To see how the concentration of acid, reacting with potassium carbonate, affects the rate of reaction
Aim: To see how the concentration of acid, reacting with potassium carbonate, affects the rate of reaction.
Intro:
This is the reaction I am using in my coursework:
2HCl + K2CO3 CO2 + 2KCL + H2O
In order for substances to react together the particles in the substances must collide with each other and the collision must have enough energy. If there isn't enough energy, no reaction occurs. If there are lots of successful collisions then a lot of CO2 will be produced. The rate of a reaction depends on how many successful collisions there are in a given unit of time.
A reaction can be made to go slower or faster by changing the concentration of a reactant.
Acid particle Water molecule Potassium carbonate tablet
1 2
In dilute acid, there are not so many acid particles (see diagram 1). This means there is not much chance of acid particles hitting a potassium carbonate particle.
In a more concentrated solution of acid, there are more acid particles (see diagram 2). There is now more chance of a successful collision occurring.
Concentration is how much of a substance there is in a certain volume and is measured in Moles per litre of solution (M). The concentration of a solution is the amount of solute, in grams or Moles that is dissolved in a litre of solution.
That is what my coursework is mainly about.
I predict that on my graphs for the results I get directly from the experiment, the curves will bend from a vertical line to a horizontal line and the faster the reaction the steeper the curve will start and the quicker it will reach a horizontal line. This is because as the reaction starts it is very quick and as the chemicals continue to react the reaction produces less CO2 per 10 seconds so it slows down gradually resulting in a curve, This is because after some time there are fewer acid and potassium carbonate particles so the reaction slows down.
I predict that the reaction will go slower when the concentration of the acid decreases. This is because the rate of reaction increases with the concentration. If the concentration of the acid is increased the reaction goes faster. In dilute acid there are not so many acid particles, so the chance of an acid particle colliding with a potassium carbonate particle is not very high.
To go with this piece of research there is another prediction I have come up with: I also predict that when the concentration is halved the rate of reaction will halve as well.
Method:
In this experiment I (by myself) am reacting potassium carbonate and hydrochloric acid together and collect the gas given off called carbon dioxide (CO2). To measure the rate of reaction I will collect the carbon dioxide produced and record the time taken every 10 seconds. This will tell me the speed of the reaction because, e.g. if a small amount of gas is produced in a long time then I will know it is slow reaction.
During the different times of the week when I will be carrying out this experiment the air temperature might change rapidly because it is summer and hot days vary a lot from cooler days. To make this fair I will attempt to cool the acid down before carrying out my experiment if I think the temperature is significantly higher by keeping the acid in an ice bath before the experiment. I will not cool the acid down while carrying out the reaction because if the reaction generates or loses its heat it will be different to cooler days because the acid won't be cooled as it is reacting.
I will drop a circular tablet of potassium carbonate (which are all the same in all the tests so there is no need to measure surface area or mass) into a conical flask containing the acid and the CO2 produced will be pushed through a delivery tube and bubble into an upside down measuring cylinder therefore allowing me to record the amount of CO2 produced every 10 seconds.
To change the concentration of the hydrochloric acid in the conical flask, in each case I will use 20cm3 acid and add certain amounts of tap water each time. ...
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I will drop a circular tablet of potassium carbonate (which are all the same in all the tests so there is no need to measure surface area or mass) into a conical flask containing the acid and the CO2 produced will be pushed through a delivery tube and bubble into an upside down measuring cylinder therefore allowing me to record the amount of CO2 produced every 10 seconds.
To change the concentration of the hydrochloric acid in the conical flask, in each case I will use 20cm3 acid and add certain amounts of tap water each time. This will give me different concentrations of the acid.
I will do a trial experiment to see if it works successfully because when I come to the experiment that counts I will know what I am doing and I will not make as many mistakes.
In this practical I might make a mistake, even if the mistake is small e.g. spilling a couple of drops of water instead of putting them in the conical flask. I will take it into account when I am writing the conclusions and results and take care when carrying out the experiment to be as accurate as possible.
I will be setting up the experiment in the same way each time so to keep it a fair experiment (see diagram).
Diagram:
As the hydrochloric acid and the potassium carbonate react, carbon dioxide is produced. The gas is collected in an upside-down measuring cylinder (the gas displaces water in the measuring cylinder). This method of collecting a gas is called collecting gas over water.
