ibhl investigations investigating acids

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]


Page | 2

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


Page | 3

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.


Page | 4

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.


Page | 5

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.


Page | 6

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


Page | 7

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


[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


[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


Page | 8

At 35oC

pH = (2.12 + 2.13 + 2.13 + 2.13) / 4 = 2.13

[H+] = 10 ^ (- 2.13) = 0.0074 mol dm-3


[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


[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


Page | 9

At 45oC

pH = (1.92 + 1.90 + 1.90 + 1.91) / 4 = 1.91

[H+] = 10 ^ (- 1.91) = 0.0123 mol dm-3


[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


[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


Page | 10

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


Page | 11

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


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


Page | 13

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." Â» New Chemistry Data Booklet (2009). Web. 15 Mar. 2012.

<http://liakatas.org/chemblog/?p=295>.

ibhl investigations investigating acids

Documents