white light).

\[2\text{Mg}(\text{s}) + \text{O}_{2}(\text{g}) \rightarrow 2\text{MgO}(\text{s})\]

2.2 Reaction with Water (H₂O)

Group 2 metals react with water to form the metal hydroxide and hydrogen gas. The vigour of the reaction increases down the group.

A. Magnesium (Mg): Reacts slowly with cold water, but reacts rapidly with steam.

• Cold Water (Slowly): \(\text{Mg}(\text{s}) + 2\text{H}_{2}\text{O}(\text{l}) \rightarrow \text{Mg(OH)}_{2}(\text{s}) + \text{H}_{2}(\text{g})\)
• Steam (Fast): \(\text{Mg}(\text{s}) + \text{H}_{2}\text{O}(\text{g}) \rightarrow \text{MgO}(\text{s}) + \text{H}_{2}(\text{g})\)

B. Calcium (Ca) onwards (Ca, Sr, Ba): React increasingly vigorously with cold water.

\[\text{M}(\text{s}) + 2\text{H}_{2}\text{O}(\text{l}) \rightarrow \text{M(OH)}_{2}(\text{aq}) + \text{H}_{2}(\text{g})\]

Example: Calcium reacts readily, producing bubbles of \({\text{H}}_2\).

2.3 Reaction with Dilute Acids (HCl and H₂SO₄)

Group 2 metals react easily with dilute acids to form a salt and hydrogen gas.

\[\text{M}(\text{s}) + 2\text{HCl}(\text{aq}) \rightarrow \text{MCl}_{2}(\text{aq}) + \text{H}_{2}(\text{g})\]

However, when reacting with dilute sulfuric acid, things get interesting due to solubility trends (which we cover fully in Section 3):

• Mg and Ca: React easily, forming soluble sulfates.
• Ba and Sr: The reaction stops almost immediately. Why? The products, \(\text{BaSO}_{4}\) and \(\text{SrSO}_{4}\), are highly insoluble. They form a protective, impenetrable white layer (a coating of precipitate) on the surface of the metal, preventing further acid attack. This is called passivation.

---

Section 3: Properties and Solubility of Group 2 Compounds

The compounds of Group 2 metals (Oxides, Hydroxides, Carbonates) are generally basic because the metal ions are large and have low charge density compared to Group 13 elements like Aluminium.

3.1 Reactions of Oxides, Hydroxides, and Carbonates

All these compounds react with acids (like \(\text{HCl}\) and \(\text{H}_{2}\text{SO}_{4}\)) in standard acid-base neutralisation reactions.

1. With Water:

• Oxides (\(\text{MO}\)): React to form hydroxides (\(\text{M}(\text{OH})_{2}\)). These reactions are strongly exothermic.

\[\text{MgO}(\text{s}) + \text{H}_{2}\text{O}(\text{l}) \rightarrow \text{Mg(OH)}_{2}(\text{s})\]

(Note: \(\text{Mg(OH)}_{2}\) is only sparingly soluble, which we cover next).

• Hydroxides (\(\text{M}(\text{OH})_{2}\)): Dissolve partially in water. The resulting solution is alkaline due to the production of \(\text{OH}^{-}\) ions.

2. With Dilute Acids:

• Oxides & Hydroxides (Neutralisation): Form salt and water.

\[\text{M}(\text{OH})_{2}(\text{s}) + 2\text{HCl}(\text{aq}) \rightarrow \text{MCl}_{2}(\text{aq}) + 2\text{H}_{2}\text{O}(\text{l})\]

• Carbonates (Acid-Base): Form salt, water, and carbon dioxide.

\[\text{MCO}_{3}(\text{s}) + \text{H}_{2}\text{SO}_{4}(\text{aq}) \rightarrow \text{MSO}_{4}(\text{s/aq}) + \text{H}_{2}\text{O}(\text{l}) + \text{CO}_{2}(\text{g})\]

3.2 Solubility Trends (The Crucial AS Level Facts)

The solubility of Group 2 compounds varies significantly down the group, and you must know the trends for Hydroxides and Sulfates (AS Level requirement 10.1.5).

Trend 1: Group 2 Hydroxides (\(\text{M}(\text{OH})_{2}\))

Solubility Trend: Solubility increases down the group (Mg to Ba).

