bioc practical session 1

8/4/2019 BIOC Practical Session 1

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BIOC 20010 28/09/11

Thomas Dodd.

Practical session 1

Wednesday 28th September 2011.

Experiment 1: Calibration.

IntroductionpH meters require calibration before they can be used to accurately detect the pH of sample solutions.

In order for the pH reader to have a scale upon which to base its readings, it must be calibrated against

two solutions of known pH concentration, the first to set a benchmark for the reader to base its

readings on and the second as a comparison benchmark, as the pH meter does not have its own storedreference pHs.

As temperature will also play a factor in pH, it must be constantly measured, using a temperature

probe which is also incorporated into the pH probe. Using this technique, we can ensure reliable pH

readings from the probe, provided it is handled with care and stored in a buffer solution of pH 7 when

not in use.

Aims

The aim of this experiment is to successfully set calibration points for a pH probe, using solutions of

known pH concentration.

Materials and Methods

This experiment incorporates the use of a pH probe, which is also fitted with a temperature probe,

deionized water, a waste beaker, phosphate pH 6.86 and calibration buffer (pH 9.18)

Procedure

The pH meter was plugged in and switched on. The first calibration solution of phosphate pH 6.86

was opened and the pH electrode and temperature probe were lowered into the container and stirred

gently. The CAL button on the pH meter console face was pushed and the detection setting was set to

6.86. The screen flashed the message NOT READY for several seconds. It then read READY andthe CFM (confirm) button was pushed allowing this calibration number of pH 6.86 to be stored.

The electrode was removed from the solution and washed down with deionized water which was

rinsed into the waste beaker to avoid cross contamination. The probe was then lowered into the

second solution containing calibration buffer pH 9.18. The screen flashed the message NOTREADY for several seconds before switching to ready. The CFM (confirm) button was pushedallowing the instrument to store the second calibration point. The pH electrode was washed down

with deionized water and the rinsing added to the waste beaker. The pH electrode was then placed

back into its storage solution.

ResultsThe pH probe was successfully calibrated and could now be used to measure the pH of solutions in

the other experiments.


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Discussion

The use of a pH meter is obviously invaluable in practical laboratory conditions, as it removes the

need to constantly take samples of a solution for the total solutions pH to be checked. The simplicityof a pH meter is also quite considerable, as it is a simple circuit that measures the voltage of the

probe, which will vary depending on the pH it is placed in and then converts the voltage to pH using

the Nernst equation

Experiment 2: The effect of buffer strength.

IntroductionThis experiment will demonstrate the ability of a buffer to drastically alter the [A

-] of a solution at a

steep point even through the volume that brought about this shift was quite minute and the pH does

not vary greatly. At the point around the pKa of the solution, the addition of a small volume of buffer

will result in a sharp shift in [A-]

This experiment will be repeated twice, with a smaller concentration of buffer the second time, using

a proportionally smaller volume of acid. In this manner, we can demonstrate that it is not the volume

of the solution that plays a role in the [A-] shift; rather it is the addition of a buffer to the solution at

around its pKa that will result in the shift of [A-].

AimsThe aim of this experiment is to successfully measure the pH shift of a solution of known

concentration and pH with the addition of an acid of known volume and pH and obtain a table of

results which can be graphed in order to indicate the sharp change in the curve where the pH shift

took place.

Materials and MethodsThis experiment incorporates the use of a pH probe, deionized water, a waste beaker, a 50mL conical

flask, a stock of 1M HCl, 100mM Tris buffer, 10mM Tris buffer, a waste beaker, a P 1,000 pipette

and a P 200 pipette.

Procedure

50 mL of 100 mM Tris Buffer solution was placed into a clean 50 mL conical flask. The pH probe

was removed from its storage solution, washed with deionized water which was added to the waste

beaker. The probe was then placed into the solution. The pH reading was allowed to settle and the

screen read an initial reading of pH 9.1

0.5 mL of 1M HCl was then added to the conical flask using the P 1,000 pipette and the solution wasgently swirled. The probe was placed back into the solution and the second reading was taken.

This process was repeated a total of 10 times, by the end of which the pH of the solution had changed

to a pH of 1.24.

