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IB HL Chemistry Investigating Acids Internal Assessment May 2012
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Claudia Braganza IBHL Chemistry Grade 12
Page | 1
IBHL Investigations: Investigating Acids
Aim
To find out the effect of different temperatures (25oC, 30oC, 35oC, 40oC, 45oC and 50oC)
on the Ka value of 1 mol dm-3 of ethanoic acid (CH3COOH)
Introduction & Hypothesis
A weak monobasic acid HA, such as ethanoic acid, reacts with water according to this
equation:
HA (aq) ⇌ H+ (aq) + A- (aq)
CH3COOH (aq) ⇌ CH3COO- (aq) + H+ (aq)
The equilibrium constant for this reaction is known as the acid dissociation constant, Ka,
and has units of mol dm-3.
Ka = [H+] [A-]/ [HA]
The acid dissociation constant is a measure of the strength of a weak acid. The larger the
value of Ka, the stronger the acid and the greater the extent of ionization or dissociation.
Since acid dissociation constants (Ka) tend to be small and vary considerably, they are
often expressed as pKa values where:
pKa = - log10 Ka (cf. [H+] and pH)
Values of pKa are also a measure of acid strength, but now the smaller the value of pKa
the stronger the acid. A change of 1 in the value of the pKa means a change in acid
strength of a factor of 10 (cf. [H+] and pH).
Acid dissociation constants are not usually quoted for strong acids because these
effectively undergo complete ionization or dissociation in water. Their dissociation
constants are very large and tend towards infinity in dilute solutions. It is difficult to
measure them accurately because the concentration of undissociated acid molecules is
so low. This is why Ka values are usually quoted only for weak acids, like ethanoic acid.
Values of Ka and pKa are equilibrium constants, and like other equilibrium constants, are
not affected by changes in concentration, only by changes in temperature. This means
that acid strengths vary with temperature and that the order of acid strengths can vary
with temperature.
The pH of a solution of a weak acid can only be calculated if the acid dissociation
constant, Ka, (or pKa) is known.
Ka = [H+] [A-]/ [HA]
Claudia Braganza IBHL Chemistry Grade 12
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But since [H+] = [A-], in a solution where only the acid is present:
Ka = [H+] / [HA]
Rearranging:
[H+] = √[HA] x Ka
And then
pH = - log10 [H+]
This approach can be reversed in order to calculate the Ka (and hence pKa, or vice versa)
of a weak acid if you know the pH of the solution and its concentration. In this
experiment, this is exactly what I will be doing as I will know the initial concentration
and pH of ethanoic acid and from there calculate the pKa value from which the Ka value
will be derived from.
The calculations that I will use are as follows:
pH = (Average of 4 trials)
[H+] = 10 ^ (- pH)
Then I will use Henderson-Hasselbalch equations to calculate the pKa and thus Ka.
pH = pKa + log [A-]/[HA]
Since it is a solution where only the acid is present, [H+] = [A-]. pKa will be calculated
from rearranging the equation to get pKa. Then, Ka will be calculated as follows:
Ka = 10 ^ (- pKa)
The reason why Ka values only vary with temperature is as follows, and I will explain it
by using pH values first.
pH is a measure of the [H+] ion concentration (potential of hydrogen ion) and is
independent of the volume of the solution. pH can indicate the acidity of a solution as it
is a measure of [H+] ion concentration. As the investigation is regarding temperature’s
effect on the Ka value of ethanoic acid, the initial pH values and subsequent ones can
indicate whether the reaction is an exothermic or endothermic one. As the [H+]
increases with temperature, we know that the reaction is endothermic, according to Le
Chatelier’s principle as illustrated below:
CH3COOH (aq) ⇌ CH3COO- (aq) + H+ (aq)
Ka = [H+] [CH3COO-]/ [CH3COOH]
As temperature increases, the particles start to collide faster and the kinetic energy of
the molecules increases. This makes the concentration of ions will increase and the
Claudia Braganza IBHL Chemistry Grade 12
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forward reaction is favored. This means the concentration of the acid itself decreases in
comparison to its ions.
