Conclusion
I can conclude that if you double the concentration of the acid the reaction rate would also double, this is because the ions are closer together in a concentrated solution. The closer together they are, the more often the ions collide. The more often they collide, the higher the chance of a reaction between the magnesium and the hydrochloric acid. Also because there are more particles in the solution which would increase the likelihood that they would hit the magnesium so the reaction rate would increase. The graph gives us a good device to prove that if you double the concentration the rate of reaction doubles. If you increase the number of particles in the solution it is more likely that they will collide more often. In the reaction, when the magnesium hit the acid, it fizzed and produced many bubbles it was silver in colour (which is one of magnesium’s physical properties silvery
white metallic element), the activation energy of a particle gets higher with heat, the particles which have to have the activation energy are those particles which are moving, in the case of magnesium and hydrochloric acid, it is the hydrochloric acid particles which have to have the activation energy because they are the ones that are moving and bombarding the magnesium particles to produce magnesium chloride. The graph for 1/time had the form of an s curve and it did have some relationship with the other graph.
Evaluation
There are many reasons why our results for the 1/time graph did not prove the point that concentration~1/time, such as
1. When the reaction takes place bubbles of H2 are given off, which might stay around the magnesium, which therefore reduces the surface area of the magnesium and so the acid cannot react properly so this affects the results.
2. We could have controlled factors in the investigation better (e.g. the stirring of the solution because if this is not done properly it can lead to incorrect results).
3. Using larger concentrations of acid would give a bigger more accurate conclusion instead of just using 10ml test tubes use 1litre test tubes, this way graphs would be more spaced out and give an accurate form or curve.
metallic element; symbol Mg, atomic number 12, found in group two in the periodic table. It is quite reactive giving vigorous reactions towards acids. It is one of the alkaline earth metals, and the lightest of the commonly used metals. It is used in alloys, flash photography, flares, fireworks and flash bulbs because it burns vigorously in air with a bright white light. Magnesium reacts with steam to release hydrogen and it also burns in carbon dioxide gas.
Hydrochloric acid, HCl, is a solution of hydrogen chloride (a colorless acidic gas) in water. The concentrated acid is about 35% hydrogen chloride and is corrosive. The acid is a typical strong, monobasic acid forming only one series of salts, the chlorides. Like most acids, it releases hydrogen ions when it is added to water and certain metals, and has a pH of less than 7. Hydrochloric acid is a common laboratory acid.
Extracts from: a chemistry coursework from an Internet source and
The HUTCHINSON Dictionary of SCIENCE second edition.
Before looking at the factors that can alter the rate of reaction, we must consider what happens when a reaction take place.
First of all, the particles of the reacting substances must collide with each other and, secondly, they need a certain amount of energy to break down the bonds of the particles and form new ones. This energy is called the activation energy or Ea. If a collision between particles can produce sufficient energy (i.e. if they collide fast enough and in the right direction) a reaction will take place. Not all collisions will result in a reaction.
The investigation could be done using one variable and therefore have a set of results which were related in some way. The variables that could be used are:
1. Concentration
2. Particle size/surface area
3. Pressure (for reactions involving gas)
4. Temperature
5. Light
6. Presence of a catalyst.
These variables can be used because:
1. The more concentrated the reactants, the greater the rate of reaction will be. This is because increasing the concentration of the reactants increases the number of collisions between particles and, therefore, increases the rate of reaction.
2. When one of the reactants is a solid, the reaction must take place on the surface area of the solid. By breaking up the solid into smaller pieces, the surface area is increased, giving a greater area of collisions to take place and so causing an increase in the rate of reaction.
3. When one or more of the reactants are gases an increase in pressure can lead to an increased rate of reaction. The increase in pressure forces the particles closer together. This causes more collisions and increases the rate of reaction.
4. An increase in temperature produces an increase in the rate of reaction. A rise of 10º C approximately doubles the rate of reaction. When a mixture of substances is heated, the particles move faster. This has two effects. Since the particles are moving faster they will travel greater distance in a given time and so will be involved in more collisions. Also, because the particles are moving faster a larger proportion of the collisions will exceed the activation energy and so the rate of reaction increases.
5. The rates of some reactions are increased by exposure to light. Light has a similar effect as temperature because it produces heat.
