There are two in telling if a reaction has speeded up. Firstly it is possible to measure the rate of products formed. Secondly, measure the rate at which the chemicals disappear.
I have some background knowledge on this subject field. I conducted an experiment to investigate the effect of temperature on the rate of reaction.
Equipment:
Tripod Sodium thiosulphate solution
Bunsen burner dilute hydrochloric acid
Conical flask
Thermometer
Stopwatch
Pen and paper
Method:
Firstly heat 50cm³ of sodium thiosulphate solution to the appropriate temperature. Then place the conical flask on the cross and add 5cm³ of the hydrochloric acid and time how long it takes for the cross to disappear. Then record the results in a table. Repeat for each of the temperatures shown in the table below.
From this experiment, I am able to show that as the temperature increases, so does the rate of reaction. At a greater temperature the sodium thiosulphate solution reacted faster with the hydrochloric acid. This is because the greater the temperature, the greater the increase of the speed and frequency at which the particles in the reacting chemicals bump into each other. Hot particles have more energy (kinetic), so they move around more quickly, so they collide more often. The collisions are also harder and more effective.
Prior to the main investigation, I decided to conduct simple experiment using different methods of altering the rates of reaction. In this particular investigation using a catalyst was not applicable. The surface area could be used, however this was fairly difficult to do, and the results would see fairly inaccurate, due to the time and equipment that we were provided with. I conducted an experiment; I changed the temperature of the acid by using a Bunsen burner and went up by intervals of 5°C. 60°C, was the maximum temperature that we could use, due to safety. I filled a conical flask with 50 cm³ of 1M hydrochloric acid and used the Bunsen burner to heat the acid up to the appropriate temperature that is was needed. I attached a delivery tube to a burette, which was placed deep into a bowl of water, in order to measure the volume of hydrogen produced. I placed a 2cm strip of magnesium ribbon into the conical flask and began to time for 30 seconds. I measured the volume of hydrogen gas that evolved. I recorded the results as follows:
From the results shown above, the accuracy of the experiment seems to be fairly poor, as the reaction occurred fairly quickly and using a burette of 0.1 increments meant that the results could be inaccurate and unreliable. It was fairly difficult to keep the temperature constant through each experiment, and by using interval of 5°C; this would mean that the results of the experiment would not be reliable enough.
I set up a similar experiment, but this time I decided to change the concentration of the hydrochloric acid. I would dilute it by water in order to obtain the concentration required. Initially I used 2M hydrochloric acid, but the reaction seemed extremely fast, and would not enable me to record accurate results, as the times were extremely close, and from using my own judgement, using a 2M acid seemed not appropriate. Then, I used 0.5M acid, however, the reaction seemed too slow in order to record fairly accurate results. I used 1M hydrochloric acid. This seemed an adequate acid to use, as the time between the concentrations were large enough in order to obtain accurate and reliable results. I will therefore use 1M acid and go down by 0.1M, until I reach a 0.1M concentrated hydrochloric acid. As well as this I used a 250ml-measuring cylinder to measure the volume of hydrochloric acid and water needed, but this seemed inaccurate as I need to measure precisely fairly low volume and such a measuring cylinder seemed inadequate. I decided to use a 10ml, and 50ml measuring cylinders.
Apparatus:
Retort stands 20 x 20mm pieces of magnesium
Delivery tube 1M hydrochloric acid
Conical flask
Measuring cylinder (10ml and 50ml)
Burette
Stopwatch
Bowl
Burette clamp
Method:
Using the measuring cylinder, measure the 50 cubic centimetres of 1M hydrochloric acid, and put into the conical flask. Fill a large bowl approximately three quarters full. Fill the burette with 50cm ³ of distilled water and inset deep into the bowl. Attach the delivery tube from the conical flask to the lower end of the burette in order for the hydrogen gas to be collected. Secure the burette safely with the clamp onto the stand. Cut 20mm strip of magnesium ribbon (approx., thickness 0.03mm, and width 3.5mm). Insert the strip of magnesium into the conical flask, secure quickly with bung and began to time every 10seconds up to 90 seconds the amount of hydrogen collected in the burette. Repeat the experiment for each concentration starting from 1M, 0.9
0.8,0.7… 0.1M. To calculate and alter the concentration mix the acid with water, as shown below:
Diluted concentration = Volume of acid x Original concentration
Total volume
For example, 45cm³ of 1M acid and 5cm³ of water.
Diluted concentration= 45/5 x 1 = 0.9M
Repeat the experiment a second time to obtain more reliable and accurate results. Then take the average.
