Rate of reaction Revision Notes

In reactions involving solutions, the concentration of reactants is the amount of solute dissolved per unit volume of solvent, usually expressed in mol/dm³. Increasing concentration means more reactant particles are present in the same volume, which raises the frequency of collisions between reacting particles.

According to collision theory, a higher collision frequency increases the chance of successful collisions that overcome the activation energy, thus speeding up the reaction.

For example, in the reaction between hydrochloric acid and sodium thiosulfate, increasing the acid concentration increases the rate at which the solution becomes cloudy due to sulfur precipitate formation:

• HCl (aq) + Na2S2O3 (aq) → S (s) + SO2 (g) + NaCl (aq)
• Higher [HCl] → more frequent collisions → faster reaction

For reactions involving gases, increasing the pressure compresses the gas molecules into a smaller volume. This increases the concentration of gas particles, leading to more frequent collisions per unit time.

For example, in the reaction between hydrogen and iodine gases:

• H2 (g) + I2 (g) → 2HI (g)
• Increasing pressure → increased collision frequency → faster reaction

When a reactant is a solid, only the particles on the surface can collide with other reactants. Increasing the surface area exposes more particles to reactants, increasing the number of collisions and speeding up the reaction.

For example, powdered calcium carbonate reacts faster with dilute hydrochloric acid than large lumps because the powder has a greater surface area:

• CaCO3 (s) + 2HCl (aq) → CaCl2 (aq) + CO2 (g) + H2O (l)
• Powdered CaCO3 → more surface area → faster reaction

Raising the temperature increases the kinetic energy of particles, causing them to move faster. This results in:

More frequent collisions between particles per unit time.

A greater proportion of collisions having energy equal to or greater than the activation energy, making them successful collisions that form products.

• For example, the reaction between sodium thiosulfate and hydrochloric acid proceeds faster at higher temperatures.
• The Arrhenius equation relates rate constant k to temperature T and activation energy Ea: k = Ae^(-Ea/RT)
• where A is the frequency factor, R is the gas constant, and T is temperature in kelvin.

Catalysts increase the rate of a chemical reaction by providing an alternative reaction pathway with a lower activation energy. They are not consumed in the reaction and do not affect the position of equilibrium.

Enzymes are biological catalysts that speed up reactions in living organisms by lowering activation energy, often with high specificity for substrates.

For example, the decomposition of hydrogen peroxide is slow but is catalysed by manganese(IV) oxide (MnO2):

• 2H2O2 (aq) → 2H2O (l) + O2 (g) (slow)
• With MnO2 catalyst → reaction is faster
• Catalysts lower activation energy → more particles have sufficient energy → increased rate

A catalyst is a substance that increases the rate of a chemical reaction without itself undergoing any permanent chemical change. It achieves this by providing an alternative reaction pathway with a lower activation energy compared to the uncatalysed reaction.

Because the catalyst is not consumed, it can participate repeatedly in multiple reaction cycles. This distinguishes catalysts from reactants, which are used up during the reaction.

For example, in the decomposition of hydrogen peroxide (H₂O₂), manganese(IV) oxide (MnO₂) acts as a catalyst to speed up the reaction without being changed chemically.

• Catalyst increases reaction rate.
• Catalyst remains chemically unchanged after reaction.
• Catalyst provides an alternative pathway with lower activation energy.

Chemical reactions require particles to collide with sufficient energy to overcome an energy barrier called the activation energy. A catalyst lowers this activation energy by providing a different mechanism or pathway for the reaction to proceed.

Lowering the activation energy means more particles have enough energy to react when they collide, increasing the frequency of successful collisions and thus speeding up the reaction.

This effect can be represented on an energy profile diagram, where the peak representing activation energy is lower for the catalysed reaction compared to the uncatalysed one.

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

Unlike reactants, catalysts do not get used up during the reaction. They participate in the reaction mechanism but are regenerated by the end of the reaction cycle in their original chemical form.

This means the mass and chemical identity of the catalyst remain the same before and after the reaction, allowing it to be used repeatedly.

For example, in the catalytic decomposition of hydrogen peroxide, MnO₂ remains as MnO₂ after the reaction and can be recovered and reused.

• Catalyst is not consumed during the reaction.
• Catalyst's chemical composition and mass remain unchanged.
• Catalyst can be recovered and reused multiple times.

