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The rate of Hydrochloric acid reacting with magnesium ribbon.
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Introduction In the reaction between hydrochloric acid and magnesium ribbon, the hydrochloric acid will dissolve the magnesium and produce hydrogen gas. All chemical reactions involve reactants which when mixed may cause a chemical reaction which will make products. In my case the reactants are hydrochloric acid and magnesium ribbon. The chemical reaction takes place when the magnesium ribbon is dropped into the hydrochloric acid. The products that are formed during this reaction are hydrogen gas and magnesium chloride. The formula equation for this experiment is: Mg + 2HCl r MgCl2 + H2 Magnesium + Hydrochloric...
be similar to the ones I already have, so I would not use this for any further experiments. The reason I could use a different type of acid for any further work is to find if the is a difference between them if their molecular build up is different.

I could also use the less reactive metals of the reactivity series zinc, aluminium, iron and lead that way I could find the initial rate of reaction at 5 seconds for the higher concentrations of acid like 3.0M or 3.5M and I could find there relative activity.

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For this investigation I am... For this investigation I am reacting magnesium ribbon Mg with Hydrochloric acid HCl I am measuring the rate of reaction between the two, and to do this am measuring the hydrogen given off by the reaction. Magnesium is a shiny silver coloured metal, an element with an atomic number of 12, which belongs to the group 2 alkali metal group; it therefore has only two orbiting electrons and is very unstable and reactive. Hydrochloric acid is a solution of hydrogen and chloride a colourless acidic gas in water, Hydrochloric Acid has a high acidity and therefore will react with an alkali metal, the magnesium will displace the hydrogen and bond with the chlorine giving off hydrogen gas and producing the salt magnesium chloride. To keep the experiment fair I must look at all the factors in the reaction, firstly the concentration of the hydrochloric acid is important because it determines how fast the reaction will happen, more HCl molecules in the hydrochloric acid solution will produce more collisions with the magnesium and dissolve it quicker this is the only variable which I am planning to vary, I will keep the different variations of acid in different containers to prevent contamination and will wash out containers before I use them too, the heat of the reaction will also determine how fast it takes place, more heat means more movement of molecules and more collisions between the HCl and the Mg particles, the amount of magnesium will also effect the reaction rate so I will use measured amounts that are all the same, and the surface area of the magnesium must be kept constant too because more surface area means more collisions and effects reaction speed, the apparatus used will be kept the same to give the reaction the same amount of room to take place in, keeping the equipment still which will reduce vibrations which in turn effect the reaction rate, the amount of hydrochloric acid used will be kept constant too because more HCl acid means more HCl molecules to collide with the magnesium, there is also the pressure at which the reaction takes place which will cause more reaction because the particles would be pushed closer together, although the gas given off would have to be measured at room pressure or the volume would be changed by the pressure, The issue of room temperature and the starting temperature of the solution will also be kept constant by keeping the acid bottles in a constant temperature and measuring the temperature at the beginning of the experiment to make sure that this temperature stays the same and therefore there will be a constant amount of activation energy and the reactions will start with the same amount of heat present and therefore will start at the same rate of reaction if no variables were changed. I am not planning to vary the pressure because it brings up to many complications and is not needed, I will start the stop watch as soon as the magnesium ribbon hits the acid and close the bung instantly, also to do this