Practical Worksheets

Part 1 - Preparation of a Standard Solution

Calibrate the weighing balance that you will be using.

Weigh approximately between 1.25 and 1.45g of anhydrous sodium carbonate.

Carefully transfer the sodium carbonate to a large beaker, accurately and precisely
recording measurements to determine the exact mass transferred.

Add 150cm3 of distilled water to the beaker, stir and completely dissolve the sodium
carbonate.

Carefully and accurately transfer all of the solution to a 250cm 3 volumetric flask, and
make up the solution to 250cm3 with more distilled water.

You must risk assess your experiment and have it checked by your supervisor before
starting.

You will need to demonstrate and evidence your skills and techniques in calibrating
equipment, transferring solids and liquids, mixing of substances and taking and
recording accurate measurements.

You will need to accurately calculate the concentration of the solution you have made.

Part 2 - Standardisation of an acid

Calibrate the pipette that you will be using.

Carefully and accurately transfer 25cm 3 of your sodium carbonate solution into a
250cm3 conical flask and add a few drops of methyl orange indicator.

Clean and fill a burette with a given solution of hydrochloric acid, which will have a
concentration of approximately 0.1M.

Titrate the sodium carbonate solution with the hydrochloric acid, until the indicator
changes colour at the end point.

Accurately and precisely record all measurements to determine the exact titre of
hydrochloric acid required to reach the end point of the titration.

You will need to decide whether the titration needs repeating and how many times.

You must risk assess your experiment and have it checked by your supervisor before
starting.

You will need to demonstrate and evidence your skills and techniques in calibrating
equipment, using volumetric glassware, transferring and mixing liquids, and taking and
recording accurate measurements.

You will need to use your results to accurately calculate the precise concentration of the
hydrochloric acid.
 Part 3 - Titrations of sodium hydroxide with hydrochloric acid

You must risk assess your experiment and have it checked by your supervisor before
starting.

You will need to demonstrate and evidence your skills and techniques in calibrating
equipment, using volumetric glassware, transferring and mixing liquids, and taking and
recording accurate measurements.

Part 3a - Titration of sodium hydroxide with hydrochloric acid (using an

indicator)

Calibrate the pipette that you will be using.

Carefully and accurately transfer 25cm 3 of sodium hydroxide solution (unknown

concentration) into a 250cm3 conical flask and add a few drops of methyl orange
indicator.

Clean and fill a burette with the standardised solution of hydrochloric acid.

Titrate the sodium hydroxide solution with the hydrochloric acid, until the indicator
changes colour at the end point.

Accurately and precisely record all measurements to determine the exact titre of
hydrochloric acid required to reach the end point of the titration.

You will need to decide whether the titration needs repeating and how many times.

You will need to use your results to accurately calculate the precise concentration of the
sodium hydroxide.

Part 3b - Titration of sodium hydroxide with hydrochloric acid (using a pH

meter)

Calibrate the pipette that you will be using.

Carefully and accurately transfer 25cm 3 of sodium hydroxide solution (unknown
concentration) into a small beaker (size 100cm 3 or 150cm3).

Calibrate the pH meter that you will be using with the buffer solutions provided.
Place the pH meter into the beaker of sodium hydroxide.

Clean and fill a burette with the standardised solution of hydrochloric acid.

Add hydrochloric acid from the burette to the sodium hydroxide solution in 1cm3
portions until all of the acid has been added. Measure the pH reading on the pH
meter every 1cm3 of hydrochloric acid added.

Accurately and precisely record all burette and pH measurements in a table.

Plot a graph of pH against volume of acid added (burette reading) / cm3.

You will need to decide whether the titration needs repeating and if the precision
needs to be improved.

You will need to use your graph to accurately calculate the precise concentration of the
sodium hydroxide.
Introduction to Learner work

The learner work that follows has been assessed accurately to national standards. This is
one example of Distinction grade achievement for Learning Aim A on an internally
assessed unit.

The learner is in Year 12 and is completing the Pearson BTEC Level 3 National Certificate
in Applied Science at the Sixth from college alongside other qualifications.

The learner has submitted Assignment Concentrate on keeping up your standards,

Learning Aim A and it has been assessed as Distinction standard.

