Determining gain a rough estimate of the

the Concentration of an Unknown Hydrochloric Acid Solution Using Sodium
Carbonate as the Alkali in a Titration

different values of pH can be crucial in various parts of the body, for example
a low pH is normal for the stomach where food is digested (Crowe and Bradshaw,
2014) however a higher pH (7.06±0.04) is expected for saliva in the mouth where
amylase is the main enzyme present (Baliga, Muglikar and Kale, 2013).

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Acidosis is an example of when pH balance is disrupted
and it occurs when there is a higher than normal concentration of Hydrogen ions
in the bodily fluids, causing pH of the blood to decrease below 7.4. This can
also be a sign of malfunctioning kidneys due to inadequate removal of acid
(, 2017).

The aim of this research was to determine the
concentration of the Unknown Hydrochloric acid (HCl) using a titration.

and Methods: A pH indicator was prepared by boiling red
cabbage in water. This was then added to solutions of various values of pH.

Sodium Bicarbonate (Na2CO3)
was used to make a 250 ml, 0.1M standard solution. Using a pipette, 20cm³ of
the solution was then transferred to a conical flask. A few drops of indicator were
added afterwards. The burette was filled with HCl acid of unknown

The first titration was used as a trial to gain a rough
estimate of the equivalence point (end point),
which was determined when the Na2CO3
changed from the colour green to purple. After this, titrations
were performed until two readings within 0.1ml of each other were accomplished.

average titre was 7.45*10-3 L. The number of moles of Na2CO3 was 0.025. This
value was then used to calculate the number of moles per 20cm³ of Na2CO3 by using the
multiplication factor 12.5, 0.025moles/12.5= 2*10-3.

Using the balanced equation below, the number of moles of
HCl was calculated using the 1:2 ratio.

Finally, the molarity was calculated by dividing the
value of moles of HCl by the average titre (L), 0.537 mol/L (4*10-3 moles/7.45*10-3

The pH of the unknown Hydrochloric acid was calculated to
be 2.51 by using the two equations below.

= Ka*HCl              pH=

Conclusion: The pH of this
solution is notably higher than that of the human stomach, pH 1.5 (Beasley et
al., 2015). This highlights the high acidity needed to digest food effectively
but also shows how different values of pH are needed for different functions in
and outside of the body.

However, the pH of this solution could have
been measured using the cell potential of the sample in association with the
standard Hydrogen electrode. By using this equipment a pH curve could have been
constructed and the equivalence point determined via this, removing
the subjective nature of this titration allowing a more accurate calculation of the unknown HCl concentration
Mullen and Roy, 2010).

On the basis of the findings in this experiment, it takes
approximately 7.45*10-3 L of 0.537 mol/L HCl to neutralise a 20cm³ solution of
0.1M Sodium Carbonate therefore the aim of this experiment has been fulfilled.






Use of Bomb Calorimetry to Determine the Enthalpy of Combustion and Therefore
the Calorific Content of Vegetable Oil

oil has many uses in today’s society, it can be used a bio-fuel when combined
with alcohol to power diesel engines or act as an additive aiding reduction in
motor emissions (, 2013). However, its omega-3-fatty acids also
provide a source of nutrition which is essential for human development (Kumar,
Sharma and C. Upadhyaya, 2016).

This particular experiment aims to conclude the enthalpy
of combustion and therefore the calorific content of vegetable oil. The way the
body uses this energy is one application of the results and is discussed
further in this report.

calibrate the calorimeter, the 1-dodecanol was lit and extinguished shortly
after to determine the size of the flame from the Bunsen burner. The mass (g)
of the lamp was determined and 100mL of ice cold water was added to the
aluminium can, the temperature was then taken (°C). The 1-dodecanol was lit and
the water was then stirred with a thermometer until the temperature had reached
10°C above room temperature. The final
temperature was recorded after extinguishing the flame. The mass of the lamp
was recorded after allowing it to cool. This was repeated with 1-dodecanol, and
then with vegetable oil.

absorbed by water during calibration using 1-dodecanol was calculated first (Q)
using the equation below.   

Q= MC?T             
M= Mass of water (100g), C= Specific heat
capacity of water 4.184J/(g°C), ?T= Change in
temperature (°C)

The heat given off by combustion of
1-dodecanol was then calculated, 1.
-55.86 kJ, 2. -59.43 kJ. The heat
given off but not used by to warm the water was also calculated and an average
of this value was determined qcal-7792.26 kJ. These values were then
used to calculate values in relation to the vegetable oils enthalpy of combustion.

Heat energy released by the combustion
of vegetable oil was calculated using the qcal of 1-dodcenaol, –qoil
= qwater + qcal. The value was negative due to the
reaction being exothermic, loosing heat to the surroundings (Experiment 1. -1827.155 kJ, Experiment 2. -1408.7504 kJ)

The enthalpy of combustion of vegetable oil was calculated
using the qoil and the moil. The average of these two
values was taken-
H -2.848 kJ/g. The calorie content was therefore 680.92
Calories / 0.6809 Kcal per 0.566 g of vegetable oil.


enthalpy of combustion for vegetable oil was negative, suggesting it is an
exothermic reaction therefore releasing energy when it reacts. In the human
body the energy released from vegetable oil is used for metabolic reactions and
physical activity (Thivel et al., 2013). However, it is important to consider
that not all energy from this oil will be absorbed and used for metabolism in
humans. Therefore calculations including energy lost via excretion would be
required to be incorporated for application to the biological aspect of this
experiment (, 2008).