The variable that I am going to change is concentration. I predict that the order of the fastest reaction to the least would be from the greatest concentration of 1M to the least concentration of 0.1M. The scientific reason behind my hypothesis is that when increasing the concentration of the hydrochloric acid, the numbers of particles present also increase and so increase the chance of collision. There would be more chance of the particles interacting, and so the reaction will proceed at a faster rate.
Analysis
From the results that I collected it is clear that the concentration with the fastest rate of reaction was that of 1M hydrochloric acid. The slowest reaction was when the concentration was 01.M.
Here is the order of the fastest to slowest reaction:
Fastest 1.0M
0.9M
0.8M
0.7M
0.6M
0.5M
0.4M
0.3M
0.2M
Slowest 0.1M
As shown in the results tables, I collected a first set of results, then a second and then averaged the results in order to obtain fairly accurate and reliable result. I was aiming to do collect a third setoff results but due to the amount of time that was available this could not be completed. My prediction that the highest concentration, 1M, would be the fastest reaction was correct, with a total amount of 22.7cm³ of hydrogen produced. The reaction was extremely fast during the first 30 seconds, however towards the end of that experiment the amount of hydrogen produced stayed the same, as he 20mm strip of magnesium was used up, and therefore no further reaction could occur. The reason why the highest concentration was the fastest reaction is because when increasing the concentration of the hydrochloric acid, the number of particles present also increases and so increases the chance of collision. There would be more chance of the particles interacting, and so the reaction will proceed at a faster rate. The second fastest reaction was when the concentration of the acid was 0.9M, with a total amount of 19.2cm³ of hydrogen produced. This again was a fairly a vigorous reaction, and the hydrogen produced in the first 30seconds was high, but again towards the 90seconds the amount of hydrogen collected was decreasing as the magnesium got used up in that time. The thirst fastest reaction was when the concentration of the hydrochloric acid was 0.8M, as the amount of hydrogen produced in 90 seconds was 13.5cm³. The fourth fastest reaction was when the concentration of the acid was 0.7M, with a total of 10.5cm³ of hydrogen produced. The fifth fastest reaction was when the concentration of the hydrochloric acid was 0.6M, with a total of 8.1 cm³ of hydrogen gas collected. The fifth slowest reaction was when the concentration of the acid was 0.5M, this total to 5.1cm³ of hydrogen produced. The fourth slowest reaction was when the concentration of the acid in the conical flask was 0.4M, this was a total of 20cm³ of hydrochloric acid, diluted with 30cm³ of water, and this resulted in 4.3cm ³ of hydrogen being collected. The third slowest reaction was that of 0.3M concentrated hydrochloric acid. This resulted in a total of 2.8cm³ of hydrogen gas being produced and collected in the burette. The second slowest reaction was when the hydrochloric acid was 0.2M; this resulted in a total of 1.5cm³ of hydrogen gas being produced. The slowest reaction in this investigation was when the concentration of the hydrochloric acid was 0.1M, this was where only 5cm³ of acid was diluted with 45cm³ of distilled water, this resulted in a total amount of 0.5cm³ hydrogen gas being produced. The rates of reactions decreases as the concentration of the hydrochloric acid decrease. As the number of particles present reduce the chance of any collision also decreases, and so there is les chance of particles interacting and so the reaction will proceed at a slower rate. The reaction rate has been multiplied by a thousand in order to plot adequate points on the graph. This can be shown by the table and graph below:
As well as comparing the concentrations and the amount of hydrogen produced throughout the experiment, I can also compare the total of amount of hydrogen gas produced at the end of the experiment (90seconds), in order to compare the reactions when different concentrations of hydrochloric acid are reacted with a 20mm strip of magnesium.
From the graph above it is clear that the concentration that produced the greatest volume of hydrogen in 90 seconds, and therefore enabled the fastest reaction was when the concentration of the hydrochloric acid was 1M, the volume of hydrogen was 22.7 cm³. The least concentration, 0.1M, produced the least volume of hydrogen at 90 seconds, and was therefore the slowest reaction throughout the investigation. It produced a total volume of 0.5cm³.
It is also possible to calculate the average rate of reaction for the period of the experiment. For example, if it takes 90 seconds to collect 22.7 cm³ of gas.
The average rate= 22.7/90
= 0.25cm³/second
From the graph above it is clear that as the concentration of the hydrochloric acid increases, so does the average rate of reaction. For example, when the concentration of the acid is 0.8M, its average rate of reaction is 0.15cm³/second, whereas when the concentration of the acid is 0.7M, its average rate of reaction is 0.12cm³/second.
