Reaction Rates (OCR AS Chemistry)

The reaction of calcium carbonate with dilute hydrochloric acid results in the formation of a gas.

The effect of increasing the concentration of the acid on the rate of reaction can be determined by measuring the amount of gas produced at regular intervals for several minutes.

The following method was used to measure the rate of the reaction of CaCO₃ with HCl:

The results for the reaction of 2.5 g of CaCO₃ with 50 cm³ of 0.050 mol dm^-3 dilute HCl are shown below. 

Using the results given in the table, plot a graph of time (s) on the x-axis and volume of CO₂ (cm³) on the y-axis.

Draw a line of best fit on the graph

Use the results and your graph from part (c) to calculate the initial rate of reaction.

Give your answer to 2 decimal places and state the correct units.

A Maxwell-Boltzmann distribution is a graph that shows the distribution of energy amongst particles within a chemical reaction. The Maxwell-Boltzmann distribution for a sample of gas at a fixed temperature, T₁ is shown.

State why a Maxwell-Boltzmann distribution curve always starts at the origin and what the area under the curve represents.

Chemical reactions take place at different speeds. For a chemical reaction to take place, particles must collide with each other in the correct orientation and with sufficient energy. 

Give one reason why a reaction may be slow at room temperature.

A platinum-rhodium catalyst is used in the reaction forming nitrogen (II) oxide from ammonia.

A group of students were completing a practical, investigating the factors which affect the rate of the following chemical reaction. 

A (s) + B (aq) → C (g)

The students collected the gas produced and plotted the graph shown below. 

Explain why the gradient of the curve in part (b) decreases as the time of the reaction progresses.

Another way to increase the rate of reaction is to increase the temperature. 

Explain why a small increase in temperature has a large effect on the initial rate of a chemical reaction.

The decomposition of hydrogen peroxide into water and oxygen is a very slow chemical reaction.

Write the equation for the decomposition of hydrogen peroxide. 

The rate of decomposition of hydrogen peroxide can be ascertained by collecting and measuring the volume of gas formed at specific time intervals.

Two students set up the practical apparatus from part (b) to measure the rate of decomposition of hydrogen peroxide. Student A states that one way to increase the rate of the decomposition, would be to increase the concentration of hydrogen peroxide. 

Is student A correct? Explain your answer.

The decomposition of hydrogen peroxide is a slow reaction, so a catalyst is often added to speed up the rate of the reaction. Catalysts are used in many chemical reactions to increase the rate. 

The following shows a two-step reaction mechanism of a chemical reaction, where a catalyst, X is used. 

   STEP 1:     W + X  →  Y + Z

   STEP 2:     Y + W → Z + A + X

   OVERALL REACTION:     2W → 2Z + A 

Give a reason, other than the rate of reaction increasing, why it can be deduced from the three equations above that X is a catalyst.

During the following reaction, A and B react together to produce C. 

A  +  2B  →  C

The production of C over time is shown below.

In the reaction in part (a), lumps of A were used.
Use collision theory to explain what would happen to the rate of the reaction if powdered A was used instead of lumps. 

Hydrogen peroxide decomposes more rapidly in the presence of aqueous hydrogen bromide. The decomposition proceeds as shown by the following equations.

H₂O₂ + HBr  →  HBrO + H₂O

HBrO + H₂O₂ →  H₂O + O₂ + HBr

Give two reasons, other than an increase in the reaction rate, why these equations suggest that hydrogen bromide is behaving as a catalyst.

In the Haber Process, nitrogen and hydrogen are reacted together in a sealed container to form ammonia, NH₃. Increasing the pressure and adding a catalyst both speed up the rate of this reaction.

This question is about the reaction between sodium thiosulfate, Na₂S₂O₃, and dilute hydrochloric acid.

When a solution containing sodium thiosulfate ions is mixed with dilute acid, sodium chloride, water, sulfur dioxide, which is readily soluble in water, and a precipitate of sulfur are formed. The rate of reaction can be investigated using a colorimeter which measures the amount of light absorbed by the solution.

