Preliminary work
For the initial experiments I decided that I would use magnesium metal and hydrochloric acid to carry out a study on the rate of reaction. I used these reactants because I know that an alkaline earth metal plus acid produces hydrogen. Therefore to measure the rate of reaction of the effect of concentration, I would measure the rate at which hydrogen was produced. This is because it is easy to record the volume of gas per selected unit of time. To evaluate what concentrations, amount of acid, equipment and time period I would use, I did some experiments. From the results I recorded, I believe that a timing interval of 20 seconds for 0.5 mole, 10 seconds for 1 mole and 5 seconds for 1.5 and 2 mole was appropriate. I will use these time periods because the 0.5 mole solution reacts slowly, so in this case I want a large time period so I can study the results more effectively. For a 2 mole solution that reacts a lot quicker than a 0.5 mole solution I want a very short time period as the reaction is spaced out over a very short period. Also I decided that 30cm cubed of acid would be used, as this would make sure the reaction happened fully and did not stop it prematurely and that 0.06g of magnesium strip would be used. In addition to cut out experimental error the experiment I decided to repeat the experiment 4 times.
Theory
The law of the collision theory, states that to have a reaction you need a sufficient energy, (activation energy), with a sufficient collision to create an efficient reaction.
Quote from chemistry for you on the law of collision theory
To explain this law on the effect of concentration on the rate of reaction, you must simply understand that as concentration of the acid increases, there are more acid particles in the same area. Therefore there is a greater chance in a more concentrated solution of acid particles colliding, and reacting, with particles on the surface of the magnesium. As to give an example of an increase in rate of reaction in modern life.
It is like a dance floor. Before a party, few people are on the dance floor, so the chance of bumping into someone is slight. When it begins, lots more people crowd on the dance floor. Therefore the concentration of people on the dance floor increases. Now there is more chance of bumping into someone than before. This is illustrated below, with two different concentrations of acid.
1 mole 2 mole
For this reaction, I would expect using the law of collision that the 2-mole experiment would be faster and double the speed. In addition, I predict that the product would be producing the same volume of gas. This is because we are using the same variables while just changing concentration. The mass of magnesium and volume of acid will also be the same and because of this, there can be no extra product produced but can only be varied in speed of reaction. We can prove this by looking at the results and where the reaction stops. This is shown below.
Most molecular collisions do not result in chemical change. Before any change takes place on collision, the colliding molecules must have a minimum kinetic energy called the Activation Energy. This is shown on the energy level diagrams below. It does not matter whether the reaction is an exothermic or an endothermic energy change. Now when heated molecules have a greater kinetic energy, a greater proportion of them have the required activation energy to react. The increased chance of higher energy collisions greatly increases the speed of the reaction.
Activation Energy simple definition
“Amount of energy that must be absorbed by reactants in their ground states to reach the transition state so that a reaction can occur.”
Prediction
Because of my theory above, my prediction for the effect of concentration on the results of reaction time would be that with greater concentration the greater the rate at which the reaction will take place. Meaning that a reaction of 2 mole would be twice as fast as a reaction of 1 mole. My reasons for this is as the theory states I know that the more concentrated solution will react quicker than a less concentrated solution if all other variables are kept the same.
Secondary sources
Internet:
Collision Theory of Reaction Rates
One of the most stable theories in the Physical Sciences is the Collision Theory of Reaction Rates. According to this theory, product formation can only take place when there are "effective" collisions between reactant molecules involved in the rate-determining step of the process. Simply having reactant molecules colliding is necessary but not sufficient in it. The collisions have to be effective.
What constitutes an effective collision?
In order for a molecular collision to be effective it must meet two conditions:
- The collision must have sufficient enough impact energy to overcome the Activation Energy. The Activation energy is the minimum energy necessary for product to form. This impact energy must be sufficient so that bonds can be broken within the reactant molecules and new bonds formed to produce the products.
- The molecules must have a proper positioning for effective collisions to occur.
If we examine the four factors that influence the rate of a reaction we can see how each of these postulates result in these factors influencing the rate.