Ka = [↑] [↑]/ [↓]
When this happens, the Ka value will increase along with the ion concentration. This is
why the acid dissociation constant is only affected by temperature. As Ka values
increase, it is known that it indicates the increase in acidity of the solution. Therefore,
my hypothesis is that as temperature increases, the Ka value will also increase, thereby
increasing the acidity of 1 mol dm-3 of ethanoic acid.
Claudia Braganza IBHL Chemistry Grade 12
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Apparatus and Materials
1) 200ml beaker x 2
2) Pipette 10ml x 1
3) pH meter x 1
4) Stir plate heater x 1
5) Thermometer x 1
6) 750ml of ethanoic acid (CH3COOH) x 1
7) 100ml calibration solution pH 4 x 1
8) 100ml calibration solution pH 7 x 1
9) pH meter screw x 1
10) Pen and paper x 1
Safety and Precautions
1) Always wear lab goggles.
2) Clean any spills immediately as some solutions can stain or be hazardous. Clean
it by wiping inwards with a paper towel. Then, immediately wash hands.
3) Handle all equipment with care.
4) Keep electrical equipment far from contact with water.
5) Always clean glassware before and after it is used. Using defective glassware can
lead to accidents as well as experimental errors during calculations.
6) Wash hands before and after lab work.
Variables
Controlled
What is controlled? How is it controlled? Why is it controlled? Concentration of ethanoic
acid It is controlled by making the concentration 1 mol
dm-3
It is controlled because although Ka is only affected by temperature, the experiment still needs to be controlled so that it doesn’t interfere with data collection.
Volume of acid for each data point
It is controlled by pipetting 30ml of the acid for each
data point
It is controlled to make the experiment a fair trial.
pH meter It is controlled by calibrating it beforehand in
pH 4 and pH 7 solutions
It is controlled to ensure that pH readings don’t differ from each other and interfere with data collection.
Claudia Braganza IBHL Chemistry Grade 12
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Pressure in the room It is controlled by conducting the experiment in a room at standard 1 atm
pressure
It is controlled because although Ka is only affected by temperature, the experiment still needs to be controlled so that it doesn’t interfere with data collection.
Independent: Temperature (25oC, 30oC, 35oC, 40oC, 45oC and 50oC)
Dependent: pH value during experiment which then determines final Ka value
Method
1) Prepare all apparatus and materials immediately. Find a clean working space
with ample space to carry out experiment safely.
2) First, prepare the stir plate heater by connecting it a plug point. Don’t turn it on
at this point.
3) Prepare the pH meter and the calibration solutions of pH 4 and 7.
4) Dip the pH meter into the calibration solution of pH 4. Wait for an unchanging
value.
5) Depending on how much the value is above or below 4, use the screw of the pH
meter to turn the bolt until the value on the meter screen displays 4.
6) Wash the pH meter before calibrating it with a solution of pH 7.
7) Dip the pH meter into the calibration solution of pH 7. Wait for an unchanging
value.
8) Depending on how much the value is above or below 7, use the screw of the pH
meter to turn the bolt until the value on the meter screen displays 7.
9) Wash the pH meter again.
10) Prepare a 200ml beaker of water and put the pH meter inside it.
11) Now, turn on the stir plate heater and turn it option 4 or 5. Leave it be.
12) Move on to preparing the solution of ethanoic acid for the trials. From the 750ml
inside the bottle, pipette out 30ml of the acid into the awaiting 200ml beaker.
13) Measure the temperature to make sure it is 25oC (RT). Then, measure the pH at
RT using the meter. Record both values.
14) Put the beaker onto the stir plate heater to heat the acid. Tilt the beaker a little
to make sure the thermometer’s tip is fully submerged in the acid.
15) Once the value reaches the next temperature value (i.e. 30oC), take the beaker
away from the heater.
16) Measure the pH using the meter. Record the value. Put the pH meter back into
the water-filled beaker so that it stays calibrated.
17) Repeat steps 12-16 for all other data points. Repeat the procedure for three
more trials.
Claudia Braganza IBHL Chemistry Grade 12
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18) At the end, don’t forget to clean up all apparatus and material.
