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IGCSE Chemistry 0620 · Chapter 6 notes · ~17 min read

Rates, equilibrium & redox

Why some reactions are fast and others slow, how to speed them up and read a rate graph, what equilibrium really means (with the Haber and Contact processes) and how to spot oxidation and reduction.

Physical vs chemical changeRate factorsCollision theoryCatalystsReversible reactionsHaber & Contact processesRedox
RSby Riyaz Sir
mentorriyaz.in
No tag = Core (everyone)EXTENDED Extended onlyhighlight = comes up again and againTest yourself at the end

In 30 seconds

  • A chemical change makes a new substance. A physical change does not.
  • Rate goes up with higher concentration, pressure, temperature, surface area or a catalyst. Particles must collide with enough energy. EXT
  • In a reversible reaction (⇌) the products can turn back into reactants. At equilibrium the forward and backward rates are equal. EXT
  • Oxidation = gain of oxygen (or loss of electrons). Reduction = loss of oxygen (or gain of electrons). EXT They always happen together.

By the end you can

  • Tell physical and chemical changes apart.
  • Describe and explain how concentration, pressure, temperature, surface area and catalysts affect rate, and interpret rate graphs.
  • Use collision theory to explain rates. EXT
  • Describe reversible reactions, including hydrated and anhydrous salts, and explain equilibrium and the Haber and Contact processes. EXT
  • Identify oxidation, reduction, oxidising and reducing agents, and work out oxidation numbers. EXT

1Physical and chemical changes

Is a new substance made?

Physical changeChemical change
new substance?no — the same particles, just rearrangedyes — new substances with new properties
reversible?usually easilyusually difficult
energy changesmalloften large (heat, light)
examplesmelting, boiling, dissolving, crushing, magnetisingburning, rusting, cooking, photosynthesis, a metal reacting with acid
  • Signs of a chemical change: colour change, gas produced (bubbles), precipitate formed, temperature change, light or heat given out.
  • A physical change can sometimes look dramatic: boiling water gives a lot of steam, but it is still water. Dissolving sugar is physical — evaporate the water and the sugar is back.
Common mistake"Dissolving salt is a chemical change because the salt disappears." It is physical: it is still sodium chloride, and you get it back by evaporating the water.

2Rate of reaction

What is rate?

Rate of reaction — how fast reactants are used up, or how fast products are made.

  • Fast: explosions, burning magnesium. Slow: rusting, ripening. A steeper graph line = a faster reaction.
concentration or pressurelowhighmore crowded → more collisions per secondsurface areaone lumppowdermore exposed surface → more collisionstemperaturecoldhotfaster + more collisions with energy ≥ Eacatalystno catalystcatalystlower Ea → more successful collisions
Five ways to speed up a reaction, and why. In each, the particles collide more often, or a bigger fraction of collisions succeed.
FactorEffect on rateWhy EXT
higher concentration (solutions)fastermore particles in the same volume, so more frequent collisions
higher pressure (gases)fastergas particles are closer together, so more frequent collisions
higher temperaturemuch fasterparticles move faster → more collisions, and a larger fraction have energy ≥ Ea
larger surface area (smaller pieces, powder)fastermore particles on the surface exposed to the other reactant, so more frequent collisions
catalystfasterprovides a different pathway with a lower Ea, so more collisions are successful
  • A catalyst speeds up a reaction and is not used up — it is chemically the same at the end. It does not change the amount of product.
  • Examples: manganese(IV) oxide for hydrogen peroxide decomposition; iron in the Haber process; vanadium(V) oxide in the Contact process; enzymes (biological catalysts) in living things.
Picture itImagine trying to bump into a friend in a crowded room, rather than an empty one. More people packed in = more bumps per second. A powder is like a lump broken into dust — many more faces get exposed.
Common mistake"At a higher temperature, particles collide more often, so the rate rises." That is only half the answer. The most important reason is that more of the collisions have enough energy to react.
Common mistake"A catalyst gives the particles more energy." It does not — it lowers the energy needed (Ea).

