pipette x 2 30cm ruler x 1
Chemicals
hydrochloric acid
magnesium ribbon
distilled water
Safety:
Safety precautions
*wear safety glasses at all times
*take care when handling hydrochloric acid as it is corrosive – wash hands immediately
in case of accidental contamination and after use
Variables:
- The Independent variable will be the concentration of Hydrochloric acid, because it will vary.
- The dependent variable will be the time taken for the rate of reaction to take place (i.e the magnesium to disappear), because it depends upon the concentration of Hydrochloric acid.
Other variables throughout this investigation which will vary are the volume of water used and the volume of hydrochloric acid.
The variables which will remain the same are the temperature (room temp.) will stay the same in order for it to be a fair test, because if the temperature changes it will effect the rate of reaction between the reactants, either by speeding it up if the temperature rises because the particles move faster and travel a greater distance in a given time and so will be involved in more collisions. The amount magnesium used will stay the same (2cm long), so that it is a fair test.
Depending on certain factors the rate that this reaction will take place will either increase or decrease. The factors that may affect the rate of reaction are as follows:
· Temperature of the Hydrochloric Acid
· Mass of the magnesium ribbon used
· Concentration of the Hydrochloric acid
· Surface area of the magnesium ribbon used
All of these factors will change the rate of reaction because of the Collision Theory. This is a theory that is used to predict the rate of a reaction. The Collision Theory is based on the idea that for a chemical reaction to take place, it is necessary for the reacting particles to collide with each other with enough energy to break or form new bonds between the other particles, which is called a successful collision. If when they collide and they do not have enough energy to break or form new bonds then they will simple bounce of each other, causing an unsuccessful collision.
Factors
The factors that could affect the rate of reaction of my experiment are as follows:
· Concentration of acid
This could affect the rate of reaction because the higher the concentration of the acid then the more acid particles per 100cm3 so more collisions per second and then there will be more successful collisions per second.
· Temperature of the acid
If the starting temperature of the acid is different each time the speed at which the acid particles collide with the magnesium ribbon will increase more the higher the temperature goes. This means the acid particles move with more energy, which means they will collide with the magnesium with more energy, which will give more successful collisions per second.
· Surface area of the magnesium
If the magnesium had a bigger surface area each time the experiment was done, then the acid particles will have a bigger area to collide with, so more collisions will occur every second and the more collisions per second than the more successful collisions per second.
· Type of acid used
If you changed the type of acid then the rate of reaction would change. Hydrochloric, Sulphuric and Nitric acid all would produce a different rate of reaction, so if I do change the type of acid then all three kinds would produce a different set of results.
Key factor
I have chosen to use the concentration of the acid as my factor that I will change. I chose this because several different concentrations can be made up before the experiment by the lab technicians and they will be able to make them accurately.
There will be several different concentrations of acid, which will give me a wide range of results, which will be reliable and reproducible.
Fair test: To ensure a fair test is carried out the time taken for the magnesium to disappear will be measured accurately using a stop clock as soon as the HCL is poured into beaker with the magnesium in it, this way most results will be accurate. It is important to keep the reactants separate while setting up the apparatus so that the starting time of the reaction will be measured accurately.
Factors which may not be easy to control are, how well the solution is mixed when it is diluted, to get the correct concentration. We will continously stir the test tube during the experiment. This factor is quite important because it determines exactly what the concentration of the solution is each time it is mixed.
To reduce the heat produced from the experiment a water bath will be used throughout the experiment. This will control the temperature.
Solutions:
I have decided to use 1 mole for the concentration of acid. This is because 2 moles was too reactive and there were not enough results to collect due to the rapid rate of reaction. I used 5 different strengths of hydrochloric acid. These strengths would determine the rates of reactions.
Method:
- Set the apparatus accordingly:
-
Once the apparatus is set out, measure out solution 1 for the first experiment. So measure out 10 cm3 of Hydrochloric acid, 40 cm3 of water and 10 strips of length 3cm of magnesium ribbon, so it totals 30cm of magnesium however the surface area is increased/reduced to help the reaction.
- After you have measured the substances for solution 1, get the stopwatch ready and place the magnesium ribbon in the colinder. Pour the water into the colinder as well.
