Acid-base titrations with citric acid: part 1

David Cash PhD, Mohawk College (retired)

Introduction

Citric acid is a user-friendly, inexpensive, water-soluble, crystalline solid carboxylic acid. In comparison to using hydro-chloric acid solutions, using solid citric acid and its solutions in water is safer and more convenient. For inexperienced students its titrations against strong base are easier to perform. This article presents a neutralization titration of a citric acid solution by sodium hydroxide solution in a format suitable for beginner titrators. A second article will suggest applications of the same experiment that are suitable for experienced titrators.

Citric acid

Citric acid1 is produced by the fermentation of sugars (> 1 M tonne per year) for use in beverages and foods (70%); in detergents (20%); and in cosmetics, pharmaceuticals and other chemicals (10%). It is produced as a crystalline solid, either anhydrous, or as a monohydrate, and is available in either form at low cost. The solid monohydrate loses water below 100 °C when heated, forming the anhydrous solid, which melts at 156 °C, and decomposes at 175 °C. Both solids are stable and remain free-flowing in the bottle after many years of storage. It is very soluble in water. Citric acid can be purchased at low cost as a reagent chemical of either form at various purity levels,2 or at even lower cost as a consumer substance of unspecified purity and usually unspecified form.3

Titrating citric acid against strong base

Citric Acid

(a triprotic acid)

192.1 g / mol

H3C6H5O7

chemical structure of citric acid demonstrating it is a triprotic acid

pKa Values (1)

3.14

4.75

6.40

Citric acid has three carboxylic acid groups, three ionizable, acidic hydrogen atoms and three Ka/pKa  values. The neutralization reaction with sodium hydroxide has 3 to 1 stoichiometry as illustrated by the balanced complete neutralization equation. The reaction goes to completion and is suitable for analytical titrations:

3 NaOH (aq)  +  H3C6H5O7 (aq)  -->  Na2C6H5O7 (aq)  +  3 H2O

When titrated by a strong base such as 0.1 M NaOH solution, a solution of citric acid traverses a buffer region during which the pH of the solution climbs gradually then more steeply. From about one added drop of the 0.1 M NaOH solution before the equivalence point of the titration to about one added drop after the equivalence point the pH of the titration solution climbs extremely steeply from slightly below 7 to above 9.
A visual acid-base indicator chosen for the titration must change from its acid colour to its base colour in the range 7 to 9. Phenolphthalein indicator is an excellent choice for this titration, changing from colourless to pink to red. In comparison to titrations with a strong acid such as hydrochloric acid, the diluted weak acid “mops up” the base in the drops of added titrant more slowly. The result is that as the equivalence point is approached, even with swirling, the pink-red colour appears where the drops of titrant enter the titration flask and persists, at first for fractions of a second, then longer and longer, until a faint pink colour can be seen throughout the solution that persists for at least one minute (see Question 3). This phenomenon makes these titrations very easy to perform, and therefore very suitable for novices.

Safety, disposal of waste, and storage

Citric acid is a relatively strong weak acid, but no special precautions are required for its use. Citric acid powder is sold for home use with no restrictions. The pH of 0.033 M citric acid is about 2.2, which is slightly higher than that of lemon juice.4 The 0.1 M sodium hydroxide and the phenolphthalein indicator are more hazardous. Handle and clean up solid citric acid as you would solid sodium hydroxide.

Titrated solutions and excess of reagent solutions may be safely disposed of in a sink. At Mohawk College we store the dropper bottles containing the sodium hydroxide solution for long periods with no apparent ill effects. Storage of citric acid solution may not be advisable, since it may well support microbiological life.

Experiment description

Here is a simple, fast and inexpensive method for performing titrations of citric acid solution with sodium hydroxide solution. It is assumed that the titrations are being performed gravimetrically using inexpensive, unbreakable, 60-mL controlled drop-dispensing polymer squeeze bottles.5a

Preparation, apparatus and supplies

  • 2-place top-loading digital balance
  • Sufficient volume of 0.1 M NaOH (4.0 g/L)
  • Sufficient volume of 0.033 M citric acid (6.4 g/L)
  • Sufficient volume of phenolphthalein indicator in small labeled dropper bottles
  • 60-mL controlled drop-dispensing polymer squeeze bottles, labeled ‘NaOH’
  • 125-mL Erlenmeyer flasks
  • 10-mL graduated cylinders*
  • small beakers
  • dropper pipets
  • distilled water in squeeze bottles
  • waste containers

*A smaller bore diameter is better for precision.

Make up the two reagents in approximate fashion, such that your advance trial titrations indicate that a 5-mL sample of the citric acid solution is neutralized by roughly 5 g of the sodium hydroxide solution. As a sample calculation, label one or other with a “fictitious” value for concentration and have the students determine the concentration of the other solution by titration and calculation. (See calculations below.)

