The index ‘a’ is called the order of the reaction in respect to the reactant A; ‘b’ is the order of the reaction with respect to B, and so on. The sum of the indices is called the overall order of the reaction. The orders of the reaction are experimental quantities they cannot be deduced from the chemical equation for the reaction.
The possible orders of a reaction
Consider the rate equation in the form:
r = k[A]
If a = 0 (zero order) the graph will be a straight line since the rate of the reaction is the gradient of the graph of the concentration against time, which is constant and therefore a is zero. In such cases the reaction precedes at the same rate, whatever the concentration, until there is no reactant left. Then the reaction stops.
r = k[A] is the same as r = k
If a = 1 (first order) the graph will be a curve such that the time it takes for the concentration of the reactant to be halved (the half-life) is constant whatever value of concentration you start from.
If a = 2 (second order) the graph will also be a curve but as a second-order curve is much deeper than a first-order one, the half-life is not constant but will increase dramatically as the reaction proceeds.
REACTION MECHANISMS
An organic reaction can occur in a number of successive steps. These steps are known as the mechanism of the reaction. It is very unlikely that these steps will take place at the same rate, knowing this, it is clear that the whole reaction goes at the rate of the slowest of the steps in the mechanism. The slowest step is called the rate-determining step.
EXPERIMENT 1
In this experiment I will be using magnesium metal and hydrochloric acid of varying concentrations. I intend to device a method to measure the rate of the reaction in order to determine the effect of different concentrations on the order of the reaction.
EXPERIMENT 2
In this experiment I will be using magnesium metal, hydrochloric acid (strong acid) and ethanoic acid (weak acid). I intend to device a method to measure the rate of the reaction at different temperatures in order to calculate the activation energy.
PREDICTION EXPERIMENT 2
When using the weak acid the rate of the reaction will be slower than when using the strong acid. This is because there will be fewer molecules to the same volume and so they will be less likely to start colliding and gain enough energy to overcome the activation energy and collide with enough force to start reacting. As a result the reaction is slower.
POSSIBLE METHODS FOR MEASURING THE RATE OF A REACTION
The rate of the reaction could be obtained by measuring the amount of gas collected at certain time intervals until the reaction is finished or by measuring the amount of time taken to produce a certain amount of gas such as 20 or 30cm³. This would give the initial rate.
Another way to obtain the rate of a reaction would be to react the magnesium with the HCl and measure the amount of time taken for all of the magnesium used to react.
POSSIBLE METHODS TO FOLLOW A REACTION AND INVESTIGATE ITS ORDER
Titration-
A reaction mixture is made up and samples are withdrawn from it, using a pipette. Some means is then found of ‘quenching the reaction – slowing it abruptly at a measured time from the start of the reaction, perhaps by a rapid cooling in ice or by removing the catalyst. The samples can then be titrated in some way depending on what is in the reaction mixture.
Measurements of electrical conductivity
If the total number of ions in a solution changes during a reaction it may be possible to follow the reaction by measuring the changes in the electrical conductivity of the solution, using a conductivity meter. This uses an alternating current so that electrolysis of the solution is avoided.
Measuring any other property which shows significant change
Possibilities include pH and chirality.
PRELIMINARY EXPERIMENTS
EXPERIMENT 1- To determine whether the rate of the reaction changes as the concentration is
varied.
METHOD
Measure 10cm³ of 0.5M HCl using a measuring cylinder and pour into conical flask.
Place 2 magnesium strips each 2cm in length and 0.5cm width into the conical flask. Immediately place bung in conical and start the stop clock. Measure the time taken to produce 30 cm³ of gas and record. Repeat for concentrations of 1, 1.3, 1.5 and 2. Do three runs for each different concentration and calculate the average.
METHOD FOR OBTAINING DIFFERENT CONCENTRATIONS OF ACID
USING 1M HCl
For 0.1M – In a 10 cm³ measuring cylinder place 1 cm³ of acid then add pure water up to 10 cm³.
For 0.3M - In a 10 cm³ measuring cylinder place 3 cm³ of acid then add pure water up to 10 cm³.
For 0.5M - In a 10 cm³ measuring cylinder place 5 cm³ of acid then add pure water up to 10 cm³.
Using 2M HCl
For 1.5M - In a 10 cm³ measuring cylinder place 7.5 cm³ of acid then add pure water up to 10 cm³.
RESULT
I found from the preliminary experiment that the amount of reactants used were too small to produce a sufficient amount of gas to measure an accurate initial rate. The concentrations of 0.1M and 0.3M were too low to produce a sufficient amount of gas. To improve the final experiment I will use concentrations ranging from 0.5M – 2M HCl I will also use three magnesium strips to ensure a sufficient reaction occurs and enough gas is produced to measure the initial rate.