Commentary

The learner has submitted Assignment Concentrate on keeping up your standards to

cover Learning Aim A: Undertake titration and colorimetry to determine the
concentration of solutions.

Learning Aim A criteria, A.P1 and A.P2, require learners to calibrate equipment listed
in the guidance in order to correctly prepare and standardise solutions for titration and
colorimetry investigations. After preparing the solutions, learners will follow
instructions to accurately calculate the concentrations of unknown solutions, using
procedures and techniques in titration and colorimetry. Throughout the practical
investigations, learners must follow instructions carefully and demonstrate safe
working practices.

For the merit criterion A.M1, learners must demonstrate skillful application of
procedures and techniques in titration and colorimetry to accurately determine the
concentration of solutions. At merit level, learners are expected to carry out the
required procedures with a high degree of accuracy and precision and under minimal
supervision.

To achieve A.D1, learners must evaluate the practical work they have carried out in
terms of the outcomes achieved and the accuracy of the quantitative analytical
procedures and techniques used. Outcomes of titrations and colorimetry investigations
must be interpreted and the accuracy of their results critically judged against the
results obtained by other learners. Learners must evaluate any problems or issues
encountered in relation to their effect on precision and accuracy and the validity and
reliability of results. Suggestions of methods to improve procedures and techniques
must be made.

Detailed feedback

Assessment criteria A.P1 and A.P2

The learner has selected appropriate glassware and calibrated the equipment to be
used in carrying out titrations and colorimetry investigations. A standard solution of
sodium carbonate was correctly prepared and the number of moles of sodium
carbonate dissolved in the prepared solution was calculated in order to titrate the
Na2CO3 against hydrochloric acid (HCl) in order to standardise the hydrochloric acid.
Using the equation Na CO + 2HCl 2NaCl + H O + CO , the concentration of HCl
was calculated. A further titration was carried out using an indicator and using a pH
meter to calculate the concentration of solution of sodium hydroxide. The learner also
calculated the concentrations of two samples of unknown concentration of copper(II)
sulfate using colorimetry. Dilutions of copper(II) sulfate (CuSO4) were prepared by the
learner and used in this investigation. A calibration curve was produced and the values
of the unknown samples calculated.
The learner has submitted sufficient evidence to justify the award of A.P1 and
A.P2.

Assessment criterion A.M1

The learner has been observed selecting and using good techniques to calibrate
equipment, measure and mix solutions and prepare dilutions in order to investigate
the concentrations of various solutions. Titrations have been carried out using an
indicator and a pH meter to determine the end point of the titrations. The assessor has
confirmed that the procedures were completed with minimal supervision through the
completion of an observation record providing a formal record of observation of the
formance during the practical activities. Formula mass has been used to
accurately calculate the number of moles in the solutions used. A titration curve and a
calibration curve have been drawn and used to identify the concentrations of unknown
samples.
The requirements of the grading criterion and additional guidance for A.M1
have been met.

Assessment criterion A.D1 has been correctly awarded. This learner has
provided a detailed evaluation of the quantitative analytical procedures and techniques
used. The problems encountered in carrying out these practical investigations have
been discussed and valid judgements regarding the accuracy and validity of the results
achieved have been made. The learner has given careful consideration to aspects of
health and safety.
Learner work
Part A.P1 Part A.M1

This assessment required me to find the concentration of an unknown sample of sodium

hydroxide (NaOH).
To do this, I first had to find the concentration of the hydrochloric acid (HCl) I was going
to use by titrating the HCl against a standard solution of sodium carbonate (Na 2CO3).
Firstly, I prepared a standard solution of sodium carbonate. In order to do this, a range of
equipment was provided. This included a beaker, 100 cm3 measuring cylinder, weighing
boat, filter paper, distilled water, 250 cm3 volumetric flask, bulb pipette, funnel, anhydrous
sodium carbonate.
A measuring cylinder could have been used but a volumetric flask is much more
accurate, this factor is important as during the process I needed to dissolve the sodium
carbonate; the volumetric flask enabled me to agitate the flask in order to ensure all of
the sodium carbonate was dissolved. Instead of using filter paper to weigh out the
sodium carbonate, a weighing boat was used because the sodium carbonate powder can
easily stick to filter paper and then not all of the sodium carbonate would be transferred
to the volumetric flask. The contents of the weighing boat were emptied into a beaker to
enable the sodium carbonate to be dissolved before putting the solution into the flask.
Approximately 50 cm3 of water was added to the powder in order to dissolve it before
placing the solution into the flask. The beaker and weighing boat were repeatedly rinsed
with distilled water to ensure that any trace of the sodium carbonate was transferred
into the volumetric flask. A glass funnel was used to transfer the dissolved sodium
carbonate into the volumetric flask; this was also rinsed with distilled water to make sure
that all the sodium carbonate had been transferred.
Part A.P1 Part A.M1

