Lab Report Sample: Acid-Base Titration: Standardising NaOH

Model Answer

Abstract

This report describes the standardisation of an approximately 0.1 mol dm⁻³ sodium hydroxide (NaOH) solution by acid–base titration against potassium hydrogen phthalate (KHP, C₈H₅KO₄), a primary standard acid with a high molar mass and known purity. Three separate samples of solid KHP were dissolved and titrated against the NaOH solution using phenolphthalein as indicator. Concordant titres (within 0.10 cm³) were obtained in all three trials, giving a mean NaOH concentration of 0.1028 ± 0.0004 mol dm⁻³. The small spread between replicate results indicates a precise and reliable standardisation, suitable for use as a working standard in subsequent quantitative analyses.

Introduction

Sodium hydroxide is a widely used base in quantitative analysis, but solid NaOH is hygroscopic and readily absorbs carbon dioxide and water vapour from the air, so a freshly prepared NaOH solution cannot be assumed to have exactly the concentration implied by the mass weighed out (Atkins and de Paula, 2014). NaOH is therefore classed as a secondary standard: its exact concentration must be determined experimentally by titration against a primary standard, a substance available in a very pure, stable, non-hygroscopic form with a precisely known molar mass (Harris, 2015).

Potassium hydrogen phthalate (KHP) is one of the most commonly used primary standards for standardising bases. It is a solid monoprotic acid, is not hygroscopic, can be dried and weighed accurately on a standard laboratory balance, and has a relatively high molar mass of 204.22 g mol⁻¹, which reduces the percentage error associated with weighing compared with standards of lower molar mass (Skoog, West and Holler, 2013). A useful general principle in analytical chemistry is that, for a fixed absolute weighing uncertainty, a standard of higher molar mass requires a larger mass to be weighed out for a given number of moles, which reduces the relative (percentage) weighing error; this is one of the main reasons KHP is preferred over lower-molar-mass acids such as oxalic acid for everyday standardisations (Harris, 2015). KHP reacts with NaOH in a simple 1:1 stoichiometric ratio: KHC₈H₄O₄ + NaOH → KNaC₈H₄O₄ + H₂O. Because the reaction is between a weak acid and a strong base, the equivalence point occurs on the alkaline side of pH 7, making phenolphthalein (colour change range approximately pH 8.2–10.0) an appropriate indicator, as its transition range closely brackets the equivalence-point pH of the resulting solution (Harris, 2015). By contrast, an indicator such as methyl orange, which changes colour in a strongly acidic range, would give an end point well before the true equivalence point in this particular reaction and would therefore be unsuitable. The aim of this experiment was to determine an accurate, precise value for the concentration of a given NaOH solution, to quantify the uncertainty associated with that value from the apparatus used, and to identify which piece of apparatus contributed the greatest source of error to the final result.

Method

Apparatus: analytical balance (±0.0001 g), 50 cm³ burette, 25 cm³ conical flasks, glass funnel, white tile, wash bottle of distilled water, phenolphthalein indicator solution, solid potassium hydrogen phthalate (KHP), and the sodium hydroxide solution to be standardised.

Procedure: the burette was rinsed with a small volume of the NaOH solution before being filled to just above the zero mark and the initial reading recorded to the nearest 0.05 cm³. Approximately 0.50 g of dried KHP was weighed accurately into a conical flask using the analytical balance, and the exact mass recorded. The KHP was dissolved in approximately 25 cm³ of distilled water (the exact volume of water used does not affect the number of moles of KHP present and therefore does not need to be measured precisely), and two to three drops of phenolphthalein indicator were added. A rough titration was performed first, adding NaOH steadily while swirling until the solution just turned permanently pale pink, to give an approximate titre. Three further accurate titrations were then carried out, adding NaOH dropwise near the expected end point until a single drop produced a permanent, faint pink colour that persisted for at least 30 seconds. The burette reading was recorded to the nearest 0.05 cm³ before and after each titration, and the process was repeated with fresh KHP samples until three concordant titres (within 0.10 cm³ of one another) were obtained.

