Two Unlabelled Bottles and a Neutralisation Problem
Two unlabelled bottles stand on a bench. Both contain colourless liquids of concentration 0.10 mol/L; one is hydrochloric acid and the other is sodium hydroxide solution. A student is allowed four tests: (i) universal indicator paper; (ii) a piece of magnesium ribbon; (iii) a few drops of bromothymol blue (BTB) solution; (iv) evaporating 10 mL of each to dryness. Results: bottle P gives bubbles with magnesium, turns BTB yellow, and leaves no residue on evaporation. Bottle Q does not react with magnesium, turns BTB blue, and leaves a white solid. The student then pours 30 mL of P into a beaker, adds 20 mL of Q, stirs, and asks the teacher whether the mixture is neutral and what colour BTB would show in it. The teacher asks for a mole-level answer rather than a guess from the colour.
1. Identify P and Q and state which observations are decisive. Explain why 'no residue on evaporation' on its own would not prove that a liquid is an acid.
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Identification. P is hydrochloric acid, Q is sodium hydroxide solution. Only P reacts with magnesium ribbon, producing hydrogen gas: a metal reacts with an acid to give a salt and hydrogen, not with an alkali. BTB is yellow in acid and blue in alkali, which fits the same assignment. Q leaves a white solid because the dissolved sodium hydroxide remains (and partly becomes carbonate in air), while P leaves essentially nothing because hydrogen chloride escapes as a gas as the water evaporates.
Judgement on the evaporation test. 'No residue' is consistent with an acid, but it is not proof: pure water, or a solution of any volatile substance, also leaves nothing behind. The evaporation test would not by itself exclude a neutral liquid. The magnesium test and the indicator are the decisive evidence because each depends on a chemical property unique to acids or alkalis.
2. Is the mixture of 30 mL of P and 20 mL of Q acidic, basic or neutral? Support your answer with mole calculations and state the BTB colour.
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Moles. n(HCl) = 0.10 mol/L x 0.030 L = 3.0 x 10^-3 mol. n(NaOH) = 0.10 x 0.020 = 2.0 x 10^-3 mol. The reaction is 1:1, so 2.0 x 10^-3 mol of each neutralise and 1.0 x 10^-3 mol of HCl is left over.
Concentration. The total volume is 30 + 20 = 50 mL, so [H+] = 1.0 x 10^-3 / 0.050 = 0.020 mol/L, giving pH = -log(0.020) which is about 1.7.
Conclusion. The mixture is clearly acidic, not neutral, and BTB would be yellow. Careful reasoning about volume matters: the larger volume of P is decisive, and 'equal concentration' does not mean 'equal amounts' once the volumes differ.
3. The student wants a neutral solution from all 30 mL of P. How much Q is needed, and what mass of solid is left after evaporating that neutral mixture? (Na = 23, Cl = 35.5) Would a green BTB colour prove the solution was exactly neutral?
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Amount needed. 30 mL of P contains 3.0 x 10^-3 mol HCl, which requires 3.0 x 10^-3 mol NaOH - exactly 30 mL of Q at 0.10 mol/L. The volumes are equal only because the concentrations are equal.
Solid remaining. Neutralisation gives 3.0 x 10^-3 mol NaCl. Molar mass = 23 + 35.5 = 58.5 g/mol, so mass = 3.0 x 10^-3 x 58.5 = 0.176 g, about 0.18 g.
Indicator judgement. No. BTB changes from yellow around pH 6.0 to blue around pH 7.6, so several mixtures that are slightly acidic or slightly alkaline still look green. Green is consistent with neutrality but does not prove it. A pH meter, or a check that adding a single drop of either solution shifts the colour in one direction only, would be a fair test.