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Aim: To investigate the rate of reaction between Hydrochloric acid and marble chips. Background Knowledge: Factors that affect the rate of reaction between hydrochloric acid and marble chips or any other reaction are called variables. They are: ¨ The surface area of the chips Solids with a large surface area react faster than solids with a smaller surface area. This is due to the fact that if there is more area on the solid to react with the substance the reaction is able to occur much faster. Page 79 "“ GCSE Chemistry revision Guide This diagram shows a large particle small surface area and lots of small particles large surface area and how the particles can react with more area on the smaller pieces. ¨ The temperature of the acid The more heat particles have, the more energy they have. So if the particles have more energy they're going to move around faster. As they're moving around faster, there's more chance of collisions. So the higher temperature increases collisions therefore speeds up the reaction. Page 79 "“ GCSE Chemistry revision Guide This diagram shows the difference in movement between hot and cold particles. ¨ The concentration of the acid As the concentration increases, the rate of reaction increases. This, like the temperature of the acid, is based on the collision theory. The higher the concentration, the more particles therefore the more collisions so the reaction takes place faster. Page 79 "“ GCSE Chemistry revision Guide This diagram shows the movement and difference between a low concentration of particles and a high concentration of particles. ¨ Catalysts A catalyst speeds up a reaction. It does this by lowering the activation energy. The activation energy is what is needed to turn reactants marble chips into products hydrogen gas. To make reactants turn into products a sufficient amount of energy is needed to make the particles collide to start the reaction. This is activation energy and it gives an exothermic reaction the energy it needs to continue the reaction. Catalysts lower the activation energy so it is easier for particles to react so a lot more particles have enough energy to react, therefore, speeding up the rate of reaction. Page 79 "“ GCSE Chemistry revision Guide This diagram shows how a catalyst gives particles something to stick to, increasing the number of collisions. Page 80 "“ GCSE Chemistry revision Guide This graph shows the effect of a catalyst on the rate of reaction. The factors that affect the rate of reaction are all based on the collision theory. The theory that all particles have to collide to cause a reaction. Preliminary work: To investigate the different concentrations of the acid would be the easiest to measure conducted over a short period of time and satisfactory results would be produced. To measure the rate of reaction, the amount of gas given off could be measured. From the equation: CaCO3 + 2HCL à CaCL2 + H20 + CO2 It is seen that a gas is produced, CO2 so the amount of gas that is produced at different concentrations could be measured. A mole calculation was used to find out how much Calcium Carbonate to use. If I use a 100cm3 measuring cylinder to measure gas: 24000cm3 of gas is 1 mole of gas 100/24000 = 1/240 = 4.2 x 10-3 moles CaCO3 : CO2 1 : 1 1 mole : 1 mole 100g : 44g 4.2 x 10-3 m : 4.2 x 10-3 m 4.2 x 10-3 m x 100g = 0.42g 0.42g of calcium carbonate should produce 100cm3 of gas. Therefore, the minimum of calcium carbonate I will use to get sufficient results is 0.5g. I will be using five different concentrations of acid: 100%, 75%, 50%, 25% and 0%. So the amount I will use will be: 100% = 20cm3 HCL 0cm3 Water 75% = 15cm3 HCL 5cm3 Water 50% = 10cm3 HCL 10cm3 Water 25% = 5cm3 HCL 15cm3 Water 0% = 0cm3 HCL 20cm3 Water This is used as a control A 0% concentration will be used as a control to see if calcium carbonate would react with water or not. This would then make sure that the reaction only takes place if HCL is present. Prediction: The higher the concentration, the faster the reaction will occur. From background knowledge, it is known that a reaction will occur when particles collide, so the more particles there are the more collisions there will be. If there are more reactant particles per set volume higher concentration more collisions will occur per second, consequently, more particles reacting per second and the rate of reaction is increased. So for a lower concentration there will be less particles, so there will be less collisions therefore the reaction will be slower. Also the higher the concentration the more gas will be produced. This is because if there's more particles higher concentration reacting with the solid marble chips then the reaction will take place quicker. Consequently, the lower the concentration, the less particles to collide and start a reaction so less gas is produced. Equipment: · HCL · Water · Marble chips · Pessel and mortar · Stopwatch · Weighing scales correct to 2d.p. · Spatula · Water trough · Measuring cylinder x2 · Boiling tube with bung and pipe · Clamp stands x2 Method: To measure the rate of reaction, time how long it takes for the marble chips to react and measure the gas given off. To do this put a 100cm3 measuring cylinder in a water trough, with water inside it, held up by a clamp stand. Then put the pipe from the boiling tube under the measuring cylinder. The boiling tube with a pipe will be held by another clamp stand opposite the measuring cylinder. Crush the marble chips into powder with a pessel and mortar and measure out 0.5g of powder for each experiment with the weighing scales. Then, measure the amount of water and HCL needed with the second measuring cylinder. For each different concentration the exact same thing will be done. Put the HCL/Water solution into the boiling tube and make sure the pipe is under the measuring cylinder. After that pour the calcium carbonate powder into the solution, then start the stop clock and put the bung on the boiling tube the same time the calcium carbonate goes in. Then, every five seconds, measure how much gas has been produced using the scale on the measuring cylinder. Repeat the experiment three times for each different concentration and then take an average. Diagram: Chemistry for you page 190 This diagram is similar to the experiment conducted except a boiling tube held by a clamp stand with a pipe and bung was used instead of a flask. Fair test: · The marble chips are crushed to make sure the surface area is the same for each experiment because a larger surface area would take longer to react than a smaller one. So if all the chips are of the same surface area, then they will all react at the same speed, making it a fair test. · All the HCL will be of the same strength, as all experiments will use the same HCL from the same bottle. Stronger acid will speed up the rate of reaction. · The water and acid will be of the same temperature each time because temperature affects the rate of reaction. · After each experiment, the boiling tube will be cleaned properly to get rid of the acid and bits of Calcium carbonate so there's no extra acid or calcium carbonate in the next experiment. · The 100cm3 measuring cylinder will always be full to the top with water so that measurements will be fair. Safety: · To ensure that no acid gets into anyone's eyes, safety goggles will be worn. · Make other persons aware of harmful chemicals. HCL · Necessary medical equipment near by, e.g. eye wash. · Have a cloth or towel near by to clean up any spilt acid so it isn't hazardous to anyone around. Results: Amount of HCL cm3 Amount of water cm3 Gas produced every 5seconds cm3 Average 1st time 2nd time 3rd time 20 0 25 24 25 24.67 45 44 40 43.00 55 57 53 55.00 61 60 57 59.33 65 68 64 65.67 67 68 66 67.00 68 69 67 68.00 69 69 68 68.67 70 71 69 70.00 70 70 70 70.00 70 70 70 70.00 70 70 70 70.00 15 5 20 19 22 20.33 38 35 40 37.67 44 40 41 41.67 47 44 42 44.33 48 45 45 46.00 51 47 46 48.00 53 50 48 50.33 55 54 49 52.67 57 56 50 54.33 58 57 55 56.67 59 58 58 58.33 61 59 59 59.67 64 60 60 61.33 64 62 63 63.00 65 63 65 64.33 66 65 66 65.67 66 67 70 67.67 66 67 71 68.00 10 10 12 10 13 11.67 26 23 20 23.00 29 25 26 26.67 31 27 27 28.33 32 28 28 29.33 33 28 29 30.00 34 29 32 31.67 35 31 34 33.33 36 32 35 34.33 36 33 37 35.33 37 35 38 36.67 38 36 39 37.67 39 37 39 38.33 39 38 40 39.00 40 38 41 39.67 41 39 41 40.33 41 40 42 41.00 43 40 43 42.00 43 41 44 42.67 44 42 44 43.33 45 43 45 44.33 45 44 45 44.67 45 44 46 45.00 5 15 12 13 11 12.00 20 19 21 20.00 23 24 22 23.00 24 24 25 24.33 25 25 25 25.00 25 25 25 25.00 25 26 26 25.67 0 20 0 0 0 0 All results will be plotted on the same graph. This will then make it easier to analyze my results. The average amount of gas measured cm3 will be plotted against time seconds. Graph to show results: The graph was produced by hand and scanned into the word document. Analysis: All concentrations produced gas rapidly to begin with but the most rapid was the 100% concentration. This happened with all the different concentrations except they all started to increase with a steady rate at different times. 