Rate of reaction Revision Notes

In reactions involving solutions or gases, concentration refers to the amount of reactant particles in a given volume. Increasing the concentration means more particles are present in the same space, which raises the likelihood of collisions between reactant particles.

For example, if hydrochloric acid concentration is increased in a reaction with magnesium, the rate at which hydrogen gas is produced increases because more acid particles collide with magnesium atoms per second.

This increase in collision frequency leads to a higher reaction rate. Therefore, concentration and reaction rate are directly proportional, assuming other factors remain constant.

• Higher concentration → more particles per unit volume.
• More frequent collisions between reactants.
• Increased chance of successful collisions → faster reaction.

When one reactant is a solid, its surface area determines how many particles are exposed and available to react. Breaking a solid into smaller pieces increases its total surface area, exposing more particles to the other reactant.

For instance, powdered calcium carbonate reacts faster with dilute acid than large chunks because the powder has a greater surface area, allowing more acid particles to collide simultaneously.

Thus, increasing surface area increases the number of collisions per second, speeding up the reaction.

• Smaller solid pieces → larger surface area.
• More exposed particles available for reaction.
• Higher collision frequency → faster reaction rate.

Temperature affects the kinetic energy of particles. When temperature rises, particles move faster and collide more often and with greater energy.

Only collisions with energy equal to or greater than the activation energy cause a reaction. Higher temperature increases the proportion of collisions that have sufficient energy to overcome this barrier.

For example, heating a solution of sodium thiosulfate and hydrochloric acid causes the reaction to proceed faster because particles collide more energetically and frequently, increasing the rate of sulfur precipitate formation.

• Higher temperature → increased particle speed.
• More frequent and energetic collisions.
• Greater proportion of successful collisions → faster reaction.

Catalysts are substances that increase the rate of a reaction without being consumed. They provide an alternative reaction pathway with a lower activation energy.

This means more collisions have enough energy to react, increasing the reaction rate without changing the temperature or concentration.

For example, adding manganese dioxide speeds up the decomposition of hydrogen peroxide by lowering the activation energy, producing oxygen gas faster.

• Catalysts lower activation energy.
• Increase proportion of effective collisions.
• Catalysts remain unchanged after reaction.
• Speed up reactions without altering reactants or products.

A catalyst is a substance that increases the speed of a chemical reaction without being consumed or permanently changed by the reaction. Unlike reactants, catalysts do not appear in the overall chemical equation as they remain chemically unchanged at the end. Their presence allows reactions to proceed faster, often under milder conditions such as lower temperature or pressure.

• Catalysts increase reaction rates.
• They are not reactants or products.
• They remain unchanged after the reaction.

Every chemical reaction requires a minimum amount of energy for reactant particles to collide successfully and form products; this is called the activation energy. A catalyst provides an alternative reaction pathway with a lower activation energy. This means more particles have enough energy to react when they collide, increasing the frequency of successful collisions and speeding up the reaction.

For example, in the decomposition of hydrogen peroxide (H₂O₂), manganese dioxide (MnO₂) acts as a catalyst by lowering the activation energy, allowing the reaction to produce oxygen gas faster at room temperature.

• Activation energy is the minimum energy needed for a reaction.
• Catalysts provide an alternative pathway with lower activation energy.
• More particles can react per unit time, increasing reaction rate.

A key feature of catalysts is that they are not consumed during the reaction. They participate in the reaction mechanism, often by temporarily forming intermediate compounds, but are regenerated by the end of the process. This means the same catalyst can be used repeatedly without being used up.

For instance, in the catalytic converter of a car, platinum acts as a catalyst to speed up the conversion of harmful gases into less harmful ones. The platinum does not get used up and continues to catalyse reactions over many miles of driving.

• Catalysts are regenerated at the end of the reaction.
• They do not appear in the overall balanced chemical equation.
• This allows catalysts to be reused multiple times.

Catalysts are essential in industrial and biological processes because they allow reactions to occur faster and often under less extreme conditions, saving energy and cost. For example, the Haber process uses an iron catalyst to produce ammonia efficiently at lower temperatures and pressures than would otherwise be possible.

In biological systems, enzymes are natural catalysts that speed up biochemical reactions necessary for life without being consumed.

• Catalysts reduce energy consumption in industry.
• They enable faster production rates.
• Enzymes are biological catalysts.

The rate of a chemical reaction indicates how quickly reactants are converted into products. Measuring this rate experimentally involves tracking a measurable change over time, such as the loss of mass or the volume of gas produced. These changes provide quantitative data to calculate the reaction speed. Accurate measurement requires careful control of variables and repeated trials to ensure reliability.

• Reaction rate = change in quantity / time
• Common measurable changes: mass loss, gas volume produced
• Control variables to isolate the effect of the independent variable

Some reactions produce a gas that escapes from the reaction mixture, causing the total mass to decrease. Using a digital balance, you can record the mass of the reaction vessel at regular time intervals. The rate is determined by how quickly the mass decreases.

For example, when hydrochloric acid reacts with marble chips (calcium carbonate), carbon dioxide gas is released, reducing the mass.

