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TITRATION OF VINEGAR
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Introduction
Objective
To practice titration technique and to measure the concentration of acetic acid in vinegar by
titration with sodium hydroxide.
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Materials
250-mL Erlenmeyer flask, 50-mL buret, 10-mL graduated cylinder, vinegar solution,
standardized sodium hydroxide (NaOH) solution (0.10M), phenolphthalein indicator solution.
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Safety
Sodium hydroxide solution is caustic. Wear safety goggles at all times.
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Waste disposal
Unless otherwise instructed, liquid solutions can be poured down the drain with lots of water.
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Introduction
Acetic acid is probably the first acid discovered by man, and was probably discovered in sour
wine. Fermented grape juice produces ethanol (CH3CH2OH), which slowly oxidizes to acetic
acid (CH3COOH). Acetic acid has a sour taste, which led to one of the earliest definitions of an
acid – a substance having a sour taste.
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Many common materials have either acidic or basic properties. Acetic acid is an example of a
weak acid – dilute aqueous solutions of acetic acid only partly ionize to form H+ and CH3CO-2
(acetate anion). The majority of acetic acid molecules remain intact in solution. Strong acids,
such as hydrochloric acid (HCl) ionize virtually 100% in dilute aqueous solutions to form H+
and Cl-, with practically no intact HCl molecules left in solution.
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There are only about a dozen strong acids you will regularly encounter: the overwhelming
majority of acids are weak acids. Weak acids dissolved in water rapidly establish equilibrium
between their ionized and unionized form. The equilibrium for acetic acid is shown below.
CH3COOH(aq) ↔ H+(aq) + CH3COO-(aq)
Ka = 1.7 x 10-5
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The equilibrium constant for weak acid ionizations is called the acid dissociation constant and
given the general symbol “Ka”. For acetic acid ionization, the equilibrium expression has the
form:
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Students should remember that square bracket notation, [ ], indicates that the concentrations of
species are molar concentrations. All concentrations are equilibrium concentrations, not starting
concentrations.
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All acids containing hydrogen ion (H+) including weak acids react stoichiometrically with strong
bases such as sodium hydroxide (NaOH). In our titration, the chemical reaction between sodium
hydroxide and acetic acid is:
CH3COOH(aq) + NaOH(aq) → Na+(aq) + CH3COO-(aq) + H2O
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By carefully measuring the volume and knowing the concentration of sodium hydroxide
solution, the concentration of acetic acid can be readily calculated.
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The normal procedure for titrating a weak acid with standard sodium hydroxide solution requires
using a buret to add the sodium hydroxide solution. A fixed volume of acid is put into a suitable
container, such as an Erlenmeyer flask. Two or three drops of phenolphthalein are added to the
acid. Standard sodium hydroxide indicator solution is put into the buret. The sodium hydroxide
solution is added in small increments, with stirring, to the acid solution. Eventually, all of the
acid has reacted with sodium hydroxide, and excess sodium hydroxide turns the phenolphthalein
pink. The pink color indicates the titration is completed, and no additional sodium hydroxide
needs to be added to the Erlenmeyer flask.
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Normally, we determine the volume of sodium hydroxide added by difference between the
starting buret volume and the final volume (once the indicator has changed color). Students must
read the volume correctly, and the proper way of reading the volume is shown in Figure 9.1
below.
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Figure 9.1: Proper method for reading volume from a buret.