Although the rate of reaction depends on the concentrations of A and B, you can’t say that the rate of reaction is proportion to the concentration of A and proportional to the concentration of B. The relationship is;
Reaction Rate ∝ [A] [B] = k[A] [B] (Physical Chemistry)
This is called a rate equation. m and n are integers, 0, 1 or 2. This is the order of the reaction, m is the order with respect to A and n is the order with respect to B. The overall order is m+n, k is the rate constant. “The rate constant for a reaction relates to the concentrations of the reactants.” (Physical Chemistry)
To find out the rate equation you need to know the rate of reaction. In this experiment rate will be measured by the time taken for the reaction. Once you know the rate of reaction you can work out the order of reaction.
Orders of reaction:
Zero Order. Rate = k[A]º
When the order is 0, the concentration of A doesn’t affect the rate.
First order. Rate = k[A]¹
Second order. Rate = k[A]²
By doing experiments involving the reaction between HCL and Na2S2O3 you can find the way the rate of reaction was related to the concentrations of HCL and Na2S2O3.
“The concentrations of HCL and Na2S2O3 have to be raised to some power to show how they affect the rate of reaction. These powers are the order of reaction with respect to HCL and Na2S2O3” (www.chemguide.co.uk)
Apparatus
Beakers – to mix solutions in.
Labels – so you can distinguish between beakers
Measuring Cylinders – to measure out volumes of water and Sodium Thiosulphate.
Pipette – to measure volume of Hydrochloric acid. As this is used in smaller amounts it needs to be measured more accurately.
Cross on a piece of paper – to know when the reaction is complete.
Stopwatch – to time the reaction
Water – to slow the reaction.
Variables
The variables that need to be controlled are:
Temperature – the temperature needs to be kept constant for the experiment to be fair. The temperature can be kept constant by doing all the reactions on the same day.
The amount of solution – Water will be used to make sure there is the same volume of liquid for every reaction.
Method 1
- Measure out volumes of water and Sodium Thiosulphate using amounts shown in table;
Use separate measuring cylinders for the Sodium Thiosulphate and water.
- Put the sodium Thiosulphate and Water in a beaker. Label the beaker.
- Measure out 5cm³ of HCl using the pipette.
- Put the beaker on top of the cross.
- Add the HCl, start the stopwatch when Half the HCL is left in the pipette.
- Stop the stopwatch when the cross is no longer visible.
- Record the results.
- Repeat for the other concentrations.
Method 2
- Measure out volumes of water, Hydrochloric acid and Sodium Thiosulphate using amounts shown in table;
- Put Sodium Thiosulphate and water in a beaker. Label the beaker.
- Put the beaker on the cross
- Add the HCl, start stopwatch when half the HCl is left in the pipette.
- Stop stopwatch when cross is no longer visible.
- Record results.
- Repeat for the other concentrations.
Results
1.
2.
Analysis of Results
The results show that, as predicted, the rate of reaction does increase with concentration. This is because of the collision theory – for a reaction to happen particles must collide. Increasing the concentration increases the number of particles. This increases the number of collisions per second, therefore increasing the rate of reaction.
Graph 1 has a straight line going through the origin showing that this is a first order reaction. There are no obviously wrong results and none are anomalous as they all follow the trend.
Graph 2 also shows that the order of reaction is between 1 and 2, (roughly ¼) with the rate increasing with concentration. All the results follow the trend, so none are anomalous, but results 1 and 3 are quite far from the line of best fit showing there is some error.
This means that, if we take the order of reaction 2 to be ¼, the rate equation is:
Rate=k[HCl] [Na2S2O3]¹[H2O]º
Or Rate=k[HCl] [Na2S2O3]
This means we can also work out the rate constant (k).
For the first set of result, experiment 1:
Rate=k[HCl] [Na2S2O3]
0.091=k[5] [50]
0.091=k74.767
k=0.091/74.767
k=0.0012
For second set of results, experiment 1:
Rate=k[HCl] [Na2S2O3]
0.074=k[5] [40]
0.074=k59.814
k=0.074/59.814
k=0.0012
For second set of results, experiment 2:
Rate=k[HCl] [Na2S2O3]
Rate=k[4] [50]
0.086=k70.712
k=0.086/70.712
k=0.0012
For fourth set of results, experiment 2:
Rate=k[HCl] [Na2S2O3]
0.071=k[2] [50]
0.071=k59.460
k=0.071/59.460
k=0.0012
This shows that the rate constant for this reaction is 0.0012
Evaluation
Overall this experiment was fairly accurate and gave the expected results, but some results were inaccurate due to errors in the experimental method. The main error was deciding when the reaction had finished. This was done by placing the beaker on a piece of paper with a cross drawn on it. When the cross disappeared, the clock was stopped.