6. A catalyst is a substance, which can alter the rate of a reaction but remains chemically unchanged at the end of the reaction. Catalysts usually speed up a reaction. A catalyst, which slows down a reaction, is called a negative catalyst or inhibitor. Catalysts speed up reactions by providing an alternative pathway for the reaction, i.e. one that has much lower activation energy. More collisions will, therefore, have enough energy for this new pathway.
Extracts from: Letts Study Guide, GCSE CHEMISTRY.
All this information is relevant to my investigation, as I now know what would happen to the molecules when using different variables. It also makes it easier to decide what variable I am going to use in this circumstance.
I decided to use the concentration of acid as my variable. I used 5 different strengths of hydrochloric acid. These strengths would determine the rates of reactions. I decided to measure the acid in millilitres. I predicted that the higher the concentration of the acid, the faster the reaction between magnesium ribbon and the hydrochloric acid. This would be because there were more acid molecules to react with the magnesium ribbon. I decided that I would do 5 experiments and the different concentrations of hydrochloric acids were:
v 1.0 molar
v 1.25 molar
v 1.50 molar
v 1.75 molar
v 2.0 molar
Before doing the actual experiment I decided to do some preliminary work. These were to tell me the details that I would need to know for my investigation to be successful. I saw from these preliminary investigations that the magnesium ribbon started to react with the hydrochloric acid the moment that I dropped it in. I decided that it would be a good idea to start timing the second that I dropped it in. When the ribbon had been eaten away by the acid, it stopped fizzing. I decided that I would stop timing the second that the fizzing stopped. I discovered from my preliminary experiments that when I used a low concentration of hydrochloric acid, for instance, 0.25 molar, it took a long time for the magnesium ribbon to be eaten away. I decided that it would be impractical to spend time on the following strengths of hydrochloric acid:
v 0.25 molar
v 0.50 molar
v 0.75 molar
This was because they were the three slowest strengths of acids available to react with the magnesium ribbon over a period of time. I also learnt from my preliminary experiments that it was sometimes quite difficult to stop timing on the exact moment that the fizzing stopped. I decided therefore that I would carry out each of the 5 experiments three times and find the average time as this would result in a more accurate figure.
When the magnesium ribbon reacts with the hydrochloric acid, magnesium chloride is formed. I wrote down the equation to show this:
Magnesium + Hydrochloric acid = Magnesium Chloride + Hydrogen
Mg + 2HCl = MgCl + H
The equipment I needed for the investigation were:
v Magnesium ribbon- 15 pieces, 1cm long and weighed 0.01g
v Hydrochloric acid – 30 ml of 1.0 molar
30 ml of 1.25 molar
30 ml of 1.50 molar
30 ml of 1.75 molar
30 ml of 2.0 molar
v Test tubes – 5
v Test tube rack
v Stop clock
v Pipette
v Measuring Cylinder
v Thermometer
v Safety goggles
I decided to do 5 experiments, three times each, using all the information that I gained while I was doing my preliminary experiments.
To ensure a safe experiment and working environment I needed to have at least 1 meter squared of working space around me, wear safety goggles at all times when using acid, use a test tube rack instead of holding the test tubes, secure all equipment and make sure that all the equipment were fully functional and not damaged.
To make the experiment a fair test I used the same amount of acid for all experiments, only changing the concentrations. I used the same size of magnesium ribbon and weight (approximately 0.01g). I also started the stop clock when the magnesium touched the acid and stopped it when the magnesium stopped fizzing for each experiment. I always washed out the test tubes when an experiment had finished so the different concentration wouldn´t get mixed together causing strange results.
First I measured out the amount of hydrochloric acid using the measuring cylinder. I used a pipette to pour the acid into the measuring cylinder as to be accurate. I needed 10 ml of acid in the cylinder and poured it into a test tube. I then put a thermometer into the test tube for 1 minute to check the temperature. I did this to see if the experiments with the same strength of acids affected the rate of reaction if there was a change in temperature. I then got a piece of magnesium ribbon about 1 cm long weighing 0.01g and dropped it into the acid and started timing the moment that the magnesium ribbon touched the acid solution. When the magnesium ribbon stopped fizzing, I stopped the clock and recorded the number of seconds (rounded up to nearest second) taken for the reaction from start to finish.
I made a table to record my results in. The table is shown below.
