Chemistry - Rates of Reaction.

Science Coursework:

Chemistry

Science Coursework; Chemistry

Rates of Reaction

Aim and background information:

The purpose of this study was to investigate the catalysts of reaction rate between substances; the elements used to test this were chips of calcium carbonate (CaCO3) and hydrochloric acid (HCl) of varying molarity. In order to complete the task I first had to examine and comprehend the basic principles that form reaction rates. Primarily, two theories must be studied, these are the Collision theory and the principles of Activation energy.

Collision Theory:-

The collision theory states that in order for a reaction to occur between two species they must come into contact with each other. They must also collide the right way and with enough energy for the bonds to break (this is known as the Activation energy which is examined below). The more particles that are required to obtain a certain reaction, the less likely the chances are that it will occur. This is mainly because most reactions necessitate a specific collision between two particles (eg. a certain molecule of one species colliding with a certain bond of the other), so the chances of reactions occurring that involve more than two species become very remote.

Activation energy:-

Even when the particles are orientated so that a reaction can occur this still will not happen unless there is a minimum amount of energy, the Activation energy. If the particles collide with less energy than the activation energy then nothing happens and they will bounce apart. Any chemical reaction involves the creation and destruction of atomic bonds, a process that releases and requires energy respectively. Where collisions are relatively gentle (or when another factor effects the collision) there isn’t enough energy available to start the bond-breaking process so no reaction occurs.

Reaction rate is something that can be affected by a multitude of factors, these are identified and explained below:

· Concentration of substance - the concentration of a substance (eg. the molarity of the HCl in my experiments) will have a substantial effect upon the rate of the reaction. A higher concentration means a higher multitude of particles, this makes the collision between particles more frequent and consequently hastens the rate of reaction.

· Surface area of solid reactants - a larger surface area of a substance increases the amount of particles exposed to collisions, this factor is normally only applicable in reactions involving solid elements.

· Temperature of substance - if the elements of a reaction are heated then the particles that compose them have a higher level of kinetic energy, they will move about faster. This means that collisions are far more likely to produce the required activation energy for successful reactions.

· Catalyst within reaction - a catalyst within a reaction will decrease the required activation energy for a reaction so more particles will be able to have successful collisions.

· Light intensity - light energy that is absorbed by reactants is transferred into chemical energy thus increasing the amount of particles that have the activation energy for a reaction to occur.

This is a preview of the whole essay

· Catalyst within reaction - a catalyst within a reaction will decrease the required activation energy for a reaction so more particles will be able to have successful collisions.

· Light intensity - light energy that is absorbed by reactants is transferred into chemical energy thus increasing the amount of particles that have the activation energy for a reaction to occur.

Reaction rates are important and can effect our daily lives a great deal. For instance, if you were to leave a dairy product out of a refrigerator for a given number of days and then compare it with an identical product that had been kept in the refrigerator, you would witness an example of faster reaction rates due to an increase in temperature.

My coursework will focus on the concentration of the reactants and its resultant effects on the reaction rate; in this case the variable will be the molarity of the hydrochloric acid.

Apparatus:

The following equipment will be needed during the experiment:

∙ 1 burette, capacity of 50cm³

∙ 2 stands

∙ 1 large beaker

∙ 1 test tube

∙ 1 bung with delivery tube

∙ Stopwatch

∙ Digital scales (to 0.01g accuracy)

Chemicals/substances needed:

∙ Hydrochloric acid (2 molar, this will be diluted appropriately to vary the concentration)

∙ 5 chips of calcium carbonate weighing at or around 0.4g

∙ Tap water

Method:

In order to investigate the reaction rates my colleagues and I must follow a set experiment that will determine the results. First the burette will be filled with 50cm³ of tap temperature water, this will be turned upside down and placed in the large beaker (also containing water) and held in place by one of the stands. This will stop water being lost and make it easier for the experiment to function. The hydrochloric acid will then be prepared, with respect to all safety precautions. As varying molarities of the HCl are not readily available, 25ml of 2 molar hydrochloric acid will have to be diluted in order to obtain the 1, 0.5 and 0.25 molar concentrations needed to complete the experiment. Five chips of calcium carbonate that are relatively the same surface area and weight will be prepared, one is then placed in the test tube which will be fixed into position next to the beaker. The 25ml HCl acid solution will be added to the test tube, the bung will then be fastened on top and the delivery tube of the bung will be submerged into the beaker and positioned under the opening of the burette. Now as hydrogen gas is released from the reaction it will travel through the delivery tube and form bubbles that will ascend the burette, the water will be displaced and we can measure the reading on the burette to obtain an insight into the reaction rate. This will be timed for ten minutes and recorded every 30 seconds. A preliminary test will be carried out prior to the experiment in order to practice and develop a technique for the experiment. The results for the preliminary and pursuing tests are shown in Fig. 1 of the results section.