Equipment:
* Stand
* Base and clamp
* Conical flask
* Bung with a delivery tube
* Large measuring cylinder
* Large plastic beaker
* Small measuring cylinder (measuring out water and acid)
* Ice
* Stop clock
* Paper and pen
In some cases I will repeat the experiment if needed. I will be testing the following mixtures of hydrochloric acid and water:
Vol. of 2M hydrochloric acid (cm3)
Volume of distilled water (cm3)
Equation for working out the mole concentration
Concentration of solution (M)
20
0
(2 x 20) ÷ 20
2.00
20
5
(2 x 20) ÷ 25
.60
20
5
(2 x 20) ÷ 35
.14
20
20
(2 x 20) ÷ 40
.00
20
40
(2 x 20) ÷ 60
0.67
20
60
(2 x 20) ÷ 80
0.50
20
80
(2 x 20) ÷ 100
0.40
20
00
(2 x 20) ÷ 120
0.33
I have chose this range of readings because for a start it all fit in the same sized conical flask and because I wanted to get quite a lot of readings so I can make a graph of the gradients from each graph from the tests.
Once I have got the results I will make graphs of each set of results. This then allows me to obtain the gradient from each curve. Having the amount of CO2 collected in cm3 on the y-axis and the time in seconds on the x-axis.
The variables I must keep the same are ones that I should not change because it would change the results and would be an unfair test. Some of these are:
* The temperature
* The amount of potassium carbonate
* The amount of acid
* The water in the measuring cylinder used for collecting gas, the same. (This means I must change it every time because the CO2 would dissolve in the water, so between results I must refresh the water).
Preliminary results:
These are a couple of sets of results from my preliminary tests (tests to find out the best way of doing the actual experiment, both tables are the results where there was 20cm3 of acid and no water, so it has a concentration of 2M
Time (seconds)
CO2 (cm3)
0
5
20
25
30
The measuring cylinder was too small, so I could only measure these results before the gas bubbled under the edge of the measuring cylinder and escaped into the air.
40
-
50
-
Time (seconds)
CO2 (cm3)
0
5
20
22
30
31
40
39.5
50
47.5
60
51.5
70
54
80
56
90
65
00
The large plastic beaker had too much water in, so I could not read the results. This is because the water made the numbers on the measuring cylinder blurred so I could not read them.
10
-
20
-
I am glad I carried out some preliminary tests because they showed that there were problems of how the experiment had been setup.
I corrected the experiment using the mistakes from the preliminary tests: I got a large measuring cylinder instead of a small one and I emptied some of the water out of the large plastic beaker.
Results:
20cm3 of hydrochloric acid with no water added:
Time (seconds)
CO2 (cm3) 1st try
CO2 (cm3) repeat
0
2
20
20
21
30
30
31
40
40
40
47
50
46
58
60
55
66
70
61
70
80
69
82
90
73
86
00
77
88
10
81
91
20
84
93
30
86
95
40
87
96
50
88
97
60
90
98
70
91
98
80
91
98
90
91
98
I carried out a repeat of this experiment just to make sure I was heading in the right direction with my results. They are all more or less the same so it is all right.
20cm3 of hydrochloric acid with 5cm3 water added:
Time (seconds)
CO2 (cm3)
0
3
20
22
30
30
40
37
50
45
60
52
70
57.5
80
63.5
90
69
00
72
10
75
20
79
30
80.5
40
82
50
82.5
60
83
70
83
80
83
90
83
20cm3 of hydrochloric acid with 15cm3 water added:
Time (seconds)
CO2 (cm3)
0
0
20
3
30
6
40
22
50
26
60
30
70
35
80
37
90
42
00
45
10
46
20
46
30
46
40
46
50
46
20cm3 of hydrochloric acid with 20cm3 water added:
Time (seconds)
CO2 (cm3)
0
8
20
4
30
9
40
24
50
29
60
34
70
39
80
44
90
47
00
51
10
55
20
58
30
62
40
66
50
69
60
71
70
73
80
74
90
75
200
76
210
76
220
76
230
76
20cm3 of hydrochloric acid with 40cm3 water added:
Time (seconds)
CO2 (cm3)
0
6
20
2
30
7
40
22
50
29
60
34
70
38
80
42
90
47
00
51
10
55
20
58
30
62
40
64
50
66
60
67
70
68
80
68
90
68
20cm3 of hydrochloric acid with 60cm3 water added:
Time (seconds)
CO2 (cm3)
0
5
20
2
30
8
40
23
50
28
60
33
70
37
80
42
90
47
00
51
10
56
20
60
30
63
40
65
50
67
60
69
70
70
80
70
90
70
20cm3 of hydrochloric acid with 80cm3 water added:
Time (seconds)
CO2 (cm3)
0
4
20
0
30
5
40
9
50
24
60
29
70
34
80
38
90
43
00
46
10
50
20
54
30
56
40
58
50
59
60
60
70
60
80
60
90
60
20cm3 of hydrochloric acid with 100cm3 water added:
Time (seconds)
CO2 (cm3)
0
4
20
8
30
2
40
6
50
20
60
24.5
70
30
80
33
90
36
00
39
10
43
20
45.5
30
49
40
51
50
52.5
60
53
70
54
80
54
90
54
Analysis
I will be making graphs for these results (pg. 7-13) and then I will be making two more graphs. A graph for all the gradients from the other curves and a graph where all the curves are on the same axis.