someone else will time for me while I handle the magnesium and close the bung up, I will also clean the jar in which the reaction was carried out to prevent inconsistencies in the strength of acid and possible side effects caused by residues left from the reaction. To carry out the experiment we will be using a conical flask which the reaction will take place in, a bung with tubes coming out of it for the gas to flow down, 4 small measuring cylinders to measure the amount of acid used the measuring cylinders measure in ml, which is equivalent to cm3 in a ration of 1:1 so we can use them to measure the volume of the gas, a margarine tub filled with water and a large measuring cylinder filled with water which will be put upside down in the water to measure gas produced by the reaction, varied strengths of hydrochloric Acid in 10ml quantities, Magnesium ribbon in4cm lengths, a stop watch to time the reaction, and safety goggles for protection from acid splashes. The hydrochloric acid will be put in the conical flask, the magnesium will be dropped in to start the reaction, the bung will be promptly placed on top of the conical flask, from the bung tubes run in to the margarine tub full of water and under the measuring tube so that when the reaction takes place gas will be pushed through the tube and collect in the measuring cylinder. There will be 4 variations of the acid strength, 2 molar, 1.5 molar, 1 molar, and 0.5 molar. I predict that the higher the concentration of the acid the quicker the reaction, but there will be a point where all the magnesium is depleted and the reaction rate will level out, some of the weaker concentrations will not reach that level, but some of the stronger ones should and there will be a point where there is no more magnesium to react and the gas is no longer produced. Before I did the actual experiments I tried some preliminary tests with some 1m acid and some 2m acid, the hydrogen was produced as soon as I dropped the magnesium ribbon in the acid, and the 2m acid reaction finished quicker than the 1m acid reaction, this determines that my original assumption was correct and the magnesium was dissolved quicker in the 2m acid, although both reactions produced the same amount of gas because I used the same amount of magnesium and therefore the reaction was limited to the amount of magnesium to react with, I tested the gas produced under a flame and it produced a high pitched squeak which indicated the presence of hydrogen and proves more of my hypothesis correct and I can determine that when the magnesium ribbon reacts with the hydrochloric acid, magnesium chloride is formed. Here is a graph to show my preliminary test results. Time 1m HCl 2m HCl 30 "“ 0:30 22 cm3 43 cm3 60 "“ 1:00 33 cm3 45 cm3 90 "“ 1:30 43 cm3 46 cm3 120 "“ 2:00 45 cm3 46 cm3 150 "“ 2:30 46 cm3 46 cm3 180 "“ 3:00 46 cm3 46 cm3 210 "“ 3:30 46 cm3 46 cm3 240 "“ 4:00 46 cm3 46 cm3 It would appear that the reaction levels out at the point where 46 cm3 of hydrogen gas is produced when using a 4cm long piece of magnesium. I wrote down the equation to show the reaction between the 2 reactents: Magnesium + Hydrochloric acid = Magnesium Chloride + Hydrogen Mgs + 2HClaq = MgClaq + Hg I will repeat each experiment 4 times to smooth out inconsistencies and be able to produce an average result. To ensure that the experiment is carried out safely I will wear protective goggles at all times when handling acid, stand up while doing experiments to get out of the way quickly in the case of acid spills, and keep a clear and tidy workspace around me to prevent things getting in the way and being damaged by acid. With the equipment set up I would drop the magnesium In to the acid, then begin timing for a set amount of time even if the reaction had finished, and measure the gas produced in the large measuring cylinder and note the volume every 30 seconds for 4 minutes. Here are my 4 results tables Time 0.5m HCl 1m HCl 1.5m HCl 2m HCl 30 "“ 0:30 7 cm3 23 cm3 39 cm3 43 cm3 60 "“ 1:00 12 cm3 34 cm3 42 cm3 45 cm3 90 "“ 1:30 19 cm3 42 cm3 45 cm3 47 cm3 120 "“ 2:00 26 cm3 46 cm3 47 cm3 47 cm3 150 "“ 2:30 30 cm3 46 cm3 47 cm3 47 cm3 180 "“ 3:00 34 cm3 46 cm3 47 cm3 47 cm3 210 "“ 3:30 37 cm3 46 cm3 47 cm3 47 cm3 240 "“ 4:00 39 cm3 46 cm3 47 cm3 47 cm3 Time 0.5m HCl 1m HCl 1.5m HCl 2m HCl 30 "“ 0:30 7 cm3 21 cm3 38 cm3 44 cm3 60 "“ 1:00 