Rate of reaction

1. Overview

The rate of reaction measures how quickly reactants are converted into products over a specific period of time. Understanding these rates is vital in industrial chemistry to maximize production efficiency and ensure safety by controlling the speed of chemical processes.

---

Key Definitions

• Rate of Reaction: The change in concentration, volume, or mass of a reactant or product per unit time.
• Catalyst: A substance that increases the rate of a chemical reaction but remains chemically unchanged at the end of the reaction.
• Enzyme: A biological catalyst (usually a protein) that speeds up metabolic reactions.
• Activation Energy ($E_a$): The minimum energy that colliding particles must possess for a reaction to occur.
• Frequency of Collisions: The number of collisions occurring between particles per unit of time.

---

Core Content

Factors Affecting the Rate of Reaction

The speed of a chemical reaction can be changed by five main factors:

1. Concentration of solutions: Increasing concentration increases the rate.
2. Pressure of gases: Increasing pressure increases the rate.
3. Surface area of solids: Decreasing the particle size (increasing surface area) increases the rate.
4. Temperature: Increasing temperature increases the rate.
5. Catalysts: Adding a catalyst increases the rate.

Investigating Rates of Reaction

There are two primary practical methods used to measure rate:

Method A: Measuring the volume of gas evolved Using a gas syringe to collect gas over time.

• Example: Magnesium reacting with Hydrochloric Acid.
• Word Equation: magnesium + hydrochloric acid → magnesium chloride + hydrogen
• Symbol Equation: $Mg(s) + 2HCl(aq) \rightarrow MgCl_2(aq) + H_2(g)$

Method B: Measuring the change in mass Placing the reaction vessel on a digital balance; mass decreases as gas escapes.

• Example: Calcium carbonate reacting with Nitric Acid.
• Word Equation: calcium carbonate + nitric acid → calcium nitrate + water + carbon dioxide
• Symbol Equation: $CaCO_3(s) + 2HNO_3(aq) \rightarrow Ca(NO_3)_2(aq) + H_2O(l) + CO_2(g)$

Interpreting Rate Graphs

• Steepness of gradient: A steeper curve indicates a faster rate.
• Plateau: When the curve goes flat, the reaction has stopped (one reactant is used up).
• Final Volume: The total amount of product formed depends on the amount of limiting reactant, not the rate.

How to Write a Collision Theory Explanation

The exam frequently asks “explain why increasing X increases the rate.” You need a chain of reasoning, not just a single sentence.

Template for concentration/pressure: More particles in the same volume → particles are closer together → more frequent collisions → more collisions exceed the activation energy → faster rate.

Template for temperature: Particles gain more kinetic energy → they move faster → collisions are more frequent AND a greater proportion of collisions have enough energy to exceed the activation energy ($E_a$) → faster rate.

The temperature answer needs both reasons (frequency and energy). Just saying “particles move faster” is worth only 1 mark out of 2–3.

Template for catalyst: A catalyst provides an alternative reaction pathway with a lower activation energy. More particles now have enough energy to react → faster rate. A catalyst does not change the amount of product formed — it only speeds up how quickly you get there.

---

Extended Content (Extended Only)

Collision Theory

For a reaction to occur, particles must collide with energy greater than or equal to the activation energy ($E_a$).

• Concentration & Pressure: Increasing these increases the number of particles per unit volume. This leads to a higher frequency of collisions, increasing the rate.
• Surface Area: Breaking a solid into smaller pieces exposes more particles. This increases the frequency of collisions between reactant particles.
• Temperature:
  • Particles gain more kinetic energy and move faster, leading to more frequent collisions.
  • Crucially, a much higher proportion of particles now have energy $\geq E_a$, leading to a higher frequency of successful collisions.
• Catalysts: A catalyst provides an alternative reaction pathway with a lower activation energy ($E_a$). Therefore, more particles have sufficient energy to react upon collision.