Collision theoryEXTENDED

not enough energyenough energy (≥ Ea)beforeafterslowfastparticles bounce apart:no reactionbonds break and new bonds form:products madeonly collisions with energy ≥ Ea are successful
Only collisions with energy of at least the activation energy are successful. Left: not enough energy, the particles bounce apart. Right: enough energy, a reaction occurs.
  • A reaction can only happen when particles collide with energy greater than or equal to the activation energy (Ea).
  • Rate depends on (1) the frequency of collisions and (2) the energy of the collisions.
  • Not every collision is successful: most particles just bounce apart.
Worked example

EXT Explain why powdered calcium carbonate reacts faster with acid than lumps of the same mass.

  1. Powder has a much larger surface area
  2. more particles of calcium carbonate are exposed to the acid
  3. so there are more frequent collisions between the particles, and the rate is faster
Worked example

EXT Explain why increasing the temperature increases the rate of reaction.

  1. The particles move faster
  2. so they collide more frequently
  3. and a greater proportion of the collisions have energy ≥ the activation energy
  4. so more of the collisions are successful and the rate increases

Catalysts and the energy profileEXTENDED

progress of reactionenergyEa withouta catalystEa withcatalystΔHno catalyst: high humpwith catalyst: analternative route witha lower humpΔH is unchanged · the catalyst is not used up · the products are the same
A catalyst offers a different route with a lower hump (Ea). ΔH is unchanged and the catalyst is not used up.

Measuring the rate

1 measure gas volume2 measure loss of massgas syringe: read thevolume every 10 secondsbalanceCO₂cotton wool lets gas escape butkeeps spray in: mass goes downdrawn for marble chips + acid: CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂
Two ways to follow a reaction that gives off a gas: collect the gas in a gas syringe, or measure the loss of mass on a balance.
  • Gas volume: use a gas syringe or an upturned measuring cylinder and read the volume at regular times.
  • Mass loss: put the flask on a balance. As gas escapes, the mass goes down. Cotton wool in the neck lets the gas out but stops spray escaping.
  • Time for something to happen: for example, the time for a cross to disappear when a precipitate forms.
  • EXT Compare methods: mass loss works well for a heavy gas such as CO₂ but poorly for hydrogen (too light to measure); a gas syringe is accurate but may stick, and some gas can escape while the bung is put in.
average rate = amount of product formed (or reactant used)time taken
for example, cm³ per second, or g per second

Reading a rate graph

timeamount of productfinal amountdotted line: steepest slope (the start)A: faster (hotter,more concentrated, powder,or catalyst added)B: slowersteeper line = faster rate · a line goes flat when a reactant has been used upboth reach the same final amount if they start with the same amount of reactant
Curve A is steeper, so the reaction is faster. Both curves flatten when a reactant has run out, and both reach the same final amount.
  • The gradient of the line is the rate. It is steepest at the start (the most reactant, the most collisions), and it falls as reactants are used up.
  • When the line becomes flat, the reaction has stopped — one reactant has been used up.
  • Faster conditions give a steeper curve that reaches the same final amount (if the amount of the limiting reactant is the same).
Time (s)0102030405060
Volume of gas (cm³)0243846505050
Worked example

Use the table to find (a) when the reaction stops, (b) the mean rate in the first 10 s, (c) the mean rate from 20 s to 30 s, and explain why they differ.

  1. (a) the volume stops rising at 40 s, so the reaction finished at 40 s
  2. (b) 24 ÷ 10 = 2.4 cm³/s
  3. (c) (46 − 38) ÷ 10 = 0.8 cm³/s
  4. The rate is lower later because the reactant concentration has fallen, so there are fewer collisions per second
Worked example

EXT In a mass-loss experiment, 0.44 g of CO₂ escapes from a flask. What volume of gas is this at r.t.p.?

  1. n(CO₂) = 0.44 ÷ 44 = 0.010 mol
  2. V = 0.010 × 24 = 0.24 dm³ (240 cm³)
Exam tipTo compare two experiments say what is the same (fair test) and what is different. A curve that ends at a higher level means more reactant was used; the same final level means only the rate changed.