- Nature of the reactants. If we crush the reactants, we are in essence increasing the total surface area that collisions can take place. This will have an enhanced effect on rate of product formation. Gaseous reactants have a higher Kinetic energy, and therefore the impact energy will be greater resulting in a higher rate of product formation
- Concentrations of the reactants in the rate-determining step. If we increase the concentration, we are, in essence, increasing the total number of collisions. This, in turn, will increase the number of effective collisions, and the rate will increase.
- Temperature- increasing the temperature increases the average Kinetic Energy of the molecules. This will increase the impact energy enough to overcome the Energy of Activation.
- Catalyst- Catalyst provides a surface whereby the reacting molecules might position themselves more favourably for collision.
The Collision Theory of Reaction Rates explains how each of the above factors affects the rate of the reaction.
Analysing Evidence and Drawing Conclusions
Explanation of what has been found out
The experiment proved that if, “concentration,” is changed so will the rate of reaction.
In the experiment I found and confirmed my results to prove my predications and theory, for the effects of concentration on the rates of reaction. My predictions were that as concentration increases so did the rate of reaction, so that an experiment of 2 moles would be faster and double the speed of a 1-mole reaction. This prediction was seen on the graph; the reaction of 2 mole was a very steep gradient, producing the largest volume of gas in the shortest amount of time: where as the reaction of 0.5 mole had a very flat gradient and produced the least amount of gas in a certain time. Also it proved the second prediction, which was that all experiments would produce the same amount of product, “HYDROGEN,” as all experiments although varying in speed all produced the same amount of gas, which were around 63cm3
- The reaction that took place
Mg(s) + 2HCl(AG) = MgCl2 (AQ) +H2 (G)
Magnesium +Hydrochloric acid = Magnesium Chloride +Hydrogen
Patterns between the graphs (Drawn on graph sheet)
The most basic patterns that can be concluded from the graph are that with a more concentrated solution, the rate of reaction increases and the gradient is steeper. Therefore the order of reaction was: 2 moles being fastest followed by 1.5 moles, than 1 mole and the slowest being 0.5 moles. In all experiments for the same mole, the rate at which the graphs gradient increased can be seen to be the same. At the start of the experiments the graphs all are faster in their reaction than at any other stage and for 1, 1.5 and 2 moles the graph increases rapidly. Than all graphs begin to slow down and the gradient was not as vertical, this can be seen to be the second stage of the experiment and than begin to level out till they have reached a certain stop in production of hydrogen, which is the final stage. Thus meaning that all graphs finish with a straight horizontal line, at the same hydrogen production.
From these results I am able to conclude that for a reaction to take place you need two reactants and if you increase the concentration of the solution, you are therefore increasing the amount of particles to react with the magnesium, which will cause a faster reaction. This conclusion proves collision theory, therefore supporting that any change to the variables will cause a change in reaction.
From the experiment I have used my prediction and the results obtained to make a conclusion that,
With a greater concentration the rate of reaction will increases, so that a reaction of 2 moles would be double the time of a 1-mole solution. When also using the same quantities of reactants, the product that is produced will be the same. Results table that was used to create a graph that resembled my theory.
To create a straight-line graph I firstly draw a tangent at the point where you wish to measure the gradient. I then used the straight-line formulae, which is:
Y=mx+c
Graph on paper
The gradient of the graph was in correlation with my predicted graph. At the beginning of the reaction the gradient was very steep or steeper than any other point of the reaction. At the second stage of the experiment, the middle, the gradient was slightly flatter and than flattened out towards end creating a horizontal line at the end. The line graphs gradient began steep and than formed a curve as it levelled off.
When I placed the results against y/x, I found that if the concentration was close to my prediction that the time of the reaction would double, when the concentration was doubled. For the reaction of 2 mole, I discovered that roughly 5.8cm3 of gas was produced every second and for 1 mole, 2cm3 of gas was produced. When I doubled the production of gas for 1 mole in 35 seconds, I got 4cm3 per second against the 2-mole solution, which were 5.8 moles. This can be explained below in processed information.
