IBHL Investigations Investigating Acids

Claudia Braganza

IBHL Chemistry

Grade 12

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 suc h 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, K a, and has units of mol dm-3. Ka = [H+] [A-]/ [HA] The acid dissociation constant is a measure of the strength of o f 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 pK a 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 pK a the stronger the acid. A change of 1 in the value of the pK a 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 K a values are usually quoted only for weak acids, like l ike 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 o f 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 | 1

Claudia Braganza

IBHL Chemistry

Grade 12

But since [H+] = [A-], in a solution where only the acid is present: Ka = [H+] / [HA] Rearranging: [H+] = √[HA] x K a And then pH = - log10 [H+] This approach can be reversed in order to calculate the K a (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 pK a 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 pK a 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 pK a. Then, Ka will be calculated as follows: Ka = 10 ^ (- pK a) The reason why K a 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 | 2

Claudia Braganza

IBHL Chemistry

Grade 12

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 K a values increase, it is known that it indicates the increase in acidity of the solution. Therefore, my hypothesis is that as temperature increases, the K a value will also increase, thereby increasing the acidity of 1 mol dm-3 of ethanoic acid.

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Claudia Braganza

IBHL Chemistry

Grade 12

Apparatus and Materials 1) 200ml beaker 2) Pipette 10ml 3) pH meter 4) Stir plate heater 5) Thermometer 6) 750ml of ethanoic acid (CH 3COOH) 7) 100ml calibration solution pH 4 8) 100ml calibration solution pH 7 9) pH meter screw 10) Pen and paper

x2 x1 x1 x1 x1 x1 x1 x1 x1 x1

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? Concentration of ethanoic acid

How is it controlled? It is controlled by making the concentration 1 mol dm-3

Why is it controlled? It is controlled because although Ka is only affected by temperature, the experiment still needs to be

controlled so that it doesn’t Volume of acid for each data point pH meter

It is controlled by pipetting 30ml of the acid for each data point It is controlled by calibrating it beforehand in pH 4 and pH 7 solutions

interfere with data collection. It is controlled to make the experiment a fair trial. It is controlled to ensure that pH readings do n’t differ from each other and interfere with data collection. Page | 4

Claudia Braganza

Pressure in the room

IBHL Chemistry

It is controlled by conducting the experiment in a room at standard 1 atm pressure

Grade 12

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 (25 oC, 30oC, 35oC, 40oC, 45oC and 50oC) Dependent: pH value during experiment which then determines final K a 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. 30 oC), 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 | 5

Claudia Braganza

IBHL Chemistry

Grade 12

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 pK a values, from which the Ka value will then be derived.

Trial

1 2 3 4

25 2.36 2.37 2.36 2.36

pH +0.01 Temperature (oC)+1oC 30 35 40 45 2.24 2.12 2.05 1.92 2.26 2.13 2.03 1.90 2.24 2.13 2.05 1.90 2.26 2.13 2.03 1.91

50 1.80 1.78 1.78 1.80

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Claudia Braganza

IBHL Chemistry

Grade 12

Data processing

At 25 oC (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

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Claudia Braganza

IBHL Chemistry

Grade 12

At 35 oC


At 40oC


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Claudia Braganza

IBHL Chemistry

Grade 12

At 45 oC


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

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Claudia Braganza

IBHL Chemistry

Grade 12

The Ka values calculated above are shown below with its respective temperatures. Temperature (oC) +1oC 25 30 35 40 45 50

K a (mol dm-3) 1.905 x 10 -5 3.162 x 10 -5 5.495 x 10 -5 8.318 x 10 -5 1.514 x 10 -4 2.630 x 10 -4

Now, the uncertainties must be calculated in order to get an idea of the errors. Instrument Thermometer Pipette pH Total random error

Percentage Error 1% (0.05/10) x 100 = 0.5% 0.01% 1 + 0.5 + 0.01 = 1.51%

Therefore, with errors the K a 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 25 30 35 40 45 50

K a (mol dm-3) 1.905 x 10 -5 +2.877 x 10-7 3.162 x 10 -5 +4.775 x 10-7 5.495 x 10 -5 +8.297 x 10-6 8.318 x 10 -5 +1.256 x 10-6 1.514 x 10 -4 +2.286 x 10-6 2.630 x 10 -4 +3.9713 x 10-6

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Claudia Braganza

IBHL Chemistry

Grade 12

Below, the graph of K a values of ethanoic acid against temperature:

Ka of ethanoic acid against temperature 0.0003

0.000263 0.00025

0.0002 ) 3 -

m d l o m0.00015 ( e u l a v a K

0.0001514

0.0001 0.00008318

0.00005495

0.00005 0.00003162 0.00001905 0 25

30

35

40

45

50

Temperature (oC)

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Claudia Braganza

IBHL Chemistry

Grade 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 K a 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 K a, 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 35 oC. This way, when moving the beaker away from the heater and readying the pH meter for measurement, the Page | 12

Claudia Braganza

IBHL Chemistry

Grade 12

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. .

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IBHL Investigations Investigating Acids

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