100% 30 seconds 75% 15 seconds 50% 10 seconds 25% 10 seconds From this we can see that the higher the concentration, the faster the reaction starts and the longer it continues rapidly. The graph indicates this in the linear gradient of the slope. As the reaction increases the gradient becomes steeper. This result supports the predictions made based on the collision theory. As there are more particles in a higher concentration, there are more collisions so the reaction is faster. When the graph became flat, it was shown that there was no more solid to react with the HCL saturation. The reactions all varied in how long the reaction took place for. 100% 60 seconds 75% 90 seconds 50% 115 seconds 25% 40 seconds The longest reaction was the 50% concentration. The graph shows this by the line leveling out for longer linear gradient. Although it was the longest reaction it didn't produce the most gas. It just produced gas very slowly as it was a low concentration, because there wasn't enough particles to react to make the reaction faster. So gas was produced but very slowly and not much of it. 100% concentration solution was over quickly again, shown by the line on the graph and produced a lot of gas; due to there being more particles to react with the solid marble chips. The 25% concentration however, took place over an even shorter time than the 100% concentration but a lot less gas was produced in the 25% concentration again, due to there not being many particles. The different concentrations also varied on how much gas was produced overall on average. 100% 70.00 cm3 75% 68.00 cm3 50% 45.00 cm3 25% 25.67 cm3 As predicted, the most gas was produced by the higher concentration and the least gas was produced by the lowest concentration. From the graph it can be seen that for different concentrations the amount of gas produced varies. This is due to there being more particles in a higher concentration to react with the solid marble chips. The results gained support the theory that the more concentration, the faster the reaction and the more gas is produced. This matches the predictions made. It is also seen that as the concentrations become less, gas is produced at a much slower, yet at a steady rate because of not having enough particles to react with the substance making the reaction slower. The conclusions and prediction are all based on the collision theory: All particles have to collide in order to react with one another. Evaluation: The method used for conducting the experiment was an effective one as: · It was easily done over the amount of time given in class to conduct the experiment. · It was simple and easy to repeat a lot of times to get enough results to calculate averages. · Produced sufficient results and were easy to present on a graph to compare. · It was a safe experiment. · It was an easy experiment to make sure everything was a fair test and accurate. If the investigation was to be done again, consideration may be given to repeating the test a few more times for each concentration to produce a better average. From the graph it can be seen that some of the concentrations don't level out. This is because for each concentration, each time the experiment was conducted; the gas stopped being produced at different times. So when the average was taken it didn't always show the gas had stopped being produced. So the graph doesn't always level off. Maybe if the gas produced every 5 seconds had been recorded more times, say 10 or 20, instead of 3, the graph would've leveled off. Another reason for this is maybe that the experiment wasn't left going for long enough and a few more bubbles of gas could've been recorded giving more accurate results. It is shown on the graph that the 50% concentration produced more than the 75% most probably because of the reason just mentioned. Even though the results weren't as accurate as they could've been for the reasons mentioned above, they still verified the predictions and conclusions made. Further experiments could be conducted to extend the work I have done. These could be to investigate the other variables in the same way I have conducted my experiment: · Surface area "“ different sizes of marble chips for each experiment. · Temperature of the acid "“ investigate a range of temperatures. · Catalysts "“ investigate the effect of a catalyst in an experiment. If then all these different factors were investigated, all the results could be put together to prove the conclusions further. Bibliography: Books: 1. Chemistry For You, National Curriculum Edition for GCSE "“ Lawrie Ryan Page 190 "“ diagram of experiment 2. Revision Guide for GCSE Double Science, Chemistry, Higher level "“ Richard Parsons Page 79 "“ diagrams to show how different variables affect the rate of reaction Page 80 "“ Graph to show the effect of a catalyst on the rate of reaction Websites: 1. http://www.revisioncentral.co.uk 2. http://www3.mistral.co.uk/cns/depts/science/sc1/GCSE/
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Aim: To investigate the rate of reaction between Hydrochloric acid and marble chips. Background Knowledge: Factors that affect the rate of reaction between hydrochloric acid and marble chips or any other reaction are called variables. They are: ¨ The surface area of the chips Solids with a large surface area react faster than solids with a smaller surface area. This is due to the fact that if there is more area on the solid to react with the substance the reaction is able to occur much faster. Page 79 – GCSE Chemistry...

