Physical Chemistry Practicals

Section 1 of 26PAG 1 — three ways of finding moles

PAG 1.1

• PAG 1.1 determination of the composition of basic copper(II) carbonate.

PAG 1.2

• PAG 1.2 determination of the relative atomic mass of magnesium.

PAG 1.3

• PAG 1.3 determination of the formula of magnesium oxide.

Common idea

• Common idea all three use experimental data to calculate moles and then the final answer.
• Other practicals may replace the suggested ones, so the same knowledge must be applied to new examples.

Section 2 of 26Measuring gas volumes or changes in mass

Measuring gas volumes or changes in mass

Gas volume method

• Gas volume method a reaction produces a gas whose volume is measured.
• The moles of gas are calculated and the stoichiometric ratio gives the moles of the other substance.
• The final answer could be the composition of copper(II) carbonate or the relative atomic mass of magnesium.

Mass change method

• Mass change method a reaction causes an increase or decrease in mass.
• Magnesium of known mass is burned and the mass of magnesium oxide formed is measured.
• The moles of magnesium and oxygen give the ratio of each element in the formula.

Hydrated salts (not a PAG practical)

• Hydrated salts (not a PAG practical) heat a salt such as CuSO₄·xH₂O to constant mass.
• The decrease in mass gives the moles of water of crystallisation.

Section 3 of 26PAG 1.1 — method and calculation

1

Set up a measuring cylinder in a water bath connected to a conical flask.

2

Weigh a known mass of CuCO₃·Cu(OH)₂ and add it to the conical flask.

3

Add excess sulfuric acid and insert the stopper immediately so the gas collects in the measuring cylinder.

4

Record the final volume of carbon dioxide.

Alternative apparatus

• Alternative apparatus a gas syringe can measure the gas volume instead.

Moles of gas

• Moles of gas n = V(dm³) / 24.0 or n = V(cm³) / 24 000.
• 24.0 dm³ is the volume of 1 mol of gas at room temperature and pressure (101 kPa).
• Under other conditions use the rearranged ideal gas equation n = PV / RT.

Equation

• CuCO₃ + H₂SO₄ → CuSO₄ + H₂O + CO₂, so CuCO₃ : CO₂ is 1 : 1.

Mass of CuCO₃

• Mass of CuCO₃ m = n × M_r.
• The sample contained CuCO₃·Cu(OH)₂, so the percentage by mass of CuCO₃ in the original sample is then calculated.

Section 4 of 26PAG 1.1 — sources of error and fixes

Source of error	Improvement
CO2 escapes before the bung is inserted	Insert the bung immediately after adding the acid.
Some copper carbonate does not react	Swirl the flask to mix the contents and use excess acid so the reaction goes to completion.
CO2 dissolves in the water	Use a gas syringe to measure the gas volume.
Solid remains on the weighing apparatus	Weigh the apparatus with and without the solid; the difference is the mass actually added.

Section 5 of 26PAG 1.2 — relative atomic mass of magnesium

1

Set up a measuring cylinder in a water bath connected to a conical flask, as in PAG 1.1.

2

Add a known mass of magnesium to the conical flask.

3

Add excess sulfuric acid and insert the stopper immediately.

4

Record the final volume of hydrogen collected.

Calculation

• Calculation n(H₂) = V(dm³) / 24, and Mg + H₂SO₄ → MgSO₄ + H₂ gives Mg : H₂ = 1 : 1.
• The relative atomic mass is then found from A_r = mass of Mg / moles of Mg.

Errors

• Errors hydrogen may escape before the bung is inserted, so insert it as soon as the acid is added.
• Some magnesium may not react, so use magnesium powder, excess acid and swirling.
• Magnesium left on the weighing apparatus is avoided by weighing the apparatus with and without it.

Mass loss cannot be used

• Mass loss cannot be used the mass of hydrogen lost is too small to measure significantly.

Section 6 of 26PAG 1.3 — formula of magnesium oxide

1

Weigh a known mass of magnesium and place it in a crucible with a lid.

2

Heat the magnesium so that it reacts with oxygen in the air.

3

Keep the lid open enough to let air in but closed enough not to lose product.

4

Cool and weigh the product.

