Overview
This lecture discusses the properties, reactions, and trends of group 2 elements (alkaline earth metals), focusing on their chemical behavior, solubility patterns, thermal stability, and practical uses. The material is designed for the Cambridge International A Level Chemistry curriculum and highlights how group 2 topics may overlap with other areas, such as topic 9.
Electronic Structure & Ion Formation
- Group 2 elements have electron configurations ending in s², meaning their outermost subshell contains two electrons.
- When reacting, group 2 metals lose these two outer electrons to form +2 ions (e.g., Mg → Mg²⁺ + 2e⁻).
- All group 2 metals form +2 ions because they lose both outer electrons.
- The atomic radius increases as you move down the group because each successive element has an additional electron shell, making the atom larger.
- The increase in atomic radius is due to the addition of extra shells as you go down the group.
Trends Down Group 2
- The first ionization energy decreases down the group. This is because:
- There are more electron shells, increasing shielding.
- The outer electrons are further from the nucleus, so the attraction between the nucleus and these electrons is weaker.
- It becomes easier to remove an electron as you go down the group.
- Reactivity increases down the group because the outer electrons are more easily lost due to increased distance from the nucleus and greater shielding.
- Although the number of protons (nuclear charge) increases down the group, the effect of increased shielding and atomic radius outweighs this, so it does not prevent the decrease in ionization energy.
- The shielding effect is strong enough to override the increased nuclear charge as you go down the group.
Reactions with Water & Oxygen
- Group 2 metals react with water to form metal hydroxides and hydrogen gas. The reaction becomes more vigorous down the group.
- Beryllium does not react with water; it is effectively inert in this context.
- Magnesium reacts only very slowly with cold water, but reacts more readily with steam to form magnesium oxide instead of magnesium hydroxide.
- Example: Strontium reacts with water to form strontium hydroxide and hydrogen gas, with visible fizzing.
- As you go down the group, the atom gets larger, and the outer electron is further from the nucleus, making it easier to remove and increasing reactivity with water.
- Group 2 metals also react with oxygen to form white solid metal oxides.
- Example: Magnesium burns with a bright white flame to form magnesium oxide (MgO).
- These reactions involve redox changes: the metal is oxidized from 0 to +2, and oxygen is reduced from 0 to -2.
- Magnesium oxide and other group 2 oxides are white solids.
Group 2 Oxides, Hydroxides & Solubility
- Group 2 oxides react with water to form alkaline hydroxide solutions.
- Example: Strontium oxide reacts with water to form strontium hydroxide, which dissociates to give Sr²⁺ and OH⁻ ions.
- The OH⁻ ion is responsible for the alkaline property.
- The alkalinity of these solutions increases down the group because the hydroxides become more soluble.
- Magnesium oxide reacts very slowly with water, and magnesium hydroxide is only slightly soluble.
- Strontium hydroxide is much more soluble, producing a more strongly alkaline solution.
- As you go down the group, the solubility of group 2 hydroxides increases, while the solubility of group 2 carbonates decreases.
- The increased solubility of hydroxides means more OH⁻ ions are produced, making the solution more alkaline.
Neutralization Reactions
- Group 2 oxides and hydroxides act as bases and neutralize acids to form salts and water.
- Example: Calcium oxide reacts with hydrochloric acid to form calcium chloride and water (CaO + 2HCl → CaCl₂ + H₂O).
- Calcium hydroxide reacts similarly: Ca(OH)₂ + 2HCl → CaCl₂ + 2H₂O.
- These are classic acid-base reactions, where the oxide or hydroxide neutralizes the acid.
- Both oxides and hydroxides of group 2 elements can be used to neutralize acids, producing a salt and water as products.
Thermal Decomposition
- Group 2 carbonates decompose upon heating (thermal decomposition) to form metal oxides and carbon dioxide gas.
- Example: CaCO₃ (heated) → CaO + CO₂.
- Group 2 nitrates decompose with heat to form metal oxides, nitrogen dioxide (NO₂), and oxygen (O₂).
- Example: Ca(NO₃)₂ (heated) → CaO + 2NO₂ + ½O₂.
- The thermal stability of both carbonates and nitrates increases down the group. This means it becomes harder to decompose them by heating as you go from magnesium to barium.
- The reason for this trend is that larger group 2 ions have lower charge density, causing less distortion of the electron cloud in the carbonate or nitrate ion, making the compound more stable.
- Smaller ions like Mg²⁺ have higher charge density and distort the electron cloud more, making the compound less stable and easier to decompose.
- Larger ions like Ba²⁺ have lower charge density, causing less distortion and greater stability.
Solubility Trends
- The solubility of group 2 hydroxides increases down the group, so solutions become more alkaline.
- The solubility of group 2 carbonates decreases down the group; they become less soluble.
- Compounds with doubly-charged anions (like CO₃²⁻) become less soluble as you go down the group, due to a decrease in the enthalpy of hydration of the metal ion.
- For singly-charged anions (like OH⁻), solubility increases down the group because the lattice dissociation enthalpy decreases, which outweighs the enthalpy change of hydration.
- In summary:
- Group 2 hydroxides: solubility increases down the group.
- Group 2 carbonates: solubility decreases down the group.
Uses of Group 2 Compounds
- Calcium hydroxide (slaked lime) is used to neutralize acidic soils, helping farmers optimize soil pH for crop growth.
- It is sprayed onto soil to adjust pH and improve crop yield.
- Calcium carbonate (limestone) is widely used in construction, for building materials and making cement.
- Limestone is used in stone buildings and as a raw material for cement.
- Group 2 compounds are important in various everyday applications due to their chemical properties, especially their ability to neutralize acids and form useful building materials.
Key Terms & Definitions
- Ionization Energy: The energy required to remove an electron from an atom.
- Shielding: The reduction in attraction between the nucleus and outer electrons caused by inner electron shells.
- Thermal Decomposition: The process of breaking down a compound by heating.
- Solubility: The amount of a substance that dissolves in a solvent.
- Charge Density: The ratio of an ion’s charge to its size, influencing how strongly it interacts with other ions or molecules.
- Redox Reaction: A chemical reaction involving the transfer of electrons, where one species is oxidized and another is reduced.
- Alkalinity: The property of a solution to neutralize acids, often due to the presence of OH⁻ ions.
Action Items / Next Steps
- Review and practice writing balanced equations for group 2 reactions with water, oxygen, acids, and during thermal decomposition.
- Memorize the trends in reactivity, solubility, and thermal stability down group 2.
- Practice questions on the uses and reactions of group 2 elements and their compounds.
- Understand the reasons behind the observed trends, especially the effects of atomic size, charge density, and shielding.
- Study diagrams showing electron cloud distortion and charge density effects for deeper understanding of stability trends.