Safety:

So as to maintain a safe working environment the following precautions will be respected:

∙ Eye protection will be worn throughout

∙ Long hair and loose clothing will be fastened

∙ Care will be taken when handling all apparatus and substances

∙ Chemicals will be placed on protective mats to avoid damage

Constant factors:

In order to ensure a fair test and accurate results the following factors must remain constant:

∙ The weight of the calcium carbonate must remain constant or relatively similar, as must the surface area of the CaCO3.

∙ The light intensity of the environment must remain constant, this will be lab conditional, preferably on the same day so the natural light intensity is the same.

∙ The temperature of the substances will be room determined at all times, no species will be heated or cooled.

∙ The removal of all catalysts must be ensured this involves the sanitation of all equipment after each test.

∙ The amount of acid (regardless of molarity) used in the experiments must remain constant.

∙ The time allocated for measurement and the point at which timing commences must remain constant.

Notes and information on ‘molarity’:

A molar substance is one with a certain amount of particles in. 1 molar of any substance has the same amount of particles as any other 1 molar solution and so on. A 1 molar solution has 6.3x 1023 particles. The molar principle was devised to give a measurement that could be used to compare against others. If I wanted to make a 1 molar solution of Hydrochloric acid I would place 36ml of the acid into 1 litre of water, because the chemical formula of Hydrochloric acid is HCl, which equates to 35 Cl and 1 H. If, as I am in this test, I were wanting to derive lesser molars of HCl from a higher molar, in this case 2, I would simply dilute the substance accordingly; for 25ml of 1 molar acid I would add 12.5ml of water to 12.5ml of 2molar HCl.

Prediction:

Founded on the information above and using a basic logic I am able to make the prediction that as the molarity of the HCl acid halves so will the speed of the rate of reaction. This is because there will be half the amount of hydrochloric acid particles present in the mixture than before, so there will be half as many successful collisions with the calcium carbonate chips. Therefore, when the concentration is doubled the speed of the reaction should also double.

Results:

Fig. 1

Fig. 2

Fig. 4

Fig. 5

Analysis:

When the experiment was complete a table was formulated to portray the results clearly, this can be seen in Fig. 1. Using this information a second table, Fig. 2, showing the reliability was formulated and a line graph, Fig. 3, was constructed. On the line graph a gradient on the 4-minute line was measured in order to calculate the rate of reaction for each molarity. These were recorded in to Fig. 4 and compared with the actual results. The reaction rates were then transformed into a bar chart, Fig. 5. It can be seen from Fig. 5 that my prediction is fundamentally correct; as the molarity halves the rate of reaction basically halves as well, this is with the exception of 0.25 molar which seems to have achieved exceptionally low liberation levels. I think that this value is very inaccurate, and would have preferred to retest the 0.25m acid now that these results are revealed. It would help to have textbook values to compare with. I consider the rate of reaction values to be acceptably accurate, as Fig. 4 portrays the theoretic and actual total volume figures are reasonably close and the inaccuracies can probably be counted for in the ones listed in my evaluation. Obviously I have to bear in mind that the actual total amount of gas liberated by the 2m and 1m acids is not known due to the 50cm³ capacity of the burette, so these can not be compared.

Evaluation:

I feel that this was a successful experiment, the results that have been produced provide a relatively reliable insight into the rate of reactions and my prediction has basically been confirmed. The preliminary test was largely a failure but this was obviously never intended to participate in the conclusion and succeeded in tutoring me upon how to conduct the practical accurately. Unfortunately there are many inaccuracies that form the margin of error within my practical:

∙ The CaCO3 chips surface areas would not have remained constant unless I had used powdered marble or accurately cut blocks, thus demanding a ridiculously low molarity or a ridiculously high amount of calcium carbonate powder.

∙ The time between placing the CaCO3 chip in the acid, fastening the bung on top and positioning the delivery tube under the burette meant an initial loss of hydrogen and an inaccuracy in the results, even though timing only began when the delivery tube was positioned.

∙ The dilution of the 2 molar acid to convert it may have lowered the temperature of the substance when compared to the experiment using the root acid, because of the temperature of the tap water. Having supplies of varying molarities of HCl acid at hand could solve this.

∙ There is obviously a small delay between looking at the stopwatch and reading the burette, this coupled with the possible inaccuracy of the readings themselves increased the margin of error in my experiment.

∙ Contrary to my plan the experiments were carried out over two separate days, this meant a variation in the room temperature and light intensity conditions of the experiment. I would have preferred to complete all the practical elements in one day.

As I commented upon in my analysis it would have been beneficial to have textbook values at hand for this experiment as they would aid me a great deal in calculating my margin of error and reliability of the results. Also mentioned in the analysis was the lack of a larger burette, the 2m HCl acid, and to an extent the 1m, released 50cm³ of hydrogen far quicker than the 10-minute termination time.

To extend upon my investigation I could try the reaction with a more varied array of molarities and at different temperatures. I could then find the optimum temperature for the reactions to take place in. I could also try different acids such as sulphuric (H2SO4) and nitric acid (HNO3); I could then compare the results and ascertain which acid reacts more quickly with the calcium carbonate.

Daniel Murphy