The Concentration Of Acetic Acid Biology Essay

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Aim: To determine the concentration of acetic acid by titrating against a standard solution of NaOH (0.1M). Using this to find the equivalence point, half the equivalence point and the neutralizing pH that will give the PKA value of acetic acid.

Data collection:

Table 1.1: The change in pH of 10 ml of acetic acid (unknown concentration) when sodium hydroxide (0.1M) was titrated over 3 trials.

pH ± 0.01

Volume of NaOH added (ml) ± 0.1ml

Trial 1

Trial 2

Trial 3

1.0

3.93

2.15

2.92

2.0

3.72

2.76

4.20

3.0

4.55

3.47

4.05

4.0

4.37

3.69

4.48

5.0

4.63

3.73

4.40

6.0

5.08

3.85

4.60

7.0

5.48

3.65

4.59

8.0

5.64

3.94

4.34

9.0

5.48

3.86

4.42

10.0

5.57

4.12

4.54

11.0

5.76

4.33

4.85

12.0

6.63

4.51

5.28

13.0

9.06

4.77

5.32

14.0

11.34

5.91

5.38

15.0

11.61

6.27

5.45

16.0

12.16

8.21

6.11

17.0

12.43

8.78

6.34

18.0

12.39

10.13

6.58

19.0

12.44

10.86

7.13

20.0

12.66

11.28

10.26

21.0

12.65

11.15

11.26

22.0

12.69

11.29

11.53

23.0

12.59

11.42

11.80

24.0

12.61

11.32

12.22

25.0

12.62

11.68

12.56

Qualitative observations: The solution turned slightly milky initially, but then turned colourless after about 12-13 ml of NaOH was added.

Data processing:

Graph 1.1: The volume of NaOH titrated into 10 ml acetic acid vs. the pH for trial 1.

Using the above graph, we can extrapolate the midpoint of the steep rise (11ml to 15 ml). The midpoint comes to 13 ml of NaOH (0.1 M) being titrated into the acetic acid and the neutralisation pH comes to about 9.06 (verified from table 1.1).

ph = 8.21, volume = 16 mlGraph 1.2: The volume of NaOH titrated into 10 ml acetic acid vs. the pH for trial 2.

Similarly we can do the same with trial 2.

ph = 7.13, volume = 19 mlGraph 1.3: The volume of NaOH titrated into 10 ml acetic acid vs. the pH for trial 3.

Similarly we can do the same for trial 3.

Table 1.2: Summary of results of the neutralizing volume of NaOH (0.1M) and the pH of neutralization for trials 1,2 and 3.

Volume of NaOH (ml) ± 0.1ml

pH ± 0.01

Trial 1

13

9.06

Trial 2

16

8.21

Trial 3

19

7.13

Average (±0.3 ml, 0.03)

16

8.13

Thus, from table 1.2 we can see that 8.13 is the equivalence point for this reaction and therefore half the equivalence point would be 4.065.

Calculations and error propagation:

Average volume:

(Trial 1 + Trial 2 + Trial 3)/3

= (12 + 16 + 19)/3

= 48/3 = 16 ml

Error:

ΔAvg. = (ΔTrial 1 + ΔTrial 2 + ΔTrial 3) x Avg.

Trial 1 Trial 2 Trial 3

= (0.1/13 + 0.1/16 + 0.1/19) x 16

= ±0.3 ml (1 sf)

Average pH:

(Trial 1 + Trial 2 + Trial 3)/3

= (9.06 + 8.21 + 7.13)/3

= 8.13 (3 sf)

Error:

ΔAvg. = (ΔTrial 1 + ΔTrial 2 + ΔTrial 3) x Avg.

Trial 1 Trial 2 Trial 3

= (0.01/9.06 + 0.01/8.21 + 0.01/7.13) x 8.13

= ±0.03 (1 sf)

Determining the concentration of the Acetic acid:

Acid + Base  Salt + Water

CH3COOH + NaOH  CH3COONa + H20

1 : 1

Concentration of NaOH = no. of moles/ volume in dm3 (using data from table 1.2)

= x/ 0.016

x = 1.6 x 10-3 moles of NaOH

No. of moles of NaOH = No. of moles of CH3COOH

Thus, concentration of CH3COOH = 1.6 x 10-3 / 0.01

= 0.16 M

Uncertainties:

Δ No. of moles of NaOH = (Δvolume/volume) x No. of moles of NaOH

(note: the uncertainty in the concentration of NaOH was unknown since it was pre-made)

Δ No. of moles of NaOH = (0.1/ 0.016) x 1.6 x 10-3

± 0.01 moles

Δ Concentration of Acetic Acid = [(ΔNo. Of moles/No. of moles) + (Δvolume/volume)] x Concentration of Acetic Acid