11 cm3 35 cm3 43 cm3 46 cm3 90 "“ 1:30 16 cm3 42 cm3 45 cm3 47 cm3 120 "“ 2:00 27 cm3 47 cm3 48 cm3 47 cm3 150 "“ 2:30 32 cm3 48 cm3 48 cm3 47 cm3 180 "“ 3:00 35 cm3 48 cm3 49 cm3 47 cm3 210 "“ 3:30 37 cm3 48 cm3 49 cm3 47 cm3 240 "“ 4:00 40 cm3 48 cm3 49 cm3 47 cm3 Time 0.5m HCl 1m HCl 1.5m HCl 2m HCl 30 "“ 0:30 6 cm3 20 cm3 38 cm3 44 cm3 60 "“ 1:00 12 cm3 33 cm3 43 cm3 47 cm3 90 "“ 1:30 17 cm3 43 cm3 46 cm3 48 cm3 120 "“ 2:00 26 cm3 48 cm3 49 cm3 48 cm3 150 "“ 2:30 29 cm3 49 cm3 49 cm3 48 cm3 180 "“ 3:00 34 cm3 49 cm3 49 cm3 48 cm3 210 "“ 3:30 36 cm3 49 cm3 49 cm3 48 cm3 240 "“ 4:00 39 cm3 49 cm3 49 cm3 48 cm3 Time 0.5m HCl 1m HCl 1.5m HCl 2m HCl 30 "“ 0:30 8 cm3 25 cm3 40 cm3 45 cm3 60 "“ 1:00 13 cm3 35 cm3 47 cm3 47 cm3 90 "“ 1:30 19 cm3 42 cm3 49 cm3 49 cm3 120 "“ 2:00 26 cm3 49 cm3 49 cm3 49 cm3 150 "“ 2:30 32 cm3 50 cm3 49 cm3 49 cm3 180 "“ 3:00 36 cm3 50 cm3 49 cm3 49 cm3 210 "“ 3:30 39 cm3 50 cm3 49 cm3 49 cm3 240 "“ 4:00 42 cm3 50 cm3 49 cm3 49 cm3 These are the 4 sets of results, I recoded 2 each lesson, the last results seem to be react slightly quicker than the others, this could be due to temperature of the room or contamination of the equipment, the reactions seem to have happened quicker, although they don't seem to be too random and I will use them in my average table set of results. Time 0.5mHCl 1mHCl 1.5mHCl 2mHCl 30 "“ 0:30 7 cm3 22.25 cm3 38.75 cm3 44 cm3 60 "“ 1:00 12 cm3 34.25 cm3 43.75 cm3 46.25 cm3 90 "“ 1:30 17.75 cm3 42.25 cm3 46.5 cm3 47.25 cm3 120 "“ 2:00 26.25 cm3 47.5 cm3 48.25 cm3 47.25 cm3 150 "“ 2:30 30.75 cm3 49 cm3 48.25 cm3 47.25 cm3 180 "“ 3:00 34.75 cm3 49 cm3 48.5 cm3 47.25 cm3 210 "“ 3:30 37.25 cm3 49 cm3 48.5 cm3 47.25 cm3 240 "“ 4:00 40 cm3 49 cm3 48.5 cm3 47.25 cm3 It was noticeable, when looking at the results table, that the more concentrated acid had a faster rate of reaction than the less concentrated acid. This was probably because there are more particles in a concentrated acid and therefore more collisions will occur, for example; the 0.5 molar acid reactions produced on average 7 cm3 of hydrogen gas in the first 30 seconds, whereas the 1.5 molar acid reactions produced on average 38.75 cm3 of hydrogen gas in the first 30 seconds. The results appear to level out at around about 48.4cm3; the concentration of the 0.5 acid reactions does not level out because we stopped the timer before the reaction had time to complete. I have made a graph of the average reaction rate for this experiment. The numbers along the bottom indicate time; the numbers along the side indicate cm3 of gas produced. The graph supports my original prediction, it shows that the higher the concentration of the acid in molars the faster the reaction occurs and hydrogen is produced quicker, therefore I can deduce that In a higher concentration there are more acid particles to react with the magnesium ribbon and therefore it is dissolved faster. Therefore if you increase the concentration of the acid you are introducing more particles into the reaction which will in turn produce a faster reaction because there will be more collisions between the particles which is what increases the reaction rate. If we would have carried on the practical for a longer time the 0.5 molar reactions would have eventually levelled out at about 48.4cm3. While performing the experiment I had to ensure that the temperature was kept constant throughout, because varying temeperature will vary the results, if the temperature increases from the start time to the finish time then the reaction speeds will get quicker at a slightly greater rate, there was also the issue of room temperature which we measured but could not do much about because the room is a large environment and has many sources of heat. The reaction could have been sped up or slowed down in many ways but the amount of hydrogen produced remains constant. There are always ways to improve an experiment like this, I could have measured the temperature of the acid to make sure that it all started at the same temperature, and could have recorded temperature results while doing the practical too, so that I have 2 sets of results for each experiment, and could compare these and analyse how they are relevant to the experiment. Also the measuring of the acid could have been improved using small measuring syringes, and the measuring of the hydrogen produced