The first two graphs I have made from my tables are quite accurate. I say this because there is a perfect example of a mistake. On the table where I have checked it (20cm3 of acid and no water), the second set of results (the check), it seems to produce more gas and more quickly too.
The graph that has all the sets of results is quite strange because I expected to find the sets of results to be near a good order (as in 2.00 M to have the steepest curve and 0.33 M have the least steep curve). Also I expected the curves to end at all the same place. I expected this to happen because there is always the same amount of potassium carbonate and hydrochloric acid so there should be the same amount of CO2 produced at the end.
This is a table of all the gradients I have collected:
Concentration of solution (M)
Initial rate of reaction (cm3 per second) (3d.p)
Amount of water added to the acid (cm3)
2.00 1st try
.174
0
2.00 (repeat)
.559
0
.60
.333
5,
.14
.600
5,
.00
0.825
20,
0.67
0.683
40,
0.50
0.750
60,
0.40
0.658
80.
0.33
0.417
00,
From looking at the table my first reaction was to notice that as the concentration decreases and the amount of water added increases, overall the Initial rate of reaction seems to decrease apart from the odd result. To back this point up I have made a graph of concentration against rate of reaction. On the graph there was a curve so the point I made was correct.
I predicted that the reaction would go slower when the concentration of the acid decreases. This is because as the reaction starts the reaction very quick and as the chemicals continue to react and because there are fewer particles in the volume there is a lower chance of a successful collision. So, if the concentration of the acid is increased the reaction goes faster. In dilute acid there are not so many acid particles, so the chance of an acid particle colliding with a potassium carbonate particle is not very high.
Evaluation
My experiment overall had unexpected results, this could have been caused
Most of the predictions I made were correct:
* On the graphs for the individual experiments, the curves looked like what I predicted.
* The reaction did go slower when the concentration decreased.
* But, the prediction of that when the concentration is halved the rate of reaction is halved to cannot be proved correct with my set of results.
I predicted that the curves would bend from a steep line to a horizontal line and the faster the reaction the steeper the curve would start and the quicker it would reach a horizontal line. This is because as the reaction starts it is very quick and as the chemicals continue to react and because there are fewer particles in the volume there is a lower chance of a successful collision, so the reaction produces less CO2 per 10 seconds so it slows down gradually resulting in a curve.
From my graphs I can say that the reaction rate does not double if concentration of acid doubles. I had predicted this because there are double the amount of acid particles so the rate of collision may double. The actual increase is not clear from my graphs, but I can say that it is not double. I think that this because the more concentrated the acid the more gas is produced therefore allowing less collisions of particles. Maybe if I tried a gain I could use a different acid or/and use a different solid to add to the acid to see if my theory that the gas produced gets in the way of the collisions is true.
There are quite a few reasons that this could have happened:
* The measuring cylinder was at an angle,
* I may have added the wrong amount of hydrochloric acid.
* Some of the gas dissolved in the water so not giving me accurate results.
* It's hard to start the timer and drop the tablet in the acid at the same time.
The reason I think would be most likely to have happened is the fact that I didn't read the measuring cylinder correctly. This is because I must not have read it incorrectly as such, just made the mistake of, in each case, having the measuring cylinder at different angles, because I did not hold the measuring cylinder, just let it lean against the side. This is a mistake I would correct next time by maybe having another clamp to hold the measuring cylinder upright, making it a fair test. To solve the problem I would maybe use a gas syringe or I would use a set of better apparatus so the measuring cylinder would stand vertically so it would be easier to read.
This is a problem because in the first attempt some of the CO2 dissolved in the water. This will have meant the gas doesn't get collected overall plus the reaction looks slow at the start because the gas is too busy dissolving to go into the measuring cylinder. This means that in the second attempt the gas did not dissolve so these problems did not occur, so the results were different. The reason for this is that - Before I started the second try I forgot to empty the water from the large plastic beaker and the large measuring cylinder. To cure this problem, next time I would consider using a gas syringe so there is no water for the gas to dissolve in.
When I started each individual test there was always a couple of second between starting the timer and dropping the tablet in the acid. This is quite a big problem because in that time you could miss an amount of gas and so changing the start in each of my results. To rid the tests of this problem I would get some help from a classmate so the time difference from dropping the tablet in to starting the timer is reduced significantly.
Bibliography:
* GCSE chemistry by Gallagher and Ingram
* Heinemann coordinated science foundation Chemistry
* Usborne dictionary of chemistry
Andy Sharpe