Making the standard solution of sodium carbonate

Firstly, the balance was calibrated to 0.1g 1.0g, 10g and 100g in mass, this was to
ensure the balance was accurate for a range of masses, to calibrate the balance a mass
was placed on the balance and a calibration screw was used to adjust the measurements
to register the exact mass.

Prior to the experiment all of the equipment was rinsed with distilled water to prevent
contamination, the beaker was placed in a drying cabinet to ensure the sodium
carbonate does not stick to it in case I put too much in and needed to remove some.

adding the sodium carbonate. 1.4g of sodium carbonate was measured out. The
weighing boat was removed from the balance and emptied it into a beaker to which
approximately 50 cm3 of distilled water was added. It was not essential to be precisely
accurate at this stage as this was not the final amount. A magnetic stirrer was placed in
the beaker and the beaker was placed on the stirring plate. When the powder had
dissolved, a funnel was used to pour the solution into a 250 cm 3 volumetric flask.
Distilled water was used to rinse the beaker and the stirrer to make sure that none of
the solution was left in the beaker. The funnel was then rinsed and the volumetric flask
filled until the bottom of the meniscus was just touching the 250 cm3 mark on the
volumetric flask. The stopper was placed on the volumetric flask and the flask was
inverted several times to ensure the solution was thoroughly mixed.

In order to use the solution to standardise the hydrochloric acid (HCl) I needed to
determine the number of moles of sodium carbonate in the solution that I had made.
The formula mass of sodium carbonate (Mr) is (23x2) + 12 + (16 x 3) which comes to
106. The number of moles of sodium carbonate dissolved in 250 cm 3 is 1.4/106 =
0.0132 moles.
Carrying out the titration to standardise the hydrochloric acid

Firstly, a 25cm3 glass pipette was calibrated. A dry, clean weighing bottle was weighed.
Next, distilled water was sucked up into the pipette so that the bottom of meniscus
rested on the 25cm3 line. The distilled water was transferred into the weighing bottle and
was weighed. The difference in mass was calculated and the mass of water was figured
out to be 24.97g. The laboratory was 18 oC so the density of water at that temperature is
0.99862 g/cm3 (according to International Critical Tables of Numerical Data, Physics,
Chemistry and Technologu, vol III). Using volume = mass x density, the volume of water
from the pipette was 24.97 x 0.99862 = 25.0045 cm3. I feel that this is close enough to
25.00cm3 to use for my experiments. The pipette was then rinsed out with the sodium
carbonate solution that had been prepared. The pipette was used to transfer 25cm 3 of
sodium carbonate solution into a conical flask and a few drops of methyl orange indicator
were added.

Secondly, a burette was calibrated. A dry, clean flask was weighed. Next, distilled water
was poured into the burette so that the bottom of meniscus rested on the 0.00cm 3 line.
The distilled water was delivered into the flask until the bottom of the meniscus rested
on the 50.00cm3 line. The difference in mass was calculated and the mass of water was
figured out to be 49.95g. The laboratory was 18oC so the density of water at that
temperature is 0.99862 g/cm3 (according to International Critical Tables of Numerical
Data, Physics, Chemistry and Technology, vol III). Using volume = mass x density, the
volume of water from the pipette was 49.95 x 0.99862 = 50.0090 cm 3. I feel that this is
close enough to 50.00cm3 to use for my experiments. The burette was then rinsed out
with hydrochloric acid, this is because the HCl was going to fill up the burette so it was a
useful way to clean the burette without contaminating the HCl with water.