To find the gradient of each graph I used the formulae:
Gradient: m=y/x
Graph on paper
- Scientific knowledge for processed information
At the beginning of the experiment the rate at which the experiment produces hydrogen is much faster than at any other stage. This is because at this point in the experiment there are more particles have not reacted in any way, and so there is more chance of an easy efficient reaction. In the middle of the experiment or at the second stage, the reaction has already been taking place and the reactants are being used and there is magnesium chloride particles being produced. This will interfere with a clean reaction. Therefore there is less reactant; therefore less chance of a reaction to take place and the experiments gradient begins to decrease. In the final stage at the end of the experiment the reactants have been completely used up, creating hydrogen and magnesium chloride and therefore there cannot be a reaction, so the experiment levels off to a horizontal gradient, (flat line).
The slight differences in the volume of gas produced each unit of time could be explained. The production of hydrogen could have insulated the surface of the strip and not allowed the acid particles to react at the same rate as other experiments but not to cause a change in the final product. The inaccuracies could have also been caused by the magnesium strip varying slightly in length and volume of hydrochloric acid not being in exactly the same quantities or concentration. Also the gas syringe throughout the experiment could have been stuck and as the pressure increased the syringe could have jumped
For the processed information I found there to be no mathematical relationship between concentration and the rate of reaction but did find that when the table is crossed examined at points in a certain experiment, that the production of hydrogen is the same. This is because at each of these points and up to them the reactants have been reacting in the same way using the same quantities of reactants, therefore producing the same volume of gas.
- How does results contrast with prediction
For the experiment no anomalous results were found, therefore concluding that all experiments were successful with little inaccuracy. My results were therefore found to be accurate in proving my prediction that the effect of a more concentrated solution on the rate of reaction will cause a faster experiment, if all other variables are kept the same. The experiment to an extent proved that the product would produce the same quantity of hydrogen if all other variables were kept the same; this proved one of my predictions. My results proved this, although my average results for the volume of hydrogen were; 63.3, 62.5, 63.7 and 62.8, which are slightly inaccurate. These inaccuracies were expected, as not every experiment can be accurate to a decimal place, although the results are accurate when rounded up to two significant figures. The graph that was obtained from the results contrasted accurately with the predicted graph and gave a similar rate of reaction.
As the experiment produced good clear results that did not stray from the line of best fit by to big a margin. The procedure was good because the experiment could be repeated easily and efficiently whilst not affecting the experiment. When these results were plotted on a graph, it was clear that the stronger the mole of an acid, the quicker there reaction was. I found the procedure to be good and efficient except for the first experiment. In this experiment the equipment and procedure failed as the gas was not being correctly collected.
- How could the experiment be improved
The experiment could always be improved to a stage were the equipment was highly professional. The equipment could be changed to which no gas could escape. The magnesium could always be measured to a more approximate point and the volume and concentrations of acid could have been better manufactured then using a measuring cylinder. It could have been supplied already prepared by a chemical factory.
For the gas syringe, it could be improved if I was to place some Vaseline in the syringe to allow it to move freely, to stop it from getting stuck.
The air pressure and temperature could be examined in more depth throughout the experiment to make certain that all factors were kept the same, but this would not make a noticeable change. You would use a monometer to check pressure and a thermometer to measure temperature.
From information obtained I know that if magnesium is left in the air it reacts with oxygen to form magnesium oxide, (silver coat,) this is how we received it. This would cause some magnesium strips to react faster than others. The magnesium in these cases could be cleaned with sandpaper to stop the acid having to eat through this coat.
Added factors to be considered: To widen the experiments range of results for concentration, I could have used all reactive metals in-group 2. With these metals I could have seen if my theory for double the concentration would double the rate of reaction and would have given me the same results. This experiment would also prove the effect of different concentrations on metals. In cohesion with this I could have broadened the range of concentrations used, to gain a better knowledge of the effect of different concentrations.