Bibliography:

Books:

1. Chemistry For You, National Curriculum Edition for GCSE – Lawrie Ryan

Page 190 – diagram of experiment

2. Revision Guide for GCSE Double Science, Chemistry, Higher level – Richard Parsons

Page 79 – diagrams to show how different variables affect the rate of reaction

Page 80 – Graph to show the effect of a catalyst on the rate of reaction

Websites:

1. http://www.revisioncentral.co.uk

2. http://www3.mistral.co.uk/cns/depts/science/sc1/GCSE/

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I will be investigating how... I will be investigating how the concentration of hydrochloric acid affects the rate of reaction between hydrochloric acid and magnesium ribbon. Equation: Mgs+2HClaq→MgCl2aq+H2g Planning: The input variables are: "¢ temperature of acid "¢ concentration of acid "¢ surface area of magnesium "¢ use of catalyst I will measure the concentration of acid and control the temperature of acid, surface area of magnesium, mass of magnesium and use of catalyst. The temperature of the acid will effect the rate of reaction because as the acid particles heat up, they gain kinetic energy, therefore move faster and collide more often and more successfully because the collisions are more energetic, therefore there are more collisions in a given time, and the rate of reaction increases. The surface area of the magnesium will effect the rate of reaction because the larger the surface area, the more collisions can take place in a given time, and the rate of reaction increases. The use of a catalyst will effect the rate of reaction because they give an alternative pathway that has a lower activation energy, therefore more of the collisions will result will result in reaction because they need less energy to be successful. This is the apparatus I will use: I will place some magnesium in a flask with an amount of acid. I will place a bung on the tube to prevent any gas from escaping. A tube will run from the flask to a graduating tube in a trough, both filled with water. I will then measure the amount of gas given off in a certain time. I will repeat the experiment 3 times in order to highlight any anomalies. In order to work safely, I will wear goggles in order to protect my eyes from acid. Prediction: I predict that the higher the concentration of the acid, the faster the rate of reaction. This is because there will be more acid particles in a given volume, therefore there are more collisions in a given time, and the rate of reaction increases. If the concentration was to double, I would expect the rate to double. This is in reference to a previous experiment with a reaction between marble chips and various concentrations of acid. It was proved in this experiment that doubling the concentration of the acid doubled the rate of the reaction, therefore the two sets of results were directly proportional. Pre-test: The concentrations we decided to use as our upper and lower values were 2M and 0.2M. From our pre-test, we discovered that 0.2M took too long to react, so the concentrations that we chose were not workable. We then decided to change our lowest concentration to 0.6M. Concentration of Acid M Time to react seconds 2 8.81 0.6 65.90 From the results of my pre-test, I will use 3cm of magnesium and measure the time to collect 10cm3 gas, as these provided workable results. Obtaining Evidence: Concentration of Acid M 1 2 3 2.0 8.52 7.84 8.42 1.8 11.52 6.49 9.12 1.6 9.35 16.02 10.01 1.4 13.95 11.58 11.21 1.2 13.25 13.78 16.05 1.0 18.81 18.24 20.72 0.8 32.78 32.81 22.02 0.6 67.51 48.28 62.46 Concentration of Acid M Temp Before ˚C Temp after ˚C Temp Before ˚C Temp after ˚C Temp Before ˚C Temp after ˚C 2 23 29 20 20 23 30 1.8 23 29 22 22 23 30 1.6 23 29 22 22 24 28 1.4 23 30 21 21 23 27 1.2 23 31 23 23 24 27 1 23 27 23 23 24 27 0.8 23 25 23 23 24 26 0.6 21 23 22 22 23 25 Concentration of Acid M Average Time seconds Rate x100 2 8.260 12.107 1.8 10.185 9.818 1.6 10.010 9.990 1.4 12.246 8.166 1.2 14.915 6.705 1.0 19.256 5.193 0.8 29.203 3.424 0.6 64.985 1.539 Results in bold are anomalies. Analysing and Conclusions From the temperatures taken before and after the reaction, I have discovered that the reaction is exothermic. My graphs show that the higher the concentration of acid, the faster the rate of reaction. This is because there are more acid particles in a given volume, therefore more collisions in a given time, therefore increasing the rate of reaction. If the concentration was doubled, there would be double the amount of acid particles in a given volume, double the amount of successful collisions in a given time, therefore the rate of reaction should double. This means that the concentration of acid and rate of reaction are directly proportional, as stated in my prediction. However, form looking at my results table, I have discovered that the rate of reaction is the concentration of the acid squared. This disagrees with my original prediction. Evaluating: The readings on my graph do not fit the best-fit line very well, as there are several anomalies, therefore the results are not very reliable. The anomalies could be due to the fact that we had to use two different bottles of acid during the experiment, so the concentrations could have been slightly different. We did parts of our experiment at different times, so the temperature in the lab could have been higher, therefore increasing the temperature of the acid, so the particles would have more energy, move faster, increasing the number of collisions, therefore altering the rate of reaction. Where the magnesium had oxidised, a substance had formed on the surface on the magnesium, which meant that the acid had to get through it to reach the magnesium, increasing the time to react. We did however try to make our data more accurate by using a syringe to measure small amounts of acid and water, and using as many decimal places as we could in our calculations. The temperature of the acid was hard to control, as it was controlled by the temperature of the room. If I was to repeat the experiment, I would use more varied concentrations in order to show more clearly the pattern, and also to attempt to clear any anomalies. I would also like to investigate larger concentrations of acid, but would have been hard to measure.   