3Reversible reactions and equilibrium

Reactions that go both ways

Reversible reaction — a reaction in which the products can react to form the reactants again. It is shown by the symbol ⇌.

reversible changes you can seehydrated copper(II)sulfateCuSO₄·5H₂Oblue crystalsheatadd wateranhydrous copper(II)sulfateCuSO₄white powder+ 5H₂O(steam)cobalt(II) chloride(dry, anhydrous)CoCl₂blueadd waterheatcobalt(II) chloride(hydrated)CoCl₂·6H₂Opinktests for water: white anhydrous copper(II) sulfate turns blue;blue cobalt chloride paper turns pink
Hydrated copper(II) sulfate (blue) loses its water on heating to become white anhydrous copper(II) sulfate. Add water and it turns blue again. Cobalt(II) chloride paper is blue when dry and pink with water.
  • Hydrated copper(II) sulfate (blue) ⇌ anhydrous copper(II) sulfate (white) + water: CuSO₄·5H₂O ⇌ CuSO₄ + 5H₂O
  • Anhydrous means "without water". Hydrated means "with water of crystallisation".
  • Heat blue crystals → white powder and steam. Add water to the white powder → it turns blue and gets hot.
  • Test for water: anhydrous copper(II) sulfate turns white → blue; blue cobalt(II) chloride paper turns blue → pink. These tests show that water is present, but not that it is pure (use boiling point for that).

EquilibriumEXTENDED

timerateforward ratereverse rateequilibrium reachedequalA ⇌ BAt equilibrium:• rate forward = rate back• amounts stay constant• it needs a closed system• the reactions have NOT stopped
In a closed system, the forward and reverse rates become equal. The amounts of reactants and products then stay constant.

Dynamic equilibrium — the state of a reversible reaction in a closed system where the forward and reverse reactions happen at the same rate, so the concentrations stay constant.

  • It is called dynamic because both reactions are still happening — it does not stop.
  • A closed system is needed so that nothing can escape.
  • At equilibrium the amounts of reactants and products are constant but not necessarily equal.

Changing the conditionsEXTENDED

  • If you change the conditions, the position of equilibrium shifts to oppose the change:
ChangePosition shifts towards
increase temperaturethe endothermic direction (absorbs the extra heat)
decrease temperaturethe exothermic direction
increase pressurethe side with fewer gas molecules
decrease pressurethe side with more gas molecules
increase concentration of a reactantthe products (uses up the extra reactant)
add a catalystno change — it speeds up both directions equally, so equilibrium is reached faster
Worked example

EXT The reaction N₂O₄ (colourless) ⇌ 2NO₂ (brown) is endothermic in the forward direction. What is seen when the mixture is heated?

  1. Heating favours the endothermic direction
  2. so the position shifts to the right
  3. more NO₂ forms, and the mixture becomes darker brown
Worked example

EXT How does pressure affect 2SO₂ + O₂ ⇌ 2SO₃?

  1. Left side: 2 + 1 = 3 gas molecules. Right side: 2 gas molecules
  2. higher pressure favours the side with fewer molecules (the right)
  3. so more SO₃ is made at high pressure

The Haber processEXTENDED

nitrogenfrom the airhydrogenfrom natural gasreactoriron catalyst450 °C200 atmcooled: ammonialiquefies (−33 °C)mixed 1:3liquid ammonia NH₃unreacted N₂ + H₂ recycledN₂ + 3H₂ ⇌ 2NH₃ (forward reaction is exothermic)
The Haber process. Nitrogen and hydrogen are mixed 1 : 3, passed over an iron catalyst, and the ammonia is liquefied. Unreacted gases are recycled.
N₂ + 3H₂ ⇌ 2NH₃
forward reaction is exothermic
  • Nitrogen comes from the air (fractional distillation of liquid air). Hydrogen comes from natural gas (methane) reacting with steam.
  • Conditions: 450 °C, 200 atm, iron catalyst.
  • Only a fraction of the nitrogen and hydrogen turns into ammonia on each pass. The unreacted gases are recycled and the ammonia is cooled until it liquefies and is removed.

