Further Organic & Inorganic Chemistry Practicals (OCR A Level Chemistry)

A student has two unlabelled samples of colourless liquid. The samples are known to be cyclohexanol, C₆H₁₁OH, and t-butanol, (CH₃)₃COH.

Describe a chemical test to distinguish between cyclohexanol and t-butanol. Your answer should explain the expected observations.

Cyclohexanol and t-butanol both undergo dehydration reactions by heating in the presence of sulfuric acid.

Identify the carbocations formed during the dehydration reactions.

Use the carbocations to explain why t-butanol requires a lower temperature for the reaction to occur.

A student has correctly identified the cyclohexanol sample.

They set up the following equipment to determine the boiling point of the cyclohexanol.

Explain three ways that the student could improve their equipment.

The purity of a liquid can be determined by measuring its boiling point. An alternative method to distillation would be to boil a tube of the sample in an oil bath.

The expected boiling point of a cyclohexanol sample is 161 °C but following distillation it was found to be 168 °C.

Suggest why the boiling point was higher than expected and why measuring the boiling point is not the most accurate method to confirm the identity of a known sample.

In a multi-step reaction, benzoic acid can be prepared by hydrolysing ethyl benzoate using sodium hydroxide.

Write an equation for this reaction.

A student used the following method to make benzoic acid.

1. 1. Place 3.0 cm³ of ethyl benzoate (density = 1.05 g cm⁻³, M_r = 150.0 g mol^-1) in a pear-shaped flask
   2. Add 25 cm³ (an excess) of sodium hydroxide solution.
   3. Set the pear-shaped flask up for reflux.
   4. Heat the mixture, under reflux, for approximately 30 minutes.
   5. Leave the mixture to cool to room temperature.
   6. Add 50.0 cm³ of 2 mol dm⁻³ hydrochloric acid
   7. Collect the crude product using suction filtration.

The crude product mixture from part b will be purified by recrystallisation. Describe how to identify a suitable solvent for the recrystallisation process.

After successful recrystallisation, the melting point of the benzoic acid was determined to assess the purity of the sample. The expected melting point was 122.3 °C.

A standard hydrogen electrode can be used to determine the standard electrode potential of a Cu (s) / Cu²⁺ (aq) electrode.

Draw a labelled diagram of the apparatus which would be used to determine the standard electrode potential of Cu (s) / Cu² (aq). 

State the necessary conditions for this cell, which would allow the standard electrode potential to be measured.

The table shows the half equations and standard electrode potentials of some half cells. 

An experiment was carried out to determine the e.m.f of this cell.

Fe (s) |  Fe³⁺ (aq)  ||  H⁺ (aq) |  H₂ (g)  |  Pt (s)

In a Daniell cell, zinc is a more reactive metal than copper and this type of cell can be used to provide electrical energy. The conventional representation for the Daniell cell is;

Zn (s) | Zn²⁺ (aq)  ||  Cu²⁺ (aq) | Cu (s)

A student prepared a sample of aspirin by refluxing salicylic acid with an excess of acetic anhydride in the presence of a phosphoric (V) acid catalyst. The reaction was quenched with water, cooled and crystallised. The crude product was collected by suction filtration, washed and then purified by recrystallisation. Thin-layer chromatography was then used to ascertain whether the product had been sufficiently purified by recrystallisation. 

Outline the method that the student would use to determine whether the product was pure, including the results that the student would expect to see. You should include a labelled diagram to support your answer.

Outline how the student could identify the positions of any compounds on the TLC plate and describe the expected results to show that the crude product had been purified by recrystallisation.

There is an extensive range of over the counter and prescription medicines, that contain a variety of chemicals. TLC is often one of the first methods used to analyse medicines.

Give three advantages of using TLC in the analysis of such medicines.

The structure of aspirin is shown below.

Standard over the counter aspirin tablets are around 85% pure. One of the common impurities found in aspirin is ethanoic acid from the manufacturing process. A student analysed a 325 mg aspirin tablet and found it to contain 1.56 x 10^-3 moles of aspirin.

Evaluate the safety of the tablet for human consumption.

Aspirin, shown below, can be synthesised by the acylation of 2-hydroxybenzoic acid using either ethanoyl chloride or ethanoic anhydride.

Write an equation to show the formation of aspirin using ethanoic anhydride.

State two reasons why ethanoic anhydride is used in preference to ethanoyl chloride for the acylation.

A student uses the following method to make aspirin (Mr 180.158 g mol^-1) in the school laboratory.