I will be investigating how the concentration of hydrochloric acid affects the rate of reaction between hydrochloric acid and magnesium ribbon. Equation: Mgs+2HClaq→MgCl2aq+H2g Planning: The input variables are: • temperature of acid • concentration of acid • surface area of magnesium ...

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Rates of reaction Plan ...Rates of reaction Plan Aim: In this experiment I will find the rate of reaction between Sodium thiosulphate NaS2o3 and Hydrochloric acid HCl. There are different variables I could use to see the change in the rate of reaction. These include temperature, concentration or catalysts. I will investigate how temperature affects the rate of reaction between Sodium thiosulphate and Hydrochloric acid. Prediction When sodium thiosulphate and hydrochloric acid react they produce a cloudy precipitate. The two chemicals are both clear solutions and will react together to form a yellow precipitate of sulphur, the equation for which is as follows: NaS2O3 aq+ HCll¨Sg+NaCls+ H2Ol+SO2s As the solution will turn cloudy, we can observe the rate of reaction by placing a black cross underneath the beaker and seeing how long it takes for it to disappear. There are factors that affect this experiment such as temperature, concentration and time. I do not think that surface area will affect the experiment, as both chemicals are liquids. For my experiment I will study temperature as this is easily observed and can be easily varied. I think that as the temperature of sodium thiosulphate increases, the amount of time taken for a reaction decreases. I know this because before two particles can react they must meet. The higher the temperature there is the more successful collisions between other particles is increased. When temperature increases the bonds in NaS2O3 break quicker because more energy is available greater than the activation energy of the reaction. As a result S2O3 2- ions are available so it takes less time to bond with H+ ions from HCL and new bonds are formed quicker and therefore sulphur precipitates quicker and the rate of reaction increases. S2o3 2- aq +2H+ aq S02aq+Sg+H2Ol When the temperature increases it causes an increase in kinetic energy so you have more chances of successful collisions between NaS2O3 particles and HCl particles so the rate of reaction increases. Also more activation energy is made available to overcome the activation energy of the reaction; the reactants have greater energy than the activation energy, so the reaction takes place quicker. I will keep the concentration of NaS2O3 constant to prevent more successful collisions as there would be more particles available if a higher concentration is fed which will increase successful collisions. I will also keep the concentration of HCl constant as an increase or decrease in concentration will affect the rate of reaction. I will change the temperature of NaS2O3 so I can see how the temperature affects the rate of reaction. I will keep the temperature of the HCl acid at room temperature as we are only concentrating on the NaS2O3 and if we heat the HCl it might affect the rate of reaction it would not be a fair test if we heat the HCl when we are observing how NaS2O3. I also predict that every time the temperature increases by 10oC the rate of reaction doubles. The preliminary results Time on heat sec Temperature of NaS2O3 0C Time taken for cross to disappear sec 0 24 60 10 34 52 Method For the preliminary experiment I heated the NaS2O3 to get it to the temperature I wanted but it was difficult to get the NaS2O3 to the right temperature so the results were not as accurate, but for my real experiment I will use a water bath to get accurate results instead of a Bunsen burner. For the preliminary experiment I only recorded the temperature of the NaS2O3 but for my real experiment I will record the temperature of the HCl as well to get more accurate results because if the NaS2O3 was high and the HCl could bring the temperature down quicker and also have to make sure all the temperature of the HCl is the same. I will also take the temperature of the mixture so I know the temperature at which the reaction took place. 