Presentation of Data and Observations

Experiment 1 is the determination of the enthalpy change of solution, $\Delta H_1$, for citric acid. This is the enthalpy change when one mole of citric acid dissolves in water.

Method

• Support one of the cups in the $250\text{ cm}^3$ beaker.
• Use the $50\text{ cm}^3$ measuring cylinder to transfer $30.0\text{ cm}^3$ of distilled water into the cup.
• Measure the temperature of the water in the cup. Record this temperature in the space for results.
• Weigh the container with FB 4. Record the mass.
• Tip all of the FB 4 into the water in the cup.
• Stir the mixture until the minimum temperature is obtained. Record this temperature.
• Weigh the container with any residual FB 4. Record the mass.
• Calculate and record the mass of FB 4 used.
• Calculate and record the temperature change.

Results

Calculate the energy change, in J, in your experiment.

Calculate the enthalpy change of solution, $\Delta H_1$, in $\text{kJ mol}^{-1}$, for dissolving $1.00\text{ mol}$ of solid citric acid, $\text{C}_6\text{H}_8\text{O}_7$, in water. Show your working.

Experiment 2 is the determination of the enthalpy change, $\Delta H_2$, for the reaction of one mole of solid citric acid with aqueous sodium hydroxide. In this experiment, aqueous sodium hydroxide, FB 6, is used in excess.

Method

• Support the second cup in the beaker.
• Add $4.80\text{--}5.00\text{ g}$ of FB 5 to the cup. Record your weighings.
• Measure and record the temperature of FB 6 in its container.
• Use the $50\text{ cm}^3$ measuring cylinder to transfer $50.0\text{ cm}^3$ of FB 6 into the cup with FB 5.
• Stir the mixture until the maximum temperature is obtained. Record the maximum temperature.
• Calculate and record the mass of FB 5 used.
• Calculate and record the temperature change.

Results

Calculations

Calculate the enthalpy change of reaction, $\Delta H_2$, in $\text{kJ mol}^{-1}$, for the reaction of $1.00\text{ mol}$ of solid citric acid, $\text{C}_6\text{H}_8\text{O}_7$, with aqueous sodium hydroxide. Show your working.

Use your values for $\Delta H_1$ and $\Delta H_2$ to calculate the enthalpy change, $\Delta H_r$, in $\text{kJ mol}^{-1}$, for the reaction of $1.00\text{ mol}$ of aqueous citric acid with aqueous sodium hydroxide.

Method

• Fill the burette with FA 4.
• Pipette $25.0\text{ cm}^3$ of FA 5 into a conical flask.
• Add FA 4 from the burette into the conical flask until the colour of the solution changes to yellow.
• Add 10 drops of FA 6 to the conical flask. Continue titrating until the blue-black colour just disappears.
• Perform a rough titration and record your burette readings in the space below.

Record, in a suitable form in the space below, all your burette readings and the volume of FA 4 added in each accurate titration.

From your accurate titration results, calculate a suitable mean value to use in your calculations. Show clearly how you obtain the mean value.

Give your answers to (c)(ii), (c)(iii), (c)(iv) and (c)(v) to an appropriate number of significant figures.

Calculate the amount, in mol, of sodium thiosulfate present in the volume of FA 4 in (b).

Calculate the amount, in mol, of iodine in $1.00\text{ dm}^3$ of FA 5.

Use the information given and your answer to (c)(iii) to calculate the amount, in mol, of iodine that reacted with sodium sulfite when solution FA 5 was prepared.

Use your answer to (c)(iv) to calculate the relative formula mass, $M_r$, of hydrated sodium sulfite.

Calculate the value of $x$.

Show your working.