• \(\text{Mg}(\text{OH})_{2}\) is virtually insoluble.
• \(\text{Ba}(\text{OH})_{2}\) is readily soluble.

Real-World Use: \(\text{Mg}(\text{OH})_{2}\) (Milk of Magnesia) is used as an antacid because its low solubility means it slowly neutralises excess stomach acid without causing a sudden, dangerous rise in pH.

Trend 2: Group 2 Sulfates (\(\text{MSO}_{4}\))

Solubility Trend: Solubility decreases down the group (Mg to Ba).

• \(\text{MgSO}_{4}\) (Epsom salts) is soluble.
• \(\text{BaSO}_{4}\) is highly insoluble.

Real-World Use: Barium sulfate (\(\text{BaSO}_{4}\)) is ingested by patients before X-rays (a 'Barium Meal'). Although Barium ions are toxic, \(\text{BaSO}_{4}\) is safe because its extreme insolubility prevents the toxic \(\text{Ba}^{2+}\) ions from entering the bloodstream.

---

Section 4: Thermal Stability of Nitrates and Carbonates (AS & A Level)

Thermal stability refers to how easily a compound breaks down (decomposes) when heated. For Group 2 compounds, thermal stability increases down the group.

4.1 Thermal Decomposition Reactions (AS Level)

A. Carbonates (\(\text{MCO}_{3}\))

Group 2 carbonates decompose upon heating to form the metal oxide and carbon dioxide gas. The temperature required for decomposition is higher down the group.

\[\text{MCO}_{3}(\text{s}) \rightarrow \text{MO}(\text{s}) + \text{CO}_{2}(\text{g})\]

• \(\text{MgCO}_{3}\) decomposes easily.
• \(\text{BaCO}_{3}\) requires very high temperatures.

B. Nitrates (\(\text{M}(\text{NO}_{3})_{2}\))

Group 2 nitrates decompose upon heating to form the metal oxide, nitrogen dioxide gas, and oxygen gas.

\[2\text{M}(\text{NO}_{3})_{2}(\text{s}) \rightarrow 2\text{MO}(\text{s}) + 4\text{NO}_{2}(\text{g}) + \text{O}_{2}(\text{g})\]

Note: The production of brown fumes of \(\text{NO}_{2}\) is a characteristic observation.

4.2 Explaining the Trend in Thermal Stability (A Level Focus)

This trend is explained by a concept called polarisation, specifically involving the ionic radius of the cation and the polarizability of the large anion (\(\text{CO}_{3}^{2-}\) or \(\text{NO}_{3}^{-}\)).

Step 1: The Role of the Cation (\(\text{M}^{2+}\))

The ability of a cation to distort the electron cloud of an anion is called polarising power.

• Small Cations (like \(\text{Mg}^{2+}\)): Have a high charge density (large charge concentrated in a small volume). They have strong polarising power.
• Large Cations (like \(\text{Ba}^{2+}\)): Have a low charge density. They have weak polarising power.

Step 2: Polarisation and Decomposition

When a small cation (like \(\text{Mg}^{2+}\)) approaches a large anion (like \(\text{CO}_{3}^{2-}\)), it pulls the anion's electron cloud towards itself, distorting (polarising) the anion.

• This distortion weakens the bonds within the anion (e.g., the C-O bonds in the carbonate ion).
• When the bonds in the anion are weakened, less energy (heat) is needed to break the molecule apart.

Conclusion for Thermal Stability

Going down Group 2, the metal ions get larger. The larger ion has a weaker polarising power, resulting in less distortion of the carbonate or nitrate anion. This means the bonds within the anion are stronger, and more energy is needed to decompose the compound.

Therefore, thermal stability increases down the group.

---

Section 5: Explaining Solubility Trends (A Level Energetics Focus)

The AS syllabus requires you to state the solubility trends. The A Level syllabus requires you to explain these trends using thermodynamics, specifically comparing two types of energy changes: Lattice Energy (\(\Delta H_{latt}\)) and Enthalpy of Hydration (\(\Delta H_{hyd}\)).