EXPERIMENT 2
METHOD
Measure 10 cm³ of HCl using a measuring cylinder. Pour into a boiling tube and place in the water bath.
Measure the temperature of the water heat if necessary to 20ºC. Measure the temp of the acid to ensure it is the same.
Once correct temperature is reached add 2 magnesium strips to the acid and immediately start the stop clock.
Measure the time taken for all the magnesium to react. Repeat for temperatures of 30, 40, 50 and 60ºC and again using CH CO H for each of the temperatures. Do three runs for each different temperature and calculate an average.
RESULT
The preliminary experiment has suggested that the quantities of reactant must be changed in order to efficiently obtain a set of sufficient results. The reactions using ethanoic acid were very slow, so to improve I will double the amount of acid used I will also use 3 magnesium strips instead of 2. I will also change the method used. Instead of measuring the time taken for all of the magnesium to react, I will measure the time taken for the reaction to produce 30 cm³ of gas. This will obtain an initial rate and I believe this will be less subjective and therefore a more accurate set of results.
Using my preliminary findings I have altered the original experiments and devised the following, practical procedures
EXPERIMENT 1
APPARATUS
DIAGRAM
METHOD
Measure 10cm³ of 0.1M HCl using a measuring cylinder and pour into conical flask.
Place 3 magnesium strips each 2cm in length and 0.5cm width into the conical flask. Immediately place bung in conical and start the stop clock. Measure the time taken to produce 30 cm³ of gas and record. Repeat for concentrations of 0.3, 0.5, 1.0 and 1.5M. Do three runs for each different concentration and calculate the average.
METHOD FOR OBTAINING DIFFERENT CONCENTRATIONS OF ACID
USING 2M HCl
For 0.5M – In a 10 cm³ measuring cylinder place 2.5 cm³ of acid then add pure water up to 10 cm³.
For 1M - In a 10 cm³ measuring cylinder place 5 cm³ of acid then add pure water up to 10 cm³.
For 1.3M - In a 10 cm³ measuring cylinder place 6.5 cm³ of acid then add pure water up to 10 cm³.
For 1.5M - In a 10 cm³ measuring cylinder place 7.5 cm³ of acid then add pure water up to 10 cm³.
RESULTS TABLE
EXPERIMENT 2
APPARATUS
DIAGRAM
METHOD
Measure 20 cm³ of HCl using a measuring cylinder. Pour into a boiling tube and place in the water bath.
Measure the temperature of the water heat if necessary to 20ºC. Measure the temp of the acid to ensure it is the same.
Once correct temperature is reached add 3 magnesium strips to the acid and immediately seal with bung and start the stop clock.
Measure the time taken to produce 30 cm³ of gas and record. Repeat for temperatures of 30, 40, 50 and 60ºC and again using CH CO H for each of the temperatures. Do three runs for each different temperature and calculate an average.
RESULTS TABLE
USING HCl
USING ETHANOIC ACID
EVALUATION AND CONCLUSIONS
EXPERIMENT 1
The results indicate that when the concentration of the HCl was doubled from 0.5-0.1M, the rate roughly halved. When the concentration was increased from 1.0-2.0M, the rate was roughly divided by eight. These results suggest that the order of the reaction increases as the concentration increases.
EXPERIMENT 2
The results suggest that when a stronger acid was reacted with magnesium the activation energy was lower than when a weaker one was used.
The results of the previous experiments may have some inaccuracies due to errors during experimental procedures. In experiments 1 and 2, a measuring cylinder inverted in water was used to measure the amount of gas given off over a certain amount of time. This was a fairly accurate procedure, however it is quite subjective, which was possibly the main source of error in measurements, throughout the investigation. In experiment 2 the temperature of the solution surrounding the test tubes containing the reactants, was heated to different temperatures in order to make comparisons of activation energy. Although the temperatures were measured several times and averages were recorded, there is a strong possibility that the temperature did not remain constant throughout the experiments, this may have caused slight inaccuracies in the final results, however this would probably have effected each run in a similar way and so have little effect on the final outcome.
The resolutions of the apparatus used play a part in how accurate the final results are. For instance the inverted measuring cylinders only measure to 0.1cm³, limiting their accuracy to this range. The thermometer used in experiment 2 only measured to the nearest degree, again limiting the accuracy to this range.
Despite some experimental errors affecting the final results, I would say they were fairly reliable, as repeats were taken and averages recorded. Also, the results seem to indicate that the initial predictions based on scientific knowledge were correct, therefore suggesting that any errors made didn’t have a major effect on the final results