could also be improved using the gas measuring syringe which would have produced much more accurate results because getting the upturned measuring cylinder in water without letting air in was difficult and the reading of the measuring cylinder could have been improved using the gas measuring syringe because the results would have been more accurate. Also weighing the magnesium ribbon, and cutting it more precisely would have helped get more accurate results, the magnesium was also covered in a white powder, some pieces more so than others, this is magnesium oxide, produced where the magnesium has been exposed to air, the pieces with more magnesium oxide on them would have less magnesium to react with the acid and the oxide may slow the reaction by getting in the way, or reacting with the acid and producing water. I could have cleaned each piece of magnesium with some emery cloth to reduce the magnesium oxide. I could have also tried varying other constants in my experiment, beginning of the year I temperature, the presence of a catalyst, the surface area of the magnesium or the pressure of the reaction chamber. These differentials in the variables would affect the reaction rates in different ways, but the tests all followed the same predicted pattern and shows that there is a level where all the magnesium is depleted, if we had used more magnesium and less hydrochloric acid we could have found a point where the amount of hydrochloric acid levels out before the magnesium, but we would have to use a lot of magnesium because using a small amount of hydrochloric acid would make it much harder to measure the results with current equipment. I also did some extra tests using 3 cm pieces of magnesium, and measuring how much gas was produced, to do these measurements we used the same equipment apart from the measuring cylinder which was replaced with the gas measuring syringe, the measuring syringe was much more accurate than the cylinder, and gave us better readings. The amount of gas produced from the 3cm piece of magnesium levels out at a point of 32 cm3 if we put this in a line graph with the maximum amount of hydrogen produced from the 5.5 cm pieces of magnesium we can predict how much the most amount of hydrogen that can be produced by a reaction between a different length of magnesium with hydrochloric acid. Here is a graph of the readings I got using the 4 cm piece of magnesium. Time 0.5m 1m 1.5m 2m 30 - 0:30 8 12 27 29 60 - 1:00 12 17 33 31 90 - 1:30 17 25 34 31 120 - 2:00 19 28 34 31 150 - 2:30 23 33 34 31 180 - 3:00 26 35 34 31 210 - 3:30 28 35 34 31 240 - 4:00 30 34 34 31 the results are much more varied in these tests, this could be due to greater accuracy and being able to note these variances, or from some sort of contaminant, the average amount of gas produced for a 5.5 cm piece of magnesium was 48.4 cm3 the average amount of gas produced for a 3 cm piece of magnesium was 42 cm3 the more magnesium means there are more magnesium atoms to react with the hydrochloric acid molecules and therefore more hydrogen is produced, if we put these in to a line graph we can use it to estimate how much hydrogen would be produced by other lengths of magnesium. I have drawn a line of best fit between the two points, with this graph we can estimate how much hydrogen would be produced if we reacted 1cm of magnesium. The graph gives a reading of around about 18 cm3 of hydrogen gas produced, if I had extra time we could test this theory, but unfortunately we do not. I also only did one set of results for the tests with the gas syringe, if I would have been able to continue this experiment further I could have produced more average results and seen if my predictions for the 1 cm piece of magnesium was correct. We could have also varied the concentration of the acid more so, and used less or more acid to get more accurate results or results for different test situations, instead of changing the strength we could have changed the amount of acid, or the temperature of the acid, or try varying these together and see how they effect each other.   

For this investigation I am reacting magnesium ribbon Mg with Hydrochloric acid HCl I am measuring the rate of reaction between the two, and to do this am measuring the hydrogen given off by the reaction. Magnesium is a shiny silver coloured metal, an element with an atomic...

