Principles Related to Practical Chemistry — Practice Questions with Answers

The equivalence point is the precise theoretical point in a titration at which the exact stoichiometric amount of titrant has been added to completely react with the analyte, according to the balanced chemical equation.

The endpoint is the experimentally observed point (typically detected via an indicator's colour change, or another physical signal) at which the titration is actually stopped in practice — ideally, a well-chosen indicator ensures the endpoint occurs very close to the true equivalence point, but a small, generally negligible difference (titration error) can exist between these two points.

A primary standard substance must be: (1) obtainable in very high purity; (2) chemically stable, both in solid form and in solution, without reacting with air, moisture, or $CO_2$; (3) non-hygroscopic (not absorbing atmospheric moisture, which would alter its precise weighed mass); and (4) possess a reasonably high molar mass, to minimise the relative error involved in precisely weighing out small quantities.

These properties collectively ensure that a precisely weighed mass of the primary standard reliably and accurately corresponds to a precisely known number of moles, allowing it to be used to reliably standardise (determine the precise concentration of) another solution.

In a strong acid-strong base titration, the equivalence point occurs at exactly pH 7 (neutral), but the pH changes extremely sharply (a steep vertical jump) over a wide pH range around this point — phenolphthalein's colour-change range (pH ~8.2-10) falls well within this steep pH jump, making it a suitable, sharp indicator here.

In a weak acid-strong base titration, the equivalence point occurs above pH 7 (basic, due to hydrolysis of the resulting conjugate base salt) — methyl orange changes colour in the acidic range (pH ~3.1-4.4), far below this basic equivalence point, meaning it would change colour prematurely, well before the true equivalence point is reached, making it unsuitable/inaccurate for this specific type of titration; phenolphthalein remains the more suitable choice here instead.

Stage 1 (preliminary/dry tests): Simple tests performed directly on the dry salt sample, such as heating it alone, or treating it with dilute and then concentrated sulfuric acid, observing for characteristic gas evolution, colour changes, or odours, which can provide preliminary clues about which anion(s) might be present.

Stage 2 (confirmatory/wet tests): Specific, targeted chemical tests performed on an aqueous solution (or extract) of the salt, using particular reagents that give a characteristic, definitive reaction (precipitate formation, colour change, or specific gas evolution) uniquely confirming the presence of a suspected specific anion, based on the preliminary test results.

A burette is a long, graduated glass tube fitted with a precision stopcock (valve) at the bottom, specifically designed to allow very precise, controllable, incremental (drop-by-drop, if needed) delivery of a measured volume of liquid, with fine graduations allowing volume readings to a high degree of precision (typically to within 0.1 mL or better, with practice).

A simple measuring cylinder, by contrast, is designed only for a single, one-time bulk volume measurement, without any mechanism for finely controlled, incremental delivery, and typically has considerably less precise graduations — this crucial ability to deliver the titrant in a slow, finely controlled, precisely measurable manner (essential for accurately locating the exact titration endpoint) is precisely why a burette, rather than a measuring cylinder, is the appropriate and necessary apparatus for titrant delivery in volumetric analysis.

A solution of the salt (acidified with dilute nitric acid, to prevent interference from other anions like carbonate) is treated with silver nitrate solution.

If chloride ions are present, a characteristic white precipitate of silver chloride ($AgCl$) forms immediately: $Cl^- + AgNO_3 \rightarrow AgCl\downarrow + NO_3^-$. This white precipitate is further confirmed as specifically chloride (rather than another halide) by testing its solubility behaviour in ammonium hydroxide solution — silver chloride characteristically dissolves readily in dilute ammonia (forming a soluble diammine silver complex), distinguishing it from silver bromide (only partially soluble) and silver iodide (essentially insoluble) in dilute ammonia.

Molarity (M) is defined as the number of moles of solute dissolved per litre of solution.

Normality (N) is defined as the number of gram-equivalents of solute dissolved per litre of solution, where the number of equivalents depends on the specific reaction the substance is undergoing (e.g., the number of replaceable H⁺ ions for an acid, or the number of electrons transferred in a redox reaction).

Relationship: $\text{Normality} = \text{Molarity}\times n$, where $n$ is the "n-factor" (number of equivalents per mole) specific to that substance's reaction in the given context (for example, $n=2$ for sulfuric acid in a reaction where both acidic protons react).

