Physical Chemistry Practicals (AQA A Level Chemistry)

Iron is extracted from iron (III) oxide in the Blast Furnace by reducing it with carbon. However, the iron which is extracted is still impure. It contains small amounts of carbon which can be removed by blowing oxygen through the impure iron.

To analyse the purity of the molten iron, a redox titration can be carried out. A sample of the iron is reacted with dilute hydrochloric acid, to form iron (II) chloride and a gas.

A sample of the formed solution is then reacted with potassium manganate (VII).

The following steps were carried out to analyse the amount of carbon impurities present in the extracted iron.

Step 1 - Take samples of the iron at different stages of the process

Step 2 - React the samples of impure iron that have been collected with an excess of dilute hydrochloric acid to form iron (II) chloride solution.

Step 3 - Make up each solution to 100 cm³.

Step 4 - From each 100 cm³ sample solution which is made, take a 50.0 cm³ portion of and reacts it with a solution of potassium manganate (VII).

At one stage of the analysis, a 1.39 g sample of impure iron was reacted with an excess of dilute hydrochloric acid and the solution was made up to 100 cm³.

A 50.0 cm³ portion of this solution was reacted with 21.7 cm³ of a 0.0962 mol dm^-3 solution of potassium manganate (VII).

Using your answers to part (c) suggest whether the scientist should stop blowing oxygen at this stage. Explain your answer.

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:

Step 1 - Support a gas syringe with a stand, boss and clamp

Step 2 - Measure 50 cm³ of 0.050 mol dm^-3 dilute HCl in a beaker and then add it to a conical flask

Step 3 - Add about 2.5 g of CaCO₃ to the flask using a spatula, immediately connect the gas syringe and start counting.

Step 4 - Record the volume of gas produced every 20 seconds for 2 minutes.

Step 5 - Repeat Steps 1-4 for different concentrations of HCl.

The results for the reaction of 2.5 g of CaCO₃ with 50 cm³ of 0.050 mol dm^-3 dilute HCl were recorded in Table 2 below.

Table 2

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.

Figure 1

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.

The enthalpy of combustion of an unknown liquid fuel can be determined by carrying out a calorimetry experiment.

A student used the following method to carry out the experiment.

Step 1 - Measure 100 cm³ of water into a calorimeter using a measuring cylinder.

Step 2 - Record the initial temperature of the water using a thermometer.

Step 3 - Heat the water using the flame from the burning fuel and record the final temperature.

Step 4 - Measure the final mass of the spirit burner with the liquid fuel.

The student obtained the following results:

• Mass of water in calorimeter = 350 g
• Initial temperature of water = 9 C
• Final temperature of water = 36 C
• Mass of liquid fuel burned = 3.15 g

The actual value of ΔH_c of the unknown liquid is greater than the value calculated in part (c). Give two reasons for this.

Lactic acid is a chemical byproduct of anaerobic respiration, with the chemical formula C₃H₆O₃. Bacteria in yoghurt will produce it, as will bacteria in our stomach. Bacteria in beer will also ferment glucose, producing lactic acid by anaerobic respiration. The increase in lactic acid decreases the pH of the beer and gives it its characteristic sour taste.

A student devised an experiment to investigate the concentration of lactic acid in a sample of beer.

Step 1 - Transfer 15 cm³ of beer to a volumetric flask using a beaker.

Step 2 - Make up the volume of the beer to 300 cm³.

Step 3 - Use a pipette to take a portion of 30.0 cm³ and add it to a conical flask.

Step 4 - Give the conical flask an initial swirl and then place on the bench.

Step 5 - Add a few drops of phenolphthalein indicator to the lactic acid in the conical flask.

Step 6 - Titrate samples of 30.0 cm³ of this solution with a 0.0750 mol dm³ solution of sodium hydroxide solution.

Complete Table 1 below.

Table 1

The results of the titration are shown in Table 2.

Table 2

The burette used by the student for their titration has an error of ± 0.05 cm³. 

Suggest the maximum percentage uncertainty in using this piece of equipment using your mean titre calculated in part (b).

Lactic acidosis occurs when there's too much lactic acid in your body. Causes of lactic acidosis may include chronic alcohol use and can lead to muscle cramps or pain, body weakness and headaches.