The conical flask was rinsed with deionized water and the rinsing were placed in the waste beaker.

The pH probe was also rinsed with deionized water and the rinsing were added to the waste beaker. 50

mL of 10 mM Tris buffer was placed in the conical flask and the pH probe was placed in the beaker.

The screen showed an initial reading of pH 9.09

0.05 mL of 1M HCl was then added to the conical flask using the P 200 pipette and the solution was

gently swirled. The probe was placed back into the solution and and the second reading was taken.

This process was repeated a total of 10 times, by the end of which the pH of the solution had changedto a pH of 2.29.


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Results

The following tables of results and accompanying graphs were taken from this experiment:

1M HCl = 1 mol/Litre

Additions made at 0.05 ml = 0.0005 Litres

Molecular weight of HCl= 36.5

36.5 X 0.0005 = 0.01825 g

Addition of 0.0005 moles HCl.

100 mM TrisReading pH mL HCl moles HCl

Initial 9.10 0.00 0

1st addition 8.83 0.50 0.0005

2nd addition 8.58 1.00 0.0010

3rd addition 8.39 1.50 0.0015

4th addition 8.17 2.00 0.0020

5th addition 7.95 2.50 0.0025

6th addition 7.67 3.00 0.0030

7th addition 7.28 3.50 0.0035

8th addition 5.30 4.00 0.0040

9th addition 1.57 4.50 0.0045

10th addition 1.24 5.00 0.0050

0

1

2

3

4

5

6

7

8

9

10

0.0000 0.0005 0.0010 0.0015 0.0020 0.0025 0.0030 0.0035 0.0040 0.0045 0.0050

pH

Moles HCl

pH


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1M HCl = 1 mol/Litre

Additions made at 0.05 ml = 0.00005 Litres

Molecular weight of HCl= 36.5

36.5 X 0.00005 = 0.001825 g

Addition of 0.00005 moles HCl.

10 mM TrisReading pH mL HCl moles HCl

Initial 9.09 0.00 0.00000

1st addition 8.64 0.05 0.00005

2nd addition 8.43 0.10 0.00010

3rd addition 8.22 0.15 0.000154th addition 8.06 0.20 0.00020

5th addition 7.81 0.25 0.00025

6th addition 7.55 0.30 0.00030

7th addition 7.14 0.35 0.00035

8th addition 5.33 0.40 0.00040

9th addition 2.65 0.45 0.00045

10th addition 2.28 0.50 0.00050

0

1

2

3

4

5

6

7

8

9

10

pH

Moles HCl

pH


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Discussion

As discussed in the aims section this experiment demonstrated the sharp shift of pH as the

concentration approaches the pKa of Tris buffer (8.06) becomes extremely pronounced, compared to

the addition of the same volume of acid at proportionally distant values from the pKa. The

equivalence point of this solution is reached when the number of moles of HCl added = 0.005/2 =

0.0025 moles, as [HA] = [A-] at this point. At the equivalence point, pH = pKa + log([A-]/[HA]), but

when [A-] = [HA], log([A-]/[HA]) will = 1. Therefore at the equivalence point, pH = pKa.

Questions

Q 3.1What does this curve tell you about the effect of the buffer concentration?

A: This curve tells us that the pH of a solution can shift rapidly before settling to a more steady

increase or decrease in pH when protons are added or taken away from the solution.

For example, if there is acid being introduced into a solution, a buffer will protonate A-to produce

HA in order to absorb the acid.

If there is base being introduced into a solution, a buffer will de-protonate HA to produce A-

in order

to absorb the base.

Q.32 Where does the pH change the least per mole of acid added?

A: Based on the data gathered from the experiment, experiment, the pH changed least upon the

addition of 0.001 moles and 0.0001 moles of HCl.

Q3.3Draw the structure of Tris at pH 10.0? at pH 2.0?

Tris buffer neither protonated nor deprotonated.

OH

OH

OHNH2

Tris buffer at pH 10.0, having been deprotonated in order

to neutralise the base in the solution.

Tris buffer at pH 2.00, having protonated in order to remove

H+

ions from the solution


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O+

O+

O+

NH2

HH

H

H

H

H Q 3.4 What is the pKa of Tris? In which pH range does it work effectively as a buffer and why?