Dilute sulfuric acid is added specifically to maintain a strongly acidic reaction medium — under these acidic conditions, permanganate ion ($MnO_4^-$) is reduced most efficiently and predictably all the way to the colourless $Mn^{2+}$ ion (a 5-electron reduction process), providing a clean, well-defined, quantitatively reliable reaction with a sharp, easily detectable colour-change endpoint.

Other acids are specifically avoided for this purpose: hydrochloric acid can itself be oxidised by permanganate (introducing an unwanted side reaction and error), while nitric acid is itself an oxidising agent that could interfere with the intended redox reaction — dilute sulfuric acid is chemically inert towards $KMnO_4$ under these conditions, making it the specifically preferred acid for this type of titration.

In the borax bead test, a small bead of molten borax (sodium tetraborate) is formed on a heated platinum/nichrome wire loop, then touched to a small amount of the salt sample and reheated in a flame.

Certain transition metal ions present in the salt form characteristically, distinctively coloured borate glass beads (due to formation of specific coloured metal borate compounds within the molten bead) — for example, a cobalt salt characteristically produces a distinctive blue bead, while a copper salt typically produces a blue-green (or sometimes red, depending on oxidising/reducing flame conditions) bead — this characteristic colouration provides a simple, useful preliminary qualitative test for identifying the presence of certain specific coloured transition metal cations in a sample.

As titrant is added dropwise from the burette, it must be given sufficient opportunity to fully and completely react with (and disperse uniformly throughout) the entire volume of the analyte solution in the flask — without adequate mixing, localised regions of excess, unreacted titrant concentration could develop near where the titrant drops enter the solution.

This could cause a temporary, localised indicator colour change that doesn't actually reflect the true overall reaction state of the complete, well-mixed solution, potentially leading to a false, premature reading of the endpoint (stopping the titration too early, before the true equivalence point has actually been reached throughout the entire solution) — continuous, thorough mixing (typically by gently swirling the flask) ensures the indicator's observed colour genuinely and accurately reflects the true, complete reaction state of the entire solution, which is especially critical for accuracy as the true endpoint is being approached.

Reaction: $NaOH + HCl \rightarrow NaCl + H_2O$ (1:1 stoichiometry)

At the equivalence point: moles of NaOH = moles of HCl, so $M_1V_1 = M_2V_2$

$M_{NaOH}\times25 = 0.1\times20$

$M_{NaOH} = \dfrac{0.1\times20}{25} = \dfrac{2}{25}$

$M_{NaOH} = \mathbf{0.08\text{ M}}$

Moles of $Na_2CO_3$ in 100 mL $= \dfrac{0.53}{106} = 0.005\text{ mol}$; Molarity of $Na_2CO_3 = \dfrac{0.005}{0.1} = 0.05\text{ M}$

Reaction: $Na_2CO_3 + 2HCl \rightarrow 2NaCl + H_2O + CO_2$ (1:2 stoichiometry)

Moles of $Na_2CO_3$ used in 20 mL aliquot $= 0.05\times0.020 = 0.001\text{ mol}$

Moles of HCl reacted $= 2\times0.001 = 0.002\text{ mol}$

Molarity of HCl $= \dfrac{0.002}{0.018} \approx \mathbf{0.111\text{ M}}$

In iodometric determination of $Cu^{2+}$: $2Cu^{2+} + 4I^- \rightarrow 2CuI + I_2$, followed by $I_2 + 2S_2O_3^{2-} \rightarrow 2I^- + S_4O_6^{2-}$

Overall, moles of $Cu^{2+}$ = moles of thiosulfate used (1:1 overall relationship, since 2 mol $Cu^{2+}$ → 1 mol $I_2$ → 2 mol thiosulfate).

Moles of thiosulfate $= 0.1\times0.020 = 0.002\text{ mol}$

Moles of $CuSO_4 = 0.002\text{ mol}$ (from the 1:1 relationship)

Molarity of $CuSO_4 = \dfrac{0.002}{0.025} = \mathbf{0.08\text{ M}}$

Using the normality equivalence principle: $N_1V_1 = N_2V_2$, where $N = M\times n$

Normality of $KMnO_4 = 0.2\times5 = 1\text{ N}$; Normality of $FeSO_4 = 0.5\times1 = 0.5\text{ N}$

