'Investigating factors that affect the rate of chemical reactions.'
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'Investigating factors that affect the rate of chemical reactions.' Matthew J Hampton Silverdale School Class 11L Chemistry Coursework Scientific Knowledge. The rate of a reaction is a measure of how fast the reaction between the reactants occurs and so how fast the products are produced. Some reactions are very slow, such as the rusting (oxidisation) of iron to ferrous oxide (rust) in the presence of water. Some reactions are very fast, such as the reaction between sodium metal and water or when dynamite explodes. Both these reactions would not be good for us to study as they occur too slowly or too quickly for us to measure accurately. A reaction that occurs at a reasonable rate should be studied, such as the reaction between dilute hydrochloric acid and calcium carbonate, in the form of marble chips or powder. Calcium carbonate + hydrochloric acid calcium chloride + water + carbon dioxide CaCO3(s) + 2HCl (aq) CaCl2 (aq) + H2O(l) + CO2(g) Another advantage of this reaction is that carbon dioxide is a product. Carbon dioxide is a gas and thus the rate of reaction can be measured by collecting, and measuring, the amount of gas given off as the reaction proceeds. This is one of three ways of measuring a reaction rate. If a gas is given off, the gas can be collected and measured. If one of the reaction products is cloudy then the rate at which the reaction becomes cloudy can be measured. The reaction between calcium carbonate and hydrochloric acid does not product a precipitate, but the reaction between sodium thiosulphate and hydrochloric acid produces a yellow precipitate of sulphur and so the reaction rate can be measured by the rate of cloudiness of the reaction products. If neither of these methods can be used, the change in mass of the reaction can be measured. It would be possible to measure the rate of reaction in the reaction between calcium carbonate and hydrochloric acid using this method, as the weight of
I can achieve this result by using the small marble chips a 1.2g each time, that is the calcium carbonate with a moderately large surface area, larger than the medium chips but not so big as the powder. If I use these chips the experiment will work in a predictable and measurable way with 20 mls of 1.0 M down to 0.2M HCl and so I can use a range of concentration in between, such as 1.0M, 0.8M, 0.6M, 0.4M and 0.2M, using the same experimental conditions as in the preliminary experiment. The Main Experiment. Method I will use the same apparatus as in the preliminary experiment and the same apparatus. I will use the same weight (1.2g) of small marble chips and the same amount (20 mls) of dilute HCl. I will use five different concentration of HCL, namely 1.0M hydrochloric acid x 3 0.8M hydrochloric acid x 3 0.6M hydrochloric acid x 3 0.4M hydrochloric acid x 3 0.2M hydrochloric acid x 3 I will use 1.2g of marble chips in my main experiment because in my preliminary experiment this amount gave me good and reliable results and I think that it will do the same in my main experiment. I recorded the amount of CO2 given off every 10 seconds for a total of 200 seconds. I performed each experiment three times and averaged the results to get more precise and reproducible results from which to plot the reaction rate graphs. Diagram Prediction for main experiment From my preliminary results and my background knowledge of collision theory, I predict that in this experiment the reaction rate, and so the rate at which the CO2 gas collects in the cylinder, will be quickest in the 1.0M HCl experiment and slowest in the 0.2M HCl experiment, and the other concentrations will range in between in an orderly fashion. This is because the concentration of the dilute aqueous hydrochloric acid is a measure of the number of particles of HCl in the liquid.
However I do not think this had much effect on my results as we measured the experiment for 200 seconds and so there were a lot of results for each concentration used and the three times we did the experiment produced results that were very similar every time. The only anomalous result we obtained was one experiment using 0.8M HCl when the reaction rate was faster than expected. I have already explained that I think this was because I accidently used 1.0M HCl and if this was so then I would expect a faster rate, as I saw in that run. Because I decided that was an anomalous result, I did another experiment at that concentration, which gave similar results to the other 0.8M experiments and I used that one in my averages. I think this was a good experiment as it gave results that were the same as my predictions and agree with what I known about reaction rates from my knowledge of collision theory, that if you increase the number of particles reacting ( the concentration of the HCl) then there will be more useful collisions and the reaction will go faster resulting in the production of more CO2 gas, in a given time. It would be possible to extend the readings I obtained using higher or lower concentrations of acid, but this would mean using different equipment, as higher concentrations of acid would be dangerous as HCl is corrosive and also a bigger cylinder would be necessary, while using lower concentrations of acid would need a more accurate way of measuring the smaller volumes of gas given off, perhaps by using a smaller, but long and thin measuring cylinder. Also other variables could be investigated that alter collision theory, such as changing surface area of the marble, or changing the temperature at which the reaction is carried out or using a catalyst to aid the reaction. A gas syringe could be used to give me more accurate results to the volume of gas given off. Books used: Chemistry for you: Lawrie Ryan GCSE Double Science: Chemistry revision guide: Richard Parsons
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