5.1 Definitions and the Energy Cycle

For a salt (\(\text{MX}_{2}\)) to dissolve, we consider the Enthalpy Change of Solution (\(\Delta H_{sol}^{\ominus}\)).

\[\text{MX}_{2}(\text{s}) + \text{water} \rightarrow \text{M}^{2+}(\text{aq}) + 2\text{X}^{-}(\text{aq})\]

According to Hess’s Law, \(\Delta H_{sol}^{\ominus}\) is determined by two opposing steps:

1. Breaking the lattice: Input energy needed to turn the solid into gaseous ions (related to Lattice Energy). \(\Delta H_{latt}^{\ominus}\) is always exothermic (negative, by syllabus definition) when forming the lattice from ions, so breaking the lattice is endothermic (positive).
2. Hydrating the ions: Energy released when the gaseous ions are surrounded by water molecules (Enthalpy of Hydration, \(\Delta H_{hyd}^{\ominus}\)). This is always exothermic (negative).

The relationship is: \[\Delta H_{sol}^{\ominus} = -\Delta H_{latt}^{\ominus} + \Delta H_{hyd}^{\ominus}(\text{M}^{2+}) + 2\Delta H_{hyd}^{\ominus}(\text{X}^{-})\] (Note: We use \(-\Delta H_{latt}^{\ominus}\) here because we are breaking the lattice, which requires energy input, while the standard definition of \(\Delta H_{latt}^{\ominus}\) is formation of the lattice, which is exothermic).

For a salt to be soluble, \(\Delta H_{sol}^{\ominus}\) should be small (ideally negative or only slightly positive).

• If \(\Delta H_{hyd}^{\ominus}\) is significantly more exothermic than the energy required to break the lattice, the salt dissolves easily.

5.2 How Ion Size Affects \(\Delta H_{latt}\) and \(\Delta H_{hyd}\)

Both Lattice Energy and Hydration Energy depend on ionic charge and ionic radius.

In Group 2, the charge (+2) is constant, so we look only at the radius (which increases down the group).

A. Effect on Lattice Energy (\(\Delta H_{latt}\)):

• Lattice energy is proportional to \(\frac{\text{charge}^2}{\text{radius of cation} + \text{radius of anion}}\).
• Since the cation radius (\(\text{M}^{2+}\)) increases down the group, the overall separation distance increases, making the electrostatic attraction weaker.
• Conclusion: Lattice Energy becomes less exothermic (less negative) down the group.

B. Effect on Hydration Energy (\(\Delta H_{hyd}\)):

• Hydration energy is proportional to \(\frac{\text{charge}}{\text{radius}}\) (charge density).
• Since the cation radius (\(\text{M}^{2+}\)) increases down the group, the charge density decreases. Weaker attraction to water molecules.
• Conclusion: Hydration Energy becomes less exothermic (less negative) down the group.

5.3 Explaining the Solubility Trends Qualitatively

For Group 2, both \(\Delta H_{latt}\) and \(\Delta H_{hyd}\) become less exothermic down the group. The key to solubility is the difference in how much they change.

Case 1: Hydroxides (\(\text{OH}^{-}\)) – Solubility Increases Down Group

• The hydroxide ion (\(\text{OH}^{-}\)) is relatively small.
• When dissolving, the overall lattice size is dominated by the size of the cation (\(\text{M}^{2+}\)).
• As you move down the group, the Lattice Energy (\(\Delta H_{latt}\)) decreases faster than the Hydration Energy (\(\Delta H_{hyd}\)).
• Since less energy is needed to break the lattice (the opposing step to \(\Delta H_{hyd}\)), the overall \(\Delta H_{sol}^{\ominus}\) becomes less positive (or more negative), promoting solubility.

In short: For small anions, the drop in Lattice Energy dominates the trend.

Case 2: Sulfates (\(\text{SO}_{4}^{2-}\)) – Solubility Decreases Down Group

• The sulfate ion (\(\text{SO}_{4}^{2-}\)) is relatively large.
• When dissolving, the lattice size is dominated by the size of the anion (\(\text{SO}_{4}^{2-}\)), which is constant.
• As you move down the group, the changing factor is the Hydration Energy (\(\Delta H_{hyd}\)) of the cation. Because the cation size increases, its charge density drops rapidly, making its Hydration Energy less exothermic faster than the Lattice Energy changes.
• Since less energy is released upon hydration, the overall \(\Delta H_{sol}^{\ominus}\) becomes more positive, hindering solubility.

In short: For large anions, the drop in Hydration Energy dominates the trend.