A student used a colorimeter to investigate how the rate of reaction varies with the concentration of thiosulfate ions. Different volumes of a solution containing thiosulfate ions were each made up to 10 cm³ with distilled water. Each mixture was placed into the colorimeter and 5 cm³ of dilute acid was added. The time taken for the intensity of the light passing through the reaction mixture to fall to 75% of the original intensity was measured for each mixture. 

The results of the experiment are given in the table below:

The rate of reaction can be calculated by .

Use the graph drawn in part (b) to:

For the solution containing 3 cm³ of thiosulfate ions, the student accidentally used an acid that was half the concentration than the acid used in the rest of the investigation.

What effect would this have on the rate of reaction?

Explain your answer.

A student is asked to plan an investigation to determine the initial rate of reaction of magnesium and hydrochloric acid.

The student plans to collect approximately 84 cm³ of hydrogen gas at room temperature and pressure and use an excess of magnesium.

They choose to use the apparatus shown in the diagram and the following:

• 100 cm³ measuring cylinder
• 2 decimal place balance
• Stop clock

Give an outline of how the student could plan to carry out this investigation. In your answer, you should explain how the results will be used graphically.

Show all working in any calculations.

The graph below shows the expected results from the investigation planned in part (a).

The student repeated the experiment using:

• The same temperature and pressure
• An acid with twice the concentration but half the volume 
• Twice the mass of magnesium  

On the graph, draw a line to show the results you would expect in this experiment. 

Explain why, using collision theory, the graph obtained from the investigation in part (a) shows that the reaction decreases over time.

Copper is a relatively unreactive metal but will react with hydroiodic acid, HI (aq), to form hydrogen gas:

Cu (s) + 4HI (aq) → 2H[CuI₂] (aq) + H₂ (g)

A student investigated the reaction by connecting a gas syringe to a conical flask. At room temperature and pressure, they placed 100 cm³ of 0.05 mol dm^-3 of dilute hydroiodic acid into the flask. An excess of copper granules were added and the flask immediately stoppered. The volume of gas collected was recorded at 20 second intervals until the reaction was complete. 

They obtained the graph shown below.

Use the graph to find:

The rate of reaction doubles approximately every 10 ^oC increase in temperature. By what factor would the rate of reaction increase by if the reaction was carried out at 70 ^oC compared with the rate of reaction at 20 ^oC?

This question is about changes in rate of reaction and catalysts.

Reaction rates can be affected by factors, such as changes in pressure and temperature.

Hydrogen iodide is used in the manufacturing of pharmaceuticals and can be broken down back into iodine and hydrogen. 

2HI (g)  H₂ (g) + I₂ (g)                     ΔH = - 52 kJ mol^-1

The activation energy when uncatalysed is +183 kJ mol^-1

• reactants and products
• curves for both the catalysed and uncatalysed reactions
• activation energy for the uncatalysed forward reaction and the catalysed forward reaction
• enthalpy change of reaction, ΔH.

Catalysts are often used in industrial processes. 

The Contact process is an important industrial multi-staged process, used in the production of sulfuric acid. During the second stage, solid vanadium (V) oxide, a heterogeneous catalyst, is used to make sulfur trioxide from sulfur dioxide and oxygen. This process is reversible.

Methanol, for use as a fuel, can be produced by the reaction of carbon monoxide with hydrogen. The reaction is typically carried out at 310 °C and 3.25 × 10⁷ Pa with a catalyst which is a mixture of zinc oxide and chromium(VI) oxide. The Boltzmann distribution for 1 dm³ of the reaction mixture is shown by sample 1 in the diagram below.

CO (g) + 2H₂ (g) → CH₃OH (g)              ΔH =  – 90 kJ mol^–1

A change was made to the reaction conditions as shown by sample 2 in diagram above.

Consider the Boltzmann distribution curve below.

For each of the changes in parts (i), (ii) and (iii) below, state and explain the effect that the change would have on:

• The area under the curve 
• The position of the peak of the curve
• The proportion of molecules with energy greater than or equal to E_a

A chemist performed a reaction at three different temperatures, 100K, 300K and 700K as shown by the Boltzmann distribution graph below.

Hydrogen will react with chlorine to form the hydrogen halide, hydrogen chloride, a colourless gas.

H₂ (g) + Cl₂ (g) → 2HCl (g)