In the experiment I’m going to carry out I’m going to investigate:
- How changes in concentration of HCl affect the rate of the reaction with magnesium.
- How the rate of reaction of magnesium varies with three different acids.
- How temperature affects the rate of the reaction of the three different acids with magnesium.
PLIMINARY WORK.
In order to find out which method I was going to use to find how the rate of the reaction alters with changes in concentration of HCl, I carried out a few pilot tests to see which method was suitable. I also did the calculations to find out the mass of magnesium used and the volume of HCl used.
CALCULATIONS.
The equation for the reaction is:
Mg(s) + 2HCl(aq) ----------> MgCl2(aq) + H2(g)
The ionic equation for the reaction is:
Mg (s) + 2H (aq) 2Clˉ(aq) ----------> Mg²(aq) 2Clˉ(aq) + H2(g)
The chlorine is a spectator ion. So the equation is:
Mg (s) + 2H (aq) ----------> Mg²(aq) + H2(g).
Mass of 10cm of Mg = 0.08g.
Mass of 1cm of Mg = 8 × 10ˉ³g.
RMM of Mg = 24g
Moles of Mg = Mass = 8 × 10ˉ³g = 0.333 ×10ˉ³ mols.
RMM 24
The ratio of Mg:HCl is 1:2.
Moles of HCl = 2 × 0.333 ×10ˉ³
= 0.667 ×10ˉ³ mols.
If the concentration of HCl is 1moldmˉ³ then:
The volume of HCl = Moles = 0.667 ×10ˉ³ = 0.667 ×10ˉ³ dm³× 1000
Concentration 1
= 0.667cm³.
To make sure that all the magnesium reacts I will use an excess of HCl. I will use 10 cm³
METHOD 1.
I used this method to collect the volume of gas given off every 10 seconds until all the magnesium disappeared.
RESULTS.
Amount of 2M HCl used = 15cm³.
Length of Mg used = 1cm.
I found this method to be unreliable because it was hard to measure the volume of gas given out every 10 seconds. When I tried to find the volume of gas given off every 10 seconds I found the results to be unreliable. During my 1st and the 2nd try I found that the volume of gas collected varied significantly. The bubbles of gas present in the burette made it hard for me to get an accurate reading of the gas collected. Since I found from this work that this method is not going to give me reliable results I’m not going to use this method.
METHOD 2.
I tried to measure the loss of mass of the solution but when I did this experiment I did not find any change in mass. This is because the loss of mass is so small that it cannot be measured by a 2 decimal place balance. And since we are only provided with 2 decimal place balances I won’t be able to follow the rate of the reaction using this method.
METHOD 3.
I used this method to find the time taken for the magnesium strip in the acid to stop fizzing. But when I to do this for different temperature I found it very hard to see the magnesium strip when the test tube was placed inside a beaker with ice cubes. This method also took a lot of time when I did it with concentrations of 0.1M. So in order to get reliable results for this experiment I’m not going to use this method.
METHOD 4.
I used this method to measure the time taken to collect 10cm³ of gas given off.
RESULTS.
Volume of acid used = 10cm³.
Concentration of acid = 2 moldmˉ³.
Length of Mg used = 1cm.
Volume of gas collected = 10 cm³.
Time taken = 9 seconds, 8 seconds and 9 seconds.
Average time taken = 8.67 seconds.
When I carried out this experiment it took me 9 seconds to collect 10 cm³ of gas given off. I found this experiment to be reliable because the time taken to collect the gas each time did not vary too much. So I’m going to use this method to find the rate of the reaction.
This method uses the initial rate method. The initial rate method is used when the initial concentrations of the reactants are known.
APPARATUS.
- Syringe.
- HCl (2M solution.)
-
H2SO4 (2M solution.)
-
CH3COOH (2M solution.)
- Pumpette * 4.
- Pipettes * 4 (50cm³).
- Pipette * 1 (10cm³).
- Measuring cylinders * 1 (100cm³)
- Marker pen.
- Magnesium.
- Water baths - 40ºC, 60ºC and 80ºC.
- Ice cubes.
- Stop clock.
- Beaker * 5 (250 cm³).
- Test tube rack.
