CHEM 43303
Experiment No. 01
The Oxidation of Formic Acid by Bromine
1. Introduction
Bromine oxidises formic (methanoic) acid to carbon dioxide according to the following
stoichiometric equation
HCOOH + Br2——>CO2+ 2H+ + 2Br In this experiment you will find the order of this reaction with respect to formic acid, bromine and H+. The
progress of the reaction is monitored electrochemically: the reaction takes place in an electrochemical cell
whose voltage (or electromotive force, EMF) depends on the concentration of bromine and bromide. This
is a very convenient way of following the reaction as all we have to do is read the EMF from a voltmeter
as a function of time.
The electrochemical cell has two electrodes. One electrode is simply a piece of platinum which dips
into the solution. The reaction that takes place at this electrode is the reduction of bromine to bromide
Br 2 +2e - ——> 2Br -
[1]
The second electrode is a Calomel electrode which consists of mercury metal in contact with a paste made
from insoluble Hg2Cl2, all surrounded by a saturated solution of KC1. The reaction which takes place at
this electrode is
Hg2Cl2(s) + 2e-——>2 Hg(m) + 2Cl[2]
The saturated KC1 solution is held in a glass container around the Hg and Hg2Cl2; a small porous plug
allows electrical contact between any solutions that the Calomel electrode is dipped into and the KC1
solution, but mixing of the two solutions is inhibited.
The EMF of a cell consisting of the Calomel electrode and the platinum electrode dipping into a
solution containing Br - and Br2 depends on the concentrations of all the species involved in the two
electrode reactions, Eqns. [1] and [2]. The Calomel electrode is in electrical contact with the solution, but
the contribution of the electrode to the EMF remains constant as all of the species in Eqn. [2] are
contained within the body of the electrode in such a way that their concentrations do not change. In this
experiment we arrange for the Br - ions to be in such large excess that their concentration is unchanged
during the experiment. Thus, any change in the EMF of the cell is attributable solely to changes in the
concentration of Br2.
The relationship between the cell EMF, E, and the concentration of Br2, [Br2], can be shown to be
E = constant + RT ln [Br2]
2F
[3]
where R is the gas constant (8.314 J K-1 mol-1), T is the absolute temperature and F is the Faraday constant
(which is the charge on one mole of electrons: 96,490 C mol-1). We are simply using the EMF as an
indicator of the concentration of bromine.
The relationship of Eqn. [3] holds provided that the cell is not producing any current. A simple voltmeter
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(i.e. one in which a pointer moves across a scale) does draw a significant current in making a
measurement; however modern electronic voltmeters draw currents of only a few nA. By using such an
instalment we can approach close to the ideal of measuring the voltage without drawing a current.
2. The Method in Principle
We will assume that the rate law can be written
rate = k[HCOOH]a[Br2]b[H+]c
where a, b and c are the orders with respect to formic acid, bromine and H+ respectively. The aim of the
experiment is to determine these orders.
As was explained above, an electrochemical cell essentially enables us to measure the concentration of
bromine, so it is convenient to express the rate in terms of the rate of loss of bromine. Expressing the rate
law using calculus we have
d[Br2]
a
b
+ c
[4]
= -k[HCOOH] [Br2] [H ]
dt
The minus sign in Eqn. [4] is to account for the fact that as the reaction proceeds the concentration of
bromine falls.
To simplify the kinetics we make up reaction mixtures in which formic acid and H + are in such large
excess compared to bromine that their concentration remains essentially constant throughout the reaction.
The rate law can then be written
d[Br2]
Where; kapp = k[HCOOH]ainitial[H+] cinitial
-kapp[Br2]b
=
dt
kapp is the apparent rate constant, called such because it is not really a true rate constant as its value
depends on the initial concentrations of formic acid and H+.