Lactic acidosis may occur when the blood levels of lactic acid are 4.0 x 10^-3 mol dm^-3 or greater.

C₆H₅COOH + NaOH → C₆H₅COONa + H₂O

Another student is required to make up 300 cm³ of an aqueous solution that contains a known mass of lactic acid. The student is provided with a sample bottle containing the lactic acid.

Describe the method, including full apparatus and practical details, that the student should use to prepare this solution.

A student investigates the rate of reaction between copper sulfate solution and dilute sodium hydroxide solution.

75 cm³ of copper sulfate solution was poured into a 250 cm³ conical flask using a measuring cylinder.

The conical flask was then placed on a cross drawn on a piece of paper. 15 cm³ of dilute sodium hydroxide was added to the conical flask and the timer was started as shown in Figure 1.

Figure 1

The timer was started, and then was stopped when the cross could not be seen anymore. The student then recorded the time.

The student wants to investigate the effect of temperature on the rate of the reaction described in part (a).

The student carries out the following steps:

Step 1 - Measure out copper sulfate using a beaker and pour it into a conical flask.

Step 2 - Record the initial temperature of the copper sulfate in the conical flask using a thermometer.

Step 3 - Place the flask with the solution onto a cross drawn on a piece of paper.

Step 4 - Heat the sodium hydroxide to a specific temperature.

Step 5 - Start the stopwatch and then add the dilute sodium hydroxide.

Step 6 - Stop the stopwatch when the cross can no longer be seen.

Step 7 - Repeat the experiment at different temperatures.

From the experiment, the student made the following conclusion:

“The rate of reaction increases with increasing temperature”

Explain whether the student’s conclusion is correct using the collision theory.

Compounds A to F, as shown in Figure 1, are a 50 : 50 mixture of polar and nonpolar covalent molecules.

Figure 1

Table 1 shows the electronegativity values for some elements.

Table 1

Use the information to suggest a maximum value for the difference in electronegativity which allows a covalent molecule to be classified as nonpolar. Justify your answer.

Give your answer to 2 decimal places.

The displayed formulae for methane, dibromomethane and tetrabromomethane are shown in Figure 2.

Figure 2

A student concluded that all three compounds were nonpolar because they are symmetrical.

Use your knowledge of structure and bonding to evaluate the student’s conclusion.

Student 1 suggests that compounds B and D, from Figure 1, are nonpolar molecules; student 2 suggests that they are polar.

Explain how both students could be considered to be correct. You may draw Lewis structures with partial charges to illustrate your explanation.

Cyclohexane, ethanol, propanone and water are placed in separate burettes.

Each liquid is tested by opening the burette to provide a steady stream of the liquid. A woolen cloth is used to transfer electrons to a plastic ruler. The ruler is then held near the stream of liquid, as shown in Figure 3, to see the deflection, if any, that the charged plastic ruler causes to the stream of liquid.

Figure 3

Place the four liquids in order from 1 - the least to 4 - the most deflected. Use information from Table 1 in part (a) to help explain your order.

A student is set the challenge of creating a sodium hydroxide solution with a concentration of exactly 0.12 mol dm^-3 using limited chemical resources. They do have access to any glassware that they require.

The chemical resources available to the student are:

• Approximately 0.21 mol dm^-3 sodium hydroxide solution
• 15 mol dm^-3 nitric acid
• Phenolphthalein

The student sets up a titration experiment as shown in Figure 1.

Figure 1

The student completes a rough titration, as shown in Figure 1 and records a value of 36.20 cm³. They then refill the burette with nitric acid.

State the errors the student has made and explain their effects on the titre values.

The student determines that the actual concentration of the sodium hydroxide solution is 97.57% of the approximate concentration.

One of the two concordant results that the student uses for their titration calculation is 34.20 cm³.

Use the information in part (a) to determine the volume of the student’s other concordant result. Show your working.

The student wants to convert 100 cm³ of a 0.18 mol dm^-3 sodium hydroxide solution into a 0.12 mol dm^-3 sodium hydroxide solution.

Calculate the volume of water that they need to add.