Key apparatus includes a conical flask, marble chips, hydrochloric acid, a digital balance, and a stopwatch.

• Place the reaction mixture on the balance and start timing.
• Record mass at fixed intervals (e.g., every 10 seconds).
• Plot mass against time to observe the rate.
• Calculate rate from the slope of the mass-time graph.

If a reaction produces a gas, its volume can be measured using a gas syringe or an inverted measuring cylinder in a water trough. This method is suitable when the gas is collected and measured directly.

For example, the reaction between zinc and dilute sulfuric acid produces hydrogen gas, which can be collected in a gas syringe.

Apparatus includes a conical flask, delivery tube, gas syringe or water trough with measuring cylinder, reactants, and stopwatch.

• Set up apparatus to collect gas safely.
• Start the reaction and record gas volume at regular intervals.
• Plot gas volume against time to determine the rate.
• Calculate rate from the gradient of the volume-time graph.

In these experiments, the independent variable might be concentration, temperature, or surface area, while the dependent variable is the measured change (mass or gas volume) over time.

Control variables such as temperature, pressure, and volume of reactants to ensure fair testing.

Repeat experiments multiple times and calculate an average to improve reliability and reduce random errors. Use graphs to visualise data and calculate rates from gradients or average changes over time.

• Identify independent, dependent, and control variables.
• Repeat trials and calculate mean values.
• Use graphs to determine reaction rates.
• Keep conditions constant except for the independent variable.

Handle acids and reactive metals with care, using gloves and eye protection.

Ensure apparatus is secure to prevent spills or gas leaks.

Limitations include measurement precision (e.g., balance sensitivity, gas syringe volume increments) and timing accuracy. To improve, use more precise instruments, smaller time intervals, and ensure thorough mixing of reactants.

• Wear appropriate personal protective equipment.
• Secure apparatus to avoid accidents.
• Use precise instruments and consistent timing.
• Repeat and average results to reduce errors.

In experiments investigating reaction rates, it is essential to identify the independent variable (the factor changed by the experimenter), the dependent variable (the factor measured), and controlled variables (factors kept constant). Common independent variables include concentration, temperature, surface area, and catalyst presence. The dependent variable often involves measuring the amount of product formed or reactant used over time, such as volume of gas produced or mass lost.

For example, in a reaction between hydrochloric acid and magnesium, the concentration of acid might be varied while measuring the volume of hydrogen gas produced every 30 seconds. Recognising these variables helps interpret how changes affect the reaction rate.

• Independent variable: factor changed (e.g., concentration).
• Dependent variable: factor measured (e.g., gas volume).
• Controlled variables: factors kept constant (e.g., temperature).

Graphs are a key tool for analysing reaction rates. Common graphs include: concentration or mass versus time, and volume of gas produced versus time. The shape and slope of these graphs provide information about how fast the reaction proceeds.

For example, a graph of gas volume against time typically rises steeply at first, indicating a fast reaction rate, then levels off as reactants are used up. The steeper the slope at any point, the faster the reaction rate at that time.

When interpreting graphs, note whether the relationship is linear or curved, and how the slope changes over time. This helps understand whether the reaction rate is constant or slowing down.

• Slope of graph = rate of reaction at that time.
• Steep slope = fast reaction; shallow slope = slow reaction.
• Curve flattening indicates reaction slowing as reactants deplete.

The rate of reaction can be calculated from the gradient (slope) of a graph showing change in concentration, mass, or volume over time. The formula is:

Rate = change in quantity / change in time = Δy / Δx

For example, if the volume of gas increases from 0 cm³ to 40 cm³ in 50 seconds, the average rate is: Rate = 40 cm³ ÷ 50 s = 0.8 cm³/s. This value represents the average rate over that time interval, not necessarily the instantaneous rate at any moment.

• Use two points on the graph to find Δy and Δx.
• Calculate slope = Δy/Δx to find rate.
• Units depend on quantity measured (e.g., cm³/s, g/s).

By comparing graphs from experiments where one variable changes, you can analyse how that factor affects the rate. For example, increasing concentration usually increases the rate, shown by steeper slopes on graphs of gas volume versus time.

Similarly, raising temperature increases particle energy and collision frequency, speeding up reactions, which is reflected in faster rises on graphs.

Interpreting these trends requires understanding collision theory: more frequent or more energetic collisions lead to faster reactions.

• Higher concentration → steeper graph slope → faster rate.
• Higher temperature → faster reaction, steeper slope.
• Catalysts increase rate without being consumed, shown by faster product formation.

When analysing data, consider the reliability and accuracy of measurements. Repeating experiments and averaging results reduces random errors.

Look for anomalies or inconsistencies in graphs that may indicate experimental errors or uncontrolled variables.

Be able to explain why a graph has a particular shape and what it reveals about the reaction mechanism or limiting reactants. For example, a plateau in gas volume indicates completion of reaction.

• Repeat experiments to improve reliability.
• Identify and explain anomalies in data.
• Use graph features to infer reaction progress and completion.