Method

• Support a cup in the $250 \text{ cm}^3$ beaker.
• Use the $50 \text{ cm}^3$ measuring cylinder to transfer $30.0 \text{ cm}^3$ of FA 7 into the cup.
• Measure and record the temperature of the solution in the cup.
• Weigh the container with FA 6. Record the mass.
• Tip all of FA 6 into the cup containing FA 7. FA 7 is in excess.
• Stir the mixture until the maximum temperature is obtained. Record the maximum temperature.
• Weigh the container with any residual FA 6. Record the mass.
• Calculate and record the mass of FA 6 used.
• Calculate and record the temperature rise.

Calculate the energy released in your experiment.

energy released = ...................................................... J

Calculate the enthalpy change of reaction, $\Delta H_1$, in $\text{kJ mol}^{-1}$ of $\text{MgO}(\text{s})$, for the reaction of magnesium oxide with hydrochloric acid.
Show your working.

Method

• Place the other cup in the beaker.
• Use the $50 \text{ cm}^3$ measuring cylinder to transfer $30.0 \text{ cm}^3$ of FA 7 into the cup.
• Measure and record the temperature of the solution in the cup.
• Weigh the container with magnesium hydroxide, FA 8. Record the mass.
• Tip all of FA 8 into the cup containing FA 7. FA 8 is in excess.
• Stir the mixture until the maximum temperature is obtained. Record the maximum temperature.
• Weigh the container with any residual FA 8. Record the mass.
• Calculate and record the mass of FA 8 used.
• Calculate and record the temperature rise.

Calculate the enthalpy change of reaction, $\Delta H_2$, in $\text{kJ mol}^{-1}$ of $\text{Mg(OH)}_2(\text{s})$, for the reaction of magnesium hydroxide with hydrochloric acid.
Show your working.

Use your answers to (a)(iii) and (b)(ii) to calculate the enthalpy change, $\Delta H_r$, in $\text{kJ mol}^{-1}$, for the reaction between magnesium oxide and water. The equation for the reaction is shown.

Show your working.

Method

• Fill a burette with FB 1.
• Pipette $25.0 \text{ cm}^3$ of FB 2 into a conical flask.
• Add approximately 10 drops of FB 3.
• Perform a rough titration and record your burette readings in the space below. The end-point is shown by the appearance of a permanent blue colour.

The rough titre is .............................. $\text{cm}^3$.

• Carry out as many titrations as you think necessary to obtain consistent results.
• Make certain any recorded results show the precision of your practical work.
• Record, in a suitable form below, all your burette readings and the volume of FB 1 added in each accurate titration.

From your accurate titration results, calculate a suitable mean value to use in your calculations. Show clearly how you obtain the mean value.

$25.0 \text{ cm}^3$ of FB 2 required .............................. $\text{cm}^3$ of FB 1.

Give all your answers to (c)(ii), (c)(iii) and (c)(iv) to an appropriate number of significant figures.

Use your answer to (b) to calculate the amount, in mol, of sodium hydroxide, FB 1, titrated.

Hence, deduce the amount, in mol, of sodium hydrogen sulfate present in $25.0 \text{ cm}^3$ of FB 2.

Use your final answer to (c)(ii) to calculate the mass of sodium hydrogen sulfate present in $1.00 \text{ dm}^3$ of FB 2.

Use your answer to (c)(iii) and the information on page 2 to calculate the percentage purity by mass of the sodium hydrogen sulfate.

• Support a cup in the beaker.
• Use the $50\text{ cm}^3$ measuring cylinder to transfer $30.0\text{ cm}^3$ of FB 6 into the cup.
• Measure and record the temperature of the solution in the cup.
• Rinse the $25\text{ cm}^3$ measuring cylinder with water and then with a little FB 7.
• Use the $25\text{ cm}^3$ measuring cylinder to add $25.0\text{ cm}^3$ of FB 7 to the FB 6 in the cup.
• Stir the mixture.
• Measure and record the maximum temperature.
• Calculate and record the temperature rise.

