Reaction types, rates, exo/endothermic

What you will learn

classify reactions as synthesis, decomposition or displacement,
use collision theory to explain how temperature, concentration, surface area and catalysts affect reaction rate,
identify activation energy on an energy profile,
distinguish exothermic (energy released) from endothermic (energy absorbed) reactions,
sketch and interpret energy profile diagrams.

Worked example 0 Real-world example: why flour mills explode

Flour stored in bags does not burn. Fine flour dust suspended in air, however, can ignite explosively. Explain using collision theory.

Flour is a combustible substance (it reacts with O $_2$ ).
A bag of flour has a small surface area relative to its mass; only the outermost grains touch oxygen.
Airborne dust has an enormous surface area — every tiny particle is exposed to O $_2$ on all sides.
Enough O $_2$ -flour collisions per second occur to sustain rapid combustion; one spark is enough to set it off.

Key idea: surface area is one of the four factors in collision theory. Increasing it can turn an inert solid into a dangerous explosive.

1. Types of chemical reactions

Synthesis (combination): two or more reactants combine. $A + B \to AB$ . Example: $2\text{Mg} + \text{O}_2 \to 2\text{MgO}$ .
Decomposition: one compound breaks into simpler substances. $AB \to A + B$ . Example: $2\text{H}_2\text{O}_2 \to 2\text{H}_2\text{O} + \text{O}_2$ .
Displacement (single replacement): a more reactive element displaces a less reactive one. Example: $\text{Zn} + \text{CuSO}_4 \to \text{ZnSO}_4 + \text{Cu}$ .

A brief activity (reactivity) series for metals (most to least reactive): K, Na, Ca, Mg, Al, Zn, Fe, Pb, (H), Cu, Ag, Au. A metal higher in the list displaces one below from its compound.

Worked example 1 Predicting displacement

Will copper displace silver from silver nitrate solution?

Copper is above silver in the reactivity series.
Yes — a displacement reaction occurs: $\text{Cu} + 2\text{AgNO}_3 \to \text{Cu(NO}_3)_2 + 2\text{Ag}$ .
A blue solution forms as Cu $^{2+}$ enters solution, and silver crystals deposit on the copper wire.

2. Collision theory and reaction rate

For a reaction to happen, reactant particles must:

Collide with each other.
Have at least the activation energy (minimum energy to react).
Collide with the correct orientation (geometry).

Rate of reaction is how fast products form (or reactants disappear) per unit time.

Four factors that increase rate:

Temperature: higher temperature gives particles more kinetic energy. More collisions per second, and a greater fraction exceed the activation energy. Roughly, a $10\,^\circ\text{C}$ rise can double the rate.
Concentration (liquids/gases): more particles per unit volume $\to$ more collisions per second.
Surface area (solids): breaking a solid into smaller pieces exposes more of its atoms.
Catalyst: a substance that provides an alternative pathway with lower activation energy. Not consumed in the overall reaction.

Worked example 2 Explaining a graph of rate

A student reacts excess zinc with dilute hydrochloric acid at two temperatures, $20\,^\circ\text{C}$ and $40\,^\circ\text{C}$ , and measures the volume of H $_2$ gas over time. Which curve is steeper, and why?

The $40\,^\circ\text{C}$ curve is steeper (gas produced faster initially).
Higher temperature means faster-moving acid particles; they collide more often with the zinc surface, and a greater fraction have enough energy to react.
Both curves reach the same final volume (same total zinc, same acid), because temperature changes the rate, not the amount produced.

3. Exothermic and endothermic reactions

Exothermic: releases energy (usually as heat). The products have less chemical energy than the reactants. Example: combustion of methane, setting-of cement, hand warmers.
Endothermic: absorbs energy. Products have more chemical energy than reactants. Example: photosynthesis, cold packs with ammonium nitrate, baking bread.

Activation energy ( $E_a$ ) is the energy “hill” that reactants must climb before products can form. This applies to both exo- and endothermic reactions.

Energy profiles for exothermic (left) and endothermic (right) reactions. The peak is the activation energy barrier.

Worked example 3 Classifying energy changes

Classify each reaction as exothermic or endothermic: (a) a methane flame on a gas stove; (b) baking soda and vinegar mixed in a beaker — the beaker feels cold.