1. I will set up my apparatus and put an X on a piece of paper and measure out 50ml of NaS2O3 and 10ml of HCl. 2. I will pour the NaS2O3 into a conical flask and measure the temperature and pour the HCl in to the same conical flask and time how long it will take for the cross on the paper to disappear. 3. I will do 4 different temperatures and I will do them three times each to get accurate results. 4. I will record the results in a table of results. Apparatus used Sodium thiosulphate NaS2O3-50ml Hydrochloric acid HCl-1M Conical flaskx2 Measuring cylinderx2 Thermometer Water bath at different temperatures Paper marked with X Stop watch Distilled water Analysis From graph 1 I can see that when temperature increases the time taken for reaction to take place decreased. In graph 2 I can see when temperature increases the rate of reaction increases. There was an anomalous result in graph 2, when the temperature was 480C and 1€time was 1.18. My results agree with my prediction because I predicted that the higher the temperature the lower the time taken for the reaction to take place and we can see this from the graphs. The graph shows this pattern taking place. For my experiment I studied temperature as this is easily observed and can be easily varied. The temperature of sodium thiosulphate increased, and the amount of time taken for a reaction decreased. When temperature increased the bonds in NaS2O3 broke quicker and more energy is available greater than the activation energy of the reaction and S2O3 2- ions are available so it takes less time to bond with H+ ions from HCl and new bonds were formed quicker and therefore sulphur precipitated quicker and the rate of reaction increased. This is why in graph 2, I had a strait line positive correlation graph. When the temperature increased it caused an increase in kinetic energy so we had more successful collisions between NaS2O3 particles and HCl particles and the rate of reaction increased. Also more activation energy was made available to overcome the activation energy of the reaction; the reactants had greater energy than the activation energy, so the reaction took place quicker. I think my results support my prediction because I predicted when temperature increases the rate at which the reaction takes place is faster. In graph 2, the theory that every time the temperature increases by 10oC, the rate of reaction will double did not work in my experiment and the results of that theory is given below: 10"¹C¨0.018 0.024€0.018=1.333 20"¹C¨0.024 0.052€0.024=2.167 30"¹C¨0.052 0.078€0.052=1.500 40"¹C¨0.078 0.086€0.078=1.103 50"¹C¨0.086 0.104€0.086=1.209 60"¹C¨0.104 0.120€0.104=1.154 70"¹C¨0.020 0.1380.120=1.1500 80"¹C¨0.138 Evaluation I think my method worked well as I repeated the experiments three times for five different temperatures and got three results which were similar. I think the experiment worked but when we used NaS2O3 with a high temperature, it was difficult for us to time the reaction as it was more rapid than we had expected. If I had the chance to repeat the experiment I would concentrate on the concentration of the NaS2O3 rather than the temperature as there are a lot of factors which could affect the temperature. I think my experiment was done reasonably well as l got similar results when I repeated them three times. There was one anomalous result in graph 2 and I think there was an anomalous result because the NaS2O3 was at a high temperature and the reactants reacted rapidly that the timing was wrong. I also think this was caused by the open window we worked next to which brought the temperature down quickly. I think my results are fairly reliable and it supports my analysis as I said, when temperature increased the time taken for the reaction to take place decreased. I could try the experiment with different methods and different reactants to get additional knowledge. I could use magnesium instead of Sodium thiosulphate and I could heat the hydrochloric acid instead of heating the NaS2O3 and to make more of a fair test I could make sure all the windows and doors are closed and no cold air comes in.   

Rates of reaction Plan Aim: In this experiment I will find the rate of reaction between Sodium thiosulphate NaS2o3 and Hydrochloric acid HCl. There are different variables I could use to see the change in the rate of reaction. These include temperature, concentration or catalysts. I will...

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