1Calculation

• Calculation the mass of oxygen is the product mass minus the magnesium mass.
• Moles = mass / M_r for magnesium and for oxygen.
• The molar ratio MgO gives the formula, so a ratio of 2 Mg : 3 O would give Mg₂O₃.

2Product escapes as smoke

• Product escapes as smoke this cannot be avoided but is reduced by careful heating and careful placement of the lid.

3Magnesium does not react

• Magnesium does not react ribbon is better than powder.
• Once ignited the ribbon keeps burning, whereas powder can lose product and leave some magnesium unreacted.
• Clean ribbon with sandpaper or glasspaper to remove any oxide coating.

4Magnesium left on the weighing apparatus

• Magnesium left on the weighing apparatus weigh the apparatus with and without the magnesium.

Section 7 of 26PAG 1 skills and PAG 2 acid-base titration

1

Skills developed in PAG 1

• Skills developed in PAG 1 accurate use of balances, gas collection by displacement or gas syringe, safe handling of heat sources and crucibles.
• Applying balanced equations and molar ratios to unknown quantities, and identifying and reducing errors.

2

PAG 2.1

• PAG 2.1 find the concentration of hydrochloric acid.
• Make a standard solution of sodium hydrogencarbonate and titrate it against the acid.

3

PAG 2.2

• PAG 2.2 find the molar mass of an acid.
• Make a solution of an unknown hydrated acid such as citric acid monohydrate and titrate it with sodium hydroxide of known concentration.
• With extra information the formula of the acid can be found.

4

PAG 2.3

• PAG 2.3 identify an unknown Group 1 carbonate.
• Titrate a solution of the carbonate with hydrochloric acid of known concentration to find its molar mass and the metal.

Section 8 of 26Volumetric analysis and standard solutions

1Volumetric analysis

• Volumetric analysis uses the volume and concentration of one reactant to find the concentration of another unknown solution.
• Titration is the most common technique and uses a volumetric or graduated pipette and a burette.

2Standard (volumetric) solution

• Standard (volumetric) solution a solution whose concentration is known precisely.
• It is made with a three-decimal-place balance and a volumetric flask to reduce measurement uncertainty.

3Apparatus

• Apparatus beaker, burette, volumetric pipette, conical flask and volumetric flask.

4Making the solution

• Weigh the solid accurately in a weighing boat or beaker and record the mass.
• Dissolve it in a little distilled water in a beaker, then transfer it with rinsings to the volumetric flask.
• Make up to the mark with distilled water, stopper and invert several times to mix.

Section 9 of 26Concentration and the burette

1Concentration

Concentration amount of solute dissolved in a solvent to make 1 dm³ of solution.

• The solute is the substance that dissolves and the solvent is often water.

2Concentrated and dilute

• Concentrated and dilute a concentrated solution has a high concentration of solute and a dilute solution has a low concentration.

3Units

• Units mol dm^-3, g dm^-3 or parts per million (ppm).

4Mass needed for a solution

• Mass needed for a solution moles = concentration × volume (dm³), then mass = moles × M_r.
• For 250 cm³ of 0.200 mol dm^-3 NaHCO₃ (M_r 84.0), 0.0500 mol is needed, which is 4.2 g.

5Burette precision

• Burette precision marked to 0.10 cm³, and as an analogue instrument its uncertainty is half the smallest marking, ±0.05 cm³.

6End point (equivalence point)

• End point (equivalence point) the point where the two solutions have reacted completely, shown by an indicator.

Section 10 of 26Steps in a titration

1

Measure a known volume (usually 20 or 25 cm³) of one solution with a volumetric pipette into a conical flask.

2

Place the other solution in the burette, usually filled to 0.00 cm³.

3

Add a few drops of indicator to the conical flask and place a white tile underneath to see the colour change.

4

Open the tap and add solution slowly until the indicator starts to change colour.

5

Near the end point, slow the flow and add the solution dropwise, closing the tap after one drop causes the change.

6

Repeat until concordant results are obtained.

• Concordant results
  • titres within 0.1 cm³ of each other.

Section 11 of 26Recording and processing titration results

1Burette readings

• Burette readings record the initial and final readings to the same precision as the uncertainty, ±0.05 cm³.