- Scissors.
- Ruler (15cm).
- Distilled water.
- Bung.
- Side–arm boiling tube.
- Paper towels.
METHOD.
The method to make the different concentrations:
1 moldmˉ³.
- Measure out 50cm³ of the 2M acid using a 50cm³ pipette.
- Transfer it to a volumetric flask.
- Add distilled water up to the mark on the volumetric flask.
0.5 moldmˉ³.
- Measure out 50cm³ of the 1M solution of the acid using a pipette.
- Transfer it to a volumetric flask.
- Add distilled water up to the mark on the volumetric flask.
0.25 moldmˉ³.
- Measure out 50cm³ of the 0.5M solution of the acid.
- Transfer it to a volumetric flask.
- Add water up to the mark on the volumetric flask.
0.1 moldmˉ³.
- Measure out 10cm³ of the 1M acid.
- Transfer it to a volumetric flask.
- Add water up to the mark on the volumetric flask.
0.2 moldmˉ³.
- Measure out 10cm³ of the 2M acid.
- Transfer it to a volumetric flask.
- Add water up to the mark on the volumetric flask.
The method for carrying out the experiment.
- Set up the apparatus as show above.
- Measure out 10cm³ of HCl using a 10cm³pipette and transfer it to the test tube.
- Measure out 1cm of the magnesium using a ruler and then cut out the 1cm piece using a scissors.
- Clean the cut piece of magnesium ribbon using a paper towel.
- Add the Mg to the test tube and immediately close the open end of the test tube with the bung.
- Start the stop clock as soon as the Mg is added.
- Measure the time taken to collect 10 cm³ of gas.
- After collecting the 10cm³ of the gas pour the solution of the acid with the remaining magnesium down the drain.
- Then wash the test tube thoroughly.
- Clean the test tube dry with a paper towel before you repeat the experiment.
- Repeat the experiment 3 times.
- Clean the gas syringe with a paper towel every 2 experiments to prevent it from sticking.
- Repeat the above steps for different concentrations of HCl.
- Find the average time for a particular reaction after doing the experiment 3 times.
- Then find the rate of the reaction using the formula:
Rate = 1
Time
- Draw a graph of rate of reaction on the y-axis against the concentration of HCl on the x-axis.
- Use this to find the order of the reaction.
Method for different temperatures.
- Measure out 10cm³ of HCl using a 10cm³and transfer it to the test tube.
- Place a thermometer inside the test tube.
- Then place the test tube inside the water bath.
- Wait till the acid inside the test tube reaches the temperature of the water bath.
- Once the acid reaches the temperature of the water bath, remove the thermometer from the test tube.
- Add 1cm of the cleaned magnesium into the test tube.
-
Measure the time taken to collect 10cm³ of H2 gas given off.
- Repeat the above method for different temperatures.
- After doing the experiment 3 times find the average time for the reaction.
- Use the average time to calculate the rate of the equation using the formula given in the earlier page.
- Draw a graph of temperature on the x-axis against the rate of the reaction on the y-axis.
- Draw another table and find 1/temperature and ln rate.
- Then draw a graph of 1/temperature on the x-axis against ln rate on the y-axis.
- Find the gradient of the graph.
- Use this gradient to calculate the activation energy of the reaction using the formula:
Gradient = - Ea where R = 8.31 J Kˉ¹molˉ¹.
R
Method for different acids.
-
Make the different concentrations of H2SO4 and CH3COOH using the method show under “method for different concentrations”.
-
Then add 10cm³ of 2M H2SO4 into a test tube.
- Add 1cm of magnesium to this test tubes.
-
Then measure the time taken to collect 10cm³ of H2 gas.
- Repeat the experiment 3 times.
-
Then repeat the experiment for different concentrations of H2SO4.
-
Repeat the above steps for different concentrations of CH3COOH.
- Then find the average time for each of these reactions.
- Use the average time to calculate the rate of the reaction.
- Then draw a graph of concentration of the acid on the x-axis against the rate of the reaction on the y-axis.
- Use these graphs to find the order of the reaction.
IMPLEMENTING.
RESULTS FOR VARYING CONCENTRATIONS AND DIFFERENT ACIDS.