Data collection
I have just collected the pH values of ethanoic acid at different temperatures as shown
below. These values will then be converted into pKa values, from which the Ka value will
then be derived.
Trial pH +0.01 Temperature (oC)+1oC
25 30 35 40 45 50 1 2.36 2.24 2.12 2.05 1.92 1.80 2 2.37 2.26 2.13 2.03 1.90 1.78 3 2.36 2.24 2.13 2.05 1.90 1.78 4 2.36 2.26 2.13 2.03 1.91 1.80
Claudia Braganza IBHL Chemistry Grade 12
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Data processing
At 25oC (RT)
pH = (2.36 + 2.37 + 2.36 + 2.36) / 4 = 2.36
[H+] = 10 ^ (- 2.36) = 0.0044 mol dm-3
pH = pKa + log [A-]/[HA]
[A-] = [H+]
2.36 = pKa + log [0.0044]/[1.00]
pKa = 4.72 (this value is very close to the data booklet value of 4.76)
Ka = 10 ^ (- 4.72)
= 1.905 x 10-5 mol dm-3
At 30 oC
pH = (2.24 + 2.26 + 2.24 + 2.26) / 4 = 2.25
[H+] = 10 ^ (- 2.25) = 0.0056 mol dm-3
pH = pKa + log [A-]/[HA]
[A-] = [H+]
2.25 = pKa + log [0.0056]/[1.00]
pKa = 4.50
Ka = 10 ^ (- 4.50)
= 3.162 x 10-5 mol dm-3
Claudia Braganza IBHL Chemistry Grade 12
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At 35oC
pH = (2.12 + 2.13 + 2.13 + 2.13) / 4 = 2.13
[H+] = 10 ^ (- 2.13) = 0.0074 mol dm-3
pH = pKa + log [A-]/[HA]
[A-] = [H+]
2.13 = pKa + log [0.0074]/[1.00]
pKa = 4.26
Ka = 10 ^ (- 4.26)
= 5.495 x 10-5 mol dm-3
At 40oC
pH = (2.05 + 2.03 + 2.05 + 2.03) / 4 = 2.04
[H+] = 10 ^ (- 2.04) = 0.0091 mol dm-3
pH = pKa + log [A-]/[HA]
[A-] = [H+]
2.04 = pKa + log [0.0091]/[1.00]
pKa = 4.08
Ka = 10 ^ (- 4.08)
= 8.318 x 10-5 mol dm-3
Claudia Braganza IBHL Chemistry Grade 12
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At 45oC
pH = (1.92 + 1.90 + 1.90 + 1.91) / 4 = 1.91
[H+] = 10 ^ (- 1.91) = 0.0123 mol dm-3
pH = pKa + log [A-]/[HA]
[A-] = [H+]
1.91 = pKa + log [0.0123]/[1.00]
pKa = 3.82
Ka = 10 ^ (- 3.82)
= 1.514 x 10-4 mol dm-3
At 50oC
pH = (1.80 + 1.78 + 1.78 + 1.80) / 4 = 1.79
[H+] = 10 ^ (- 1.79) = 0.0162 mol dm-3
pH = pKa + log [A-]/[HA]
[A-] = [H+]
1.79 = pKa + log [0.0162]/[1.00]
pKa = 3.58
Ka = 10 ^ (- 3.58)
= 2.630 x 10-4 mol dm-3
Claudia Braganza IBHL Chemistry Grade 12
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The Ka values calculated above are shown below with its respective temperatures.
Now, the uncertainties must be calculated in order to get an idea of the errors.