Hydrogen peroxide decomposes slowly into water and oxygen:

2H₂O₂(aq) → 2H₂O(l) + O₂(g) (uncatalysed) — slow reaction

When manganese(IV) oxide (MnO₂) is added as a catalyst, the reaction speeds up significantly: MnO₂ provides an alternative pathway with lower activation energy, increasing the rate of oxygen gas production without being consumed: MnO₂(s) → MnO₂(s) (unchanged)

• Uncatalysed reaction is slow.
• MnO₂ acts as a catalyst, speeding up oxygen evolution.
• MnO₂ remains unchanged after reaction.

In practical experiments, the presence of a catalyst can be observed by a faster rate of product formation or reactant disappearance.

For example, when MnO₂ is added to hydrogen peroxide, rapid bubbling of oxygen gas is observed compared to the slow reaction without catalyst.

Catalysts are widely used in industry to increase efficiency and reduce energy consumption by speeding up reactions without being consumed, lowering costs and waste.

• Catalyst increases observable reaction rate.
• Gas evolution or colour change may be faster with catalyst.
• Catalysts improve industrial reaction efficiency.

One practical method to measure the rate of a reaction is by monitoring the change in mass of the reaction mixture over time. This is especially useful when a gas is produced and escapes from the reaction vessel, causing a decrease in mass.

A typical setup involves placing the reaction vessel on a digital balance and recording the mass at regular time intervals. For example, when magnesium reacts with hydrochloric acid, hydrogen gas is produced and escapes, reducing the mass.

The balanced chemical equation for this reaction is: Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)↑ (gas escapes)

• Use a digital balance with suitable precision (e.g., 0.01 g).
• Record mass at fixed time intervals (e.g., every 10 seconds).
• Ensure the reaction vessel is open or connected to allow gas to escape.
• Plot mass against time to determine the rate of reaction.
• Control variables such as temperature and concentration for valid comparisons.

Another common method is to measure the volume of gas produced during a reaction over time. This is applicable when a gas is evolved, such as hydrogen, oxygen, or carbon dioxide.

A gas syringe or an inverted measuring cylinder over water can be used to collect and measure the volume of gas produced at regular intervals.

For example, when calcium carbonate reacts with hydrochloric acid, carbon dioxide gas is produced: CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + CO₂(g)↑ + H₂O(l)

• Set up a gas syringe connected to the reaction vessel to collect gas.
• Record the volume of gas at fixed time intervals (e.g., every 15 seconds).
• Ensure the apparatus is airtight to prevent gas leakage.
• Plot gas volume against time to find the rate of reaction.
• Keep temperature, concentration, and surface area constant during experiments.

In any rate measurement experiment, it is essential to identify and control variables to ensure valid and reproducible results.

The independent variable is the factor deliberately changed, such as concentration of reactants or surface area of a solid.

The dependent variable is the quantity measured, such as mass loss or volume of gas produced over time, which reflects the rate of reaction.

• Independent variable examples: concentration, temperature, surface area.
• Dependent variable examples: mass change, gas volume produced.
• Control variables include temperature, pressure, catalyst presence, and apparatus setup.
• Repeat experiments to check reliability and calculate mean values.

Accurate timing and precise measurement are crucial for reliable rate data. Use a stopwatch to record time intervals precisely.

Plotting data such as mass or gas volume against time allows determination of the rate from the gradient of the curve or line.

Repeating the experiment and calculating an average rate reduces random errors and increases confidence in results.

• Use a stopwatch for accurate timing.
• Record measurements at consistent intervals.
• Plot graphs of mass or volume against time.
• Calculate rate from the initial slope of the graph.
• Repeat experiments and calculate mean values for accuracy.
• Ensure apparatus is clean and calibrated to reduce systematic errors.

When measuring reaction rates, safety precautions must be observed. Handle acids and reactive metals with care, using gloves and eye protection.

Ensure good ventilation when gases are produced, especially if toxic or flammable.

Avoid spills and clean apparatus after use to prevent contamination or hazards.

• Wear safety goggles and gloves when handling chemicals.
• Conduct experiments in a well-ventilated area or fume cupboard.
• Handle gas syringes carefully to avoid damage or leaks.
• Dispose of chemical waste according to safety guidelines.

In experiments investigating the rate of reaction, key variables include concentration of reactants, temperature, surface area of solids, and presence of catalysts. Data collected typically involve measuring changes over time, such as the volume of gas produced, mass loss, or concentration changes.

Accurate data collection requires controlling variables other than the one being tested to ensure valid conclusions. Repeating measurements and calculating averages improve reliability.

Common apparatus includes gas syringes for measuring gas volume, balances for mass changes, and colorimeters for concentration changes in coloured solutions.

• Identify independent variable (e.g., concentration) and dependent variable (e.g., gas volume).
• Control other variables to isolate effects on reaction rate.
• Repeat measurements to reduce random errors.
• Use appropriate apparatus for precise measurements.