Evaluating Practical Methods

• Mass Loss: Most accurate for heavy gases like $CO_2(g)$. Not suitable for $H_2(g)$ as it is too light to produce significant mass changes on standard balances.
• Gas Collection: Good for all gases, but if the gas is soluble (like $SO_2$), some may dissolve in the water if collected over a trough, leading to inaccurate results.

---

Key Equations

1. Average Rate of Reaction $$\text{Rate} = \frac{\text{Change in mass or volume}}{\text{Time taken}}$$

• Change: measured in $g$ or $cm^3$
• Time: measured in $s$ or $min$
• Units: $g/s$ or $cm^3/s$

2. Gradient (from a graph) $$\text{Gradient} = \frac{\Delta y}{\Delta x}$$

• Used to find the rate at a specific point in time (draw a tangent).

---

Common Mistakes to Avoid

• ❌ Wrong: Saying "Higher temperature makes particles vibrate more."
• ✓ Right: Temperature increases the kinetic energy and speed of particles in fluids (liquids/gases).
• ❌ Wrong: Saying "Catalysts take part in the reaction and are used up."
• ✓ Right: Catalysts may react intermediate steps, but they are chemically unchanged and their mass is the same at the start and end.
• ❌ Wrong: Confusing "more collisions" with "more frequent collisions."
• ✓ Right: Always use the word frequency or rate of collisions (collisions per second).

---

Exam Tips

• Command Word "Explain": If a question asks you to "Explain" using collision theory, you must mention energy and frequency. Mentioning only "more collisions" will often lose marks.
• Typical Values: Be prepared to handle numerical data like $113.0g$ (mass) or $5.2 cm^3/s$ (rate). Always check your units.
• Contexts: Exam questions often use the "Marble Chip" ($CaCO_3$) experiment or the decomposition of hydrogen peroxide ($H_2O_2$):
  • $2H_2O_2(aq) \rightarrow 2H_2O(l) + O_2(g)$
• Graphing: If a catalyst is added, the curve should be steeper but must end at the same horizontal level (if the amount of reactants is unchanged).

---

Exam-Style Questions

Practice these original exam-style questions to test your understanding. Each question mirrors the style, structure, and mark allocation of real Cambridge 0620 Theory papers.

Exam-Style Question 1 — Short Answer [5 marks]

Question:

A student investigates the rate of reaction between magnesium ribbon and hydrochloric acid. The hydrogen gas produced is collected in a gas syringe.

(a) State two observations, other than the volume of gas collected, that would indicate that the reaction is proceeding at a faster rate. [2]

(b) The student repeats the experiment using the same mass of magnesium ribbon, but this time the magnesium ribbon is powdered. Explain, using collision theory, why the reaction rate increases. [3]

Worked Solution:

(a)

1. Faster bubbling/effervescence. [An increase in the speed of bubbling indicates a faster production of hydrogen gas]
2. The magnesium ribbon disappears faster. [The reaction is consuming the magnesium at a faster rate]

How to earn full marks:

• One mark for each valid observation.
• Do not accept "the reaction happens faster" as it is too vague.

(b)

1. Powdering the magnesium increases the surface area. [This is the initial observation]
2. This increases the frequency of collisions between magnesium particles and hydrochloric acid particles. [More frequent collisions lead to a greater chance of a reaction occurring]
3. Therefore, there are more successful collisions per unit time, increasing the rate of reaction. [Successful collisions are those with sufficient energy to overcome the activation energy]

How to earn full marks:

• 1 mark for stating that the surface area increases.
• 1 mark for linking increased surface area to increased collision frequency.
• 1 mark for explaining that more successful collisions per unit time increases the rate of reaction.

Common Pitfall: Students often forget to link the increased collision frequency directly to the increased rate of reaction. Make sure to explicitly state that more frequent and successful collisions lead to a faster reaction. Also, be specific in your observations; avoid vague statements like "the reaction is faster."

Exam-Style Question 2 — Short Answer [6 marks]

Question:

Hydrogen peroxide decomposes slowly at room temperature according to the following equation:

$2H_2O_2(aq) \rightarrow 2H_2O(l) + O_2(g)$

Manganese(IV) oxide, $MnO_2$, acts as a catalyst for this reaction.

