Chemistry 102
EXPERIMENT 11
Determination of the Equilibrium Constant for a Reaction
Background
Organic esters are formed in an acid-catalyzed condensation reaction between a carboxylic acid
and an alcohol. Water is also produced. The reaction, called Fisher esterification, is an
equilibrium process whose equilibrium constant can be calculated if one can determine the
equilibrium concentrations of all species.
H+
O
'
R-COH + R OH
alcohol
acid
O
R-C-O-R' + H2O
ester
In this experiment, we will mix known amounts of the 2 reactants with a catalytic quantity of
strong mineral acid. The mixture must be heated to reach equilibrium in a reasonable time.
Periodically during the heating, we will withdraw a homogeneous sample of the reaction
mixture and will find the concentration of acid in the solution by titration. When the
concentration becomes constant (unchanged on 2 successive analyses) we assume the mixture
has reached equilibrium. From the results of the titration, we can determine the equilibrium
concentration of the carboxylic acid. Knowing all initial reactant concentrations and the
equilibrium concentration of one component allows us to calculate all equilibrium
concentrations. Then, from the equilibrium expression, we can solve for the equilibrium
constant value of the reaction.
One simplification makes use of the fact that the reaction mixture's volume is essentially
constant during the course of the reaction. Since the stoichiometry shows equal moles of
reactants and products, volume cancels from the equilibrium expression. Thus, the Keq
expression simplifies to the mole ratio of each component at equilibrium.
PROCEDURE:
You may work in pairs for this experiment. An individual lab report should be submitted by
each student.
Part 1:
Stock bottles of ~0.2M NaOH will be provided. This base must be standardized. Use
potassium hydrogen phthalate, KHC8H4O4 (KHP for short) as the primary standard.
KHC8H4O4 (aq) + NaOH (aq) → KNaC8H4O4 (aq) + H2O (l)
 First, calculate the appropriate mass of KHP required to titrate ~30 mL of the ~0.2M
base. Obtain a KHP sample and dissolve it in 25 mL of deionized water. Rinse and fill
Los Angeles City College
1
Chemistry 102
a buret with the base and, using phenolphthalein as the indicator, titrate to the endpoint.
Repeat for at least 3 good trials.
 Calculate the molarity of the NaOH solution for each trial then average the values.
Part 2:
Each pair of students will find an assembled reaction apparatus consisting of a 200-mL roundbottomed flask, a condenser with hoses attached to a water supply, a heating mantle and a
power controller. The round-bottomed flask can be removed from the heating mantle by
loosing the clamp on its neck and by also lowering the ring supporting the mantle. DO NOT
move the condenser or the water hoses! DO NOT turn on the power to the heating mantle until
the reactants have been mixed and the flask is secure once more.
Pipet into the round-bottomed flask 20.0 mL of ethanol (d=0.789 g/mL, MW = 46.1 g/mol) and
20.0 mL of the glacial acetic acid/H2SO4 mixture (0.20 mL conc. H2SO4 per 20.0 mL mixture;
for H2SO4, d=1.84 g/mL and MW=98.1 g/mol). Mix these thoroughly to form a homogeneous
solution then immediately pipet 1.00 mL of the mixture into a 125-mL Erlenmeyer flask. Add
about 20 mL of deionized water, 2 drops of phenolphthalein indicator and titrate to a pale pink
endpoint. Record the volume of base consumed.
Add a boiling stone to the liquid in the flask. Attach the round-bottomed flask securely to the
bottom of the water-filled condenser, resting on the heating mantle, and turn on the power
controller. Heat the flask so that the liquid turns to vapor but the vapor condenses in the bottom
third of the condenser and drips back into the flask. This procedure is known as "reflux".
Reflux the mixture for 45 minutes. Do not leave the mixture unattended; do not let the water
flow through the condenser stop and do not dislodge the drain hose from the sink.
During this time, refill the buret and prepare for another titration trial. Then, turn off the power
controller and drop the heating mantle. Cool the reaction mixture in an ice bath with the
condenser still attached to room temperature. Swirl to mix the mixture then pipet 1.00 mL of
the cooled reaction mixture to an Erlenmeyer flask, add indicator, and titrate. Record the base
volume consumed. Reflux the reaction mixture for an additional 15 minutes; cool, remove a
1.00-mL aliquot and titrate as before. If the last two base volumes differ by more than 0.50 mL,
reflux for an additional 15 minutes then titrate yet again.
CALCULATIONS:
(Assume that volumes are additive and that, as mentioned before, the total room temperature
volume of the reaction mixture does not change during the reaction)
Calculate the initial number of moles of ethanol in 1.00 mL of reaction mixture. Also, calculate
the number of moles of H+ contributed by H2SO4 in 1.00 mL of reaction mixture (assume both
protons are titrated).
From the average NaOH molarity and the volume used to titrate the initial 1.00 mL sample of
reaction mixture, calcuate the total moles of H+ in the reaction mixture. Subtract the moles
Los Angeles City College
2
Chemistry 102
contributed by H2SO4 to find the moles of H+ contributed by acetic acid (or, moles acetic acid
initially present).
Using the smallest base volume obtained in titrating the refluxed mixture, calculate the total
moles H+ at equilibrium and the moles of acetic acid at equilibrium. Use the stoichiometry of
the reaction to calculate the equilibrium number of moles of each other participant. Then,
calculate the Keq of the reaction.
QUESTIONS:
1. Why is it critical to keep water flowing through the condenser during the entire
experiment?
2. Estimate the H of this reaction using reference data either from your text or from the
CRC Handbook of Chemistry and Physics. Why do we heat the reaction mixture before
titrating samples? Will the heating then cooling have any effect on the equilibrium
position and on our calculated value for the Keq?
3. Briefly explain why the addition of 20 mL of water to each 1.00-mL of reaction mixture
does not affect the volume of NaOH consumed in the titration.
4. What is the purpose of the H2SO4? Why does the H+ contributed by the sulfuric acid
remain constant during the reaction while the total moles of H+ decreases?
Los Angeles City College
3