Calculate the energy released in your experiment.
(Assume that $4.2\text{ J}$ change the temperature of $1.0\text{ cm}^3$ of solution by $1.0^\circ\text{C}$.)

energy released = .............................. J

Calculate the enthalpy change of reaction, $\Delta H_1$, in $\text{kJ mol}^{-1}$, for the neutralisation of $\text{NH}_3(\text{aq})$ with $\text{HCl}(\text{aq})$.
Show your working.

• Support a cup in the beaker.
• Rinse the $50\text{ cm}^3$ measuring cylinder with distilled water.
• Use the $50\text{ cm}^3$ measuring cylinder to transfer $30.0\text{ cm}^3$ of distilled water into the second cup.
• Measure and record the temperature of the water in the cup.
• Weigh the container with FB 8. Record the mass.
• Tip all of the FB 8 into the water in the cup.
• Stir until all FB 8 dissolves and record the minimum temperature observed.
• Calculate and record the temperature change.
• Weigh and record the mass of the container with any residual FB 8.
• Calculate and record the mass of FB 8 used.

Calculate the enthalpy change of solution, $\Delta H_2$, in $\text{kJ mol}^{-1}$, for FB 8, ammonium chloride.
(Assume that $4.2\text{ J}$ change the temperature of $1.0\text{ cm}^3$ of solution by $1.0^\circ\text{C}$.)

The values for the enthalpy changes of solution of ammonia and hydrogen chloride are given.

From your answers to (a)(iii), (b)(ii) and the data above, use Hess' Law to calculate the enthalpy change, $\Delta H_r$, in $\text{kJ mol}^{-1}$, for the reaction below.

Method

Preparing a solution of FA 1

• Weigh the stoppered container of $\text{FA 1}$. Record the mass in the space below.
• Tip all the $\text{FA 1}$ into the beaker.
• Reweigh the container with its stopper. Record the mass.
• Calculate and record the mass of $\text{FA 1}$ used.
• Add approximately $100\text{ cm}^3$ of $\text{FA 2}$ to the $\text{FA 1}$ in the beaker.
• Stir the mixture until all the $\text{FA 1}$ has dissolved.
• Transfer this solution into the $250\text{ cm}^3$ volumetric flask.
• Rinse the beaker and glass rod with distilled water and transfer the washings to the volumetric flask.
• Make up the solution in the volumetric flask to the mark using distilled water.
• Shake the flask thoroughly.
• This solution of the hydrated salt is $\text{FA 4}$. Label the flask $\text{FA 4}$.

Titration

• Fill the burette with $\text{FA 3}$.
• Pipette $25.0\text{ cm}^3$ of $\text{FA 4}$ into a conical flask.
• Use the $25.0\text{ cm}^3$ measuring cylinder to add $10\text{ cm}^3$ of $\text{FA 2}$ to the $\text{FA 4}$ in the conical flask.
• Perform a rough titration and record your burette readings in the space below.

The rough titre is .............................. $\text{cm}^3$.

• Carry out as many accurate titrations as you think necessary to obtain consistent results.
• Make sure any recorded results show the precision of your practical work.
• Record in a suitable form below all of your burette readings and the volume of $\text{FA 3}$ added in each accurate titration.

Keep FA 3 and FA 4 for use in Question 3.

From your accurate titration results, obtain a suitable value for the volume of $\text{FA 3}$ to be used in your calculations.
Show clearly how you obtained this value.

$25.0\text{ cm}^3$ of $\text{FA 4}$ required .............................. $\text{cm}^3$ of $\text{FA 3}$.

Calculate the number of moles of potassium manganate(VII) present in the volume of $\text{FA 3}$ calculated in (b).

moles of $\text{KMnO}_4$ = .............................. mol

$1\text{ mol}$ of $\text{KMnO}_4$ reacts with $5\text{ mol}$ of the hydrated salt, $\text{FA 1}$.