The full method and a sample titration should be demonstrated for novices. Specify a precision criterion of successful completion. A reasonable standard would be “three titration masses within ± 3% of the mean mass”. If using 50-mL burets, increase the volume of the citric acid samples from 5 mL to 10 mL to have roughly the same precision in the results.

a squeeze bottle over a flask and squeeze bottle on a balance

Gravimetric titration with a polymer controlled drop-dispensing squeeze-bottle and a 2-place digital balanceGravimetric titration with a polymer controlled drop-dispensing squeeze-bottle and a 2-place digital balance

Student instructions

  1. Transfer about 10 mL of the citric acid solution into a small beaker (100 or 250 mL). Use this portion of solution to rinse the inner surface of the small beaker, the 10-mL graduated cylinder and the dropper pipet. This rinse portion is waste. Repeat the rinsing twice more. DO NOT RINSE the 125-mL Erlenmeyer flask, it may be left wet, but only with distilled water. (See Question 1.)
  1. Transfer about 50 mL of the citric acid solution into the rinsed small beaker. This amount will be needed for the three trials. For trial 1 transfer some of the solution as carefully as you can into the 10 mL graduated cylinder, up to the 5.0 mL line, using the dropper pipet to adjust the bottom of the meniscus to the line.
  1. Transfer the 5.0-mL sample of citric acid solution for trial 1 from the cylinder into a 125-mL Erlenmeyer flask. Empty the cylinder totally, by waiting until the last drops fall. Add distilled water from the squeeze-bottle so that the total volume in the flask is between 20 mL and 30 mL. Add 4 to 5 drops of phenolphthalein indicator solution to the flask. Swirl the flask gently to completely mix the contents.
  1. Press the zero/tare button on the 2-place digital balance. Place the drop-dispensing bottle containing NaOH solution on the balance pan. The balance pan, the outside of the dropper bottle and your fingers must be kept dry at all times. (See Question 2.)
  1. Record the mass of the bottle and its contents.
  2. Titrate the solution in the Erlenmeyer flask by adding drops of solution from the drop-dispensing bottle. Hold the bottle upside down over the mouth of the flask. Count drops. Swirl the flask gently. The end point is reached when, on adding one drop of the base, the solution in the flask changes from colourless to pink or red and the colour remains for at least a minute.

If, after a multiple drop addition of base the solution is deeply red, the end-point may have been passed, and the titration result must not be used for calculations. If you “lose” a drop during a titration, the titration must not be used for calculations.

  1. Repeat Steps 4 and 5. Record the new mass of the bottle and its contents. Subtract your mass values to get the titration mass of the 0.1 M NaOH solution.
  1. Repeat the titration procedure. The drop count can serve as a guide to speed up the repeat titrations. Continue until the criterion of successful completion is achieved.

Calculations

The density of a 0.125 M NaOH solution at 20 °C is 1.0039 g/mL.6 For student calculations the density of a 0.1 M NaOH solution is so close to unity in g/mL units that the mass values of titrations in g units can be used as volumes in mL units without significant error.

My preferred method for introductory students is a 3-part calculation. Suppose that repeat titrations of 5-mL samples of citric acid solution produced a mean titration result of 4.87 g of 0.0989 M NaOH (= 4.87 mL of 0.0989 M NaOH):

1.

Calculate the moles of the known reagent (NaOH):

4.82 × 10-4 mol NaOH

2.

Calculate the moles of the unknown reagent (citric acid) using the balanced equation: 

1.61 × 10-4 mol citric acid

3.

Calculate the molarity of the citric acid solution:

0.0322 M citric acid

* Three significant digits (5.00 mL) are used in this sample problem. The precision of the volume measurement of the graduated cylinder lies between 5.0 mL and 5.00 mL. If 10.0 mL had been used, there would be three significant digits. It is up to you as a teacher whether you enter into this discussion with your students. Some find significant digits so confusing that it might be easier to have students use 5.0 mL.
 

Questions for students

  1. You were instructed to rinse the beaker, the graduated cylinder and the dropper pipet, but not the Erlenmeyer flask, with the citric acid solution. Explain.
  1. What errors occur if your fingers, the balance pan or the outside of the dropper bottle are wet with water?
  1. The very slow fading of the colour of the phenolphthalein indicator is blamed on carbon dioxide in the air reacting with the hydroxide ion in the solution. Write a balanced chemical equation for this reaction and explain why it causes the colour to fade.

Acknowledgements

The author thanks Randy Travis, technologist of the Department of Chemical, Environmental, and Biotechnology of Mohawk College, for his invaluable help.

References

  1. http://www.wikipedia.org for acid dissociation constants of citric acid.
  1. Sigma-Aldrich Citric acid (99 %) Citric acid monohydrate (98 %)

    3.  Boreal Science: Citric acid monohydrate (unspecified %):

Fibre Garden: Citric acid (unspecified form or %)

4.   Engineering Toolbox: Acids - pH values  Reference 4 lists values of pH for solutions whose concentrations are given in Normality units. For a citric acid solution 0.033 M = 0.100 N. There are some areas of technology where normality is still in use. Industrial advisors to college programs insist that this topic is important to our graduates.

5.   Articles available from the Web page:

Or contact David Cash for editable Word® versions of these articles.

6.   CRC Handbook (1973-74): Concentrative properties of aqueous
solutions - sodium hydroxide.    ∎