- Lastly, pour the Hydrochloric acid in there and replace the stopper, the stopwatch should immediately be started as soon as the HCL is poured in. Also, the colinder should be stirred in the water bath.
- One person should read out when the time reaches every 20 seconds, the person holding the buirette will say the point at witch the gas has reached, so the last person can record the results, every 20 secs.
- This method should be repeated 5 times with different concentrations.
- Then the entire experiment must be repeated and recorded again for a fair test.
Table of Results:
Solution 1
Solution 2
The following solutions were measured every 10 secs:
Solution 3
Solution 4
Analysis:
Graphs
Conclusion: In conclusion I think our results did prove my hypotheses that as the concentration of the Hydrochloric acid increases, the time taken for the magnesium to disappear decreases. However, the experiment did not go very well which may have altered some of the results.
Evaluation: The experiment could be improved by:
- Smaller particles have a bigger surface area than larger particle for the same mass of solid.
- Insulating the experiment prevents heat loss therefore it keeps more energy within the experiment having the same effect as raising the temperature of the experiment. At higher temperatures, particles are moving faster, so there are more collisions. Also (and more importantly), the collisions are more energetic.
If the stirring factor is not carefully controlled it could lead to inaccurate results if the concentration of the acid solution is not what we calculated it to be.
The surface area made no apparent difference, although it is hard to tell with the small amount of data that we collected for surface area.
I think it would have been better to concentrate on one or two factors and repeat the same experiment more than once, rather than doing 28 separate ones. This would have allowed us to be surer of the results.
Information from:
- The HUTCHINSON Dictionary of SCIENCE second edition.
- Internet: www.studentcentral.co.uk
- Letts Study Guide, GCSE CHEMISTRY
To investigate the factors that affect the rate of reaction between magnesium and hydrochloric acid.
We placed 25 cm3 of Hydrochloric acid in a conical flask. We then added 4 cm of magnesium ribbon. We immediately connected a gas syringe and ran the experiment for 1 minute. The amount of gas given off was recorded every 10 seconds, in cm3.
We repeated the whole procedure with four different concentrations of HCl., one, two, three and four molar. With each concentration we also used three different surface areas for the magnesium. However, since it comes as a strip, we could only alter the area slightly. For the low surface area we cut it into eight parts. For the high surface area we folded it tightly. For the medium surface it was kept as a strip. This does not really change the surface area, as the liquid can still get to the centre, although it would be replaced at a reduced rate.
I have chosen to repeat the experiment 3 times because it therefore allows me to calculate an average rate of reaction. This will ensure that there are no abnormal results and it will increase accuracy. I have decided to start readings at 20OC and increase by 10OC each time until 60OC is reached, since it will allow me to see the increase in rate of reaction and 5 results should be enough to identify any trends.
* Rates of Reaction
Increasing the temperature increases the speed of the particles. The faster the particles move, the greater the number of collisions, and therefore the rate of the reaction increases. A 10OC rise in temperature almost doubles the rate of most reactions.
Chemical reactions take place by chance. Particles need to collide with enough velocity so that they react. As the temperature is increased the particles move faster since they have more energy. This means that they are colliding more often and more of the collisions have enough velocity to cause a reaction. Since there are more collisions the chemical reaction takes place faster.
A tangent was drawn at the beginning of each curve and its gradient calculated, the gradients are shown in the table below.
Evaluating Evidence
I believe that the experiment was successful but some of the results were unexpected/unreliable. The lines on the graphs for 20OC and 30OC cross, this doesn’t affect my results since I am only concerned with the initial rate or reaction but it was unexpected. In graph 3 the rate of reaction for 60OC is lower than that of 50OC - this result is anomalous and has been ignored in this investigation.
I believe that the experiment was designed well but there were a few problems. Although the initial rate of reaction (which is what I am concerned with in this investigation) seemed to fit a trend, the rate of reaction curves of some temperatures on the graphs crossed. This could have been because some of the magnesium had corroded forming a magnesium oxide layer which would have affected the rate of reaction.
Other factor which could have given me unreliable results could have been that the gas syringes were wet causing them to jam and so not giving correct results or that the bung was not placed on the top of the side arm tube fast enough which allowed gas to escape. I conducted all three experiments for each temperature at the same time to save time. An error in my graphs (plotting, drawing curves or calculating gradients) could have also affected the calculated rates of reaction.