Part A.P1 Part A.M1

A rough titration was carried out in order to establish some understanding of where the
end point would be. The rough titration showed that the end point would come between
25cm3 and 30cm3. Following the rough titration, the process was repeated more
accurately. 25cm3 of sodium carbonate was measured into a conical flask using the
pipette and a few drops of methyl orange indicator was added to the solution, which
turned yellow. The start reading was recorded from the burette and HCl was poured from
the burette into the sodium carbonate. The concial flask was swirled to mix the contents
and distilled water was sometimes used to squirt any acid on the sides into the flask. As
the end point was getting close, the HCl was added a drop at a time until the solution
just turned pink and the final volume was recorded. The titration was repeated until
there were concordant results and the results were averaged to improve the accuracy by
reducing the margin of human error.

Learner evidence of investigation

Results:

Rough 1st 2nd 3rd

Initial volume / cm3
0.00 0.00 0.00 0.00
Final volume / cm3 27.50 29.00 28.90 28.80
Volume of HCl used / cm 3 27.50 29.00 28.90 28.80
Average volume of HCl used / cm3 (29.00 + 28.90 + 28.80) / 3 = 28.90
The equation for the reaction is

Na CO + 2HCl 2NaCl + H O + CO

The number of moles of sodium carbonate dissolved in 250 cm3 is 1.4/106 = 0.0132
moles.
The number of moles of sodium carbonate in 25 cm 3 is 0.0132/10 = 0.00132 moles
I can now use this figure to calculate the concentration of the hydrochloric acid
Na CO + 2HCl 2NaCl + H O + CO

1 mole of Na CO reacts with 2 moles of HCl. This means from my earlier calculation
the number of moles of sodium carbonate is 0.00132 so the number of moles of HCl are
0.00132 * 2 which equals 0.00264 moles in 25 cm3

The following equation was used:

concentration = number of moles/ volume in dm 3

Concentration = 0.00264 * 1000 / 28.9 = 0.0913

Therefore, the concentration of the hydrochloric acid was 0.0913 mol/dm³

Carrying out the titration to find the concentration of the sample of sodium
hydroxide.
Part A.P2 Part A.M1

To find out the concentration of sodium hydroxide, two different titration methods were
used. First with methyl orange indicator and then secondly using a pH meter to monitor
the change in pH over the titration.

The first titration was carried out using exactly the same technique as was used to
standardise the hydrochloric acid. Once again the hydrochloric acid was poured into the
burette so this was conditioned with the hydrochloric acid before use. The conical flask
needed to contain exactly 25cm3 of sodium hydroxide so after being rinsed with distilled
water to remove any contamination it was filled with 25cm3 of sodium hydroxide using
the 25cm3 pipette.

A rough titration was carried out in order to establish some understanding of where the
end point would be. The rough titration showed that the end point would come between
25cm3 and 28cm3. Following the rough titration, the process was repeated more
accurately. 25cm3 of sodium hydroxide was measured into a conical flask using the
pipette and a few drops of methyl orange indicator was added to the solution, which
turned yellow. The start reading was recorded from the burette and HCl was poured from
the burette into the sodium hydroxide. The concial flask was swirled to mix the contents
and distilled water was sometimes used to squirt any acid on the sides into the flask. As
the end point was getting close, the HCl was added a drop at a time until the solution
just turned pink and the final volume was recorded. The titration was repeated until
there were three concordant results and the results were averaged to improve the
accuracy by reducing the margin of human error.
Learner evidence of investigation
Results:

Rough 1st 2nd 3rd 4th

Initial volume / cm3
0.00 0.00 1.00 0.00 5.50
Final volume / cm3 27.00 26.00 27.10 26.50 31.70
Volume of HCl used / cm 3 27.00 26.00 26.10 26.50 26.20
Average volume of HCl used /
(26.00 + 26.10 + 26.20) / 3 = 26.00
cm3
The equation for the reaction is

HCl + NaOH NaCl + H O

For the standardisation of HCl, the concentration was found to be 0.0913 mol/dm³.

The number of moles of hydrochloric acid used in 26.00 cm3 was 0.0913 * 26.00 / 1000
= 0.002374 moles

I can now use this figure to calculate the concentration of the sodium hydroxide
HCl + NaOH NaCl + H O

1 mole of HCl reacts with 1 mole of NaOH.