A According to Wikipedia, Tris has a pKa of 8.06. Tris would work best as a buffer in more

basic environments, as it would easily de-protonate in order to release free H+

ions which would be

used to neutralize basic compounds being introduced into the system. It would not work as

proportionally well in an acid solution, as it will not willingly accept H+

ions, compared to loosing

them.

Experiment 3: Determination of pKa values.

Introduction

Different compounds will have the ability to cope with pH changes to different extents. In this

experiment, the compounds : Glycine, Cysteine and Imidazole will be used to demonstrate this fact.

Additions of Sodium Hydroxide (NaOH) will be made at intervals of o.5 mL and the pH of eachsolution will be measured after each addition.

By taking these readings, we will be able to determine the pKa values of each solution. As Glycine

exists in a zwitterionic form as both a cationic carboxylic acid and an anionic amine when in solution,

it will require two titrations in order to determine it two pKa values. Both forms are indicated below.

Cationic carboxylic acid H3N+CH2CO2H Anionic amine H2NCH2CO2

-


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AimsThe aim of this experiment is to successfully measure the pH shift of solutions of known

concentration and pH with the addition of a base of known volume and pH and obtain a table of

results which can be graphed in order to extrapolate the pKa values of each solution. A second

titration will be required for Glycine during which acid will be added, in order to obtain both

zwitterionic pKa values.

Materials and MethodsThis experiment incorporates the use of a pH probe, deionized water, a waste beaker, a 50mL conical

flask, a stock of 1M NaOH, 50 mL 0.01M Glycine, 50 mL 0.01M Cysteine, 50 mL 0.01M Imidazole ,

a waste beaker, a P 1,000 pipette and a stock of 1M HCl.

Procedure

Glycine

50 mL of 0.01 M Glycine was placed into a clean 50 mL conical flask. The pH probe was removed

from its storage solution, washed with deionized water which was added to the waste beaker. The

probe was then placed into the solution. The pH reading was allowed to settle and the screen read an

initial reading of pH 1.8

0.5 mL of 1M NaOH was then added to the conical flask using the P 1,000 pipette and the solution

was gently swirled. The probe was placed back into the solution and the second reading was taken.


to a pH of 11.21

Cysteine

50 mL of 0.01 M Cyseine was placed into a clean 50 mL conical flask. The pH probe was removed







to a pH of 11.50

Imidazole

50 mL of 0.01 M Imidazole was placed into a clean 50 mL conical flask. The pH probe was removed







to a pH of 11.01

Results

The following tables of results and accompanying graphs were taken from these experiments:


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In order to obtain [A-] values, the total concentration of NaOH added was divided by the total volume

of the solution at this concentration.

For example, at total volume = 50.5 ml, [NaOH] = 5x10-5

5x10-5/ 0.0505 L = 9.901 X 10

-4= [A

-].

This calculation was repeated for all addition in order for find [A-] for all three titrations.

The [HA] values were then obtained and the log of [A-]/[HA] was graphed agains the pH in order to

find the intercept of the two axes. This would allow us to obtain the pKa values of the solutions.

The following table was obtained in this manner in order to calculate the log [A-]/[HA] values. [HA]

values and log values are not included where [HA] = 0.

The following tables of information were obtained from measuring the pH values of each solution

after each addition of NaOH.

Glycine Cystine Imidazole

Reading

Moles

NaOH pH Reading Moles NaOH pH Reading

Moles

NaOH pH

Initial 0.00000 1.80 Initial 0 5 Initial 0 5

First 0.00005 2.13 First 0.00005 7.93 First 0.00005 5.1

Second 0.00010 2.25 Second 0.0001 8.4 Second 0.0001 6.02

Third 0.00015 2.40 Third 0.00015 8.74 Third 0.00015 6.45

Fourth 0.00020 2.50 Fourth 0.0002 9.07 Fourth 0.0002 6.6

Fifth 0.00025 2.90 Fifth 0.00025 9.44 Fifth 0.00025 6.83

Sixth 0.00030 3.70 Sixth 0.0003 9.55 Sixth 0.0003 7.14

Seventh 0.00035 8.45 Seventh 0.00035 10.52 Seventh 0.00035 7.28

Eighth 0.00040 9.45 Eighth 0.0004 10.95 Eighth 0.0004 7.49

Ninth 0.00045 9.85 Ninth 0.00045 11.24 Ninth 0.00045 7.73

tenth 0.00050 10.11 tenth 0.0005 11.5 tenth 0.0005 8.07

eleventh 0.00055 10.32 eleventh 0.00055 8.41

twelfth 0.00060 10.52 twelfth 0.0006 11.01

thirteenth 0.00065 10.73

Fourteenth 0.00070 10.96

Fifteenth 0.00075 11.21

Table of log[A-]/[HA]

