Chemistry (9701) Study Notes: Group 2 Metals (Mg to Ba)
A Friendly Introduction to Group 2: The Alkaline Earth Metals
Hello Chemist! This chapter is all about Group 2, the **Alkaline Earth Metals** (Magnesium, Calcium, Strontium, and Barium). These elements are like siblings in the Periodic Table family—they share many similarities, but their properties change in predictable ways as you move down the group.
Understanding these trends is crucial for predicting chemical behavior and mastering inorganic chemistry. Don’t worry if the energetic explanations seem tricky at first; we will break them down step-by-step!
1. Physical Trends of Group 2 Elements
Group 2 metals all have a characteristic outer electron configuration of **\([Noble \ Gas] \ ns^2\)**. Since they easily lose these two valence electrons, they always form ions with a **\(+2\) charge** (\(M^{2+}\)).
1.1 Key Physical Trends Down the Group (Mg to Ba)
As we descend Group 2, the elements have more electron shells. This fundamental change drives all the key trends:
- Atomic Radius and Ionic Radius (\(M^{2+}\)):
Trend: Increases.
Explanation: More electron shells are added, making the atom/ion physically larger. - First and Second Ionisation Energies (IE):
Trend: Decreases.
Explanation: The outer electrons are further from the positively charged nucleus, and there is increased shielding (repulsion) from the inner electron shells. Less energy is required to remove these electrons. - Reactivity:
Trend: Increases.
Explanation: Because the ionization energy decreases, it becomes easier for the metals to lose their two electrons and form stable \(M^{2+}\) ions.
Quick Review: Moving down Group 2 means bigger atoms, easier electron loss, and higher reactivity.
2. Chemical Reactions of Group 2 Metals (Syllabus 10.1.1)
2.1 Reaction with Oxygen (\(O_2\))
All Group 2 metals react with oxygen to form a metal oxide (\(MO\)).
Magnesium needs to be heated strongly (it burns with a dazzling white flame), while Barium reacts much more vigorously.
General Equation:
\(2M(s) + O_2(g) \to 2MO(s)\)
2.2 Reaction with Water (\(H_2O\))
The reactivity with water increases down the group.
- Magnesium (Mg): Reacts very slowly with cold water, but rapidly with steam to form magnesium oxide and hydrogen gas.
\(Mg(s) + 2H_2O(l) \to Mg(OH)_2(s) + H_2(g)\) (Slow, cold water)
\(Mg(s) + H_2O(g) \to MgO(s) + H_2(g)\) (Fast, steam) - Calcium (Ca), Strontium (Sr), Barium (Ba): React readily with cold water to form the metal hydroxide and hydrogen gas. The reaction becomes faster and more vigorous down the group.
General Equation: \(M(s) + 2H_2O(l) \to M(OH)_2(aq) + H_2(g)\)
2.3 Reaction with Dilute Acids (HCl and H₂SO₄)
The metals react with dilute acids to produce a salt and hydrogen gas.
General Equation (e.g., with HCl):
\(M(s) + 2HCl(aq) \to MCl_2(aq) + H_2(g)\)
⚠️ A Common Mistake to Avoid: Sulfuric Acid Reaction (Passivation)
When Mg reacts with dilute sulfuric acid (\(H_2SO_4\)), the reaction is fast because magnesium sulfate (\(MgSO_4\)) is soluble.
However, when Calcium (Ca), Strontium (Sr), or Barium (Ba) react with dilute \(H_2SO_4\), the reaction slows down or stops quickly.
Why? The sulfates of Ca, Sr, and Ba are insoluble. An insoluble layer of sulfate forms on the metal surface, protecting it from further reaction. This effect is called passivation.
3. Reactions of Group 2 Compounds (Syllabus 10.1.2)
3.1 Reactions with Water
The metal oxides (\(MO\)) and hydroxides (\(M(OH)_2\)) are alkaline.
- Oxides with Water: The oxides react to form hydroxides.
\(MO(s) + H_2O(l) \to M(OH)_2(aq)\)
The alkalinity increases down the group because the solubility of the hydroxide increases (See Section 5). \(Mg(OH)_2\) is sparingly soluble, making the resulting solution only weakly alkaline, while \(Ba(OH)_2\) is much more soluble, making the solution strongly alkaline.
3.2 Reactions with Acids (Neutralisation)
Group 2 oxides, hydroxides, and carbonates are basic and thus react with dilute acids (HCl or \(H_2SO_4\)) in classic neutralisation reactions.
- Hydroxides + Acid: (Salt and Water)
\(M(OH)_2(s) + 2HCl(aq) \to MCl_2(aq) + 2H_2O(l)\) - Carbonates + Acid: (Salt, Water, and Carbon Dioxide)
\(MCO_3(s) + H_2SO_4(aq) \to MSO_4(s/aq) + H_2O(l) + CO_2(g)\)
Did you know? Calcium hydroxide, Ca(OH)₂ (slaked lime), is often used in agriculture to neutralise acidic soil, showing its strong basic properties!
4. Thermal Stability of Carbonates and Nitrates (Syllabus 10.1.3 & 27.1.1)
4.1 The Trend in Thermal Stability
Thermal stability refers to how easily a compound breaks down when heated. For Group 2 carbonates and nitrates, the stability increases as you move down the group (from Mg to Ba).
- Magnesium compounds decompose easily at low temperatures.
- Barium compounds require very high temperatures to decompose.
