Apply volumetric analysis using acid-base titrations to determine unknown concentrations.

What this dot point is asking

SACE expects you to perform the full calculation reliably, justify your choice of indicator, and explain the role of standard solutions, primary standards and the difference between equivalence point and endpoint.

Lead worked calculation

Key terms

The equivalence point is where stoichiometrically equivalent amounts of acid and base have reacted, as defined by the balanced equation. The endpoint is where the chosen indicator changes colour. A well-chosen indicator makes the endpoint coincide with the equivalence point as closely as possible, so the titre is a true measure of the reacting amounts.

The central relationship

The general procedure for any titration calculation:

1. Write the balanced equation and read off the mole ratio.
2. Calculate moles of the standard reactant from $n = cV$ using the titre.
3. Use the mole ratio to find moles of the analyte.
4. Account for any dilution or aliquot scaling.
5. Divide by the relevant volume to get the concentration (or multiply by $M$ for a mass).

Choosing the indicator

The pH at the equivalence point depends on the salt formed, so the indicator must change colour within the steep part of the titration curve at that pH.

• Strong acid with strong base (e.g. $\text{HCl}$ and $\text{NaOH}$): equivalence pH $= 7$; the vertical region is wide, so methyl orange or phenolphthalein both work.
• Weak acid with strong base (e.g. ethanoic acid and $\text{NaOH}$): the salt is basic, equivalence pH $> 7$; use phenolphthalein (range $8.3$ to $10.0$).
• Strong acid with weak base (e.g. $\text{HCl}$ and $\text{NH}_3$): the salt is acidic, equivalence pH $< 7$; use methyl orange (range $3.1$ to $4.4$).

Back-titration

When an analyte reacts slowly, is insoluble, or is a solid such as an antacid, you add a known excess of one reagent, then titrate the leftover excess with a second standard. The amount that reacted with the analyte is the difference. For example, an antacid containing $\text{CaCO}_3$ is dissolved in a measured excess of standard $\text{HCl}$; the unreacted $\text{HCl}$ is then titrated with standard $\text{NaOH}$, and $n(\text{HCl reacted with carbonate}) = n(\text{HCl added}) - n(\text{HCl titrated})$.

Technique and reliability

• Rinse the burette and pipette with the solution they will hold; rinse the conical flask with distilled water only (residual water does not change the moles delivered into it).
• Perform a rough trial titre, then repeat for concordant titres within $0.10\ \text{mL}$ of each other and average only those.
• Read the burette to $\pm 0.05\ \text{mL}$ at the bottom of the meniscus at eye level to avoid parallax.
• The final answer can be no more precise than the least precise data; match significant figures accordingly.

Why it matters for monitoring

Acid-base titration delivers quantitative water-quality data, such as total acidity, alkalinity or the concentration of a dissolved species, using simple, low-cost equipment. It is the benchmark technique against which instrumental methods such as AAS and chromatography are calibrated, and it underpins quality control across food, environmental and industrial laboratories.