Reading

Moles

NaOH Total volume [A-] (mol/L) [HA] (mol L-1) [A-]/[HA] log[A-]/[HA]

Initial 0.00000 50 0 0.01 0 N/A

First 0.00005 50.5 9.901x10-4

8.911x10-3

0.11111 -0.95425

Second 0.00010 51 1.961x10-3

7.843x10-3

0.25003 -0.60201

Third 0.00015 51.5 2.913x10-3

6.796x10-3

0.42864 -0.36791

Fourth 0.00020 52 3.846x10-3

5.769x10-3

0.66667 -0.17609

Fifth 0.00025 52.5 4.762x10-3

4.762x10-3

1 0

Sixth 0.00030 53 5.66x10-3

3.774x10-3

1.49974 0.17602

Seventh 0.00035 53.5 6.542x10-3 2.804x10-3 2.3331 0.36793

Eighth 0.00040 54 7.407x10-3

1.852x10-3

3.99946 0.602


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Ninth 0.00045 54.5 8.257x10-3

9.173x10-4

9.00142 0.95431

tenth 0.00050 55 9.091x10-3

0 N/A N/A

We can plot these values against the pH, and the point at which they cross the y axis will give us the

pKa:

2.03 2.12 2.252.4

2.682.9

3.7

8.45

9.44

7.93

8.48.74

9.079.44

9.95

10.5210.93

11.24

5.1

6.026.45

6.7 6.837.14 7.28

7.497.73

0

2

4

6

8

10

12

-1.5 -1 -0.5 0 0.5 1 1.5

pH

log [A- ]/[HA]

pKa values

Glycine

Cysteine

Imidazole


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Discussion

By graphing the results we took from each titration, we can calculate approximate pKa values of each

solution .

As Glycine exists in a zwitterionic form, it would require two titrations, one with acid and one with

base, in order to determine the pKa values of both forms. Due to time constraints, only the titration ofan acidic solution of Glycine could be carried out. In actuality, the cationic and anionic pKa values are

2.34 and 9.6 respectively.

As Imidazole and Cysteine exist in one form only in solution, they will return a single pKa values of

6.95 and 9.58 respectively.

An interesting thing to note was that all of the solutions pKa values returned in this experiment wereslightly higher than what they are in actuality. This suggests that perhaps there was some error in the

pH calibration in, as the result was never lower than it should have been.

The true pKa values for Glycine and cysteine were obtained from

http://www.cem.msu.edu/~cem252/sp97/ch24/ch24aa.html

Imidazoles pKa value was found from:http://en.wikipedia.org/wiki/Imidazole.

Questions

Q 4.1 Look up the pKa values for the compounds. How well do your pKa values agree with those

found in literature?A The calculated values for each compounds pKa correlate quite well to the known values foreach compound. A 2.4 pKa value was obtained for Glycine, when its actual pKa values is 2.35. 6.75

was obtained for Imidazole, whose actual pKa is 6.95. Cysteine returned a value of 9.58, with the true

value being 10.25. While this is reasonably accurate, it is the least accurate of the three titrations. The

result is probably down to human error, possibly impatience while waiting for the pH reading to

settle, or incorrect cleaning of the pH probe between sample readings.

Q4.2: Some of the compounds have multiple titrations. Which?

A Glycine has multiple titrations due to being zwitterionic, as discussed before. Therefore it

would require two titrations to detect both pKa values, as discussed before.

Q4.3: For the compounds with multiple titrations: Are you able to observe the multiple titrations

individually or are you getting an overall average?

A

When the full graph is plotted, an overall pKa value can be obtained, however when the graph is

divided into its two components, the pKa for each functional group can be calculated.