4.2 Decomposition Equations
- Carbonates: They decompose into the metal oxide and carbon dioxide.
\(MCO_3(s) \to MO(s) + CO_2(g)\) - Nitrates: They decompose into the metal oxide, brown nitrogen dioxide gas, and oxygen gas. (This produces a lot of gas and is a common test in the lab!)
\(2M(NO_3)_2(s) \to 2MO(s) + 4NO_2(g) + O_2(g)\)
4.3 🧠 A2 Explanation: The Role of Polarization
This trend is explained by considering the size of the cation and its effect on the large anion (carbonate, \(CO_3^{2-}\), or nitrate, \(NO_3^-\)).
1. Cation Size: Down the group, the \(M^{2+}\) cation radius increases (\(Mg^{2+}\) is tiny; \(Ba^{2+}\) is large).
2. Polarizing Power: Small, highly charged cations have a high polarizing power. They can distort (polarise) the electron cloud of a large anion (like \(CO_3^{2-}\) or \(NO_3^-\)) towards themselves.
3. Weakening the Anion: When the anion is polarised, the electron density within the anion is shifted, which weakens the bonds within the anion (specifically the C-O bonds in carbonate or N-O bonds in nitrate).
4. Stability:
- Magnesium (\(Mg^{2+}\)): Small radius, high polarizing power. It severely distorts the carbonate/nitrate ion, weakening its structure. This means it has low thermal stability and decomposes easily.
- Barium (\(Ba^{2+}\)): Large radius, low polarizing power. It causes minimal distortion. This means it has high thermal stability and requires much more energy to decompose.
Key Takeaway for Stability: Smaller cation \(\to\) more polarisation \(\to\) weaker anion bonds \(\to\) lower thermal stability.
5. Solubility Trends of Group 2 Compounds (Syllabus 10.1.5 & 27.1.2)
The solubility trends for Group 2 hydroxides and sulfates are opposite, and you must know both!
5.1 Hydroxides (M(OH)₂): Solubility Increases Down the Group
- \(Mg(OH)_2\) is nearly insoluble.
- \(Ca(OH)_2\) (limewater) is sparingly soluble.
- \(Ba(OH)_2\) is readily soluble.
Memory Aid: Hydroxides (OH) are Happy at the bottom (soluble).
5.2 Sulfates (MSO₄): Solubility Decreases Down the Group
- \(MgSO_4\) (Epsom salts) is highly soluble.
- \(CaSO_4\) is slightly soluble.
- \(BaSO_4\) is highly insoluble (used in barium meals for X-rays).
Memory Aid: The Ba-rge of Sulfates Sinks (Barium Sulfate is insoluble).
5.3 🧠 A2 Explanation: Hydration and Lattice Energy
To explain solubility, we look at the energy changes involved in dissolving a solid: the Enthalpy Change of Solution (\(\Delta H_{\text{sol}}\)).
(Note: \(\Delta H_{\text{latt}}\) is the energy required to break the lattice (endothermic, positive value); \(\Delta H_{\text{hyd}}\) is the energy released when ions are surrounded by water (exothermic, negative value).)
For a compound to dissolve easily (high solubility), \(\Delta H_{\text{sol}}\) needs to be small (or ideally negative), meaning the energy released by hydration must overcome the energy required to break the lattice.
Key Principle: How size affects \(\Delta H_{\text{latt}}\) and \(\Delta H_{\text{hyd}}\)
As you move down Group 2, the \(M^{2+}\) ion gets larger.
Both the Lattice Energy (\(\Delta H_{\text{latt}}\)) and the Hydration Enthalpy (\(\Delta H_{\text{hyd}}\)) become less exothermic (less negative) because the charge density decreases and the electrostatic attraction between the ions/water decreases.
- For small anions (like \(OH^{-}\)):
The difference in size between the small \(Mg^{2+}\) and the large \(Ba^{2+}\) is significant relative to the small \(OH^{-}\) anion. The hydration energy decreases more rapidly down the group than the lattice energy. This difference makes dissolution favorable for the larger cations (Ba) \(\to\) Solubility Increases. - For large anions (like \(SO_4^{2-}\)):
The anion itself is already very large. As the cation size increases down the group, the overall change in lattice energy dominates and decreases dramatically. The hydration energy change is relatively similar across the group. The decrease in lattice energy dominates, making it harder to break the lattice than hydrate the ions \(\to\) Solubility Decreases.
Struggling Student Tip: Think of it like this:
The small \(Mg^{2+}\) ion fits nicely with the small \(OH^-\) ion (strong lattice) AND it can be surrounded easily by water (strong hydration). When the cation gets bigger (Ba), its ability to be hydrated drops off quickly, but its ability to pack tightly in a lattice doesn't drop off as fast (relative to the small OH⁻). This favors solubility for Ba(OH)₂.
But with the huge \(SO_4^{2-}\), when the cation grows, the ability to pack in the lattice drops drastically, making \(BaSO_4\) the most stable (hardest to break up) and therefore least soluble.
6. Summary of Key Group 2 Trends
6.1 Trends Down the Group (Mg \(\to\) Ba)
- Reactivity: Increases
- Thermal Stability (Carbonates/Nitrates): Increases
- Hydroxide Solubility: Increases
- Sulfate Solubility: Decreases
- Alkalinity of Hydroxides/Oxides: Increases
Try to link the solubility trend to a real-world application: Barium sulfate is so insoluble it can be safely ingested for medical imaging (barium meal) without dissolving in the body and causing barium poisoning.