Graphing data is essential for visualising how reaction rates change over time. Common graphs include concentration or mass versus time and volume of gas produced versus time.

The gradient (slope) of a curve on these graphs at any point represents the rate of reaction at that time. A steeper gradient indicates a faster rate.

For example, a graph of volume of gas produced against time often shows a curve that rises steeply at first and then levels off as reactants are used up.

• Plot dependent variable against time to observe reaction progress.
• Calculate gradient to find instantaneous rate.
• Compare gradients under different conditions to assess effects on rate.
• Use tangent lines for rates at specific times on curves.

The rate of reaction can be calculated using the formula:

Rate = change in quantity of reactant or product / time taken

For example, if 20 cm³ of gas is produced in 50 seconds, the rate is: Rate = 20 cm³ ÷ 50 s = 0.4 cm³/s. This represents the average rate over that time interval.

• Use consistent units for volume, mass, and time.
• Calculate average rate over time intervals.
• Determine instantaneous rate from gradients of tangent lines on graphs.
• Check units and convert if necessary (e.g., cm³ to dm³).
• Example calculation: If 15 cm³ of gas is produced in 30 s, rate = 15 ÷ 30 = 0.5 cm³/s.

By comparing data sets or graphs, one can determine how changes in variables affect reaction rate. For instance, increasing concentration usually increases the rate, shown by steeper gradients or faster gas production.

Data may show that rates decrease over time as reactants are consumed, indicated by curves flattening on graphs.

Understanding these trends helps explain reaction mechanisms and supports collision theory, where more frequent or more energetic collisions increase rate.

• Identify how rate changes with variables from data.
• Explain decreasing rate as reactants are used up.
• Relate data trends to particle collision frequency and energy.
• Use data to support or refute hypotheses about reaction conditions.

An experiment measures the volume of hydrogen gas produced when magnesium reacts with hydrochloric acid. The data show 30 cm³ of H₂ collected in 60 seconds.

Calculate the average rate of reaction: Rate = volume of gas / time = 30 cm³ ÷ 60 s = 0.5 cm³/s.

If a second experiment with higher acid concentration produces 45 cm³ in 60 seconds, the rate is 45 ÷ 60 = 0.75 cm³/s, indicating an increased rate due to concentration increase.

• Use formula: Rate = change in quantity / time.
• Substitute values carefully with units.
• Compare rates to assess effect of variables.
• Common error: forgetting to convert units or mixing units.

Collision theory states that for a chemical reaction to occur, reacting particles must collide with sufficient energy and proper orientation. Not all collisions lead to a reaction; only those where particles have enough energy to overcome the energy barrier and are aligned correctly result in product formation.

The theory helps explain why reaction rates vary under different conditions by focusing on the frequency and energy of collisions between particles.

• Particles must collide to react.
• Collisions must have sufficient energy (activation energy).
• Particles must be oriented correctly during collision.

Particle density refers to the number of particles in a given volume. Higher particle density increases the likelihood of collisions because more particles are present to collide within the same space.

Collision frequency is the number of collisions occurring per unit time. Increasing particle density raises collision frequency, which generally increases the rate of reaction, assuming other factors remain constant.

For example, increasing the concentration of a solution or the pressure of a gas increases particle density and thus collision frequency.

• Higher particle density → more collisions per second.
• Increased collision frequency → faster reaction rate.
• Concentration and pressure affect particle density.

Kinetic energy is the energy of motion possessed by particles. At a given temperature, particles have a range of kinetic energies.

Only collisions where particles have kinetic energy equal to or greater than the activation energy can result in a reaction. Increasing temperature raises the average kinetic energy, increasing the proportion of particles that can overcome the activation energy barrier.

Thus, higher kinetic energy increases the number of effective collisions, speeding up the reaction.

• Particles have varying kinetic energies at a given temperature.
• Higher kinetic energy increases effective collisions.
• Temperature affects kinetic energy distribution.

Activation energy is the minimum energy required for a reaction to occur during a collision. It represents the energy barrier that reactant particles must overcome to form an activated complex or transition state before converting into products.

If colliding particles have less energy than the activation energy, they will not react and will simply bounce apart.

Activation energy depends on the nature of the reactants and the reaction pathway.

• Activation energy is the minimum collision energy needed.
• Only collisions with energy ≥ activation energy cause reaction.
• Activation energy determines reaction feasibility.

Consider the reaction between hydrogen gas and iodine gas to form hydrogen iodide:

H₂(g) + I₂(g) → 2HI(g) (balanced equation) At higher temperatures, the average kinetic energy of H₂ and I₂ molecules increases, resulting in more collisions with energy exceeding the activation energy.