Experiment 4: The analysis of Vinegar.

IntroductionIn this experiment, a sample of vinegar will be titrated against 1M of NaOH in order to determine he

concentration of acetic acid in the sample. In order to detect the end point of the titration; the

equivalence point, we will require the use of a colour indicator. Phenolphthalein will be used in this

instance. Phenolphthalein is colourless in acidic solutions, but turns pink above a pH of 8.2. At, or

very close to the equivalence point the indicator will turn the solution from colourless to pink.

Aims
http://www.cem.msu.edu/~cem252/sp97/ch24/ch24aa.htmlhttp://www.cem.msu.edu/~cem252/sp97/ch24/ch24aa.htmlhttp://en.wikipedia.org/wiki/Imidazolehttp://en.wikipedia.org/wiki/Imidazolehttp://en.wikipedia.org/wiki/Imidazolehttp://www.cem.msu.edu/~cem252/sp97/ch24/ch24aa.html


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The aim of this experiment is to conduct two titrations on a sample of vinegar with additions of

NaOH, the first to establish a general equivalence point and the second to establish a more accurate

equivalence point. By observing the colour change that takes place at the equivalence point and

noting the amount of NaOH that had been added to the vinegar sample, we can calculate the amount

of acetic acid in the solution, since this reaction occurs in a 1:1 ratio.

Methods and Materials

25 mL vinegar,a graduated cylinder,a conical flask, a sample of phenolphthalein, a waste beaker, a

pipette, a stock of deionised water, a burette containing the 1 M NaOH and a white sheet of A4

paper.

Procedure25 ml of vinegar was placed in a clean conical flask and two drops of phenolphthalein were placed in

the beaker using a pipette. The beaker was placed underneath the burette containing NaOH. It was

ensured that the meniscus was at the graduation mark at eye level before proceeding. The tap of the

burette was opened, allowing a steady stream of NaOH to enter the conical flask. When 10 ml of

NaOH had been placed in the beaker, the tap was closed and the beaker was refilled back to its

starting conditions. The tap was opened again, allowing the NaOH to flow through, although this timeat a slower pace, so as to detect the end point. Traces of pink were detected in the conical flask, but

the solution returned to colourless after a gentle swirl of the conical flask. A final end point of roughly

19.3 was detected as the point when the solution had reached its equivalence point.

The contents of the conical flask were placed into the waste beaker and the burette was refilled to its

starting point. As before, 25 ml of vinegar was placed into the conical flask and two drops of

phenolphthalein were placed into the solution. This time, a more cautious addition of NaOH was

practiced as the volume reached 19 ml. However the second reading gave a very sharp change of pink

to colour less at almost exactly 21.2 ml NaOH. This was the figure chosen to calculate the amount of

acetic acid present in the solution.

Results

As this reaction occurs in a 1:1 ratio, we can use the equation M1 X V1 = M2 X V2 to solve for the

molarity of the acetic acid.

The molar mass of acetic acid CH3CO2H is 60.05280 g/mol

Titre 1: 19.3 ml

M1 X V1 = M2 X V2

1 X 19.3 = x x 0.25

19.3 = 0.25x

X = 0.772 M

0.772 M X 60.0528 = 46.3608 g/L

46.3608 g/L = 4.63608 g/100ml present

= 4.6% w/v


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Titre 2: (accurate titration): 21.2 ml

1 X 21.1 = x x 0.25

21.1 = 0.25x

X=0.848 M.

60.0528 x 0.848 = 50.9247 g/L

50.92g/L = 5.092g/100ml present

= 5.1% (w/v)

Discussion

One possible reason for the second titre value being higher than the first in this experiment may

have to do with the use of a graduated cylinder being used to measure the vinegar solution, which

would have been somewhat inaccurate.

Another factor may have been that myself and my lab partner took turns determining the end point

each time.

Both of these errors can be rectified by repeating the titration with a more cautios approach.

Even though two different acetic acid concentrations were obtained, they both fall withinthe limits

of regular diluted vinegar : (4% to 8% acetic acid). However, the second value is still to be taken as

the true value of the sample tested, as a much more precise colour change was detected at this

point.

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