2Titre

• Titre the volume delivered, V_final − V_initial, with an uncertainty of ±0.10 cm³.
• The uncertainty is doubled because two burette readings are used, following the rules for propagating uncertainties.

3Averaging

• Averaging average only the concordant titres and discard the others, including the rough run.

4Calculation steps

• Find the moles of the solution in the burette from concentration × volume (dm³).
• Use the stoichiometric ratio of the equation to find the moles of the other reactant.
• Divide moles by the pipetted volume (dm³) to give its concentration.

Section 12 of 26PAG 2.1 — concentration of hydrochloric acid

1Reaction

• NaHCO₃(aq) + HCl(aq) → NaCl(aq) + H₂O(l) + CO₂(g), a 1 : 1 ratio.

2Method

• Method a standard solution of NaHCO₃ is titrated against a known volume of HCl.

3Processing

• Average the concordant titres, convert to moles of NaHCO₃ with c × V / 1000, and set moles of HCl equal to this.
• Concentration of HCl = moles / (volume of acid in cm³ / 1000).

Section 13 of 26PAG 2.2 — molar mass of an acid

1Method

• Method dissolve a known mass of an unknown hydrated or monoprotic acid in a volumetric flask and titrate with sodium hydroxide of known concentration.

2Processing

• Moles of NaOH = c × V / 1000.
• Use the stoichiometry (for a monoprotic acid HX : NaOH is 1 : 1) to find moles of acid in the sample titrated.
• M_r = mass / moles, scaled up to the whole flask if an aliquot was used.

3Identifying a halogen acid HX

• Identifying a halogen acid HX subtract the A_r of hydrogen from M_r and match the result to the periodic table.
• A value near 80 gives bromine, so the acid is HBr.

4Other acids

• Other acids monoprotic, diprotic and triprotic acids can all be used and the stoichiometry of neutralisation must be accounted for.
• Extra information is given to identify an unknown acid.

Section 14 of 26PAG 2.3 — identifying an unknown Group 1 carbonate

1Reaction

• X₂CO₃ reacts with HCl in the ratio 1 X₂CO₃ : 2 HCl.

2Method

• Method dissolve a known mass in water to make a standard solution and titrate aliquots (e.g. 25.0 cm³ of 250.0 cm³) with HCl of known concentration.

3Processing

1

Calculate moles of HCl from c × V / 1000.

2

Halve it to give moles of X₂CO₃ in the aliquot.

3

Find the mass in the aliquot (one tenth of the total mass for 25.0 of 250.0 cm³).

4

M_r = mass / moles.

5

Subtract the mass of CO₃ (12.0 + 3 × 16.0) and divide by 2 to give A_r of X.

6

Match to the periodic table; for example 23.0 is sodium.

• Identification pitfall
  • values near 86 can be wrongly assigned to francium.
  • Use relative atomic mass, not atomic number: rubidium has A_r 85.5 whereas francium has atomic number 87.

Section 15 of 26Skills developed by the titration practicals

1Measuring volumes

• Measuring volumes accurate use of volumetric pipettes, burettes and volumetric flasks.

2Solutions

• Solutions preparing standard solutions and minimising measurement uncertainty.

3Recording

• Recording burette readings to the correct precision with instrument uncertainty.

4Technique

• Technique using indicators and controlling the flow rate near the end point safely.

5Calculation

• Calculation concentrations, molar masses and formulae from balanced equations and volume data.

Section 16 of 26PAG 3 — calorimetry basics

PAG 3 — calorimetry basics

Calorimetry

• Calorimetry a technique to measure enthalpy changes of chemical reactions.
• A calorimeter can be a polystyrene cup, a vacuum flask or a metal can.

Specific heat capacity

• Specific heat capacity the energy needed to raise the temperature of 1 g of a substance by 1 K.
• For water it is 4.18 J g^-1 K^-1.

Energy transferred

• Energy transferred q = mcΔT.

  What the symbols mean: E, Q or q = energy transferred (J) · m = mass · c = specific heat capacity (J kg⁻¹ °C⁻¹; water 4.18 J g⁻¹ °C⁻¹) · ΔT or Δθ = temperature change (°C or K)
  • (kg; g in chemistry calorimetry)
• q is heat energy in J, m is mass of water in g, c is specific heat capacity in J g^-1 K^-1 and ΔT is the temperature change in K.