This leads to an increased rate of formation of HI, as more effective collisions occur per unit time.

• Balanced equation: H₂(g) + I₂(g) → 2HI(g)
• Higher temperature → higher kinetic energy → more effective collisions
• Reaction rate increases with temperature due to collision theory.

Collision theory states that for a chemical reaction to occur, reacting particles must collide with sufficient energy (equal to or greater than the activation energy) and with the correct orientation. The rate of reaction depends on the frequency of such effective collisions.

Particles in gases and solutions move randomly; increasing the number of collisions or the energy of collisions increases the reaction rate.

• Effective collision requires sufficient energy and proper orientation.
• Reaction rate depends on frequency of effective collisions.

Increasing the concentration of reactants in solution increases the number of particles per unit volume. This raises the frequency of collisions between reactant particles, thus increasing the rate of reaction.

For example, in the reaction between hydrochloric acid and sodium thiosulfate, increasing acid concentration speeds up the formation of sulfur precipitate.

No change in activation energy occurs; only collision frequency increases.

• Higher concentration → more particles per volume → more collisions per second → faster reaction.
• Example: HCl + Na2S2O3 reaction rate increases with acid concentration.

Increasing the pressure of reacting gases compresses the gas particles into a smaller volume, increasing their concentration.

This leads to more frequent collisions per unit time and thus a faster reaction rate.

For example, increasing pressure speeds up the reaction between hydrogen and iodine gases to form hydrogen iodide: H2(g) + I2(g) → 2HI(g).

• Higher pressure → higher gas concentration → increased collision frequency → faster reaction.
• Example: H2 + I2 reaction rate increases with pressure.

For reactions involving solids, increasing the surface area exposes more particles to react.

Powdered solids react faster than large lumps because more particles are available to collide with other reactants.

For example, powdered calcium carbonate reacts faster with dilute hydrochloric acid than large marble chips.

• Greater surface area → more exposed particles → more collisions → faster reaction.
• Example: powdered CaCO3 reacts faster with HCl than large pieces.

Increasing temperature increases the average kinetic energy of particles, causing them to move faster.

Faster particles collide more frequently and with greater energy.

More collisions have energy equal to or greater than the activation energy, increasing the proportion of effective collisions and thus the reaction rate significantly (often doubling for every 10 °C rise).

• Higher temperature → increased kinetic energy → more frequent and more energetic collisions → faster reaction.
• Activation energy barrier overcome by more particles at higher temperature.

Catalysts increase the rate of a reaction without being consumed.

They provide an alternative reaction pathway with a lower activation energy, so a greater proportion of collisions result in a reaction.

For example, manganese(IV) oxide (MnO2) catalyses the decomposition of hydrogen peroxide: 2H2O2(aq) → 2H2O(l) + O2(g).

• Catalysts lower activation energy → more effective collisions → faster reaction.
• Catalysts remain chemically unchanged after reaction.
• Example: MnO2 catalyses H2O2 decomposition.

Every chemical reaction requires a minimum amount of energy to start, called the activation energy (Ea). This energy is needed to break bonds in the reactants so that new bonds can form to produce products.

Activation energy can be visualised as an energy barrier that reactant particles must overcome for a reaction to occur.

The higher the activation energy, the slower the reaction tends to be at a given temperature because fewer particles have sufficient energy to react when they collide.

• Activation energy is the minimum energy required to start a reaction.
• It corresponds to the energy needed to break bonds in reactants.
• Reactions with high activation energy proceed more slowly at the same temperature.

A catalyst is a substance that increases the rate of a chemical reaction by providing an alternative reaction pathway with a lower activation energy.

Importantly, a catalyst is not consumed during the reaction; it remains chemically unchanged and can be used repeatedly.

By lowering the activation energy, a catalyst allows more reactant particles to have enough energy to react when they collide, thus increasing the reaction rate without changing the overall energy change (enthalpy) of the reaction.

• Catalysts provide an alternative pathway with lower activation energy.
• They increase reaction rate without being consumed.
• They do not alter the position of equilibrium or the overall energy change of the reaction.

Energy profile diagrams plot the potential energy of reactants and products during a reaction against the progress of the reaction.

Without a catalyst, the diagram shows a higher peak representing the activation energy barrier.

With a catalyst, the peak is lower, indicating the reduced activation energy required for the reaction to proceed.

• Activation energy is represented by the height of the energy barrier on the diagram.
• Catalysts lower this barrier, shown as a lower peak on the energy profile.
• The energy difference between reactants and products remains unchanged.