Instrument Percentage Error Thermometer 1% Pipette (0.05/10) x 100 = 0.5% pH 0.01% Total random error 1 + 0.5 + 0.01 = 1.51%
Therefore, with errors the Ka values will be as follows:
Example calculation
At 30oC Ka = 3.162 x 10-5 mol dm-3 + 1.51%
Ka = 3.162 x 10-5 + 4.774 x 10-7 mol dm-3
Temperature (oC) +1oC Ka (mol dm-3) 25 1.905 x 10-5 30 3.162 x 10-5
35 5.495 x 10-5 40 8.318 x 10-5
45 1.514 x 10-4
50 2.630 x 10-4
Temperature (oC) +1oC Ka (mol dm-3) 25 1.905 x 10-5 +2.877 x 10-7
30 3.162 x 10-5 +4.775 x 10-7 35 5.495 x 10-5 +8.297 x 10-6
40 8.318 x 10-5 +1.256 x 10-6 45 1.514 x 10-4 +2.286 x 10-6 50 2.630 x 10-4 +3.9713 x 10-6
Claudia Braganza IBHL Chemistry Grade 12
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Below, the graph of Ka values of ethanoic acid against temperature:
0.00001905
0.00003162
0.00005495
0.00008318
0.0001514
0.000263
0
0.00005
0.0001
0.00015
0.0002
0.00025
0.0003
25 30 35 40 45 50
Ka v
alu
e (
mo
l dm
-3)
Temperature (oC)
Ka of ethanoic acid against temperature
Claudia Braganza IBHL Chemistry Grade 12
Page | 12
Conclusion and Evaluation
From the graph that I have constructed above using the Ka values that I calculated, it can
be seen that the trend is that as temperature increases, the Ka value of 1 mol dm-3 of
ethanoic acid also increases.
The explanation for this was mentioned in the hypothesis. As the calculations in data
processing have shown that the [H+] ions increase in concentration as the temperature
increases, this indicates that the reaction is an endothermic one. This is according to Le
Chatelier’s Principle, where we know that the reaction will try to reduce the increase in
temperature by favoring the temperature-reducing endothermic part of the equilibrium.
If the acid dissociation is endothermic, as in ethanoic acid, the reaction favors the
dissociation of the acid into its ions, as shown below:
HA (aq) ⇌ H+ (aq) + A- (aq)
As a result, with higher temperature, more of the acid dissociated into its ions, which
then increased the ions’ concentration. Based on the formula for Ka, this would increase
the Ka value. Therefore, according to this theory, my hypothesis that Ka values for
ethanoic acid would increase with temperature is correct.
This experiment could be improved in several ways. Firstly, as always, with more trials
a better average would be given for the pH of ethanoic acid at the different
temperatures. This would reduce the total random error of the experiment. A good
number of trials would be 6.
Also, the use of the pH meter may have caused some limitations when reporting the
displayed pH value. As the meter was manual and had to be calculated, the calibration
values were not always exactly 4 or exactly 7, because it was difficult to get an exact
value and rather the values were slightly above or below. This may have led to a slight
increase in random error, which then translates to pH readings at each data point which
were slightly above or slightly below the actual pH. Basically, there is no way to know if
the random error may have been slightly above or slightly below the calculated one. To
reduce the total random error of the experiment, using a Vernier machine with an
automatically calibrated pH meter would be better. This way, if there are any errors,
they would stay minimal and would be the same for each trial.
Also, the temperature measured may have been slightly below the needed data point.
This is because the pH had to be measured away from the heater in order to prevent the
ethanoic acid from heating up too much above the required temperature, and in the
time spent moving, the actual temperature that the pH was measured in may have
dropped slightly. To counter this in the next experiment, the ethanoic acid can be heated
up to 5 points above the temperature needed, for example if the needed temperature is
30oC then the solution should be heated till it reaches 35oC. This way, when moving the
beaker away from the heater and readying the pH meter for measurement, the
Claudia Braganza IBHL Chemistry Grade 12
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temperature would slowly drop to the required temperature. The pH meter can then be
quickly inserted and the value measured would be close enough. This would then
reduce the random error.
Another way to reduce the random error due to temperature can be using the Vernier
machine to measure it as well. The uncertainty would be significantly less than a manual
thermometer, and the machine can be programmed to measure the pH value at the
exact temperature.
Works cited
Harwood, Richard, and Christopher Coates. "Acids and Bases." Chemistry for the IB
Diploma. By Christopher Talbot. London: Hodder Education, 2010. 490-92. Print.
"IB Chemistry Blog." » New Chemistry Data Booklet (2009). Web. 15 Mar. 2012.
<http://liakatas.org/chemblog/?p=295>.