The Contact processEXTENDED

1 burn sulfurS + O₂ → SO₂(sulfur and air)2 converter2SO₂ + O₂ ⇌ 2SO₃V₂O₅ catalyst450 °C, 1–2 atm3 absorb SO₃ inconc. H₂SO₄ tomake oleumH₂S₂O₇4 add waterH₂S₂O₇ + H₂O→ 2H₂SO₄sulfuric acidH₂SO₄SO₃ is not added straight to water: the reaction gives off so much heatit makes a mist of acid. Oleum is diluted instead.450 °C: a compromise between rate (hotter is faster) and yield (colder is better)1–2 atm: the yield is already high, so high pressure would waste money
The Contact process makes sulfuric acid in four steps. Sulfur dioxide is oxidised to sulfur trioxide over a vanadium(V) oxide catalyst.
2SO₂ + O₂ ⇌ 2SO₃
forward reaction is exothermic
  • Sulfur dioxide comes from burning sulfur (S + O₂ → SO₂) or from roasting sulfide ores. Oxygen comes from the air.
  • Conditions: 450 °C, 1–2 atm, vanadium(V) oxide (V₂O₅) catalyst.
  • SO₃ is absorbed in concentrated sulfuric acid to make oleum (H₂S₂O₇), then diluted with water to give H₂SO₄. SO₃ is not added directly to water because the reaction is so violent it makes a mist of acid.

Why those conditions?EXTENDED

Haber processContact process
exothermic?yesyes
temperature450 °C450 °C
pressure200 atm1–2 atm
catalystironvanadium(V) oxide
  • Temperature — a compromise. Exothermic reactions give more product at low temperature, but the rate is too slow when it is cold. So a moderate temperature (450 °C) is chosen to get an acceptable rate and yield.
  • Pressure — Haber: 4 molecules → 2 molecules, so high pressure gives more ammonia and a faster rate. But very high pressure costs more (strong equipment, energy) and is dangerous, so 200 atm is a compromise.
  • Pressure — Contact: 3 molecules → 2, but the yield is already high at 1–2 atm, so high pressure would waste money.
  • Catalyst: speeds up the reaction so equilibrium is reached quickly at a lower temperature. It has no effect on the yield.
Exam tipWhen explaining industrial conditions always mention rate, yield (position of equilibrium), safety and cost. "A compromise" is the key idea.

4Redox

Oxidation and reduction

Oxidation — gain of oxygen. EXT Also loss of electrons, or an increase in oxidation number.

Reduction — loss of oxygen. EXT Also gain of electrons, or a decrease in oxidation number.

Redox reaction — a reaction in which oxidation and reduction happen at the same time.

Mg + Cu²⁺ → Mg²⁺ + CuMgCu²⁺2e⁻Mg²⁺(2,8)CuMg loses electrons= OXIDATIONMg is the reducing agentCu²⁺ gains electrons= REDUCTIONCu²⁺ is the oxidising agentOIL RIGOxidation Is LossReduction Is Gaincolour-change tests (for oxidising and reducing agents)acidified KMnO₄ turns purple → colourless:MnO₄⁻ is reduced, so it is an oxidising agentpotassium iodide turns colourless → brown:I⁻ is oxidised to I₂, so it is a reducing agent
Magnesium + copper(II) ions. Magnesium loses two electrons (oxidation); the copper ion gains them (reduction).
  • Example: Fe₂O₃ + 3CO → 2Fe + 3CO₂. Fe₂O₃ loses oxygen (reduced). CO gains oxygen (oxidised).
  • Example: 2Mg + O₂ → 2MgO. Magnesium gains oxygen: it is oxidised.
Memory trickOIL RIG — Oxidation Is Loss (of electrons); Reduction Is Gain.
  • EXT Mg + Cu²⁺ → Mg²⁺ + Cu: Mg loses 2 electrons → oxidised. Cu²⁺ gains 2 electrons → reduced.
  • EXT An oxidising agent oxidises another substance and is itself reduced. A reducing agent reduces another substance and is itself oxidised.

