Coaching Centre
Reactive chemistry review
Summary of Module 3 Reactive Chemistry: reaction types and net ionic equations, the activity series and ionisation energy, redox and half-equations, galvanic cells and E°, and rates with collision theory, followed by a mixed HSC-style test.
This module keeps asking three questions about any pair of substances: will they react, what will they make, and how fast? Use this page to check you can answer all three, then try the mixed test. Questions in the HSC often join these ideas, for example a displacement reaction that has to be predicted with E°, written as half-equations and then speeded up.
How the module fits together
| Lesson | Big idea | Must be able to |
|---|---|---|
| 3A | Predicting products | classify synthesis, decomposition and combustion; balance with states; explain First Nations detoxification (crushing, leaching, fermenting, heating) |
| 3B | Reactions in solution | use solubility rules; write molecular, full ionic and net ionic equations; acid–base and acid–carbonate reactions |
| 3C | Metal reactivity | describe reactions with water, steam, acid, O₂ and metal ions; build and use the activity series; relate to ionisation energy and its limits |
| 3D | Redox | assign oxidation numbers; identify oxidant/reductant; balance half-equations in acid; combine them |
| 3E | Galvanic cells | label anode (−, oxidation), cathode (+, reduction), electron and ion flow; E°cell = E°cathode − E°anode; predict spontaneity |
| 3F | Measuring rate | calculate average and instantaneous rates from graphs; design a rate investigation |
| 3G | Explaining rate | collision theory; energy profiles (Ea, ΔH, reverse Ea); energy distributions; catalysts |
Key relationships
| Quantity | Relationship |
|---|---|
| Cell potential | $E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$; positive means spontaneous |
| Rate | $\text{rate} = \dfrac{\Delta(\text{quantity})}{\Delta t}$; instantaneous rate = gradient of tangent |
| Energy profile | $\Delta H = H_{\text{products}} - H_{\text{reactants}}$; $E_a(\text{reverse}) = E_a(\text{forward}) - \Delta H$ |
| Gas moles (from Module 2) | $n = \dfrac{V}{24.79\text{ L mol}^{-1}}$ at 25 °C and 100 kPa |
| Oxidation numbers | sum = 0 for a compound, = charge for an ion; O = −2, H = +1 |
Putting it together
A student places a clean strip of nickel in 1.0 mol L⁻¹ silver nitrate solution. (a) Use E° values to predict whether a reaction occurs. (b) Write the half-equations and the net ionic equation. (c) The student repeats the experiment with nickel powder. Explain the effect on the rate.
(a) Ag⁺ would be reduced (E° = +0.80 V) and Ni oxidised (E° = −0.24 V): E° = 0.80 − (−0.24) = +1.04 V. Positive, so the reaction is spontaneous. This agrees with the activity series: nickel is above silver.
(b) Oxidation: Ni(s) → Ni²⁺(aq) + 2e⁻. Reduction: Ag⁺(aq) + e⁻ → Ag(s), doubled. Net ionic: Ni(s) + 2Ag⁺(aq) → Ni²⁺(aq) + 2Ag(s). Nitrate ions are spectators.
(c) Powder has a much larger surface area, so many more nickel atoms are exposed to Ag⁺ ions. There are more collisions per second, so more successful collisions per second, and the rate increases. The final amount of silver is unchanged if the same mass of nickel is limiting.
Revisit any sim from this module to check an answer: the spectator ion sieve (3B) for net ionic equations, metal reactivity trials (3C), the half-equation balancer (3D), build a galvanic cell (3E), rate tangents (3F) and the activation barrier (3G).
Key words
| Term | Meaning |
|---|---|
| Precipitate | an insoluble solid formed when solutions are mixed |
| Spectator ion | an ion present in solution that takes no part in the reaction |
| Net ionic equation | an equation showing only the species that change |
| Activity series | metals ranked by how readily they react (are oxidised) |
| First ionisation energy | energy to remove one mole of electrons from one mole of gaseous atoms |
| Oxidation / reduction | loss / gain of electrons; increase / decrease in oxidation number |
| Oxidising agent | the species that is reduced, taking electrons from another |
| Anode / cathode | electrode where oxidation / reduction occurs |
| Salt bridge | a path for ions between half-cells, keeping each neutral |
| Standard reduction potential, E° | voltage of a half-cell against the standard hydrogen electrode at 25 °C, 100 kPa, 1.0 mol L⁻¹ |
| Activation energy, Ea | the minimum energy colliding particles need to react |
| Transition state | the highest-energy arrangement on the reaction path |
| Catalyst | a substance that increases the rate by providing a lower-Ea pathway, without being consumed |
Test yourself
Which pair of reactants would produce a gas?
Acid + carbonate gives CO₂. Copper doesn't react with dilute acid.
What is the oxidation number of manganese in MnO₂?
Mn + 2(−2) = 0.
In a Zn/Zn²⁺ ‖ Ag⁺/Ag galvanic cell, which statement is correct?
Zinc is oxidised (anode), so electrons leave it. E°cell = 1.56 V; anions move to the anode.
Which is the balanced net ionic equation for magnesium reacting with dilute sulfuric acid?
Sulfate is a spectator; B is unbalanced in charge and atoms.
A reaction has Ea(forward) = 85 kJ mol⁻¹ and ΔH = +25 kJ mol⁻¹. What is Ea(reverse)?
Ea(reverse) = 85 − 25 = 60 kJ mol⁻¹.
Which observation supports the claim that tin is more reactive than copper but less reactive than iron?
Displacement shows the order directly: Fe > Sn > Cu.
Is each reaction a redox reaction?
A metal will displace another metal from solution only if it is in the activity series. The electrode where reduction occurs is the . Increasing temperature increases the fraction of collisions with energy greater than or equal to the . Ions that do not take part in a reaction are ions. A catalyst does not change the value of .
A student adds 0.30 g of zinc powder to 50.0 mL of 0.50 mol L⁻¹ copper(II) sulfate. (a) Use standard reduction potentials to show that a reaction will occur, and write the net ionic equation. (b) Determine which reactant is limiting and the mass of copper formed. (c) Explain, using collision theory, how the rate would differ if zinc granules were used instead.
(a) Cu²⁺ is reduced (E° = +0.34 V) and Zn is oxidised (E° = −0.76 V). E° = 0.34 − (−0.76) = +1.10 V, which is positive, so the reaction is spontaneous. Net ionic: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s).
(b) n(Zn) = 0.30 ÷ 65.38 = 4.59 × 10⁻³ mol. n(Cu²⁺) = 0.50 × 0.0500 = 0.0250 mol. The ratio is 1 : 1, so zinc is limiting. n(Cu) = 4.59 × 10⁻³ mol; m(Cu) = 4.59 × 10⁻³ × 63.55 = 0.29 g.
(c) Granules have a smaller surface area than the same mass of powder. Only zinc atoms at the surface can collide with Cu²⁺ ions, so fewer collisions occur per second, giving fewer successful collisions per second and a slower rate. The same mass of copper still forms in the end, because the amount of zinc is unchanged.