Calculate the concentration of the hydrated salt, in $\text{mol dm}^{-3}$, in $\text{FA 4}$.

concentration of $\text{FA 4}$ = .............................. $\text{mol dm}^{-3}$

Use your answer to (c)(ii), and your data on page 2, to calculate an experimentally determined value for the relative formula mass of the hydrated salt, $\text{FA 1}$.
Show your working.

$M_r$ of $\text{FA 1}$ = ..............................

Method

Dilution of FB 2

• Pipette $25.0\text{ cm}^3$ of FB 2 into the $250\text{ cm}^3$ volumetric flask.
• Make the solution up to the mark using distilled water.
• Shake the solution in the volumetric flask thoroughly.
• This solution of hydrochloric acid is FB 3. Label the volumetric flask FB 3.
• Rinse the pipette thoroughly.

Titration

• Fill the burette with FB 3.
• Pipette $25.0\text{ cm}^3$ of FB 1 into a conical flask.
• Add a few drops of bromophenol blue.
• Perform a rough titration and record your burette readings in the space below.

The rough titre is .............................. $\text{cm}^3$.

• Carry out as many accurate titrations as you think necessary to obtain consistent results.
• Make sure any recorded results show the precision of your practical work.
• Record in a suitable form below all of your burette readings and the volume of FB 3 added in each accurate titration.

From your accurate titration results, obtain a suitable value for the volume of FB 3 to be used in your calculations.
Show clearly how you obtained this value.

$25.0\text{ cm}^3$ of FB 1 required .............................. $\text{cm}^3$ of FB 3.

Give your answers to (ii), (iii), (iv) and (v) to the appropriate number of significant figures.

Calculate the number of moles of hydrochloric acid, $\text{HCl}$, in the volume of FB 3 calculated in (b).

Give the equation for the reaction of calcium hydroxide with hydrochloric acid.

............................................................................................................................................

Deduce the number of moles of calcium hydroxide that reacted with the hydrochloric acid in (c)(ii).

Calculate the concentration, in $\text{mol dm}^{-3}$, of calcium hydroxide in FB 1.

Calculate the mass of calcium hydroxide dissolved in $1.00\text{ dm}^3$ of limewater, FB 1.

Method

• Fill the burette with FA 3.
• Pipette $25.0\text{ cm}^3$ of FA 1 into a conical flask.
• Use the $25\text{ cm}^3$ measuring cylinder to add approximately $20\text{ cm}^3$ of FA 4 to the conical flask.
• Perform a rough titration and record your burette readings in the space below.

The rough titre is ............................ $\text{cm}^3$.

• Carry out as many accurate titrations as you think necessary to obtain consistent results.
• Make certain any recorded results show the precision of your practical work.
• Record, in a suitable form below, all of your burette readings and the volume of FA 3 added in each accurate titration.

From your accurate titration results, obtain a suitable value for the volume of FA 3 to be used in your calculations. Show clearly how you obtained this value.

$25.0\text{ cm}^3$ of FA 1 required ............................ $\text{cm}^3$ of FA 3.

Calculate the number of moles of manganate(VII) ions present in the volume of FA 3 calculated in (b).

moles of $\mathrm{MnO}_4^-$ = ............................ mol

Calculate the number of moles of hydrogen peroxide present in $25.0\text{ cm}^3$ of FA 1.

moles of $\mathrm{H}_2\mathrm{O}_2$ = ............................ mol

Using your answer to (ii) calculate the concentration, in $\text{mol dm}^{-3}$, of hydrogen peroxide in FA 1.

concentration of $\mathrm{H}_2\mathrm{O}_2$ in FA 1 = ............................ $\text{mol dm}^{-3}$

Method

• Fill the burette with FA 4.
• Pipette $25.0\text{ cm}^3$ of FA 1 into a conical flask.
• Use the measuring cylinder to add approximately $10\text{ cm}^3$ of FA 2 to the same conical flask.
• Use the measuring cylinder to add approximately $20\text{ cm}^3$ of FA 3 to the mixture in the conical flask. The mixture will now be a brown colour, due to iodine produced in the reaction.
• Begin your rough titration by adding FA 4 from the burette until the mixture becomes light brown.
• Add 10 drops of starch indicator. The mixture will become darker.
• Continue titrating until the mixture becomes an off-white colour. This is the end-point.
• Add one drop of starch indicator to check that no traces of dark colour are produced. If the mixture stays off-white, the titration is finished. If some dark colour is produced, because iodine is still present, continue the titration.
• Record your burette readings and the rough titre in the space below.