Two types of experiment

• Two types of experiment enthalpy changes of reaction in solution and enthalpy changes of combustion.
• For both you must outline the experiment and process the data by calculation or graph.

Section 17 of 26PAG 3.1 — enthalpy change of neutralisation

1

Place 25.0 cm³ of 1.0 mol dm^-3 HCl in one polystyrene cup and 25.0 cm³ of 1.0 mol dm^-3 NaOH in another, using measuring cylinders.

2

Draw a table for the initial temperatures and for the mixture's temperature every half minute up to 9.5 minutes.

3

Put a thermometer or probe in each cup, stir, and record the temperature every half minute for 2½ minutes.

4

At precisely 3 minutes add the two solutions together; do not record the temperature at 3 minutes.

5

Keep stirring and record the temperature for 6 minutes.

• Assumptions
  • the solution has the specific heat capacity of water (4.18 J g^-1 K^-1) and density 1 g cm^-3.
  • The heat capacity of the container is ignored, the reaction is complete and heat losses are negligible.

Section 18 of 26Temperature correction graphs

1Why needed

• Why needed if a reaction is not instantaneous there is a delay before the maximum temperature, during which heat is lost to the surroundings.
• The true maximum temperature is therefore never reached.

2Steps

1

Record the temperature for a few minutes before mixing to get a steady value.

2

Add the second reactant and continue recording temperature and time.

3

Plot temperature against time.

4

Extrapolate the cooling section back to the time the second reactant was added.

1Equal initial temperatures

• Equal initial temperatures both solutions plot as one line and the extrapolation at 3 minutes gives ΔT for q = mcΔT.

2Different initial temperatures

• Different initial temperatures plot the HCl and NaOH temperatures as separate lines.
• Extrapolate separately to the mixing time and calculate q = mcΔT for each.
• The average of the two gives the overall enthalpy change of neutralisation.

Section 19 of 26PAG 3.2 — enthalpy change by Hess' law

Purpose

• Purpose an extended version of PAG 3.1 using two reactions to find an enthalpy change that cannot be measured directly.

Exam board suggestion

• Exam board suggestion potassium carbonate and potassium hydrogencarbonate each reacting with hydrochloric acid.
• This gives the enthalpy change for decomposition of KHCO₃ to K₂CO₃.

Another example

• Another example hydrated and anhydrous copper sulfate each forming copper sulfate solution.

Section 20 of 26PAG 3.2 — sample method

1

Place 25.0 cm³ of 1.0 mol dm^-3 HCl in a polystyrene cup with a measuring cylinder.

2

Draw a table for the initial temperature and for the temperature every half minute up to 9.5 minutes.

3

Weigh 3.25 g of KHCO₃ on a top-pan balance.

4

Put in a thermometer or probe, stir, and record the temperature every half minute for 2½ minutes.

5

At precisely 3 minutes add the solid to the acid; do not record at 3 minutes.

6

Keep stirring and record the temperature for 6 minutes.

7

Plot a temperature correction graph to find the temperature change.

8

Repeat using 2.50 g of K₂CO₃, and repeat both to find average temperature changes.

9

Use q = mcΔT to find ΔH₁ (KHCO₃ with HCl) and ΔH₂ (K₂CO₃ with HCl).

10

Apply Hess' law: ΔH_r = ΔH₁ − ΔH₂ for decomposition of KHCO₃ to K₂CO₃.

Section 21 of 26PAG 3.3 — enthalpy change of combustion method

• Principle
  • the heat released by burning a fuel raises the temperature of a known mass of water in a simple combustion calorimeter.

1

Measure a known mass of water into a copper calorimeter and clamp it above a spirit burner.

2

Record the starting temperature of the water and the starting mass of the spirit burner with its lid.

3

Light the burner and stir the water gently until the temperature has risen by a suitable amount.

4

Extinguish the flame and record the final temperature and the final mass of the burner.

• Calculation chain
  • Mass of fuel burned = starting mass − final mass.
  • Temperature change = final − starting temperature, then q = mcΔT.
  • Moles of fuel = mass burned / M_r, and ΔH = q / n with the sign negative for combustion.