$N_{KMnO_4}\times V_{KMnO_4} = N_{FeSO_4}\times V_{FeSO_4}$

$1\times V_{KMnO_4} = 0.5\times25$

$V_{KMnO_4} = \dfrac{12.5}{1} = \mathbf{12.5\text{ mL}}$

Total moles of HCl taken $= 0.2\times0.050 = 0.01\text{ mol}$

Moles of excess HCl (neutralised by NaOH) $= 0.1\times0.010 = 0.001\text{ mol}$

Moles of HCl reacted with $CaCO_3$ = Total − Excess $= 0.01 - 0.001 = 0.009\text{ mol}$

Reaction: $CaCO_3 + 2HCl \rightarrow CaCl_2 + H_2O + CO_2$ (1:2 stoichiometry)

Moles of $CaCO_3$ = moles of HCl reacted / 2 $= 0.009/2 = 0.0045\text{ mol}$

Mass of $CaCO_3 = 0.0045\times100 = 0.45\text{ g}$

% purity $= \dfrac{0.45}{0.4}\times100 = \mathbf{112.5\%}$

(Note: A value exceeding 100% would indicate an inconsistency in the given data in a real experiment, but the calculation method demonstrated here — back-titration analysis — remains the standard, correct approach.)

Procedure: A solution of the salt is first acidified with dilute hydrochloric acid (to dissolve/exclude interference from carbonates, sulfites, and similar anions that would also give precipitates with barium but are soluble in dilute HCl). Barium chloride solution is then added to this acidified salt solution.

Observation: Formation of a white precipitate, insoluble in dilute hydrochloric acid (or dilute nitric acid), confirms the presence of sulfate: $SO_4^{2-} + BaCl_2 \rightarrow BaSO_4\downarrow + 2Cl^-$

The specific insolubility of this white precipitate in dilute mineral acids (unlike, for example, barium carbonate or barium sulfite, which would dissolve/decompose upon acid addition) is the key distinguishing confirmatory feature specifically identifying sulfate, differentiating it from other anions that might also initially give a superficially similar white precipitate with barium chloride under neutral conditions.

Reaction: $H_2SO_4 + 2NaOH \rightarrow Na_2SO_4 + 2H_2O$ (1:2 stoichiometry)

Moles of NaOH $= 0.6\times0.030 = 0.018\text{ mol}$

Moles of $H_2SO_4$ required $= \dfrac{0.018}{2} = 0.009\text{ mol}$ (since 1 mol $H_2SO_4$ reacts with 2 mol NaOH)

Volume of $H_2SO_4$ required $= \dfrac{\text{moles}}{\text{molarity}} = \dfrac{0.009}{0.5} = 0.018\text{ L}$

Volume $= \mathbf{18\text{ mL}}$

Procedure: A freshly prepared solution of the salt is treated with freshly prepared iron(II) sulfate ($FeSO_4$) solution. Concentrated sulfuric acid is then very carefully poured slowly down the side of the tilted test tube, allowing it to sink to the bottom without immediately mixing with the upper aqueous layer.

Observation: If nitrate is present, a distinctive brown ring forms precisely at the interface (junction) between the two liquid layers — this brown ring is a complex compound, $[Fe(NO)(H_2O)_5]SO_4$ (a nitroso-iron(II) sulfate complex), formed from the reaction between nitric oxide (NO, generated locally by the reduction of nitrate by $Fe^{2+}$ in the presence of the concentrated acid) and excess unreacted $Fe^{2+}$ ions.

The careful, slow addition of concentrated sulfuric acid specifically to form two distinct layers (rather than mixing immediately) is essential, since this concentrates the necessary reaction conditions specifically at the narrow interfacial boundary, producing the sharp, characteristic, easily observable brown ring.

Using normality equivalence: $N_1V_1 = N_2V_2$

Normality of oxalic acid $= 0.1\times2 = 0.2\text{ N}$

$N_{oxalic}\times V_{oxalic} = N_{KMnO_4}\times V_{KMnO_4}$

$0.2\times20 = N_{KMnO_4}\times24$

$N_{KMnO_4} = \dfrac{4}{24} = 0.1667\text{ N}$

Molarity of $KMnO_4 = \dfrac{N}{n} = \dfrac{0.1667}{5} \approx \mathbf{0.0333\text{ M}}$

Moles of NaOH required $= M\times V(\text{in L}) = 0.1\times0.250 = 0.025\text{ mol}$

Mass of NaOH required $= \text{moles}\times\text{molar mass} = 0.025\times40$

Mass $= \mathbf{1\text{ g}}$

(In practice, this 1 g of NaOH would be accurately weighed, dissolved in a small amount of distilled water, then carefully transferred to a 250 mL volumetric flask and made up precisely to the 250 mL mark with additional distilled water.)