Experiment Strength of hydrochloric acids in 10 ml No. of Mg ribbon pieces (0.01g) Test 1 secs. Temp Test 2 secs. Temp Test 3 secs. Temp AverageSecs.
1 1.0 molar 1 151.0 18 ºc 141.0 18 ºc 117.0 18.4ºc 136.3
2 1.25 molar 1 81.0 17 ºc 76.0 17.5ºc 74.0 18 ºc 77.0
3 1.50 molar 1 58.0 17.5ºc 56.0 18 ºc 56.0 18 ºc 56.3
4 1.75 molar 1 50.0 18 ºc 41.0 18º c 31.0 18 ºc 40.7
5 2.0 molar 1 20.0 19 ºc 22.0 19 ºc 16.0 21 ºc 19.3
To calculate the average time that it took for the magnesium to be eaten away by the acid, I did the following calculation:
Test 1 + Test 2 + Test 3 = Average time
3
As I already have mentioned, I used a measuring cylinder to make the measurements and used a pipette for further accuracy. I did each experiment three times so I would be able to calculate averages and thereby get more accurate results. I recorded the results in seconds instead of minutes in order to obtain more precise results. I used a stop clock instead of a 24-hour clock so I could look at the milliseconds and round it up to the nearest second, which made the results more exact.
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 instance, 1.0 molars´ average time, 136.3 seconds, is longer than 2.0 molars´ average time, which was 19.3 seconds.
I made a graph to show the results.
The graph above supports my original prediction of: the more concentrated the acid the faster the rate of reaction because it shows the time difference between the different strengths of acids. In a higher concentration there are more acid particles to react with the magnesium ribbon and therefore it is eaten away faster.
I conclude that changing one factor does have a significant effect on the rate of reaction as we have seen.
Looking at the set of results obtained, you can clearly see that they all follow the expected pattern. This is pattern suggests that the reaction rate increase when the concentration of the acid increases because 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.
The evidence I have been able to gather from this investigation seems to lead to a quite firm conclusion. I might not have been able to find the exact speed of the reactions but the pattern seems to be correct as I have repeated readings three times and as it agrees with the information I have researched.
I used the variable of concentration, which seemed to be of a good choice as it would show the results of how more acid molecules reacting with magnesium, would result in a faster reaction.
There will always be ways in which you can improve your investigations and the same thing goes to my investigation.
I found it very hard to measure out the exact number of millilitres for the acid even though I used a pipette and I was also in a hurry. If I was to redo this investigation I would put some more effort into measuring the acid. I could have used the wrong concentration of acid by accident and that would have affected the speed because there would have been fewer or more acid particles to react with the magnesium ribbon. Next time I do this experiment I would try to remember which acid I am using so it doesn´t get mixed up. Every time I washed a test tube or a measuring cylinder, I did not dry it before using it. This may have affected the rate of reaction, as water would dilute the acid. To improve my results, I could dry the test tubes and the measuring cylinder after they are washed to prevent diluted acids. The size and weight of the magnesium would have affected the rate of reaction. The experiment could be improved by measuring, adjusting and weighing the magnesium ribbons so they all are the same size and weight. I also found out from background information, that the magnesium ribbon is covered with a whitish deposit. This deposit was magnesium oxide where the magnesium had reacted with the air. I would imagine that some pieces had only a little of this oxide and some had a lot. The pieces of magnesium ribbon that did not have much oxide on them reacted faster than those with a lot. To improve my results, I could clean the magnesium oxide of all the magnesium pieces using some sandpaper, and this would mean that the acid would not have to eat through the magnesium oxide before reacting with the magnesium. In my investigation I also measured the temperature to see if there would be any change in the rate of reactions. From my results it was noticeable to look at each experiment and see how the temperature had affected the rates of reactions. For instance, if we refer back to the table on page 5 and look at experiment 5, test three is slightly faster than test 1 or test 2 because it has a temperature of 21º, where as the other two tests both have a temperature of 19º. To improve my investigation I could assure that the temperature was constant all the time. I could also do more readings to get nearer to a more accurate result.
In my investigation I used concentration as my variable. To improve my investigation further, I could use other variables such as, surface area, temperature, pressure for gas, and a presence of a catalyst.
These variables would hopefully prove that they all help speed up a chemical reaction.