(a) The flame gives out heat and light. Exothermic.
(b) The beaker feels cold because the reaction absorbs heat from its surroundings. Endothermic.

Key idea: “feels hot” = exothermic (heat flows out of the system); “feels cold” = endothermic (heat flows into the system).

Worked example 4 Catalyst on an energy profile

A catalyst is added to a slow reaction. Describe what happens to (a) the activation energy, (b) the energy of reactants and products, (c) the rate.

(a) Activation energy decreases — the energy hill is lowered by an alternative pathway.
(b) Reactant and product energies are unchanged. The catalyst does not alter how exothermic or endothermic the reaction is.
(c) Rate increases — more collisions now exceed the lower activation energy.

Practice: Year 10

Fluency

Reaction types

Classify: (a) $2\text{KClO}_3 \to 2\text{KCl} + 3\text{O}_2$ ; (b) $2\text{Na} + \text{Cl}_2 \to 2\text{NaCl}$ ; (c) $\text{Fe} + \text{CuSO}_4 \to \text{FeSO}_4 + \text{Cu}$ .
Predict whether the reaction occurs: (a) $\text{Ag} + \text{CuSO}_4$ ; (b) $\text{Mg} + \text{ZnCl}_2$ . Justify using the reactivity series.
Write a balanced equation for the synthesis of magnesium oxide from magnesium metal and oxygen.
Write a balanced equation for the decomposition of calcium carbonate on heating.

Fluency

Rate factors

State the three conditions for a successful collision between reactant particles.
List four factors that increase reaction rate.
Does increasing concentration of a solid reactant affect the rate? Justify.
Explain why powdered sugar burns much faster than a sugar cube.
What is a catalyst? Does it get used up in the reaction?

Fluency

Energy changes

Define exothermic and endothermic with one example of each.
On an energy profile, how can you tell if the reaction is exothermic?
What does activation energy mean, and where is it on an energy profile?
True or false: a catalyst changes the amount of energy released by a reaction. Justify.

Reasoning

Explain using collision theory

Cold food lasts longer in the fridge. Explain in terms of reaction rate.
A student grinds a marble chip into powder before adding to acid. Predict the effect on the rate of CO $_2$ production and explain.
Why is a $2$ mol/L acid solution more vigorous than a $0.5$ mol/L solution with the same metal?
Explain why a rise of $10\,^\circ\text{C}$ can roughly double a reaction’s rate, yet the activation energy of the reaction has not changed.
A catalyst is “not consumed” in the reaction. Explain what this means using the idea of an alternative pathway.

Problem solving

Apply and interpret

A sparkler burns brightly while a lump of iron does not. Explain using surface area and temperature.
Photosynthesis stores energy from sunlight in glucose. Is this reaction exothermic or endothermic? How is combustion of glucose (respiration) related?
A chemist measures CO $_2$ released when CaCO $_3$ reacts with HCl. At $20\,^\circ\text{C}$ the reaction finishes in $120$ s; at $40\,^\circ\text{C}$ it finishes in $50$ s. (a) Which has the faster average rate? (b) Is the total mass of CO $_2$ produced the same in both cases?
Hydrogen peroxide decomposes slowly, but adding manganese dioxide powder produces rapid bubbling. (a) What role does MnO $_2$ play? (b) What happens to its mass over the reaction?

Challenge

Reasoning

Harder reasoning

A reaction has activation energy $E_a = 50$ kJ/mol. At $300$ K, about $1$ in $10^8$ collisions has enough energy to react. Explain qualitatively why doubling the concentration does not change the fraction of successful collisions but does change the reaction rate.
Describe how a catalytic converter in a car exhaust speeds up the conversion of CO and unburnt fuel into CO $_2$ , and why precious metals (Pt, Pd, Rh) are used despite the cost.
Sketch energy profile diagrams for: (a) an exothermic reaction with and without a catalyst; (b) an endothermic reaction. Label activation energy on each.
Industrial chemists run the Haber process at $400$ - $500\,^\circ\text{C}$ even though higher temperatures reduce yield. Use collision theory and an economic argument to explain this compromise.