Conduct investigations to measure the concentration of cations and anions in solution using gravimetric analysis and precipitation titrations

What this dot point is asking

NESA wants you to use a precipitation reaction to measure concentration two ways: by weighing the dried precipitate (gravimetric analysis) or by titrating a known precipitant to a colour-change endpoint (precipitation titration, typically the Mohr method for chloride). You should know the workflow, the calculations, and the most common sources of error.

The answer

Gravimetric analysis: the workflow

1. Weigh the sample accurately on an analytical balance.
2. Dissolve in a measured volume of water, and acidify with the appropriate dilute acid to remove carbonate and other interferents.
3. Add excess precipitating reagent to drive the precipitation to completion.
4. Digest the precipitate by warming, which grows crystals and reduces co-precipitation.
5. Filter through pre-weighed filter paper (ashless) or a sintered glass crucible.
6. Wash the precipitate with a small volume of distilled water (and a dilute solution of a common ion to suppress dissolution).
7. Dry to constant mass in an oven (or ignite to a known oxide).
8. Weigh the dried precipitate.

The mass of the precipitate gives moles, which converts back to moles (and mass) of the original ion via the stoichiometry of the precipitation equation.

The canonical example: sulfate as barium sulfate

For a sample of mass $m$ that gives a precipitate of mass $m_{BaSO_4}$:

Acidify with dilute $HCl$ first; this removes carbonate (which would otherwise precipitate as $BaCO_3$) but does not dissolve $BaSO_4$.

Other common gravimetric pairs

Precipitation titration: the Mohr method

Use when you want speed and do not need part-per-billion precision. The classic Mohr titration measures $Cl^-$:

• Titrant: standardised $AgNO_3$ (commonly 0.1 mol/L).
• Indicator: a few drops of $K_2CrO_4$.
• End-point: the first persistent red-brown colour of $Ag_2CrO_4$.

The chemistry has two stages. While free chloride remains:

When chloride is exhausted, the next drop of $Ag^+$ reacts with chromate to give a red-brown precipitate, signalling the end-point:

$Ag_2CrO_4$ has a higher $K_{sp}$ than $AgCl$, so $AgCl$ precipitates first. The chromate stays in solution until all chloride is consumed.

Calculation pattern for a precipitation titration

For volume of titrant $V_t$ and concentration $c_t$:

Convert to g/L by multiplying by 35.5 g/mol.

pH window for the Mohr method

The titration must be run at pH 7 to 9.5.

• Below pH 6.5: chromate protonates to dichromate, which is soluble with silver. No coloured end-point.
• Above pH 10: $AgOH$ and brown $Ag_2O$ precipitate, consuming titrant and falsifying the result.

Adjust with $NaHCO_3$ or a phosphate buffer if needed.

Sources of error

Gravimetric:

• Incomplete precipitation if insufficient precipitant is added (use a clear excess).
• Co-precipitation of impurities (acidify to remove carbonate; dilute the sample to reduce inclusion).
• Particle loss through the filter (digest the precipitate first to grow larger crystals).
• Incomplete drying (dry to constant mass; reweigh after a second drying cycle).
• Hygroscopic precipitates absorb moisture during weighing (cool in a desiccator).

Precipitation titration:

• pH out of range distorts the end-point.
• Slow precipitate formation makes the end-point hard to spot; swirl thoroughly.
• Indicator concentration: too much chromate masks the white-to-red change.
• Coloured samples (sea water with biological matter, for example) hide the end-point.

When to choose which

Examples in context

Example 1. Sulfate in mine drainage at Cobar. EPA contractors monitoring acid mine drainage from the Cobar copper mine in central NSW use gravimetric analysis to measure sulfate. A 50.0 mL aliquot is acidified with HCl and treated with excess $BaCl_2$ to precipitate $BaSO_4$. The precipitate is filtered through pre-weighed filter paper, washed, dried at 110 degrees C overnight and weighed. A typical mass of 0.282 g corresponds to $n(SO_4^{2-}) = 0.282 / 233.39 = 1.21 \times 10^{-3}$ mol, giving $[SO_4^{2-}] = 24$ mmol L$^{-1}$ or 2300 mg L$^{-1}$, far above the 250 mg L$^{-1}$ aesthetic guideline. The result feeds discharge licence compliance.

Example 2. Mohr titration of chloride in NSW pool water. Aquatic centre managers at the Sydney Olympic Park Aquatic Centre run weekly Mohr titrations to confirm chloride concentration in saltwater pools. A 25.0 mL pool sample is titrated against standardised 0.100 mol L$^{-1}$ silver nitrate using potassium chromate as indicator. The first persistent red tint of $Ag_2CrO_4$ marks the end point. A titre of 12.50 mL gives $n(Cl^-) = 0.100 \times 0.01250 = 1.25 \times 10^{-3}$ mol, hence $[Cl^-] = 0.0500$ mol L$^{-1}$ or 1775 mg L$^{-1}$, sitting in the operating range for salt chlorination. The HSC Mohr titration framework is the same chemistry the pool operator runs in the plant room.

Try this

Q1. Outline the steps of a gravimetric analysis for sulfate ions in water. [3 marks]

• Cue. Acidify, add excess $BaCl_2$, filter the $BaSO_4$ precipitate, wash, dry to constant mass, weigh, calculate moles and concentration.

Q2. A 100 mL water sample yields 0.247 g of $AgCl$ after addition of excess $AgNO_3$. Calculate $[Cl^-]$ in mol L$^{-1}$. [3 marks]

• Cue. $n(AgCl) = 0.247 / 143.32 = 1.72 \times 10^{-3}$ mol; $n(Cl^-) = 1.72 \times 10^{-3}$ mol; $[Cl^-] = 0.0172$ mol L$^{-1}$.

Q3. Mohr's method titrates $Cl^-$ with $AgNO_3$. (a) Write the equation for the indicator end point. (b) Explain why the pH must be between 7 and 9.5. (c) State one source of error in the Mohr titration. [2+2+1 marks]

• Cue. (a) 2Ag^+ + CrO_4^{2-} \rightarrow Ag_2CrO_4_{(s, red)}. (b) Acidic conditions protonate chromate to dichromate; basic conditions precipitate AgOH. (c) Co-precipitation of bromide or iodide; impure water blank.