However, the aim of this experiment was to determine the
enthalpy of combustion and calorific content of vegetable oil so although the
biological comparisons are not fully achievable the results can be used for
other purposes, for example choosing which oil is best to cook with to maintain
a healthy diet and lower calorie consumption.


Synthesis and Characterisation of Acetylsalicylic Acid (Aspirin) Produced from
Salicylic Acid

is a pain-killer and anti-inflammatory frequently used in today’s society. The
major component of Aspirin, Salicylic Acid, is found in willow plants and was
previously used to treat pain relief from headaches and other minor ailments
prior to the vast scientific knowledge to the present day of its properties (Arias,
2017). Aspirin inhibits both the COX-1 and COX-2 enzymes which therefore allow
aspirin to act as an anti-inflammatory and can potentially reduce development
of some cancers (Choi et al., 2016).

The aim of the current experiment was to synthesise
Aspirin (Acetylsalicylic acid) and produce an adequate percentage yield from
Salicylic acid.

of Salicylic Acid was weighed and 4.0ml of Acetic Anhydride was then added to
this in a fume cupboard. Four drops of concentrated Sulphuric Acid were swirled
into this solution to ensure mixing. The solution was placed in a water bath
for 10 minutes, cooled via an ice bath and then filtered with a Buchner Funnel.
The crystals were washed with 3mL of cold, distilled water and then filtered

Crystals formed after around 10 minutes of air drying.
These were placed in an oven for 30 minutes at 100°C and then weighed shortly

Ferric Chloride was added to two test tubes, each containing
distilled water, one with commercial aspirin and one with the current
experiment’s aspirin.

The produced aspirin, in a capillary tube, was placed in
a heating block. The temperature at which it melted was recorded.

percentage yield of the current experiment’s aspirin was 81.23%. The actual yield was 2.119g and the theoretical yield was
2.6086g which was calculated by multiplying the number of moles of
Acetylsalicylic acid (0.01448) by its molecular mass (180.157).

After the addition of Ferric Chloride the commercial
aspirin solution maintained a yellow colour whereas the aspirin produced from
this experiment induced a purple solution, see Appendix II. The Fe3+ ions react with –OH groups
attached to salicylic acid, potentially suggesting the Aspirin produced
contained un-reacted salicylic acid (Lewis, 2003).

The temperature at which the aspirin melted was 122.1°C,
which in comparison is much lower than Lewis (2003) whereby the melting point
was 138-140°C. The Aspirin produced could have therefore contained impurities,
for example water.

previously mentioned, the Aspirin produced is likely to have contained
impurities. The Aspirin produced in the current experiment is unlikely to be
acceptable for use for commercial purposes due to the lack of clarity regarding
exactly what impurities it contained.

Lewis (2003) used thin layer chromatography in order to
test the purity further and compare the extent of similarity between the
commercial aspirin and the produced aspirin. This could have removed or
decreased the issue of lack of knowledge as to how pure the Aspirin produced
was in the current experiment. If it had been used it could have potentially allowed
commercial use if the two compounds had similar Rf values on the chromatography
plate. This therefore is a possible improvement of this experiment.

The results from this experiment show that Aspirin was
synthesised successfully with an acceptable percentage yield however, and
therefore the aim of this experiment was fulfilled.








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Accessed 8 Dec. 2017.

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of Periodontology, online 17(4), p.461. Available at: Accessed 26 Dec. 2017.

Beasley, D., Koltz, A.,
Lambert, J., Fierer, N. and Dunn, R. (2015). The Evolution of Stomach Acidity
and Its Relevance to the Human Microbiome. PLOS ONE, 10(7), p.e0134116.

Choi, J., Ghoz, H., Peeraphatdit, T., Baichoo, E.,
Addissie, B., Harmsen, W., Therneau, T., Olson, J., Chaiteerakij, R. and
Roberts, L. (2016). Aspirin use and the risk of cholangiocarcinoma. Hepatology, online 64(3), pp.785-796.
Available at:
Accessed 8 Dec. 2017.

Crowe, J. and Bradshaw, T. (2014). Chemistry for the Biosciences – The Essential Concepts. 3rd ed. Oxford:
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Basics. online Available at: Accessed 12 Dec.

P., Mullen, C. and Roy, M. (2010). Titration and pH Measurement. eLS.
online Available at: Accessed 15 Dec.

Lewis, D. (2003). Aspirin.
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Accessed 8 Dec. 2017. (2017). Metabolic
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Accessed 6 Dec. 2017.

Kumar, A., Sharma, A. and C. Upadhyaya, K. (2016).
Vegetable Oil: Nutritional and Industrial Perspective. Current Genomics, online 17(3),
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Accessed 6 Dec. 2017.

Thivel, D., Aucouturier, J., Metz, L., Morio, B. and
Duché, P. (2013). Is there spontaneous energy expenditure compensation in
response to intensive exercise in obese youth?. Pediatric Obesity, online 9(2), pp.147-154. Available at:  Accessed 1 Dec. 2017.