Section 22 of 26Errors and improvements in combustion calorimetry

Heat losses

• Heat losses not all heat from the flame passes to the water.
• Some is lost to the surroundings and some is absorbed by the calorimeter.

Incomplete combustion

• Incomplete combustion a main source of error that gives less heat than complete combustion.

Reducing heat loss

• Do not place the copper calorimeter too far above the flame.
• Put a lid on the calorimeter.
• Use shielding to reduce draughts.

Section 23 of 26Combustion calculation pattern

Combustion calculation pattern

Calculate q = mcΔT using the mass of water, 4.18 J g^-1 K^-1 and the temperature rise.

Calculate the moles of fuel burned as mass / M_r.

Calculate ΔH = q / n, convert to kJ mol^-1 and make it negative because heat is released.

• Check
  • 200 g of water heated by 30 K gives q of 25 080 J, and burning 1.023 g of propan-1-ol (M_r 60.11) gives about −1.5 × 10³ kJ mol^-1.

Section 24 of 26Skills developed by calorimetry

1Measuring

• Measuring temperature changes with a thermometer or probe, mass with a top-pan balance and liquid volumes with measuring cylinders.

2Data handling

• Data handling recording and plotting time–temperature data to make graphical corrections.

3Calculation

• Calculation enthalpy changes from q = mcΔT, and Hess' law to derive an enthalpy change indirectly from two reactions.

4Evaluation

• Evaluation identifying and explaining sources of error such as heat loss and incomplete combustion.

Section 25 of 26Key terms

Term	Meaning
Standard solution	A solution whose concentration is known precisely.
Volumetric analysis	Finding the concentration of an unknown solution using the known volume and concentration of another.
End point	The point at which the reactants have reacted completely, shown by an indicator.
Concentration	The amount of solute dissolved in 1 dm3 of solution.
Water of crystallisation	Water molecules chemically bonded in a hydrated crystal, found from the mass lost on heating.
Molar gas volume	24.0 dm3 mol−1 at room temperature and pressure.

Term	Meaning
Titre	The volume delivered from the burette, final reading minus initial reading.
Concordant results	Titres within 0.1 cm3 of each other.
Aliquot	A measured portion of a solution taken for titration.
Specific heat capacity	The energy needed to raise 1 g of a substance by 1 K.
Calorimetry	Measuring enthalpy changes from temperature changes.
Temperature correction graph	A graph extrapolated back to the mixing time to find the true temperature change.

Term	Meaning
Hess' law	The enthalpy change of a reaction is independent of the route taken.
Enthalpy change of combustion	The enthalpy change when one mole of a substance burns completely in oxygen.
Extrapolation	Extending a graph line beyond the data to find a value, such as the true temperature change.
Incomplete combustion	Burning in insufficient oxygen, giving less heat energy.

Section 26 of 26Exam technique

1

Using mass loss to find the relative atomic mass of magnesium — the lost hydrogen is too small to measure, so collect the gas volume instead.

2

Using 24.0 dm³ per mole at any conditions — it is valid only at room temperature and pressure, otherwise use n = PV/RT.

3

Not sealing the flask at once — gas escapes before the bung is inserted, so add acid and stopper immediately.

4

Reading burette uncertainty as ±0.10 cm³ — each single reading is ±0.05 cm³.

1

Averaging all titres including the rough run — average only concordant titres within 0.1 cm³.

2

Quoting titre uncertainty as ±0.05 cm³ — it is ±0.10 cm³ because two readings are subtracted.

3

Using atomic number to identify an unknown element — use relative atomic mass.

4

Forgetting the 1 : 2 ratio for a Group 1 carbonate with HCl — halve the moles of HCl.

1

Recording the temperature at the moment of mixing — the reading at 3 minutes is omitted and the graph is extrapolated back to it.

2

Forgetting that combustion is exothermic — ΔH is negative.

3

Writing ΔHr = ΔH₂ − ΔH1 — for KHCO₃ to K₂CO₃ use ΔH1 − ΔH₂ as in the method.

4

Using the mass of fuel in q = mcΔT — m is the mass of water heated.