Collision theory states that for a chemical reaction to occur, reacting particles must collide with sufficient energy and proper orientation. The minimum energy required for a successful collision is called the activation energy. Only collisions with energy equal to or greater than this threshold lead to product formation.

The rate of reaction depends on how often particles collide (collision frequency) and how many of these collisions have enough energy to overcome the activation energy barrier. Factors that increase collision frequency or the proportion of effective collisions will increase the reaction rate.

• Activation energy is the minimum energy needed for a reaction to proceed.
• Collision frequency is the number of collisions per unit time.
• Only collisions with energy ≥ activation energy cause reactions.

Increasing the concentration of reactants in a solution means more particles are present in the same volume. This raises the likelihood of collisions between reactant particles because they are closer together and more numerous.

For example, doubling the concentration of hydrochloric acid in a reaction with magnesium increases the number of acid particles per unit volume, leading to more frequent collisions with magnesium and a faster reaction.

This effect can be quantitatively observed by measuring the rate of gas production or mass loss over time as concentration changes.

• Higher concentration → more particles per volume → increased collision frequency.
• Example: More concentrated acid reacts faster with metal.
• Reaction rate ∝ collision frequency.

For reactions involving solids, only particles on the surface can collide with other reactants. Increasing the surface area exposes more particles to reactants, increasing the number of collisions.

For instance, powdered calcium carbonate reacts faster with acid than large chunks because the powder has a greater total surface area, allowing more acid particles to collide simultaneously.

This principle explains why grinding solids speeds up reactions and is important in industrial processes to optimise reaction rates.

• Larger surface area → more exposed particles → more collisions.
• Example: Powdered solids react faster than lumps.
• Surface area affects collision frequency at the solid interface.

Raising temperature increases the average kinetic energy of particles, causing them to move faster. Faster particles collide more frequently and with greater energy.

Because more collisions have energy exceeding the activation energy, the proportion of effective collisions rises, significantly increasing the reaction rate.

For example, heating a solution of sodium thiosulfate and hydrochloric acid causes the reaction to complete faster due to more energetic collisions.

• Higher temperature → faster particles → more frequent and energetic collisions.
• More collisions exceed activation energy → higher reaction rate.
• Example: Reactions speed up noticeably with temperature increase.

Catalysts increase the rate of a reaction without being consumed. They provide an alternative reaction pathway with a lower activation energy.

This means a greater proportion of collisions have enough energy to react, even though the temperature and collision frequency remain unchanged.

For example, adding manganese dioxide speeds up the decomposition of hydrogen peroxide by lowering the activation energy, allowing more collisions to be successful.

• Catalysts lower activation energy → more effective collisions.
• Catalysts remain chemically unchanged after the reaction.
• Example: Manganese dioxide catalyses hydrogen peroxide decomposition.

Collision theory provides a framework for understanding how and why chemical reactions occur. It states that for a reaction to happen, reacting particles must collide. However, not all collisions lead to a reaction; certain conditions must be met for a collision to be successful. These conditions involve the energy and orientation of the particles during collision.

Particle density refers to the number of particles present in a given volume. When particle density increases, there are more particles packed into the same space, which raises the likelihood of collisions between them.

Collision frequency is the number of collisions occurring per unit time. A higher particle density naturally leads to a higher collision frequency because particles are more likely to encounter each other.

For example, increasing the concentration of a solution increases the number of dissolved particles in a fixed volume, thus increasing collision frequency and speeding up the reaction.

• Higher particle density means more particles in the same volume.
• More particles increase the chance of collisions.
• Increased collision frequency generally increases reaction rate.

Particles in a substance are constantly moving and possess kinetic energy. The amount of kinetic energy depends on factors such as temperature; higher temperatures mean particles move faster and have more kinetic energy.

For a collision to result in a chemical reaction, the particles must collide with enough kinetic energy to overcome the energy barrier known as activation energy.

If particles collide with insufficient kinetic energy, they will simply bounce off each other without reacting.

• Kinetic energy is the energy of motion of particles.
• Higher kinetic energy increases the chance of successful collisions.
• Temperature affects kinetic energy and thus reaction rates.

Activation energy is the minimum energy that colliding particles must have for a reaction to occur. It represents the energy barrier that must be overcome to break bonds in reactants and form new bonds in products.

Only collisions where the combined kinetic energy of the particles equals or exceeds the activation energy will lead to a reaction.

For example, in the reaction between hydrogen and oxygen gases, molecules must collide with sufficient energy to break the H–H and O=O bonds before new bonds form to create water.

• Activation energy is the minimum energy needed for a reaction.
• Collisions below this energy do not cause reactions.
• Activation energy explains why some reactions are slow at room temperature.

Collision theory links particle density, collision frequency, kinetic energy, and activation energy to explain reaction rates.

Increasing particle density or temperature increases collision frequency and kinetic energy, respectively, leading to faster reactions.

Catalysts work by lowering the activation energy, allowing more collisions to be successful without increasing kinetic energy or particle density.

• Reaction rate depends on both frequency and energy of collisions.
• Adjusting conditions can control reaction speed.
• Understanding these concepts helps in designing safer and more efficient chemical processes.