To improve the experiment I would find a way of attaching and releasing the magnesium inside the side arm tube above the acid (with a bung at the top of the side arm tube) so that the magnesium could be dropped into the acid without any gas being lost.
Additional work, which could be carried out, is to repeat the experiment using, a wider range of temperatures. The investigation could also be extended to investigate other factors affecting the rate of reaction such as catalysts, concentration of the acid or particle size of the magnesium.
Ways that I could record my experiment
There are several way which I could record my experiment. The possible ways are as follows:
Amount of gas evolved
I could use a gas syringe to collect the gas that will evolve from my experiment. I could use these results to calculate the initial rate of reaction.
The weight before and after the experiment
I could put the conical flask with the chosen volume of hydrochloric acid onto a set of accurate electronic scales and record the weight of it. I could then drop a piece of magnesium into the conical flask and measure the decrease in weight at chosen intervals. The weight of the experiment will decrease because as the hydrogen in the hydrochloric acid is being displaced it is being released and will float up and out the conical flask, the weight change will not be very big, but there will be one.
How long the magnesium takes to dissolve
I could measure the length of time it takes for the magnesium to dissolve. The only problem with recording my experiment this way is that I could only calculate the average rate of reaction and not the initial rate of reaction
Ways to measure the rate of the reaction:
Average rate of reaction
Initial rate of reaction
I will be using the initial rate to calculate the rate of reaction as it can calculate the true rate and not the average rate of reaction.
Question
My question is to see that if I change the concentration of the Hydrochloric acid for each experiment I will see an increase or decrease in the rate of reaction between Hydrochloric acid and Magnesium ribbon.
COULD BE USED IN EVAL.
Fair Test
In order to keep my experiment a fair test I will have to make sure that I keep the following factors the same:
· Starting temperature of the acid
· Volume of acid used (cubic centimetres)
· Surface area of the magnesium
· Clean the magnesium with emery paper before experiment
· Length of magnesium
I will also have to make sure that the gas syringe is correctly connected and that it is placed quickly and tightly enough so that no hydrogen gas escapes.
Conclusion
My results table and graph show me that when I increase the concentration of the hydrochloric acid, the initial rate of reaction also increases.
Altogether I tested 5 different concentrations of hydrochloric acid. 0.0M, which was the lowest concentration of acid that I used, there was no reaction. 2.5M hydrochloric acid, which was the highest concentration that I used, produced the fasted rate of reaction. I repeated all 5 concentrations twice to be sure that they were reliable results and in all cases the higher the concentration the higher the rate of reaction. I had stated this in my prediction.
However I also stated in my prediction that if I doubled the concentration from 1M to 2M hydrochloric acid then the rate of reaction will also double. I have discovered that this is not the case.
As you can see from this table as the concentration doubles then the rate of reaction approximately quadruples. My graph also shows that as the concentration doubles then the initial rate of reaction approximately quadruples.
I therefore conclude that:
1. The initial rate of reaction increase as the concentration of the acid increases
2. There is a fourfold increase in the rate of reaction as the concentration increases
3. As the concentration doubles the initial rate of reaction is approximately squared
My original prediction was that the initial rate of reaction would double as the concentration doubles. This was incorrect because it was not based on experimental evidence.
Evaluation
Was I precise in my measurements?
I feel that I was precise and accurate in recording measurements.
I accurately cut the magnesium ribbon to the nearest millimetre with the ruler provided.
I measured the starting and the end temperature accurately with a thermometer to the nearest .5OC.
I accurately measured the volume of gas evolved to the nearest .5cm3.
I accurately measure the time that had elapsed to the nearest second with the stop clock.
Did I take enough readings?
Not enough concentrations were used for a good graph. I should have used concentration 3M acid ought to have been done but it was not known then that it was necessary and required.
The results that I did collect were reliable as I repeated the experiments twice to obtain good average results. My results were not only reliable they were reproducible.
Anomalous results
These errors both occurred in the first run, when the technique and practical skill had not been perfected over a large number of experiments. The more times an experiment is preformed then the more accuracy and skill the student acquires. At higher concentrations there is more heat evolved which gives a larger volume error, since gases have a larger volume at higher temperature.