This means from my earlier calculation the number of moles of hydrochloric acid is
0.002374 so the number of moles of NaOH is also 0.002374 moles which was in 25 cm3

The following equation was used:

concentration = number of moles/ volume in dm 3
Part A.P2
Concentration = 0.002374 * 1000 / 25.00 = 0.09496

Therefore, the concentration of the sodium hydroxide was 0.095 mol/dm³

The second titration was carried out using a pH meter instead of using an indicator.
Firstly, the pH meter was calibrated; pH buffers of pH7 and pH10 were used for
calibration as the early readings would be alkaline. The pH meter was turned on and
allowed to warm up and the buffer solutions and electrode were allowed to reach room
temperature.
The electrode was rinsed in distilled water and gently blotted with tissue before placing
the electrode into the pH7 buffer solution. The reading was allowed to stabilise. The
sensitivity control was adjusted until the pH meter read pH7. The electrode was rinsed
well in distilled water, gently blotted and the procedure repeated using the pH10 buffer
solution.
The burette was rinsed out with hydrochloric acid, this is because the HCl was going to
fill up the burette so it was a useful way to clean the burette without contaminating the
HCl.
Part A.P2

A rough titration was carried out in order to establish some understanding of where the
end point would be. This was completed by adding 5cm3 of HCl at a time to 25cm3 of
sodium hydroxide in a beaker with a pH probe inserted into the solution. After 25cm 3 of
HCl was added to the sodium carbonate, the pH was recorded as 9.92, after 30cm 3 was
added the pH reading was 1.66. This showed that the end point would be reached after
adding between 25cm3 and 30cm3 of HCl.
Part A.M1

Following the rough titration, the process was repeated more accurately. 25 cm 3 of
sodium hydroxide was measured into a conical flask using the pipette and a pH meter
probe was inserted into the solution. 1cm 3 volume of HCl was poured from the burette
into the sodium carbonate. A magnetic stirrer was used to continually mix the contents
and a pH reading was taken. This was repeated by adding another 1cm 3 of HCl at a time
and reading the pH. The pH was recorded after each subsequent 1 cm 3 volume of HCl
was added.
1cm3 of HCl was added at a time until 25cm3 of HCl had been added (the rough titration
showed that the end point would come after adding between 25cm 3 and 30cm3 of HCl).
The pH values were recorded each time after the reading had stabilised. The process was
then slowed by adding the HCl at only 0.5cm3 at a time whilst constantly using the
magnetic stirrer.
I continued the titrations until 3 concordant results under the same conditions. The
titration was repeated 3 times and the results averaged by adding the 3 results together
and dividing the total by 3, this would improve the accuracy of the results by reducing
the margin of human error.

Results: Learner evidence of investigation

pH - Run pH - Run 2 pH - Run 3 Mean

Volume HCl/cm3 1
0 13.97 13.97 13.97 13.97
1 13.98 13.73 14.13 13.95
2 13.97 13.72 14.12 13.94
3 13.96 13.71 14.11 13.93
4 13.94 13.69 14.09 13.91
5 13.9 13.65 14.05 13.87
6 13.86 13.61 14.01 13.83
7 13.84 13.59 13.99 13.81
8 13.78 13.53 13.93 13.75
9 13.72 13.47 13.87 13.69
10 13.67 13.42 13.82 13.64
11 13.63 13.38 13.78 13.60
12 13.55 13.3 13.7 13.52
13 13.48 13.23 13.63 13.45
14 13.42 13.17 13.57 13.39
15 13.35 13.1 13.5 13.32
16 13.25 13 13.4 13.22
17 13.14 12.89 13.29 13.11
18 13.05 12.74 13.14 12.98
19 12.9 12.42 13.01 12.78
20 12.75 12.36 12.89 12.67
21 12.55 12.54 12.58 12.56
21.5 12.6 12.25 12.65 12.50
22 12.32 12.07 12.59 12.33
22.5 11.99 11.74 12.14 11.96
23 9.65 9.4 9.8 9.62
23.5 7 6.75 7.15 6.97
24 3.1 3.3 2.95 3.12
25 1.47 1.67 1.32 1.49
26 1.31 1.51 1.16 1.33
From the graph drawn from the results, I found that it took 23.5cm3 of hydrochloric acid
to neutralise the sodium hydroxide.