Oxidation numbers

  • Roman numerals in names show the oxidation number of a metal: iron(II) chloride FeCl₂ has Fe²⁺; iron(III) chloride FeCl₃ has Fe³⁺; copper(II) oxide CuO has Cu²⁺.
  • EXT Rules: (1) uncombined elements = 0; (2) a simple ion = its charge; (3) in compounds, hydrogen is usually +1 and oxygen is usually −2; (4) the oxidation numbers in a compound add up to 0, and in an ion add to its charge.
  • EXT Oxidation = the oxidation number increases. Reduction = the oxidation number decreases.
Worked example

EXT Find the oxidation number of (a) S in H₂SO₄, (b) Mn in KMnO₄, (c) Cr in Cr₂O₇²⁻.

  1. (a) 2(+1) + S + 4(−2) = 0 → S = +6
  2. (b) (+1) + Mn + 4(−2) = 0 → Mn = +7
  3. (c) 2 × Cr + 7(−2) = −2 → 2 × Cr = +12 → Cr = +6
Worked example

EXT In Fe₂O₃ + 3CO → 2Fe + 3CO₂, identify what is oxidised and reduced, and the oxidising and reducing agents.

  1. Fe: +3 → 0, the number decreases, so Fe₂O₃ is reduced
  2. C: +2 → +4, the number increases, so CO is oxidised
  3. Fe₂O₃ is the oxidising agent; CO is the reducing agent

Colour tests for oxidising and reducing agentsEXTENDED

  • Acidified potassium manganate(VII) (KMnO₄): purple → colourless when it is reduced. So it shows a reducing agent is present.
  • Potassium iodide (KI): colourless → brown (iodine) when I⁻ is oxidised. So it shows an oxidising agent is present.
  • It looks backwards at first: the test chemical changes colour, and the thing it detects does the opposite job.
Common mistake"Oxidation is gain of electrons." No — Oxidation Is Loss.
Common mistakeForgetting that redox needs both processes. Something that is oxidised must be balanced by something that is reduced.

5Test yourself

Test yourself — cover the answers first

Quick check

  1. Give two signs of a chemical change.
  2. Name five factors that change the rate of a reaction.
  3. What does the symbol ⇌ mean?
  4. What colour change shows that water is present using anhydrous copper(II) sulfate?
  5. What is oxidation, in terms of oxygen?
  6. What is the oxidation number of iron in iron(III) oxide?
  7. EXT What does a catalyst do to the activation energy?
  8. EXT Why does a higher pressure increase the yield of ammonia?

Exam-style

  1. Explain, using particles, why increasing concentration increases the rate of reaction. [2]
  2. EXT Explain why the Contact process uses a pressure of only 1–2 atm. [2]
  3. EXT State the conditions used in the Haber process and explain why 450 °C is used. [4]
  4. EXT Use oxidation numbers to show that the reaction Mg + CuSO₄ → MgSO₄ + Cu is redox. [3]
Show answers
  1. any two from: colour change, gas given off, precipitate, temperature change, light
  2. concentration, pressure (gases), temperature, surface area, catalyst
  3. the reaction is reversible
  4. white → blue
  5. gain of oxygen
  6. +3
  7. lowers it
  8. there are 4 gas molecules on the left but 2 on the right, so the position shifts right
  9. more particles per unit volume
    so more frequent collisions and so a higher rate
  10. the yield of SO₃ is already high at this pressure
    high pressure would increase costs and safety risks without much more product
  11. 450 °C, 200 atm, iron catalyst
    low temperature gives a better yield (the reaction is exothermic)
    but the rate is too slow
    450 °C is a compromise between rate and yield
  12. Mg: 0 → +2, oxidised (increase)
    Cu: +2 → 0, reduced (decrease)
    oxidation and reduction occur together, so it is redox
Practise this chapter
Physical & chemical changesRate of reactionReversible reactions & equilibriumRedox
Next chapterAcids, bases & salts →
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