The rough titre is ............................. $\text{cm}^3$.

• Carry out as many accurate titrations as you think necessary to obtain consistent results.
• Make sure any recorded results show the precision of your practical work.
• Record in a suitable form below all of your burette readings and the volume of FA 4 added in each accurate titration.

Keep FA 3 and starch indicator for use in Question 3.

From your accurate titration results, obtain a suitable value for the volume of FA 4 to be used in your calculations.
Show clearly how you obtained this value.

The iodine produced required ............................. $\text{cm}^3$ of FA 4.

Calculate the number of moles of sodium thiosulfate, $\text{Na}_2\text{S}_2\text{O}_3$, in the volume of FA 4 calculated in (b).

moles of $\text{Na}_2\text{S}_2\text{O}_3$ = ............................. mol

Using your answer to (ii), calculate the number of moles of iodine that reacted with the number of moles of $\text{Na}_2\text{S}_2\text{O}_3$ calculated in (i).

moles of $\text{I}_2$ = ............................. mol

Using your answer to (iv) and the information on page 2, calculate the relative formula mass of the copper compound in FA 1.

$M_r$ of copper compound = .............................

Method

Read through the method before starting any practical work.
In the space below prepare a single table for your results of Experiments 1 and 2.

Experiment 1

• Weigh a crucible with its lid and record the mass.
• Add between $2.5\text{ g}$ and $3.0\text{ g}$ of FA 5 to the crucible. Weigh the crucible with FA 5 and lid and record the mass.
• Place the crucible on the pipe-clay triangle.
• Heat the crucible and contents gently for about two minutes, with the lid on.
• Remove the lid and continue heating gently for about three minutes.
• Replace the lid and leave the crucible and residue to cool for at least five minutes. Then reweigh the crucible and contents with the lid on. Record the mass.
• While the crucible is cooling, you may wish to begin work on Question 3.
• Calculate and record the mass of FA 5 used and the mass of residue obtained.
• State the observation(s) you made while the reaction was taking place.

.............................................................................................................................................

.............................................................................................................................................

Experiment 2

Repeat the method used in Experiment 1, using between $1.5\text{ g}$ and $2.0\text{ g}$ of FA 5 in the second crucible.

Results

Use your results from Experiment 1 to calculate the number of moles of copper oxide, $\text{CuO}$, obtained as residue.
Use the Periodic Table on page 12 for any data you may require.

moles of $\text{CuO}$ obtained in Experiment 1 = ............................. mol

Use your answer to (i), the equation on page 4 and the mass of FA 5 you used in Experiment 1, to calculate the relative formula mass, $M_r$, of malachite.

$M_r$ of malachite (from Experiment 1) = .............................

Use your results from Experiment 2 to calculate another value for the relative formula mass, $M_r$, of malachite.

$M_r$ of malachite (from Experiment 2) = .............................

Use data from the Periodic Table to calculate the relative formula mass, $M_r$, of malachite from its accepted formula, $\text{CuCO}_3\cdot\text{Cu(OH)}_2\cdot\text{H}_2\text{O}$.

$M_r$ of malachite (from formula) = .............................

If the relative formula mass of malachite obtained from either of your experiments is within $2.5\%$ of the answer in (iv), this is good evidence that the accepted formula, $\text{CuCO}_3\cdot\text{Cu(OH)}_2\cdot\text{H}_2\text{O}$, is correct.

Show by calculation whether either of your experiments supports the accepted formula.