In the first part of the experiment we test the hypothesis that the order with respect to bromine is one,
i.e. b = 1. If this is the case then the rate law can be integrated as follows
d[Br2]
dt
1
∫ [Br ]
2
= -kapp[Br2]
∫
d [Br2] = -kapp dt
ln [Br2] = -kapp t + A
[5]
where A is the constant of integration. The form of Eqn. [5] tells us that, if the reaction is first order in
Br2, a plot of In [Br2] against t should give a straight line with slope -kapp. We have already seen above,
Eqn. [3], that the cell EMF is directly proportional to ln[Br2]
E = constant + RT ln [Br2]
2F
2
[3]
This can be rearranged to give
ln [Br2] =
2F (constant)
RT
2FE RT
and this expression for ln[Br2] is then substituted into Eqn. [5] to give
2FE
RT
-
2F (constant)
= - kappt + A
RT
This equation can be tidied up to the simple form
RT
E k app
t B
2F
[6]
where B is a constant at a given temperature. Equation [6] implies that a plot of E against time will
give a straight line of slope -kapp((RT/2F). If E is given in volts and t in seconds, the usual SI units for
all the other quantities will give an apparent first order rate constant, kapp, in s-1.
You should find that your data give good straight lines when plotted in this way, confirming that the
reaction is first order in bromine.
The second part of the experiment is to determine the orders with respect to H+ and formic acid. This
is done by varying the initial concentrations of these two species (still keeping them in excess, though)
and then determining kapp . By comparing values of kapp determined for reaction mixtures with different
concentrations of one species and the same concentration of the other we can determine the order with
respect to each in the following way:
For a particular reaction mixture, let the initial concentration of formic acid be [HCOOH]initial and the
initial concentration of H+ be [H+]initial. The apparent rate constant, kapp, is thus
kapp = k[HCOOH]ainitial[H+]cinitial
taking natural logarithms of both sides gives
ln kapp= ln k + a ln {[HCOOH]initial}+ c ln {[H+]initial}
[7]
This relationship implies that if kapp is measured for a series of reaction mixtures with different values of
[HCOOH]initial, but the same value of [H+]initial , a plot of lnkapp against ln{[HCOOH]initial} will have slope a,
the order with respect to formic acid. Likewise for a series of reaction mixtures with different values of
[H+]initial, but the same value of [HCOOH]initial, a plot of lnkapp against ln{[H+]initial} will have slope c, the
order with respect to H+.
In this experiment you will determine the kapp for six different reaction mixtures, enabling you to obtain
kapp for three different initial concentrations of formic acid and the same of H+.
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3. The Apparatus
Work in pairs when recording the data, but analyze the data separately.
The apparatus is basically an electrochemical cell formed by dipping the Calomel electrode and the
platinum electrode into the reaction mixture. A magnetic stirrer is used to keep the solution mixed. An
electronic voltmeter is used to measure the cell EMF.
Ideally the reaction mixture should be kept in a thermostat so as to keep the temperature constant.
However, the heat evolved in this reaction is not sufficient to raise significantly the temperature of the
reaction mixture. It is possible, therefore, to dispense with the thermostat.
4. Procedure
You are provided with a solution which is 6.0 x 10-3 in Br2 and 0.2 M in NaBr, a 2.0 M solution of formic acid
and a 2 M solution of HC1. You will record the progress of six different kinetic runs using the following
volumes of solution:
run
vol. Br2+NaBr
solution / cm3
vol. formic acid
solution / cm 3
vol. HC1 solution / cm3
vol. H2O / cm3
1
2
50
50
5
5
2.50
6.00
42.50
39.00
3
50
5
12.50
32.50
4
50
6
7.50
36.50
5
50
8
7.50
34.50
6
50
11
7.50
31.50
Note that the total volume of each reaction mixture is 100 cm3, 50 cm3 of which is always made up from the
bromine stock solution. The reaction mixtures for runs 4, 5 and 6 have the same concentration of H but
different concentrations of formic acid. The reaction mixtures for runs 1, 2 and 3 have the same concentration
of formic acid but different concentrations of H+.
Use pipettes (either fixed volume or graduated, as appropriate) to measure out the solutions. As the
bromine tends to evaporate from the stock solution, pipette this directly from the bottle and keep the bottle
stoppered when not in use. Do not pipette the other solutions directly from the stock bottles - transfer some
first to a beaker or conical flask. Also, take care not to pollute the solutions by using the same pipette for
different solutions.
*Always use pipette fillers - do not suck up any solution by mouth. Take care when fitting a
pipette filler to the pipette: hold the pipette by the end to which you are attaching the filler.
*You must record your raw data and observations as you make them on the experimental report
form.
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Preliminary steps
1. Pipette 50 cm3 of the Br2 + NaBr solution into a clean 150 cm3 tall-form beaker. Put the magnetic
"flea" into the beaker.
