12.14-12.15 - pH Experiments And Ka

pH Data As Chemical Evidence

Two solutions can have the same concentration but very different $\mathrm{pH}$ values. That is not a measuring accident: $\mathrm{pH}$ is evidence about how many $\ce{H+}$ or $\ce{OH-}$ ions are actually present after particles have dissociated or reacted with water.

The example patterns below refer to about 25 ${}^{\circ}\mathrm C$. When the solutions are equimolar, concentration has been controlled. Differences in $\mathrm{pH}$ can then be linked to the chemical behaviour of the solute.

The $\mathrm{pH}$ values in real data depend on concentration and temperature, so questions usually provide the values or enough information to calculate them. Your job is to analyse the pattern, not memorise one table of numbers.

The salt explanations can be written as proton-transfer equilibria:

For weak ammonia, $\ce{NH3(aq) + H2O(l) <=> NH4+(aq) + OH-(aq)}$. These equations identify the source of acidity or alkalinity; the salt itself is normally fully dissociated into its ions.

Common-error contrast. A low $\mathrm{pH}$ does not automatically mean "concentrated acid". In this experiment the strong acid and weak acid may be equimolar, so the $\mathrm{pH}$ difference is mainly evidence about degree of dissociation, not about how much acid was put into the flask.

Dilution Patterns For Strong And Weak Acids

Dilution is a clean way to test whether $\mathrm{pH}$ follows the total acid concentration or the equilibrium position. A strong monoprotic acid is almost fully dissociated, so diluting it by a factor of $10$ reduces $[\ce{H+}]$ by a factor of $10$.

$\mathrm{pH} = -\log_{10}[\ce{H+}]$

If $[\ce{H+}]$ becomes ten times smaller, $\mathrm{pH}$ increases by 1.00.

The weak acid behaves differently because the acid equilibrium responds to dilution:

For weak-acid solutions where $[\ce{H+}]$ remains much smaller than the initial acid concentration $c$, and water ionisation is negligible:

$[\ce{H+}] \propto \sqrt{c}$

So a 10-fold dilution makes $[\ce{H+}]$ fall by about $\sqrt{10}$, not by $10$, and $\mathrm{pH}$ rises by about 0.5 rather than 1.0. The fraction of acid molecules dissociated increases on dilution, even though the total hydrogen ion concentration still decreases. The half-unit rule is approximate: once dissociation is a substantial fraction of c, retain $c - [\ce{H+}]$ rather than $c$ in $K_{\mathrm a}$. At extremely low acid concentrations water ionisation also matters; neither simple dilution rule should be extrapolated indefinitely.

Data-analysis habit. Do not identify the weak acid only because its $\mathrm{pH}$ is higher. Use the whole pattern: the strong acid has $\mathrm{pH}$ increases of roughly one unit for each 10-fold dilution; the weak acid has smaller $\mathrm{pH}$ increases because the equilibrium shifts as it is diluted.

Core Practical 9 Measurement Route

A dissociation constant connects a measured pH to equilibrium concentrations. The $\mathrm{pH}$ number is only useful if the solution concentration is known from a reliable preparation method.

[DIAGRAM: asset_name: Cp9 Ka From pH Apparatus; asset_slug: edexcel_a_level_chemistry_l056_cp9_ka_from_ph_apparatus; recommended_method: deterministic_chart; description: Core Practical 9 sequence showing a solid weak acid being weighed, dissolved and made up to a known volume in a volumetric flask, then measured with a calibrated pH meter or pH probe.]

A good route is:

1. Weigh a known mass of a pure, dry monobasic weak acid.
2. Dissolve the acid in deionised water in a beaker.
3. Transfer the solution quantitatively into a volumetric flask.
4. Rinse the beaker, stirring rod and funnel into the flask so all acid reaches the flask.
5. Make up to the calibration mark with deionised water, stopper and invert to mix.
6. Calibrate the $\mathrm{pH}$ meter or $\mathrm{pH}$ probe using buffer solutions near the expected $\mathrm{pH}$.
7. Rinse the probe with deionised water between solutions, gently blot it, then measure the $\mathrm{pH}$ of the weak acid solution at a consistent temperature.

Read the bottom of the meniscus at eye level and add the last drops with a dropping pipette. Wait for a stable probe reading; repeat measurements and investigate an outlier before deciding which results to average. A written account can assess planning, calculation and evaluation, but practical competence also requires supervised performance.

The final volume matters because mass alone gives amount of acid, not concentration. The $\mathrm{pH}$ meter matters because $\mathrm{pH}$ paper is usually too coarse for a reliable $K_{\mathrm a}$ calculation: an error of 0.1 $\mathrm{pH}$ unit changes $[\ce{H+}]$ by a noticeable factor.