1. 1. Place 1.50 g of 2-hydroxybenzoic acid (M = 138.121 g mol^-1) and 3.00 cm³ of ethanoic anhydride (M = 102.089 g mol^-1) in a 50.0 cm³ pear-shaped flask.
   2. Add 4 drops of phosphoric (V) acid and swirl to mix.
   3. Set the pear-shaped flask up for reflux.
   4. Heat the mixture with a boiling water bath, for approximately 5 minutes.
   5. Carefully, add 2.00 cm³ of water down the reflux condenser, without cooling.
   6. When the reaction is complete, pour the mixture into a 100 cm³ beaker containing 30.0 cm³ of cold water.
   7. Stir the mixture and rub the sides of the beaker with a stirring rod to promote crystallisation.
   8. Stand the mixture in an ice bath to complete crystallisation.
   9. Collect the product using suction filtration.

Calculate the expected mass of aspirin, assuming the student achieves an 85% yield in their crude product. You may assume that the density of ethanoic anhydride is 1.00 g cm^-3.

In your answer, you should identify the limiting reagent.

Describe how the student could purify the aspirin after it has been filtered.

This question is about the preparation of aspirin.

The salicylic acid required to prepare aspirin must be pure. Describe how the salicylic acid can be purified and explain why each step is required. 

A student prepared a sample of aspirin by following the two stages outlined:  

• Add a known mass of the purified salicylic acid into a dry pear shaped flask
• Add an excess of 5 cm³ of ethanoic anhydride
• Add by 8 drops of concentrated phosphoric acid
• Attach a condenser to the flask
  In a fume cupboard,  warm the mixture with swirling until the solid has dissolved
• Continue to warm for another 5 minutes

• Carefully, add 5 cm³ of cold water to the solution
• Stand the flask in cold water to promote precipitation
• Filter off the product and wash the product with a little cold water
• Transfer the product to a watch glass and leave to dry overnight
• Once dry, weigh the product and record the mass.  

Suggest four improvements to the method the student followed.

Draw a labelled diagram of the apparatus used to filter the mixture in Stage 2 outlined in part (b).

The following reaction scheme is for the preparation of aspirin.

The average percentage yield for the industrial synthesis of aspirin is around 79.2%. A student uses 1.05 g of salicylic acid with an excess of ethanoic anhydride for their aspirin synthesis. 

Calculate the minimum mass of aspirin that they would need to produce to meet industrial standards.

This question is about the preparation of methyl 3-nitrobenzoate.

Outline a method for the nitration of methyl benzoate. Include safety considerations and equations for the reaction. You do not need to discuss purification of the crude product. 

Describe the purification process of the product using a suitable solvent.

A student tested the methyl 3-nitrobenzoate product produced in part (a) and found the melting point to be lower than expected 78 ℃ data book value.

 Explain how the student can obtain a more accurate value.

Figure 1 shows a student's diagram of the standard hydrogen electrode.

Figure 1

Identify the error that the student has made and explain why it is incorrect.

The approximate cost of platinum in the UK is £22 500 per kg.

Explain how the design of the platinum electrode minimises the cost while maximising its efficiency. 

A student set up an electrochemical cell as shown in Figure 2.

Figure 2

The student made the silver electrode by adding 7.78 g of silver sulfate to 25.0 cm³ of water at 298 K. 

The solubility of silver sulfate in water at 298 K is 4.73 g dm^-3.

Explain if the setup in Figure 2 can be used to measure the standard electrode potential of the Ag / Ag⁺ half cell. Show your working.

The cell represented in Figure 3 was set up under standard conditions.

Figure 3

Explain why this electrochemical cell cannot measure the standard electrode potential of the Fe²⁺ / Fe³⁺ half cell.

Tollens’ reagent or diammine silver hydroxide, [Ag(NH₃)₂]⁺OH^-, is a mixture of liquid ammonia and silver nitrate. Tollens’ reagent is prepared in the laboratory in two steps. 

The first step is to react silver nitrate and sodium hydroxide to form a brown precipitate of silver(I) oxide. This precipitate is then dissolved in concentrated ammonia to form the reagent. 

Write two equations, including state symbols, to show the laboratory formation of Tollens’ reagent.

Write an equation, including state symbols and give the observation for the reaction of Tollens’ reagent with benzaldehyde. 

Suggest why the student should only use clean and dry glassware for this practical. 

Potassium dichromate is less expensive than Tollens' regent and can also be used to test for aldehydes.

Explain why Tollens' reagent is specifically used. 

A mixture of ethyl benzoate 5 cm³ and sodium hydroxide solution 15 cm³ was refluxed in a round bottomed flask fitted with a water condenser on a water bath at a temperature of 90 °C  for 30 minutes. A salt was formed in the reaction.

The solution was then cooled and acidified with hydrochloric acid. The resultant acidified solution was then cooled in an ice bath forming a precipitate of benzoic acid.

State the equation to form the salt. 

Describe the purification process of the product using a suitable solvent.

According to a data book, the melting point of benzoic acid is 122.3 °C. 

The product formed in part (a) was found to have a lower melting point than expected.

Explain how the student can obtain a more accurate value.

The mass of benzoic acid produced in the reaction was 1.2 g and the percentage yield of benzoic acid formed was calculated to be 80.53%.

Calculate the mass of ethyl benzoate that was used in the reaction.