Another student suggests that an alternative method to make the 0.12 mol dm^-3 solution is to evaporate the water from the original, approximately 0.21 mol dm^-3 solution.

Once the sodium hydroxide has been recovered from the original solution, then the required mass of sodium hydroxide to make a 0.12 mol dm^--3 solution can be weighed out and dissolved in a standard flask.

Explain how the hygroscopic nature of sodium hydroxide could affect the final concentration of the solution.

The reaction between sodium thiosulfate solution and hydrochloric acid is commonly investigated by measuring the time taken for a cross placed underneath the reaction mixture to disappear, as shown in Figure 1.

Figure 1

A student completes an investigation into the effect of temperature on the rate of reaction for sodium thiosulfate with hydrochloric acid. Their results are shown in Table 1.

Table 1

Plot a graph of reaction rate against temperature on Figure 2.

Figure 2

Use your knowledge of Maxwell-Boltzmann distribution curves to explain the effect of temperature on the reaction between sodium thiosulfate and hydrochloric acid, as shown in your answer to part (b).

A student suggests that the initial rate of the reaction between sodium thiosulfate and hydrochloric acid could be measured by collecting the sulfur dioxide produced in a 250 cm³ gas syringe.

The student’s suggested method is as follows:

1. Place 100 cm³ of 0.05 mol dm^-3 of sodium thiosulfate solution into a conical flask
2. Add 10 cm³ of 3.0 mol dm^-3 hydrochloric acid into the conical flask
3. Connect the conical flask to the gas syringe with a delivery tube and start the timer
4. Record the volume of gas produced every 30 seconds until the volume of gas has remained constant for 2 minutes
5. Plot and use a graph of the results to estimate the initial rate of reaction

Using your equation from part (a), justify whether or not this method would be effective in producing valid results in a school laboratory.

An experiment is set up, as shown in Figure 1, to determine the formula of an unknown sample of hydrated copper sulfate, CuSO₄.xH₂O.

Figure 1

The hydrated copper sulfate is carefully heated, until constant mass, to drive off the water of crystallisation. The water is then condensed in a test tube held in a beaker of ice water.

Explain why the mass of water in the test tube cannot be relied upon to determine the mass of water for a water of crystallization calculation.

An 8.0 g sample of the hydrated copper sulfate is heated to constant mass, using the equipment shown in part (a).

Student A predicts that the formula of the hydrated copper sulfate will be CuSO₄.5H₂O.

Student B predicts that the hydrated copper sulfate will contain 7 waters of crystallisation.

Calculate the constant mass that each student would expect to see to prove their prediction.

When the 8.0 g of hydrated copper sulfate is heated, it achieves a constant mass of 4.77 g. Both students from part (b) observe thin white fumes from the top of the test tube while the water of crystallisation is condensing.

Suggest a formula for the hydrated copper sulfate and explain which student is likely to be correct in their prediction.

The reaction scheme for the formation of hydrated copper sulfate from its anhydrous form is shown in Figure 2.

Figure 2

Ethanoic acid and ethanol can undergo a condensation reaction to form ethyl ethanoate and water according to the following equation.

CH₃COOH + C₂H₅OH ⇋ CH₃COOC₂H₅ + H₂O

12.0 g of ethanoic acid and 11.5 g of ethanol were mixed in a stoppered bottle. The reaction mixture was left to reach equilibrium at room temperature. The mixture was then poured into a volumetric flask and made up to a total volume of 250 cm³.

A 25.0 cm³ sample of the resulting mixture required 26.5 cm³ of 0.200 mol dm^-3 sodium hydroxide to neutralise the remaining ethanoic acid.

Calculate the number of moles of ethanoic acid in the reaction mixture at equilibrium.

Use the information in part (a) to calculate a value for K_C for the reaction, including units.

Give your answer to the appropriate number of significant figures.

The reversible reaction of ethanol with ethanoic acid is exothermic in the forward direction. Explain the effect that an increase in temperature will have on the value of K_C.

There is an equilibrium between chemicals X and Y is as shown:

X ⇋ Y

At 100 C, the value of K_C for the reaction is 39. At the start of the reaction there are 2.0 moles of X. Calculate the number of moles of X at equilibrium.