Safety
In order to make the experiment safe goggles and lab coats must be worn at all times as a protection in case of a spillage.
Clearing away chemicals and equipment
Action required for anticipated accidents
Fair testing
The independent variable in my investigation will be the surface area of the Marble Chips, however their mass will need to be controlled as well as the amount of hydrochloric acid and its concentration (2moles) used in all experiments in order to make it a fair test. The amount of gas produced will be measured (by using scales) as a dependent variable. The temperature will be kept constant as the room temperature.
Analysing the results.
From the obtained results I was able to work out the rate of reaction in all 3 test (12 experiments), they are as followed:
Test 1:
- Marble Chip powder took 2 minutes to produce 0.45g of gas in total; therefore the rate of reaction was 0.45 ¸ 2 = 0.23g per minute (to 2 decimal places)
- 2-6mm Marble Chips took 9 minutes 20 seconds to produce 0.48g of gas in total; therefore the rate of reaction was 0.48 ¸ 9.3 = 0.05g per minute (to 2 decimal places)
- 6-9mm Marble Chips took 10 minutes to produce 0.46g of gas in total; therefore the rate of reaction was 0.46 ¸ 10 = 0.05g per minute (to 2 decimal place)
- 9-12mm Marble Chips took 9 minutes 40 seconds to produce 0.45g of gas in total; therefore the rate of reaction was 0.45 ¸ 9.7 = 0.05g per minute (to 2 decimal places)
Test 2:
- Marble Chip powder took 1 minute 40 seconds to produce 0.47g of gas in total; therefore the rate of reaction was 0.47 ¸ 1.7 = 0.27g per minute (to 2 decimal places)
- 2-6mm Marble Chips took 8 minutes 40 seconds to produce 0.46g of gas in total; therefore the rate of reaction was 0.46 ¸ 8.7 = 0.05g per minute (to 2 decimal places)
- 6-9mm Marble Chips took 9 minutes 40 seconds to produce 0.46g of gas in total; therefore the rate of reaction was 0.46 ¸ 9.7 = 0.05g per minute (to 2 decimal place)
- 9-12mm Marble Chips took 10 minutes to produce 0.45g of gas in total; therefore the rate of reaction was 0.45 ¸ 10 = 0.05g per minute (to 2 decimal places)
Test 3:
- Marble Chip powder took 3 minute 40 seconds to produce 0.42g of gas in total; therefore the rate of reaction was 0.42 ¸ 3 = 0.14g per minute (to 2 decimal places)
- 2-6mm Marble Chips took 9 minutes 40 seconds to produce 0.48g of gas in total; therefore the rate of reaction was 0.48 ¸ 9.7 = 0.05g per minute (to 2 decimal places)
- 6-9mm Marble Chips took 10 minutes to produce 0.43g of gas in total; therefore the rate of reaction was 0.43 ¸ 10 = 0.04g per minute (to 2 decimal place)
- 9-12mm Marble Chips took 10 minutes to produce 0.43g of gas in total; therefore the rate of reaction was 0.43 ¸ 10 = 0.04g per minute (to 2 decimal places)
I have chosen to repeat the experiment 3 times in order to allows me to calculate an average rate of reaction. This increased the accuracy of the obtained results.
From the table of results and the graphs I can deduce that by increasing the surface area the rate at which the gas was produced also increased. The line which represented the amount of gas given off on the graph was steeper in the test using powder therefore the gas was produced at a faster rate. Roughly the same amount of gas was produced in all tests. Therefore my hypothesis was correct.
Evaluation.
As I repeated the experiment 3 times and worked out the average rate of reaction in each test, I think my results are quite reliable. However, I could make further improvement by increasing the number of tests and taking the results with a smaller gap such as 10 seconds instead of 20. I could also increase the amount of Calcium Carbonate to more than 2g and keep the amount of acid as 50ml; which would slow down the reaction and will allow me to take more accurate results, as it would decrease the number of anomalous results.
Conclusion.
Increasing the surface area increases the chance of the particles hitting the reactant (colliding) and allows more particles to react with the reactant all at once. The bigger the surface area of Calcium Carbonate, the greater the number of collisions, Carbon Dioxide was given off at a quicker rate. Therefore the rate of the reaction slowed down as the size of the particles was increased from powder to 9-12mm pieces.
The theory of how reactions happen
Reactions can only happen when the reactant particles collide, but most collisions are not successful in forming product molecules. The reactant molecules must collide with enough energy to break the original bonds so those new bonds in the product molecules can be formed. All the rate-controlling factors are to do with the frequency of reactant particle collision. In the case of temperature, the energy of the collision is even more important than the frequency effect. The particle theory of gases and liquids and the diagrams below will help you understand what is going on.
The effect of Concentration
If the concentration of any reactant in a solution is increased, the rate of reaction is increased. Increasing the concentration, increases the probability of a collision between reactant particles because there are more of them in the same volume. Examples …..
Increasing the concentration of acid molecules increases the frequency at which they hit the surface of marble chips to dissolve them (slower ➔ faster).
Increasing the concentration of reactant A or B will increase the chance of collision between them and increase the speed of product formation (slower ➔ faster). The effect of Surface Area
If a solid reactant or a solid catalyst is broken down into smaller pieces the rate of reaction increases. The speed increase happens because smaller pieces of the same mass of solid have a greater surface area compared to larger pieces of the solid. Therefore, there is more chance that a reactant particle will hit the solid surface and react. The diagrams below illustrate the acid – marble chip reaction, but they could also represent a solid catalyst in a solution of reactants.
The effect of Temperature
When gases or liquids are heated the particles gain kinetic energy and move faster (see diagrams below). The increased speed increases the chance of collision between reactant molecules and the rate increases. However this is not the main reason for the increased reaction speed.
➔
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.
Results