Year 11 Chemistry – Precipitation, Decomposition & Displacement

Vocabulary flashcards covering key terms and definitions from the Year 11 Chemistry booklet on precipitation, decomposition and displacement reactions.

Precipitation Reaction

A reaction in which two aqueous solutions form an insoluble solid (precipitate).

Precipitate

The insoluble solid produced when two solutions are mixed in a precipitation reaction.

Solubility (of ionic compounds)

The ability of an ionic compound to dissolve in water to form aqueous ions.

Soluble Ionic Compound

An ionic substance that dissociates in water to give aqueous ions, e.g., NaCl → Na⁺(aq) + Cl⁻(aq).

Insoluble Ionic Compound

An ionic substance that does not dissolve in water and remains as a solid.

Spectator Ion

An ion present in solution that does not participate in the overall chemical change.

Ionic Equation

A chemical equation that shows only the ions and molecules directly involved in the reaction, omitting spectator ions.

Thermal Decomposition Reaction

A reaction where a compound breaks down into simpler substances upon heating.

Metal Carbonate Decomposition

Most metal carbonates decompose on heating to give a metal oxide and CO₂: MCO₃ → MO + CO₂.

Metal Hydrogen Carbonate Decomposition

All metal hydrogen carbonates decompose on heating to give a metal carbonate, CO₂, and H₂O.

Metal Hydroxide Decomposition

Most metal hydroxides decompose on heating to form a metal oxide and water; Group I hydroxides are exceptions.

Activity Series of Metals

A ranked list of metals ordered by their tendency to lose electrons (reactivity), e.g., Ca > Mg > Al > Zn > Fe > Pb > (H) > Cu > Ag.

Displacement Reaction (Metals)

A reaction where a more reactive metal displaces a less reactive metal from its compound in solution.

Reactivity (Metals)

The ease with which a metal loses electrons; higher reactivity means it more readily forms positive ions.

Test for Sulfate Ions

Add Ba²⁺(aq); a white BaSO₄(s) precipitate confirms SO₄²⁻. Ionic eqn: Ba²⁺ + SO₄²⁻ → BaSO₄(s).

Test for Chloride Ions

Add Ag⁺(aq); a white AgCl(s) precipitate confirms Cl⁻. Ionic eqn: Ag⁺ + Cl⁻ → AgCl(s).

Test for Iodide Ions

Add Ag⁺(aq); a yellow AgI(s) precipitate confirms I⁻. Ionic eqn: Ag⁺ + I⁻ → AgI(s).

Example Precipitation (MgCO₃)

Mg²⁺(aq) + CO₃²⁻(aq) → MgCO₃(s); MgCO₃ forms as a white solid when mixing MgCl₂ and Na₂CO₃ solutions.

Reaction of Metal with Water

Reactive metals (e.g., Ca) react with H₂O to form a metal hydroxide and H₂ gas: Ca + 2 H₂O → Ca(OH)₂ + H₂.

Reaction of Metal with Acid

Metals above hydrogen in the activity series react with acids to form a salt and H₂ gas, e.g., Zn + 2 HCl → ZnCl₂ + H₂.

Spectator Ion Example

In MgCl₂ + Na₂CO₃, Na⁺ and Cl⁻ remain unchanged; they are spectator ions.

Exception – Alkali Metal Carbonates

Na₂CO₃ and K₂CO₃ do not thermally decompose under normal laboratory heating.

Exception – Group I Hydroxides

Hydroxides of Na⁺ and K⁺ are stable to heat and do not decompose easily.

Colour Change in Cu²⁺ Displacement

Cu²⁺(blue) → Cu(s)(pink) when Zn displaces Cu²⁺; solution turns colourless as Zn²⁺ forms.

Definition: Solution

A homogeneous mixture of a solute (e.g., an ionic solid) dissolved in a solvent (commonly water).

Balanced Ionic Equation (BaSO₄ ppt)

Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s); illustrates removal of spectator ions to show actual change.