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The aim of this experiment is to answer the following question: What is the effect of temperature on the equilibrium constant for the hydrolysis of an ester?

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Introduction

Scott Mabbutt Individual Investigation Aim: The aim of this experiment is to answer the following question: What is the effect of temperature on the equilibrium constant for the hydrolysis of an ester? The reaction I aim to investigate is a reversible reaction where an ester (an organic compound with RCOOR group) is produced and. Esters are generally insoluble in water but are soluble in other solvents. Esters are formed in an esterification reaction where and carboxylic acid (RCOOH group) and an alcohol (ROH), react to form an ester. Below is the esterification reaction/ester hydrolysis reaction. Alcohol + Carboxylic acid Ester + Water The Equipment and apparatus I will be using in my investigation are as follows: * Safety Goggles * Test Tube x 5 * Test tube stoppers x 5 * Water Bath with thermostatic control * Test tube holder (suitable for use in water baths) * Burette * Clamp stand with clamp fixings * Safety Mat * Funnel * 250ml Beaker * 250ml Conical Flask * 10ml Measuring cylinder * Thermometer * White tile Safety goggles will be worn so that the risk of chemicals coming into contact with the eyes is lessened (see risk assessment). Test tubes are being used as the environment for the reaction to take place, they are glass so can be monitored and can be easily stored in a test tube holder in the water bath. Test tube stoppers will be used so that none of the water or ethanol can evaporate out of the tube, thus giving inaccurate results. Water baths will be used so that the temperature, can be maintained at different temperatures. Test tube holders will be used so that all the test tubes can be held safely without risk of tipping over. A burette will be used in the titration because it will less small amounts of liquid at a time, and also has an accurate scale up the side, so volumes will be easier to obtain. ...read more.

Middle

This does not seem to be highly correlated, but is still a relationship that can be useful in making my prediction. However I feel there is room for improvement which I will explain more below. Modifications to Method From my preliminary results and observations I noticed some flaws that could affect my investigation if they are done in the real experiments. They are listed below with ways to change them and make the experiment more accurate. * One of the main flaws with my experiment is that when using a bunged boiling tube, I had to shake the mixture to homogenise it. The problem this caused is that after a while the reaction mixture turned an orange tint. This was because the H+ in the ethanoic acid reacted with the rubber bung causing it to mix with the reaction. The problem this caused was that some of the acid was used up in reacting with the rubber therefore the correct amount was not available to react with the alcohol. The way I have chosen to overcome this is by changing the equipment holding the reaction. I will use glass 30cm3 sample tubes with screw-in plastic tops. This will make sure that all of the acid is available to react with the alcohol in the reaction. * The second problem I found is that the titres I found were rather small. This means that the alkali solution I was using in the titration was too strong. This is because it had too many hydroxyl ions in a given volume, and therefore a small volume neutralised the acid. A way to combat this is to change the concentration of NaOH in the alkali solution. I will use 0.4 moldm-3. The benefit that this has is too make the titres larger and therefore easier to spot trends and patterns. * Another problem I found when using the water bath is that at 60oC, after 3days much of the water in the water bath had evaporated. ...read more.

Conclusion

Temperature (degrees Celsius) Initial Moles of Acid Moles of Acid Equilibrium Constant 25 0.00900 0.00392 1.68 25 0.00900 0.00396 1.62 25 0.00900 0.00392 1.68 25 0.00900 0.00396 1.62 25 0.00900 0.00396 1.62 Average 0.00900 0.00396 1.62 30 0.00900 0.00404 1.51 30 0.00900 0.00404 1.51 30 0.00900 0.00404 1.51 30 0.00900 0.00404 1.51 30 0.00900 0.00404 1.51 Average 0.00900 0.00404 1.51 35 0.00900 0.00412 1.40 35 0.00900 0.00412 1.40 35 0.00900 0.00412 1.40 35 0.00900 0.00412 1.40 35 0.00900 0.00412 1.40 Average 0.00900 0.00412 1.40 40 0.00900 0.00416 1.35 40 0.00900 0.0042 1.31 40 0.00900 0.00424 1.26 40 0.00900 0.0042 1.31 40 0.00900 0.0042 1.31 Average 0.00900 0.0042 1.31 45 0.00900 0.00428 1.22 45 0.00900 0.00428 1.22 45 0.00900 0.00428 1.22 45 0.00900 0.00428 1.22 45 0.00900 0.00428 1.22 Average 0.00900 0.00428 1.22 50 0.00900 0.00304 3.84 50 0.00900 0.0044 1.09 50 0.00900 0.00436 1.13 50 0.00900 0.00436 1.13 50 0.00900 0.00436 1.13 Average 0.00900 0.00436 1.13 Anomalous Result is ignored From this graph I can deduce that as the Temperature increased the value of Kc decreases. As I said this in my prediction, it goes some way to proving that my prediction was at least to some extent correct. Another conclusion I can draw from my graphs is that as the temperature increased the equilibrium constant changed exponentially. This is because on the graph it shows a slight curve. The reason for this is that as the temperature rises the reaction favours the production of the reactants. Therefore as the reactants increase the products decrease so the equilibrium constant goes down. There is a constant fractional change in the value of Kc as the temperature increases. The reason that the reaction favours the production of the reactants, ethanoic acid and propan-1-ol is because this direction of reaction takes in energy, i.e. its energy level is higher than the products that it makes. As the temperature increases the system tries to remove the extra heat energy being put in. The way it does this is by making the reaction that removes heat energy go faster. Therefore the reaction constant goes down.. ...read more.

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