Controlled variables and precautions: the same batch of NaOH solution and the same burette were used for every titration to avoid introducing systematic differences between trials. The conical flask was swirled continuously during each titration to ensure the KHP solution mixed thoroughly with the added NaOH, and the inside of the flask was rinsed periodically with distilled water from a wash bottle so that no reagent adhering to the glass was left unreacted. The burette was read at eye level with the meniscus aligned to the graduation mark to avoid parallax error, and the bottom of the meniscus was used consistently for both the initial and final readings. Standard laboratory safety precautions were followed throughout: safety spectacles were worn at all times, as both the NaOH solution and solid KHP can irritate the eyes and skin, and any spillages were reported and cleaned immediately with plenty of water.

Results

Table 1 shows the mass of KHP used, the titre volume of NaOH required, and the resulting NaOH concentration calculated for each of the three concordant trials.

Table 1. Mass of KHP, titre volume and calculated NaOH concentration for three concordant trials.

Worked calculation (Trial 1). The number of moles of KHP used was calculated from its mass and molar mass (204.22 g mol⁻¹):

n(KHP) = 0.509 ÷ 204.22 = 2.4926 × 10⁻³ mol

Since KHP and NaOH react in a 1:1 molar ratio, the moles of NaOH consumed equal the moles of KHP present. The concentration of NaOH was then found by dividing the moles of NaOH by the titre volume, converted to dm³:

c(NaOH) = 2.4926 × 10⁻³ mol ÷ 0.02425 dm³ = 0.1028 mol dm⁻³

The same calculation was repeated for Trials 2 and 3, giving concentrations of 0.1028 and 0.1028 mol dm⁻³ respectively (Table 1), and the mean of the three trials was 0.1028 mol dm⁻³, with a maximum deviation from the mean of only 0.0001 mol dm⁻³.

Uncertainty analysis. Each burette reading carries an uncertainty of ±0.05 cm³, so a titre calculated from two readings (initial and final) carries a combined uncertainty of ±0.10 cm³. For Trial 1, this represents a percentage uncertainty of:

(0.10 ÷ 24.25) × 100 = 0.41%

The analytical balance contributes a much smaller relative uncertainty, since a reading uncertainty of ±0.0001 g on a mass of 0.509 g gives (0.0001 ÷ 0.509) × 100 = 0.02%. Combining these in quadrature gives a total percentage uncertainty of approximately 0.41%, equivalent to an absolute uncertainty of about ±0.0004 mol dm⁻³ on the final concentration, so the standardised concentration is reported as 0.1028 ± 0.0004 mol dm⁻³.

Discussion

The three titres obtained were highly concordant, with calculated concentrations agreeing to within 0.0001 mol dm⁻³ of one another, and the resulting standardised concentration of 0.1028 ± 0.0004 mol dm⁻³ is close to the nominal 0.1 mol dm⁻³ that the solution was intended to have. This small but real difference between the intended and measured concentration illustrates precisely why standardisation is necessary: NaOH pellets absorb atmospheric moisture and carbon dioxide during weighing and storage, so a solution prepared by weighing out a target mass of NaOH will typically be slightly more dilute, and occasionally slightly more concentrated through weighing error, than intended (Atkins and de Paula, 2014).

The dominant source of uncertainty in this experiment was the burette reading, contributing roughly twenty times more relative uncertainty than the analytical balance. This is a common feature of titration-based standardisations and explains why good technique near the end point – adding titrant dropwise and rinsing the flask walls with distilled water to ensure all reagent reaches the bulk solution – is emphasised so strongly in analytical chemistry teaching (Skoog, West and Holler, 2013). Using a larger mass of KHP per trial, and therefore a larger titre volume, would reduce the relative contribution of the fixed ±0.05 cm³ burette reading uncertainty; for example, doubling the titre volume while keeping the reading uncertainty fixed would approximately halve the percentage uncertainty from the burette.

