Background:
Reactions happen at many different speeds, explosions are extremely fast reactions. Other reactions are slow, like iron rusting.
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 that has to be 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 percentage 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 un-powdered solid. Only particles on the surface of the solid will be able to undergo collisions with the particles in a solution or gas.
The reaction taking place will be:
Sodium Thiosulphate and Hydrochloric acid > Sodium Chloride + Water + Sulphur + Sulphur Oxide.
Method:
Firstly A Large X shape is drawn on a piece of paper and placed on the table
5cm of Hydrochloric Acid (at concentration 1 mol./dm3) and 15 cm of sodium thiosulphate 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 (30C to 70C) 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 at different temperatures.
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.
Results:
°C Test 1 Test 2
25 58 68
35 32 32
45 24 19
55 12 12
65 08 10
75 05 05
From these results I can clearly see a pattern emerging. As I predicted, when the temperature of the Sodium Thiosulphate is increased the rate of the reaction increases. The graph shows this pattern taking place as well. The energy given to power the reaction increases as the Sodium Thiosulphate is heated, causing the reaction to take place faster and the sulphur to precipitate colouring the solution.The end product of this experiment is From these results we now know that this reaction was a chemical reaction, and the details given from the collision theory and the kinetic theory explain why the reaction speeds up, with an increase in temperature. To justify both of these theories there are several other experiments that give similar results e.g. Hydrochloric Acid and Magnesium.
Analysis:
In this experiment I have found that as the temperature increased the time taken for the reaction to take place decreases. This means the rate of reaction increases as it takes less time for a reaction to take place, so more take place per second. In the experiment the time taken for a reaction to take place decreased by roughly 10 to 15 seconds for every 10C increase in temperature, with the one anomaly being the 30C 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 graphs (that plot temperature and concentration against time taken for the reaction to take place) the graph has negative correlation, meaning that as the temperature increased the time taken for the reaction to take place decreases.
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.