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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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Rate Of Reaction in... Rate Of Reaction in Sodium Thiosulphate and HCl Plan in this experiment I am investigating the rates of reaction, and the effect different changes have on them. The rate of reaction is the rate of loss of a reactant or the rate of formation of a product during a chemical reaction. A rate of reaction is measured by dividing 1 by the time taken for the reaction to take place. There are five factors which affect the rate of a reaction, by the collision theory of reacting particles: temperature, concentration of solution, pressure in gases, surface are of solid reactants, and catalysts. Aim: - my aim is to see the effects of a change in temperature and concentration on the rate of a reaction. The reaction that will be used is: Sodium Thiosulphate + Hydrochloric Acid Na2S2O3 aq + 2HCl aq Sodium Chloride + Water + Sulphur Dioxide + 2NaCl aq + H2O l + SO2 g + Sulphur S s Two experiments will be carried out "“ one changing the temperature while everything else remains constant and one varying the concentration while keeping everything else constant. Both the sodium thiosulphate and the Hydrochloric acid are soluble in water, so the concentration of either can be changed. temperatures and concentrations to use during my preliminary series of experiments "“ 50cm3 of sodium thiosulphate solution and 5cm3 of hydrochloric acid as the experiment one con 1 mol/dm3 of HCl acid concentration will be fixed 10-35g/dm3 of sodium thiosulphate all of these concentrations will be tested in turn going up in steps of 5g/dm3 20-70°C temperature all of these temperatures will be used going up in steps of 10°C Concentrations of 5, and 40 g/dm3 of thiosulphate were available to me but my preliminary work showed that the 5 g/dm3 and 40g/dm3 were too slow and fast respectively in reacting to be worth testing. Similarly any temperature below 20°C reacted too slowly, and 80°C and 90°C reacted too quickly to be worth including in my final results. Using my preliminary experiments I decided on using the following apparatus: 1 thermometer 1 beaker 2 measuring cylinders 1 conical flask 1 tripod 1 gauze 1 heatproof mat 1 stopwatch 1 Bunsen burner X board 1 pair of tongs 1 pair of goggles 1 apron Method: - Experiment 1 - Changing the concentration 5 cm3 of HCl at concentration 1 mol./dm3 and 15 cm3 of sodium thiosulphate at varying concentrations "“ 10 to 35 g/dm3 are poured out into two measuring cylinders and then poured into a conical flask, which is placed on top of a board marked with letter X. The stopwatch will now be started. When the mixture has turned sufficiently cloudy so that the letter X can no longer be seen the stopwatch will be stopped and the time will be recorded. The experiment is repeated with all the concentrations. The whole procedure is then repeated. Experiment 2 "“ Changing the temperature 5 cm of HCl at concentration 1 mol./dm3 and 15 cm of sodium thiosulphate at varying concentrations "“ 10 to 35 g/dm3 are poured out into two measuring cylinders. A beaker is half filled with hot water from a tap. The water is placed on top of a Bunsen on a blue flame and the two measuring placed inside the water bath. The water is heated to the necessary temperature 30°C to 70°C then the two measuring cylinders are taken out and the contents of both are poured into a conical cylinder. The time it takes for the X to disappear is timed and recorded. The experiment is repeated using all the temperatures. The entire procedure is the repeated. Repeat results and averages will be taken to improve the credibility of the findings, and present solid grounding for the final conclusion. The repeat results will help to iron out any anomalies and the average will give a good summary of the results of the experiment. However if one set of results is entirely different to the other, a third experiment will be performed to replace the anomalous set of results. Safety "“ A pair of goggles will be worn during the heating part of the experiment in order to protect the eyes. An apron will also be worn to protect the skin and clothing. When handling hot beakers and measuring cylinders a pair of tongs will be used. A gauze and heatproof mat will be used while heating to avoid any damage to the equipment. Fair Test - In order for my findings to be valid the experiment must be a fair one. I will use the same standard each time for judging when the X has disappeared. I will make sure that the measuring cylinders for the HCl and thiosulphate will not be mixed up. The amount of HCl will be 5 cm3 each time, and the amount of thiosulphate will be fixed at 15 cm3. During the heating stage of the experiment, a blue flame will be used throughout. Also the same Bunsen burner and gas tap will be used to maintain continuity. All