Consider the decomposition of hydrogen peroxide (H₂O₂) into water and oxygen:

2H₂O₂ (aq) → 2H₂O (l) + O₂ (g) (catalysed by manganese(IV) oxide, MnO₂)

Without a catalyst, this reaction is slow because of a high activation energy. Adding MnO₂ provides an alternative pathway with lower activation energy, speeding up the reaction without being consumed.

• MnO₂ acts as a catalyst, lowering activation energy.
• The reaction proceeds faster but the catalyst remains unchanged.
• No change in the overall enthalpy of the reaction.

In practical reactions, catalysts can be solids, liquids, or gases. For example, enzymes are biological catalysts that speed up reactions in living organisms.

Observations of catalysed reactions often show faster gas evolution, colour change, or temperature change compared to uncatalysed reactions.

Catalysts must be handled carefully; some may be toxic or require specific conditions to function effectively.

• Catalysts increase reaction rate observable by faster product formation.
• They remain chemically unchanged after the reaction.
• Examples include MnO₂ for H₂O₂ decomposition and enzymes in biological systems.

One practical method to measure the rate of a reaction is by recording the change in mass of the reaction mixture over time. This method is especially useful when a gas is produced and escapes from the reaction vessel, causing a decrease in mass.

The apparatus typically includes a conical flask containing the reactants placed on a digital balance. As the reaction proceeds, gas formed escapes, reducing the mass measured by the balance.

For example, when hydrochloric acid reacts with marble chips (calcium carbonate), carbon dioxide gas is produced: CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + CO₂(g) + H₂O(l). The mass decreases as CO₂ escapes. Recording mass at regular intervals (e.g., every 10 seconds) allows calculation of the rate of reaction as the change in mass per unit time (g/s).

• Use a digital balance with suitable precision (e.g., 0.01 g).
• Place the reaction flask securely on the balance.
• Start timing immediately after mixing reactants.
• Record mass at fixed time intervals until reaction slows or stops.
• Plot mass against time to analyse rate.
• Calculate rate as Δmass/Δtime (e.g., g/s).

Another common method involves measuring the volume of gas produced during a reaction over time. This is suitable when a gas is evolved and can be collected.

The apparatus includes a conical flask connected to a gas syringe or an inverted measuring cylinder in a water trough to collect gas by displacement.

For example, the reaction of zinc with dilute sulfuric acid produces hydrogen gas: Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g). The volume of hydrogen collected at regular intervals indicates the reaction rate. Recording gas volume every 10 seconds enables calculation of rate as volume per unit time (cm³/s).

• Set up a gas syringe or water displacement apparatus.
• Ensure all connections are airtight to prevent gas leaks.
• Start timing immediately after mixing reactants.
• Record gas volume at fixed intervals.
• Plot volume against time to determine rate.
• Calculate rate as Δvolume/Δtime (e.g., cm³/s).

To obtain reliable and reproducible results, it is essential to control variables such as temperature, concentration, and surface area of reactants.

Repeat measurements and calculate averages to reduce random errors.

Use appropriate apparatus to minimise systematic errors, e.g., a digital balance with fine resolution or a gas syringe with clear graduations and minimal friction.

• Keep temperature constant using a water bath if necessary.
• Use reactants of known and consistent concentration and mass.
• Use similar-sized pieces of solid reactants to ensure consistent surface area.
• Repeat experiments at least three times and calculate mean values.
• Avoid gas leaks by checking apparatus connections.
• Record data promptly and accurately.

Plotting the measured quantity (mass or gas volume) against time produces a curve that shows how the reaction progresses.

The initial rate of reaction can be estimated from the gradient (slope) of the curve at the start, where the reaction is fastest.

As the reaction proceeds, the rate decreases as reactants are used up, causing the curve to level off.

• Use graph paper or software to plot data points.
• Draw a tangent to the curve at t = 0 to estimate initial rate.
• Calculate gradient as change in y (mass or volume) over change in x (time).
• Compare rates under different conditions by comparing initial gradients.

Change in mass is straightforward and quick but requires that a gas is produced and escapes. It is less suitable if the gas dissolves in solution or if the reaction produces no gas.

Measuring gas volume provides direct quantification of gas produced and can be more accurate if apparatus is well set up. However, it requires careful sealing and can be affected by temperature and pressure changes.

Both methods require careful timing and control of variables to ensure valid comparisons.

• Mass change method is simple and requires minimal apparatus.
• Gas volume method allows direct measurement of gas quantity.
• Mass method may be affected by gas solubility or evaporation.
• Gas volume method requires airtight apparatus and consideration of gas laws.
• Both methods provide quantitative data to calculate reaction rates.