Crystals that weigh more than the chemistry allows
A Grade 10 student prepares copper(II) sulphate crystals. She is given 25.0 cm³ of dilute sulphuric acid of concentration 1.0 mol/dm³ and adds copper(II) oxide in small spatula loads, stirring and warming gently, until no more will dissolve. She adds 4.0 g of black copper(II) oxide in total. She filters the mixture and a black solid remains on the filter paper. The filtrate is clear blue. She pours it into an evaporating basin and heats it over a Bunsen, with the gauze removed for the last few minutes, until the liquid has almost gone and crystals are just appearing. She leaves it to cool overnight, filters off the blue crystals, presses them between filter papers and weighs them: 6.8 g. She also weighs the black solid left after filtering and reports 1.2 g. Her teacher asks her to check whether both masses make sense before writing up the experiment. [Mᵣ: CuO = 79.5, CuSO₄ = 159.5, CuSO₄·5H₂O = 249.5]
1. Calculate the mass of copper(II) oxide that just reacts with all the acid, the maximum mass of CuSO₄·5H₂O crystals obtainable, and the mass of black solid you would expect to be left over. Comment on the reported 1.2 g.
Read the worked answer
Moles of acid = 0.0250 dm³ × 1.0 mol/dm³ = 0.0250 mol.
CuO + H₂SO₄ → CuSO₄ + H₂O is a 1:1 reaction, so 0.0250 mol of CuO reacts.
- CuO needed = 0.0250 × 79.5 = 1.99 g ≈ 2.0 g
- Maximum CuSO₄·5H₂O = 0.0250 × 249.5 = 6.24 g ≈ 6.2 g
- Expected leftover CuO = 4.0 − 2.0 = 2.0 g
Adding copper(II) oxide in excess is correct practice: it is insoluble and is added until no more reacts, which guarantees that all the acid is used up and the filtrate is not acidic.
The reported 1.2 g cannot be right. If only 1.2 g were left over, 2.8 g of CuO (0.0352 mol) must have reacted - but that needs more acid than she was given. So either acid was lost or the solid was lost. The likeliest explanation is physical loss: black solid stuck to the beaker wall or passed through the filter paper during transfer. A mass balance that does not close is a finding in itself, not a detail to tidy away.
2. Evaluate the 6.8 g mass of crystals against your calculation. Give the most likely explanation and describe a test she could carry out.
Read the worked answer
The theoretical maximum is 6.24 g, so 6.8 g corresponds to a yield of 6.8 ÷ 6.24 × 100 = 109%.
A yield above 100% is impossible for a pure product from a fixed amount of acid - no calculation error changes that, because 0.025 mol of copper is all the copper she started with once the acid is limiting. The conclusion must be that the 6.8 g is not all CuSO₄·5H₂O.
The most likely reasons:
- Wet crystals. Surface water and mother liquor trapped between the crystals; pressing between filter papers removes some but not all of it.
- Impurity from the evaporation. She heated almost to dryness, so any soluble impurity crystallised out with the product.
- Weighing the paper. Crystals transferred with pieces of filter paper give a falsely high mass.
Test: dry the sample to constant mass - press it between fresh filter papers, leave it in a warm place (never a hot oven) and weigh at intervals until two consecutive masses agree. The mass will fall towards 6.2 g or below; if it keeps falling and the crystals turn white, water of crystallisation is being lost.
3. She removed the gauze and heated the basin strongly in the final minutes. State one risk of this step and how it would show up in what she weighs. She also says the filtrate "looked clean", so she assumes no acid was left. Assess that reasoning.
Read the worked answer
Removing the gauze and heating strongly risks driving off the water of crystallisation. Hydrated CuSO₄·5H₂O (blue) loses its four coordinated water molecules on heating to about 150 °C to give white anhydrous CuSO₄, and hotter still it decomposes to black copper(II) oxide. If that happened, the mass would be much lower, not higher: 0.0250 mol of anhydrous CuSO₄ would weigh 0.0250 × 159.5 = 4.0 g. So a mass of about 4.0 g of white powder, or a black residue, would be the fingerprint of overheating rather than a good yield. Strong heating also causes spitting, which loses solution and lowers the yield.
Her reasoning about the filtrate is weak. Colour tells her about the copper ion, not about hydrogen ions; a copper(II) sulphate solution is blue whether or not acid is left over. To test for excess acid she should add a spatula of copper(II) oxide or a small piece of carbonate to a sample of the filtrate and look for a reaction, or check with pH paper. "It looked clean" is an observation about appearance, not evidence about composition.