of these precautions will make my final results more reliable and keep anomalies at a minimum so thus make the entire investigation more successful. Prediction "“ I predict that as the temperature is increased the rate of reaction will increase. I also predict that as the concentration of the sodium thiosulphate increases the rate of reaction will increase. This means that both graphs drawn up in my analysis will have positive correlation, and will probably be curved as the increase in rate of reaction will not be exactly the same as the concentration temperature is increased. This can be justified by relating to the collision theory. When the temperature is increased the particles will have more energy and thus move faster. Therefore they will collide more often and with more energy. Particles with more energy are more likely to overcome the activation energy barrier to reaction and thus react successfully. If solutions of reacting particles are made more concentrated there are more particles per unit volume. Collisions between reacting particles are therefore more likely to occur. All this can be understood better with full understanding of the collision theory itself: For a reaction to occur particles have to collide with each other. Only a small percent result in a reaction. This is due to the energy barrier to overcome. Only particles with enough energy to overcome the barrier will react after colliding. The minimum energy that a particle must have to overcome the barrier is called the activation energy, or Ea. The size of this activation energy is different for different reactions. If the frequency of collisions is increased the rate of reaction will increase. However the percent of successful collisions remains the same. An increase in the frequency of collisions can be achieved by increasing the concentration, pressure, or surface area. Concentration "“ If the concentration of a solution is increased there are more reactant particles per unit volume. This increases the probability of reactant particles colliding with each other. Pressure - If the pressure is increased the particles in the gas are pushed closer. This increases the concentration and thus the rate of reaction. Surface Area "“ If a solid is powdered then there is a greater surface area available for a reaction, compared to the same mass of unpowdered solid. Only particles on the surface of the solid will be able to undergo collisions with the particles in a solution or gas. The particles in a gas undergo random collisions in which energy is transferred between the colliding particles. As a result there will be particles with differing energies. Maxwell-Boltzmann energy distribution curves show the distribution of the energies of the particles in a gas. The main points to note about the curves are: 1. There are no particles with zero energy. 2. The curve does not touch the x-axis at the higher end, because there will always be some particles with very high energies. 3. The area under the curve is equal to the total number of particles in the system. 4. The peak of the curve indicates the most probable energy. The activation energy for a given reaction can be marked on the distribution curve. Only particles with energy equal or greater than the activation energy can react when a collision occurs. Although Maxwell-Boltzmann distribution curves are for the particles in a gas, the same distributions can be used for the particles in a liquid or solid. Effects of a temperature change - The graph below shows Maxwell-Boltzmann distribution graphs for a fixed mass of gas at two temperatures "“ T1 and T2, where T2 is roughly 10°C higher than T1. The total area under the curve remains the same, since there is no change in the number of particles present. A small increase in temperature causes significant changes to the distribution energies. At the higher temperature: 1. The peak is at a higher energy. 2. The peak is lower. 3. The peak is broader. 4. There is a large increase in the number of particles with higher energies. It is the final change that results increase in rate, even with a relatively small increase in temperature. A small increase in temperature greatly increases the number of particles with energy greater than the activation energy. The shaded areas on the energy distribution curves show this. Effect of a catalyst - A catalyst works by providing an alternative reaction pathway that has lower activation energy. A catalyst does not alter the Maxwell-Boltzmann distribution. Because a catalyst provides a reaction route of lower activation energy, however, a greater proportion of particles will have energy greater than the activation energy. Secondary Sources Used: AS Level Chemistry Textbook kinetics module The Internet Dr. Jones's Chemistry Lessons Information sheets from Dr. Jones Obtaining Evidence Temp.