2. Set up the cell with the beaker placed on the magnetic stirrer and with a thermometer, the Calomel
electrode (carefully remove the rubber cap) and the platinum electrode dipping into the solution (the
end of the Calomel electrode only needs to be just immersed). The electrodes and the thermometer
should be clamped lightly using retort clamps which can themselves be attached to the bench rack.
Set the stirrer going gently, checking that the flea does not foul the electrodes.
With ingenuity you should be able to set the electrodes at such a height that you can remove the
beaker simply by sliding away the stirrer rather than having to move the electrodes up.
3. Connect the platinum electrode to the positive (red) terminal on the voltmeter and the
Calomel electrode to the negative (black). Ask the demonstrator if you are uncertain how to use or
read the voltmeter. You should find that the voltage is about 0.8 V; if it is not, consult the
demonstrator. Over the course of time the voltage may vary by a few mV - this is quite normal.
4. When you are happy with the set-up, move onto making the kinetic runs. You can use the 50 cm3
of bromine solution you already have in the beaker for your first run.
5. Make sure that you know how to operate the stop-watch.
Kinetic Runs
For each run the Br + NaBr solution is put into the beaker and the cell set up as above. The formic acid, HC1
and distilled water are pre-mixed in a 50 cm3 volumetric flask. As the total volume of this mixture is always 50
cm3 all you need to do is measure the hydrochloric acid and formic acid into the volumetric flask and then
make it up to the line with distilled water. Stopper the flask and invert it a few times to mix the contents.
The reaction is initiated by pouring the contents of the volumetric flask into the cell. The voltage is then read
off as the reaction proceeds. As the reaction is first order it is not necessary to start the stopwatch at the precise
time of mixing (which is ill-defined in any case). Start the stop-watch once you have added all of the pre-mixed
solution.
6. Pipette 50 cm3 of the Br2+NaBr solution into a clean 150 cm3 tall-form beaker. Put the magnetic
"flea" into the beaker and set up the cell as above; set the stirrer going. Check that the voltage is
around 0.8 V.
7.
Mix the quantities of formic acid, HC1 and distilled water in a 50 cm3 volumetric flask, as
described above.
.
•
8. Record the initial voltage from the cell
9. Add the contents of the flask to the beaker (a small funnel may help in doing this), start the stopwatch and commence recording the voltage at 1 minute intervals until the voltage has dropped by
about 0.03 V from its starting value or after 10 minutes, whichever is the sooner.
10. At some point in the run, read and record the temperature.
11. After making the measurements stop the stirrer, remove the beaker and discard the solution being
careful to retain the flea. Rinse the beaker and the flea well in tap water and then in distilled water;
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wipe them both dry with tissue (you will need both for subsequent runs).
12. Rinse the electrodes and thermometer with distilled water (use a wash bottle and let the waste run
into a separate beaker). Gently dab any excess water from the electrodes.
13. Follow the procedure, i.e. steps 6 to 12, for kinetic runs 2 through to 6.
Finishing Off
14. Rinse all of the glassware you have used and leave it to drain. Having cleaned the electrodes
(step 12), replace carefully the rubber cap onto the Calomel electrode.
5. Data Analysis
Use the tables in the experimental report form. For each kinetic run plot the cell voltage against time (in
seconds), draw the best fit straight line and determine its slope. At the start of the reaction the voltage may rise
before falling steadily for the remainder of the time (we are unsure as to why this happens). If this happens,
ignore these early points when drawing the best fit straight lines.
Convert the slopes to first order rate constants, kapp , using Eqn. [6]. Fill in the table of
concentrations for runs 1-5 on page 1 of the report form, and the summary table on page 5; determine the order
with respect to formic acid and H+ by plotting two graphs:
(1) lnkapp against ln{[HCOOH]initial} for constant [H+]initial
(2) lnkapp against ln{[H+]initial} for constant [HCOOH]initial
Answer the following questions at the end of the report form.
(1) The following mechanism has been proposed for this reaction:
First, a rapid pre-equilibrium involving the dissociation of formic acid to formate (the H+ is, of course,
present as H3O+)
H+ + HCOO-
HCOOH
Equilibrium constant Ka
Then the slow irreversible rate determining step in which Br2 is consumed
HCOO - + Br2
slow
H+ + CO2 + 2Br +
rate constant K
To what order with respect to Br2 , HCOOH and H would you expect this reaction scheme to give
rise? Explain your reasoning and, if you can, show the overall rate law can be derived.
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