The equation for the reaction is HCl + NaOH NaCl + H O

The concentration of the hydrochloric acid was standardised as being 0.0913 mol/dm³

Using the same equation which is:

number of moles (HCl) = volume (HCl) in dm3 x concentration (HCl)

This gives - number of moles = 23.5 x 0.0913 /1000 = 0.00215 moles.

HCl + NaOH NaCl + H O

1 mole of NaOH reacts with 1 mole of HCl. Using the earlier calculation, this means that
the number of moles of hydrochloric acid is 0.00215 moles therefore the number of
moles of NaOH is also 0.00215 moles.

This is the number of moles in 25 cm3 of sodium hydroxide. This gives a concentration of
sodium hydroxide as 0.00215 x 1000/25 = 0.086 mol/dm3

Part A.P2

I have determined the concentration of the sodium hydroxide from two different titration
techniques. Although the values are similar (0.095 mol/dm³ and 0.086 mol/dm³), they
are not that close to each other.
Health and safety considerations involved with making a standard solution

Pass, merit and disctinction guidance met

Because I was making up solutions using chemicals that were in powder form I had to
make sure that I wore goggles at all times to protect my eyes.
I made sure that all glassware was in a position whereby it could not fall off the bench.
When the balance was calibrated I made sure that the masses were safe and would not
fall on my toes.
When I used hot water to make up the copper(II) sulfate I used a kettle that had been
rinsed with distilled water prior to use. I allowed it to boil and then let it cool. If the
water had been too hot it could have broken the flask or I cause burns or scalds.

Health and safety considerations involved with performing a titration

I made sure that I placed the burette with the top below eye level when I filled it that
way there was no chance of the acid splashing onto my face. Once filled, the meniscus
could then be read at eye level.
Safety goggles and a lab coat were worn at all times.
When placing the pipette filler onto the pipette I made sure to hold the glass tube firmly
and close to the end being inserted as if force is applied to the centre of the glass tube,
it could break and cause serious injury.
I made sure that electrical equipment was kept away from any liquids as liquids and
electricity are a dangerous combination.

Health and safety considerations involved with colorimetry

During the colorimetry I wore safety goggles because although the copper sulfate was
diluted if a cuvette was accidently dropped, the solution could splash into my eyes.
I made sure that the cuvettes were placed in the bowl when I was finished with them,
this was to ensure they were not knocked over spilling their contents, or knocked off the
bench and broken.

Evaluation of techniques involved in making up a standard solution for titration

and colorimetry

Part A.D1
After completion of the titration of sodium hydroxide, I was informed that the
concentration of the unknown sample of sodium hydroxide used, was in fact 0.1
mol/dm3. The results of my investigation found that the concentration of the sodium
hydroxide to be 0.095 and 0.086 mol/dm3 so there were obviously errors made in the
practical technique.
The percentage difference between my results and the actual value was
(0.1 0.095/0.1) x 100 = 5% and (0.1 - 0.086/0.1) x100 = 14%
This shows that there were errors made during the preparation of the standard solution
and/or the titrations of the hydrochloric acid and the sodium hydroxide.
The method used to ensure good techniques were to use a beaker for dissolving the solid
sodium carbonate. This meant that I was only putting a liquid into the flask rather than a
solid so there was less chance of losing some of the solid by not washing it into the flask
properly. The weighing boat and funnel were also washed out to ensure that all of the
sodium carbonate was washed into the flask to ensure accuracy.
Despite using a beaker, an improvement to my method would have been to use a more
accurate balance that found the mass to more than one decimal place as 1.4 grams
could be any mass between 1.36 and 1.45 grams. I ensured that when filling the
volumetric flask, I was at eye level with the mark on the flask and that I filled the flask
until the bottom of the meniscus was on the required volume on the flask. I also shook
the flask a number of times by inverting it to ensure good mixing of the solution.

During the dilution of the copper(II) sulfate solution for the colorimetry, I felt that I was
reasonably accurate when I made up the serial dilutions as I used a bulb pipette to put
the correct amount of the standard solution into the measuring cylinder, I also used a
small measuring cylinder which is much more accurate than using a large one, 2cm3 would
get lost in a large cylinder and would not be an accurate measurement.