°C Time 1 s Time 2 s Average s 20 110.67 107.42 109.045 30 100.13 103.34 101.735 40 64.20 65.92 65.06 50 45.34 37.73 41.535 60 30.12 33.18 31.65 70 18.92 16.34 17.63 Concen.g/dm3 Time 1 s Time 2 s Average s 10 222.63 224.38 223.505 15 150.90 147.03 148.965 20 105.25 105.97 105.61 25 66.04 68.75 67.395 30 55.63 56.1 55.865 35 27.32 25.96 26.64 Temp.°C Rate of Reaction 1s-1 Rate of Reaction 2 s-1 Average s-1 20 0.00904 0.00931 0.00917 30 0.00999 0.00968 0.00983 40 0.01558 0.01517 0.01537 50 0.02206 0.02650 0.02428 60 0.03320 0.03014 0.03167 70 0.05285 0.06120 0.05703 Concen.g/dm3 Rate of Reaction 1s-1 Rate of Reaction 2 s-1 Average s-1 10.00000 0.00449 0.00446 0.00447 15.00000 0.00663 0.00680 0.00671 20.00000 0.00950 0.00944 0.00947 25.00000 0.01514 0.01455 0.01484 30.00000 0.01798 0.01783 0.01790 35.00000 0.03660 0.03852 0.03756 Temp.°C Rate of Reaction 1s x1000 Rate of Reaction 2 s x1000 Average s 20 9.04 9.31 9.17 30 9.99 9.68 9.83 40 15.58 15.17 15.37 50 22.06 26.50 24.28 60 33.20 30.14 31.67 70 52.85 61.20 57.03 Concen.g/dm Rate of Reaction 1s x1000 Rate of Reaction 2 s x1000 Average s 10 4.49 4.46 4.47 15 6.63 6.80 6.71 20 9.50 9.44 9.47 25 15.14 14.55 14.84 30 17.98 17.83 17.90 35 36.60 38.52 37.56 Analysis In this experiment I have found that as the temperature and concentration is increased the time taken for the reaction to take place decreases. This means the rate of reaction increasers as it takes less time for a reaction to take place, so more take place per second. In the temperature experiment the time taken for a reaction to take place decreased by roughly 10 to 15 seconds for every 10°C increase in temperature, with the one anomaly being the 30°C reading. There is also a trend in the increase in rate of reaction as the temperature increases. The difference is always more or less 0.02 s-1, with the same exception. Using the graphs, with lines of best fit, I can draw a conclusion from my experiment. Firstly I can see that with the "time" graphs that plot temperature and concentration against time taken for the reaction to take place the graphs have negative correlation in both cases, meaning that as the temperature concentration increased the time taken for the reaction to take place decreases. The time graph for the temperature experiment has a much steeper curve than the one for the concentration experiment, meaning that the decrease in time taken for the reaction was far more rapid. Naturally, the above means that the both the graphs plotting rate against temperature and concentration have positive correlation "“ as the temperature and concentration are increased so does the rate of reaction. This is because when the temperature is increased the particles will have more energy and thus move faster. Therefore they will collide more often and with more energy. Particles with more energy are more likely to overcome the activation energy barrier to reaction and thus react successfully, and when solutions of reacting particles are made more concentrated there are more particles per unit volume. Collisions between reacting particles are therefore more likely to occur. The graph for concentration shows that when the concentrations were relatively low 10, 15, 20 g/dm3, the increase of rate x1000 was also fairly small increasing from 4.47 to 6.71 to 9.47. There was then a gradual increase in the difference, and between 30 and 35 g/dm3 the rate more than doubled from 17.90 to 37.56s-1. This shows that there are far more collisions at a concentration of 35 g/dm3 than at 30 g/dm3. The graph plotting time against the rate of reaction x1000 shows that the difference of rate between increasing temperatures excluding the anomaly of 30°C was pretty much regular, increasing in steps of 6-10 9.17 to 15.37 to 24.28 to 31.67. However, once again there is a giant gap in the last temperature increase "“ at 60°C the RoR x1000 is 31.67 s-1, and at 70°C it is 57.03 s-1. For this to fully make sense it is necessary to recap the collision theory briefly: For a reaction to occur particles have to collide with each other. Only a small percent result in a reaction. This is due to the energy barrier to overcome. Only particles with enough energy to overcome the barrier will react after colliding. The minimum energy that a particle must have to overcome the barrier is called the activation energy, or Ea. The size of this activation energy is different for different reactions. If the frequency of collisions is increased the rate of reaction will increase. However the percent of successful collisions remains the same. An increase in the frequency of collisions can be achieved by increasing the concentration, pressure, or surface area.   

Rate Of Reaction in Sodium Thiosulphate and HCl Plan in this experiment I am investigating the rates of reaction, and the effect different changes have on them. The rate of reaction is the rate of loss of a reactant or the rate of formation of...

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