Δ Concentration of Acetic Acid = [(0.01/1.6 x 10-3) + (0.1/0.016)] x 0.16

= ± 2M

Determining the PKA value:

PKA = pH at half equivalence point

Trail 1:

Equivalence pH = 9.06 at 13 ml

Half equivalence pH= 5.28 at 6.5 ml (from graph 1.1)

PKA = 5.28 (As per Henderson-Haselbalch equation)

Trial 2:

Equivalence pH = 8.21 at 16 ml

Half equivalence pH= 3.94 at 8 ml (from graph 1.2)

PKA = 3.94 (As per Henderson-Haselbalch equation)

Trial 3:

Equivalence pH = 7.13 at 19 ml

Half equivalence pH= 4.48 at 9.5 ml (from graph 1.3)

PKA = 4.48 (As per Henderson-Haselbalch equation)

Average:

(5.28 + 3.94 + 4.48)/3

= 5.57 (3 sf.)

Results:

Table 1.3: The table shows the PKA value obtained from half the equivalence point each trial and the average.

Trial Number

PKA Value

1 (± 0.01)

5.28

2 (± 0.01)

3.94

3 (± 0.01)

4.48

Average (± 0.04)

5.57

Conclusion:

A weak acid only partially dissociates from its salt. The pH will rise normally at first, but as it reaches a zone where the solution seems to be buffered, the slope levels out. After this zone, the pH rises sharply through its equivalence point and levels out again like the strong acid/strong base reaction.

There are two main points to notice in the curves we have obtained in graph 1.1, 1.2 and 1.3. The first is the half-equivalence point. This point occurs halfway through a buffered region where the pH barely changes for a lot of base added. For trial 1 this occurs near a point where 6.5 ml of NaOH is added, for trial 2 it occurs near 8 ml and for trial 3 it occurs near 9.5 ml. The half-equivalence point is when just enough base is added for half of the acid to be converted to the conjugate base. When this happens, the concentration of H+ ions equals the KA value of the acid. This can be further extended to give, pH = pKa. The second point is the higher equivalence point; this point is usually above a pH of 7 as seen by the trend in the data obtained.

The determination of the KA, of a weak acid can be difficult. However, its PKA can be easily estimated by analysis of its titration curve. The, PKA is the pH value at the half equivalence point, that is, the point at which only half of the volume of alkali needed to reach the equivalence point has been added. The reason for this is that, at the half-equivalence point, the concentrations of the conjugate base, A-, and that of the non-dissociated acid, HA, are almost equal. Therefore, they cancel out in the expression for KA.  

We observed the PKA value to be 5.57 (average) and the range for PKA values in an aqueous medium usually range from -2 (strong acid) to 12 (strong base). Our value of 5.57 indicates that acetic acid is a weak acid as it has not completely disassociated into ions. The fluctuations between the three trials could be accounted for by the errors mentioned later, but we are sure that acetic acid is a weak acid and our experiment proved it. The theoretical value for PKA of acetic acid as per the IB data booklet is 4.76. This is 0.81 less than the experimental value we calculated meaning that there was about 17.02% experimental error in this value obtained. The systematic error accounted for only 0.84% of the total error this we can attribute the rest of the deviation to the random errors, which are later, discussed in the evaluation. We can see that our third trial was the most accurate out of all the three as it is closest to the literature value.

Evaluation:

I think this lab was conducted with adequate proficiency, however some sources of error could be:

It was quite difficult to manage the pH probe, magnetic stirrer, flow of NaOH from the burette and record the readings at the same time as we did our practice labs in groups. Thus there could have been errors in what we recorded.

There might have been a parallax error while taking the readings off the burette.

The burette could have dripped the base directly onto the pH probe and thus there may have been extra basicity detected by it giving inaccurate readings.

The acid was not used in equal amounts for every trial, as the pipette holders were not working properly. Thus there were varied amounts of acid during every trial.

The beakers we used were not clean and had some residue stuck at the bottom, which could have contributed for some mistakes in our data.

I dripped 1 ml of base into the acid at a time; this was too large an interval since acetic acid is a weak acid. This means I did not record some of the changes in the pH, which would have occurred with slightly less amounts of acid.

Some improvements that I can suggest to this lab are:

Do the lab in pairs as it becomes easier to handle all the equipment that is required to complete the lab.

Take more readings to rule out the effect of random errors.

Wash all equipment thoroughly before using and ensure they all function correctly.

Use smaller intervals for adding the base (perhaps 0.5 ml) to get a more gradual and accurate graph.

Keep the burette at eye level to reduce parallax error.

Use 2 pH probes to check the change in pH to be completely sure of the readings obtained.

 

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