Another reason for these anomalous results is that the plunger on the syringe may have been sticking on the barrel after a lot of runs, which would have given me lower readings.
Improvements to my procedure
Further work
There are several things that I could do for further work. The first is to find out what the initial rate of reaction would be for concentrations 1.25M hydrochloric acid and 1.75M hydrochloric acid as it would be interesting to see how the rate increases compared to 1M hydrochloric acid.
Another experiment that I could do is change the type of acid that I use.
I could also use the less reactive metals of the reactivity series (zinc, aluminium, iron and lead) that way I could find the initial rate of reaction at 5 seconds for the higher concentrations of acid like 3.0M or 3.5M and I could find there relative activity.
BACKGROUND INFORMATION
What affects the rate of reaction?
1) The surface area of the magnesium.
2) The temperature of the reaction.
3) Concentration of the hydrochloric acid.
4) Presence of a catalyst.
In the reaction when the magnesium hits the acid when dropped in,
it fisses and then disappears giving of hydrogen as it fisses and it
leaves behind a solution of hydrogen chloride. The activation energy of a particle is increased with heat. The particles which have to have the activation energy are those particles
which are moving, in the case of magnesium and hydrochloric acid, it is
the hydrochloric acid particles which have to have the activation energy because they are the ones that are moving and bombarding the magnesium
particles to produce magnesium chloride.
We can measure reactions in two ways:
1) Continuous:- Start the experiment and watch it happen; you can use a
computer “logging” system to monitor it. I.e. Watching a colour fade or
increase.
2) Discontinuous:- Do the experiments and take readings/ samples from
the experiment at different times, then analyse the readings/samples to
see how many reactants and products are used up/ produced.
For particles to react:-
a) They have to collide with each other.
b) They need a certain amount of energy to break down the bonds of the
particles and form new ones.
When we increase the temperature we give the particles more energy which:
1) Makes them move faster which In turn makes them collide with each other more often.
2) Increases the average amount of energy particles have so more particles have the “activation energy” Both of these changes make the rate of reaction go up so we see a decrease in the amount of time taken for the reaction and an increase in time taken. Reflects the rate of reaction. Because temperature has an effect on both the speeds at which the particles react and the activation energy they have a greater effect on the rate of reaction than other changes.
A change in concentration is a change in the number of particles in a given volume.
If we increase the volume:-
a) The particles are more crowded so they collide more often.
b) Although the average amount of energy possessed by a particle does not change, there are more particles with each amount of energy;- more particles with the activation energy.
a) is a major effect which effects the rate, but b) is a minor effect which effects the rate very slightly. In this experiment we are not concerned with whether the reaction is exothermic or endothermic because we are concerned with the activation energy needed to start and continue the reaction.
During a chemical reaction the particles have to collide with enough energy to first break the bonds and then to form the new bonds and the rearranged electrons, so it is “safe” to assume that some of the particles do not have enough energy to react when they collide.
The minimum amount of energy that is needed to break down the bonds is called the activation energy (EA). If the activation energy is high only a small amount of particles will have enough energy to react so the reaction rate would be very small, however, if the activation energy is very low the number of particles with that amount of energy will be
high so the reaction rate would be higher. An example of a low EA would be in explosives when they need only a small input of energy to start their exceedingly exothermic reactions. In gases the energy of the particles is mainly kinetic, however in a solid of a given mass this amount of energy is determined by their velocities. determined by the fraction of molecules with sufficient amounts of energy to react. Putting the probability theory and the kinetic theory together this now gives us a statement which accounts for the 100% increase in the rate of reaction in a 10K rise.
Reaction Rate and Concentration.
The reaction rate increases when the concentration of the acid increases because:
If you increase the concentration of the acid you are introducing more particles into the reaction which will in turn produce a faster reaction because there will be more collisions between the particles which is what increases the reaction rate.
Temperature.
When we did the experiment changing temperature we used both of the sets of apparatus. To get a fair reaction we had to keep the amount of magnesium the same and the concentration of the acid. In the experiment we used 0.1g of magnesium and the concentration of the acid was 50cm3 of acid to 50cm3 of water. This is because if we used 100cm3 of acid the reaction would be too fast. Still we had an excess amount of
acid, so one mole of magnesium can react with two moles of HCl. Concentration.