A further consideration is the choice of indicator. Phenolphthalein was appropriate here because the KHP/NaOH reaction produces a weakly alkaline solution at the equivalence point, but it would be unsuitable for titrations producing an acidic or near-neutral equivalence point, where an indicator such as methyl orange would be preferred instead (Harris, 2015). A small systematic error can also arise from adding one drop of NaOH too many past the true equivalence point before the colour change is judged to be permanent; this ‘indicator error’ is difficult to eliminate entirely by eye but is minimised here by titrating dropwise near the end point and confirming the colour persisted for at least 30 seconds before recording the final reading.

Overall, the closeness and consistency of the three concordant results, together with the small calculated uncertainty of ±0.4%, suggest that this standardisation was carried out to a good analytical standard and that the resulting NaOH solution could be used with confidence as a secondary standard in further quantitative work, such as determining the concentration of an unknown acid or the purity of a solid acidic sample. It is worth noting that concordance between repeats indicates good precision – that the method is repeatable – but does not on its own guarantee accuracy, since a systematic error such as an incorrectly calibrated burette or a contaminated KHP sample could produce results that are precise but consistently offset from the true value. Cross-checking the standardised NaOH solution against a second, independently prepared primary standard, or comparing the result with a certified reference material, would provide stronger evidence that the value obtained is both precise and accurate.

Comparing this result with typical undergraduate teaching values, a well-executed KHP standardisation is generally expected to agree with the nominal concentration to within about 1–2%, and to show a spread between replicate titres of well under 1% (Vogel, 2013). The results obtained here, with replicate concentrations agreeing to within 0.01% of one another and a combined percentage uncertainty of 0.41%, comfortably meet this benchmark and indicate that the technique used – slow, dropwise addition near the end point, thorough swirling, and consistent meniscus reading – was applied effectively throughout the three trials.

Conclusion

Titration of three concordant KHP samples against the sodium hydroxide solution gave a standardised concentration of 0.1028 ± 0.0004 mol dm⁻³, close to the intended 0.1 mol dm⁻³ value. The burette reading was identified as the dominant source of uncertainty, contributing roughly twenty times more relative error than the analytical balance, and could be further reduced in future work by using a larger mass of KHP per trial to increase the titre volume. This standardisation demonstrates the standard analytical approach for determining the true concentration of a hygroscopic base and provides a reliable working solution for subsequent titrations.

References

• Atkins, P. and de Paula, J. (2014) Atkins’ Physical Chemistry. 10th edn. Oxford: Oxford University Press.
• Harris, D.C. (2015) Quantitative Chemical Analysis. 9th edn. New York: W.H. Freeman.
• Skoog, D.A., West, D.M. and Holler, F.J. (2013) Fundamentals of Analytical Chemistry. 9th edn. Belmont: Cengage Learning.
• Housecroft, C.E. and Constable, E.C. (2010) Chemistry: An Introduction to Organic, Inorganic and Physical Chemistry. 4th edn. Harlow: Pearson.
• Vogel, A.I. (2013) Vogel’s Textbook of Quantitative Chemical Analysis. 6th edn. Harlow: Pearson.
• Miller, J.N. and Miller, J.C. (2010) Statistics and Chemometrics for Analytical Chemistry. 6th edn. Harlow: Pearson.
• Christian, G.D., Dasgupta, P.K. and Schug, K.A. (2014) Analytical Chemistry. 7th edn. Hoboken: Wiley.
• Royal Society of Chemistry (2016) Titration Techniques: A Practical Guide. London: Royal Society of Chemistry.
• Day, R.A. and Underwood, A.L. (1991) Quantitative Analysis. 6th edn. Englewood Cliffs: Prentice Hall.