It is also possible to calculate the gradient, by using the line of best fit. The steeper the slope of the line, the faster the reaction.
For example the gradient for when the concentration of the hydrochloric acid is 1M is
Gradient= Vertical change
Horizontal change
=0.6
For when the concentration of the hydrochloric acid is 0.9M:
Gradient = 0.4
From this example, I can show that there is a steeper slope on the graph for when the concentration is 1M than there is for when the concentration is lower, 0.9M. As the concentration increases so does the gradient of line on the graph. It is also possible to predict, by using the line of best fit what the volume of hydrogen gas would be if the concentration were 0.75M. By using the line of best I can show that it would be approximately 12.6cm³ at 90 seconds.
Evaluation
The aim of this investigation was to investigate one factor, which affects the rate of reaction between magnesium and hydrochloric acid. I chose to do concentration, as it seemed the most adequate one to do after conducting preliminary work.
Accuracy of the measurement was the hardest factor to keep constant because it is almost impossible to get completely accurate results in an experiment like this with the equipment we were provided with. Human experimental error is a problem because factors like reaction times, eyesight and our own judgement cannot be changed and they do affect the end result considerably. Such problems could not be fully controlled with the equipment available but steps were taken to avoid them. As the same person’s vision was used to determine the volume of gas that was produced. Also the same person time, and filled the burette in order to ensure a fair test and produce reliable results. To ensure that this investigation was a fair I ensured that the same length of magnesium ribbon was used, the same volume was used and also the temperature at which the acid was kept was the same throughout the investigation. If this was not done then the accuracy and reliably of the investigation would be fairly poor, as particularly temperature is a factor that affects the rate of reaction. The acid was kept at room temperature. However, a thermometer could have been used in order to ensure the exact temperature was used for each experiment. In all two tests were carried out.
However, the accuracy of this experiment could be improved. Firstly, a more accurate burette with 0.01 increments could have been used as oppose to 0.1. The bung on the conical flask may have let come of the hydrogen gas escape, this could be improved. For example, Vaseline could be applied between the bung and the conical flask in order to limit the amount of gas escaping, producing more reliable results. Also to improve the accuracy and reliability of the investigation more tests could be done, in this experiment 2 tests were carried out for each concentration, but to ensure a more accurate and reliable final result five tests could be done, then from this an average could be taken. I could also use emery paper to clean the magnesium before reacting it with the hydrochloric acid, this would ensure a more reliable result. As well as this, weighing the magnesium ribbon by using a balance may be more accurate than cutting 2cm strips.
There were not any major anomalous results. However, the amount of hydrogen produced at a concentration for 0.4M, at the times of 50 and 60 seconds were exactly the same at 2.4cm³. This could be because, some of the hydrogen gas may have escaped between the bung and the conical flask, rather than going into the burette, or the magnesium strip that reacted with the hydrochloric acid may have been used up and no further reaction could take place, so therefore the amount of hydrogen produced would have stayed the same. Such results appeared frequently in acids that had a higher concentration, as these reacted quickly with magnesium and used up the magnesium before the 90 seconds, and the amount of hydrogen produced stayed the same. This was especially so for the highest concentration in my investigation of 1M. This did not occur in the low concentration, as the magnesium was slowly reacting with the low concentration acid, and the hydrogen produced was steadily increasing.
If I were to repeat this experiment again I would certainly ensure that I would use a substance such as Vaseline to limit the amount of hydrogen escaping. I would also use a more accurate burette, than with 0.01cm³ increments, rather than 0.1, as this would enable me to obtain more accurate results, as the results were fairly close together. As well as this I would ensure than more test were done, for example five as oppose to only two. This would then give me more accurate and reliable results.
I could take this experiment further by using a different metal such as Zinc and aluminium. I could investigate other different factors that affect the rate of reaction such as temperature or surface area. I could use a fixed amount of hydrochloric acid, and could alter the surface area of the magnesium strip. I could also use the other method of measuring the rate of reaction. Measure the rate at which the chemicals disappear, rather than measuring the rate of products formed. However, this may be an unreliable method, with the equipment available, at this would be based on ones own judgement. A more advanced and technological way, could be to use a particular program that will enable me to collect data that will instantly be put onto graph, and will show how changing the concentration affects the reaction rate throughout the investigation. The investigation could also be extended to investigate other factors affecting the rate of reaction such as catalysts, temperature of the acid or particle size of the magnesium. Such ideas could give more evidence, or extend my investigation in this subject field.