When we did the reaction changing the concentration we changed the concentration until we had just enough for 1 mole of magnesium to react with two of HCl. To get a fair reaction we had to keep the amount of magnesium the same and the temperature. We used 0.1g of magnesium.
RESULTS
Temperature
From this graph you can see that if we do increase the temperature the rate of reaction also increases, but it does not show that if you increase the temperature the rate of reaction doubles.
This graph shows that there is an increase in the rate of reaction as the temperature increases. This shows a curve, mainly because our results were inaccurate in a number of ways. This is because the concentration is changed during the experiment because at high temperatures the acid around the magnesium is diluted. If this experiment was accurate it would be also a curve but if you made it into 1/time the result would be a straight line showing a clear relationship.
Even though I changed it to 1/time it still does not show a clear relationship because of the factors mentioned in the conclusion.
Concentration
This graph shows an increase in the amount of gas given off and the speed at which it is given off. This graph also does not show the rate increase, it just shows how it increases with a change in concentration.
This graph shows that if you increase the concentration of the molar solution of the acid the time in which the Mg takes to disappear becomes a lot slower. This does not show the rate at which this happens, the graph of rate vs. conc. would show a straight line.
This shows a straight line, thus proving that there is a relationship between the time it takes the magnesium to disappear and the concentration of the acid. If we take a gradient of it, it would show the rate at which the reaction was happening.
Because this shows a straight line we can say that it is a second order reaction.
This graph shows a nearly straight line which shows that there is a relationship between the temperature and the rate of reaction, as the gradient shows the rate of reaction. If you look at this graph it comes out to show that if you increase the temperature by 100C the
gradient of the line is doubled. This shows that rate µ temp.
Even though there is a greater increase in the amount of H2 given off in each of the different reactions you can see that there is a change in the amount given off, but between the temperatures 30 and 400C there is not much of a change, this could be because of our human error, there should be a big change in the amount given off.
This table shows a nice spread of results throughout the range of concentration. It clearly shows that the reaction is at different stages so is therefore producing different amounts of H2. This shows also that the reaction is affected by the concentration of the acid.
CONCLUSIONS
I conclude that if you increase the temperature by 10oC the rate of reaction would double, this is because of using the kinetic theory and the probability theory. Even though our results did not accurately prove this, the theory that backs it up is sufficient. the kinetic theory explains that if you provide the particles with a greater amount of kinetic energy they will collide more often, therefore there will be a greater amount of collisions per unit time. The probability theory explains that there is only a number of particles within the reaction with the amount of Ea to react, so if you increase the amount of kinetic energy there will be more particles with that amount of Ea to react, so this will also increase the reaction rate. If you double the concentration of the acid the reaction rate would also double, this is because there are more particles in the solution which would increase the likelihood that they would hit the magnesium so the reaction rate would increase. The graph gives us a good device to prove that if you double the concentration the rate would also double. If you increase the number of particles in the solution it is more likely that they will collide more often. There should be more H2 given off if we compare it across the range of temperatures because the reaction is going quicker and so more H2 is given off in that amount of time. There is more H2 given off if you compare it to the range of
concentrations that you are using, this shows that the reaction is at different stages and so is therefore producing different amounts of H2. Also our results were not accurate but this could be because of a number of reasons.
There our many reasons why our results did not prove this point accurately.
· At high temperatures the acid around the magnesium starts to starts to dilute quickly, so if you do not swirl the reaction the magnesium would be reacting with the acid at a lower concentration which would alter the results.
· Heating the acid might allow H Cl to be given off, therefore also making the acid more dilute which would also affect the results.
· When the reaction takes place bubbles of H2 are given off which might stay around the magnesium which therefore reduces the surface area of the magnesium and so the acid can not react properly with it so this affects the results.
To get more accurate results, we could have heated the acid to a lower temperature to stop a large amount of H Cl being given off. The other main thing that could have helped us to get more accurate results is we cold have swirled the reaction throughout it to stop the diluting of the acid and the bubbles of H2 being given off.
If I had time I could have done the reactions a few more times to get a better set of results. This would have helped my graphs to show better readings