(a) Define the term catalyst. [2]

(b) Describe a simple experiment to show that $MnO_2$ is a catalyst in this reaction. Your description should include:

(i) the observations you would make [2] (ii) how you would confirm the $MnO_2$ is unchanged at the end of the reaction. [2]

Worked Solution:

(a)

1. A catalyst is a substance. [It is a chemical substance]
2. that increases the rate of a chemical reaction without being used up in the reaction. [It speeds up the reaction but is not consumed]

How to earn full marks:

• 1 mark for stating that a catalyst increases the rate of a reaction.
• 1 mark for stating that the catalyst is not used up in the reaction.

(b) (i)

1. Add $MnO_2$ to $H_2O_2$. [Mix the reactants with the catalyst]
2. Observe increased bubbling/effervescence (oxygen production). [The rate of oxygen gas production increases visibly]

How to earn full marks:

• 1 mark for stating to add $MnO_2$ to $H_2O_2$.
• 1 mark for observing increased bubbling or effervescence.

(ii)

1. At the end of the reaction, filter the mixture to recover the $MnO_2$. [Separate the solid catalyst from the liquid]
2. Dry and weigh the recovered $MnO_2$. The mass should be the same as the mass added at the start. [Confirm that the mass of the catalyst remains constant]

How to earn full marks:

• 1 mark for describing filtration to recover the $MnO_2$.
• 1 mark for describing drying and weighing to confirm the mass remains constant.

Common Pitfall: Many students forget to mention the crucial step of drying the recovered catalyst before weighing it. If the catalyst is wet, the mass measurement will be inaccurate, invalidating the conclusion that the catalyst's mass remains unchanged.

Exam-Style Question 3 — Extended Response [8 marks]

Question:

A student investigates the rate of reaction between zinc granules and dilute sulfuric acid. The student measures the volume of hydrogen gas produced at regular time intervals. The experiment is carried out at $25^\circ C$ using $1.0 \text{ g}$ of zinc granules and $50 \text{ cm}^3$ of $1.0 \text{ mol/dm}^3$ sulfuric acid.

The experiment is then repeated with different conditions:

• Experiment 2: $2.0 \text{ g}$ of zinc granules and $50 \text{ cm}^3$ of $1.0 \text{ mol/dm}^3$ sulfuric acid at $25^\circ C$.
• Experiment 3: $1.0 \text{ g}$ of zinc granules and $50 \text{ cm}^3$ of $2.0 \text{ mol/dm}^3$ sulfuric acid at $25^\circ C$.
• Experiment 4: $1.0 \text{ g}$ of zinc granules and $50 \text{ cm}^3$ of $1.0 \text{ mol/dm}^3$ sulfuric acid at $40^\circ C$.

(a) Sketch a graph on the axes below to show the expected results for all four experiments. Label each line clearly with the experiment number (1, 2, 3, and 4). [4]

(b) Explain why the rate of reaction is faster in Experiment 3 compared to Experiment 1, using collision theory. [4]

Worked Solution:

(a)

How to earn full marks:

• 1 mark for all 4 lines starting at (0,0)
• 1 mark for Experiment 4 being the steepest and reaching the same final volume as experiment 1.
• 1 mark for Experiment 3 being steeper than Experiment 1 and reaching a higher final volume.
• 1 mark for Experiment 2 being steeper than Experiment 1 and reaching a higher final volume, and completing faster than experiment 1.

(b)

1. Experiment 3 has a higher concentration of sulfuric acid than Experiment 1. [This is the key difference between the two experiments]
2. This means there are more sulfuric acid particles per unit volume. [Increased concentration means more particles in the same space]
3. Therefore, there is a higher frequency of collisions between zinc and sulfuric acid particles. [More particles lead to more collisions]
4. This results in more successful collisions per unit time, leading to a faster rate of reaction. [More frequent and successful collisions increase the rate of product formation]

How to earn full marks:

• 1 mark for stating that Experiment 3 has a higher concentration of sulfuric acid.
• 1 mark for linking higher concentration to more particles per unit volume.
• 1 mark for explaining increased frequency of collisions.
• 1 mark for explaining more successful collisions leading to a faster rate.