Evaluation of techniques involved in performing a titration

Part A.D1

During the making of the standard solution and the titrations, the overall difference
between the concentration that I calculated for the sodium hydroxide and the actual value
was 5% and 14%. The concentration calculated from the titration with the indicator was
more accurate than the titration with the pH meter. I did find that it was harder to pH
readings that match closely and this could have affected the graph. Although this was
quite inaccurate, it does show that, on the whole my technique was quite good but that
there are areas that need improvement to get a more accurate result.

To ensure good technique when filling the burette, I placed the top of it at eye level so
that I could see the meniscus, this meant I could be positive that the liquid used started
at an exact amount with the bottom of the meniscus registering on the 0 of the burette. I
chose a burette that started at 0 rather than 50 and counts down because it is easier to
work out the volumes used without making a mistake.

I used a magnetic stirrer when doing the titration, this meant that the solutions were
constantly mixed to reduce the time when the pH meter could be taking inaccurate results
due to the poor mixing of the two solutions. To ensure better results I could have increased
the time between each reading to make sure that the solutions were fully mixed before
more acid was introduced.

I did a rough titration first that allowed me to have an idea of where the end point was,
this allowed me to then record every 0.5 cm3 when I was near the end point as I considered
that this would be accurate enough to show the endpoint. This in turn allowed me to draw
a more accurate graph.

When using the pipette to measure out 25cm 3 I found it difficult to measure out exactly
25 cm3 as due to the thinness of the tube near the mark on the pipette, the volume of
solution would shoot up over the mark and I would then have to empty some solution out
and retry to get the exact amount. I found that I had to do this a number of times before
I got the right amount and that air bubbles would sometimes enter the pipette and be
difficult to remove which means that I would not have the correct amount of solution.
Before attempting this again I would get further training and practice for this particular
technique as I felt that this was one area that may have contributed to the amount of
difference between my results and the actual measurement.

Evaluation of techniques involved in performing colorimetry

Part A.D1

Prior to using the colorimeter, I ensured it was properly calibrated, this meant that I
could be confident that my readings were accurate.
When I took the readings I gave the machine a few minutes with each sample to settle
down as the initial readings fluctuate.
I made sure that I only held the cuvette on the ridged side not the plain side so that my

Future improvements required for making standard solutions.

Part A.D1

The sodium carbonate was a fairly standard procedure that is used widely, however with
the copper sulfate because it is hard to get the crystals to dissolve it would be good to
be able to use a hot plate / magnetic stirrer combined
When making the dilutions I could use an electronic pipette filler that could be set to
exact amounts e.g. 2cm3 as this would eliminate human error in measuring small
volumes accurately using a measuring cylinder.

Future improvements required for titrations

The graph was very difficult to draw accurately, in the future I might have the
opportunity to use an installed laboratory graph drawing programme linked to the
computer and produce a really accurate graph. This would allow me to be absolutely
certain of where neutralisation occurred.
From research, I have discovered that there is a monitor that can be fitted onto the
bottom of a burette to count the number of drips of liquid used. This would be a very
accurate way of calculating exactly how much acid was used, especially if it links directly
to the pH meter so I could know how many drips it takes to neutralise the solution.
To carry out the titrations I had to put 25cm3 of the alkali in the conical flask. I used a
pipette filler to do this, to improve I could use an electronic pipette filler. This would be
much more accurate than trying to make sure the pipette filler keeps a tight seal on the
pipette whilst also trying to determine if the meniscus is in line with the required line on
the pipette.

Future improvements required for colorimetry

When I took the colorimetry reading I had to wait for the reading to settle before I
recorded it. It is possible to get software for the computer that takes an average of

seeing and give a far more accurate result.

Again as for the titrations I could use graph drawing software.

Sources of information used:

http://chemwiki.ucdavis.edu/Core/Analytical_Chemistry/Quantitative_Analysis/Titration/
Acid-Base_Titrations
http://www.chemguide.co.uk/physical/acidbaseeqia/phcurves.html
http://chemistry.about.com/od/chemistryquickreview/a/titrationcalc.htm