Common Pitfall: When sketching rate curves, many students focus only on the initial rate and forget about the final volume of gas produced. Remember to consider the limiting reactant in each experiment; a higher mass of zinc or a higher concentration of acid will lead to a greater overall volume of hydrogen gas.

Exam-Style Question 4 — Extended Response [9 marks]

Question:

A student investigates the decomposition of aqueous potassium iodide solution, $KI(aq)$, using iron(III) chloride, $FeCl_3(aq)$, as a catalyst. The equation for the reaction is:

$2KI(aq) \rightarrow 2K(aq) + I_2(aq)$

The student performs two experiments. In both experiments, the concentration of $KI(aq)$ is $0.10 \text{ mol/dm}^3$. The colour of the iodine, $I_2(aq)$, is measured using a colorimeter. The higher the concentration of $I_2(aq)$, the higher the reading on the colorimeter.

Experiment 1: $50 \text{ cm}^3$ of $0.10 \text{ mol/dm}^3 KI(aq)$ and $5 \text{ cm}^3$ of $0.01 \text{ mol/dm}^3 FeCl_3(aq)$ are mixed.

Experiment 2: $50 \text{ cm}^3$ of $0.10 \text{ mol/dm}^3 KI(aq)$ and $10 \text{ cm}^3$ of $0.01 \text{ mol/dm}^3 FeCl_3(aq)$ are mixed.

The colorimeter readings are recorded every minute for 10 minutes. The results are shown in the table:

Time (minutes)	Experiment 1 (colorimeter reading)	Experiment 2 (colorimeter reading)
0	0	0
1	5	9
2	10	17
3	14	25
4	18	32
5	22	39
6	25	45
7	28	50
8	31	54
9	33	57
10	35	60

(a) Plot the data on a graph, with time on the x-axis and colorimeter reading on the y-axis. Draw two separate best-fit curves for the two experiments. [4]

(b) Describe how the rate of reaction changes with time in both experiments. [2]

(c) Suggest two ways, other than using a colorimeter, to measure the rate of this reaction. [3]

Worked Solution:

(a)

How to earn full marks:

• 1 mark for correctly labeling the axes with units (Time in minutes, Colorimeter Reading in arbitrary units).
• 1 mark for plotting the data points for Experiment 1 correctly.
• 1 mark for plotting the data points for Experiment 2 correctly.
• 1 mark for drawing smooth best-fit curves for both experiments.

(b)

1. The rate of reaction is fastest at the start of the experiment. [Initially, the concentration of reactants is highest, leading to a higher rate]
2. The rate of reaction decreases with time as the reactants are used up. [As the reaction proceeds, the concentration of reactants decreases, slowing down the rate]

How to earn full marks:

• 1 mark for stating that the rate is fastest at the start.
• 1 mark for stating that the rate decreases with time.

(c)

1. Titration: Take samples of the reaction mixture at regular time intervals and titrate with sodium thiosulfate solution to determine the concentration of iodine. [Titration is a quantitative method to determine the concentration of iodine]
2. Measure the change in mass: As $I_2$ is produced, it can be extracted into an organic solvent. The mass of $I_2$ extracted over time can be measured. [Measuring change in mass offers another way to quantify product formation]
3. Measure the change in electrical conductivity: As the reaction proceeds, the concentration of ions changes, and thus the electrical conductivity of the solution will change. Measure the conductivity of the solution over time. [Changes in conductivity relate to changing ion concentrations]

How to earn full marks:

• 1 mark for each valid method suggested, including a brief explanation of how it relates to the reaction rate.
• Accept other valid methods that measure the change in concentration of reactants or products over time.

Common Pitfall: When suggesting alternative methods for measuring reaction rate, students often provide vague answers. Be specific about what you would measure and how that measurement relates to the change in concentration of reactants or products. For example, simply saying "measure the mass" is insufficient; you need to specify what mass you would measure and how it changes over time.