RESULTS FOR VARYING TEMEPRATURES.
The acid used = HCl.
Concentration of acid = 1M.
The volume of acid used = 10cm³.
The length of magnesium used = 1cm.
The volume of gas collected = 10cm³.
ANALYSIS.
RESULTS.
Acid used = HCl.
Acid used = H2SO4
Acid used = CH3COOH.
ANALYSIS OF THE GRAPHS AND CONCLUSION.
After carrying out the experiment, I used the time I found to collect 10cm³ of H2 gas when magnesium reacted with HCl to find the rate of the reaction. I found the rate of the reaction using the formula shown in the tables above.
The graph of the concentration of the HCl against the rate of the reaction shows that as the concentration of HCl increases then the rate of the reaction increases.
I drew a graph of concentration against the rate of the reaction instead of the time against concentration graph because I did not measure the volume of hydrogen gas given off with time for each different concentrations of HCl. But since I only measured the time it took to collect a fixed volume of gas I cannot draw this graph for the experiment I carried out. I can still use this graph to find out the order of the reaction.
The graph of concentration of H2SO4 against the rate of the reaction shows that the rate of the reaction increases as the concentration of the acid increases. The same results can be seen for the graph of concentration of CH3COOH against the rate of the reaction.
Generally the rate of the reaction increases as the concentration of the reactants increase. This is because according to the kinetic theory the molecules in a liquid are constantly moving and bumping into each other. When these molecules collide there is a chance that they will react. So the more the reactants there are the greater the rate of reaction. This is because increasing the concentration increases the number of molecules present and as the number of molecules increase, there will be a greater number of collisions and so the rate of reaction increases.
The results from the experiment also show that out of the 3 acids I used the rate of reaction is the fastest for H2SO4 and the rate is the slowest for CH3COOH. This proves that my predictions were right when I said that the rate of the reaction was going to be the fastest for H2SO4 and the slowest for CH3COOH in my plan. CH3COOH is the slowest because CH3COOH is a weak acid so it does not dissociate completely in water. So the resulting solution only has a few hydrogen ions to react with the magnesium. So the rate of the reaction is going to be slow.
HCl -----> H + Clˉ
H2SO4 -----> 2H + SO4²ˉ
According to the equation it shows that H2SO4 completely dissociates in water to form 2H and SO4²ˉ. But when HCl dissociates there is only one mole of H produced. This shows that a solution of H2SO4 will have more hydrogen ion concentration than a solution of HCl. So the more the hydrogen ions present the greater the rate of reaction.
The equation also shows that the rate of the reaction for H2SO4 should be double the rate of the reaction of that of HCl. But when I carried out the experiment I did not actually find this to be right. This could be due to anomalous results and errors in the apparatus I used.
RESULTS FOR TEMPERATURE.
Acid used = HCl.
Concentration = 1M
Gradient = Δ ln Rate
Δ1/temperature
= - 3.25
0.8 × 10ˉ³
= - 4062.5
The gradient of the graph = - Ea
R
- 4062.5 = - Ea
8.31
Ea = 33759.4 J molˉ¹.
= 33.8 KJ molˉ¹.(3s.f).
ANALYSIS OF THE GRAPHS AND CONCLUSION.
The graph of temperature against the rate of the reaction shows that as the temperature increases the rate of the reaction increases. The results I got form an almost smooth curve on the graph.
If the molecules collide with each other softly no reaction will take place because the molecules will just bounce off each other. The molecules must collide with each other with enough energy to break the chemical bounds for a reaction to take place. When a mixture is heated the kinetic energy of the molecules will increase. This causes the molecules to move faster. Since the molecules have more kinetic energy now, a larger proportion of the collisions will exceed the activation energy of the reaction and so the rate of the reaction increases.
This is show in the graph because when the reaction is allowed 2 take place at 0ºC the molecules of the reactants will have less energy so most of the collisions will not exceed the activation energy of the reaction and so the rate of reaction is slow. Where else if the reaction is carried out at 80ºC then most of the molecules will have a high kinetic energy so most of the collisions will exceed the activation energy of the reaction and so the rate of the reaction increases.
The activation energy for this reaction can be found using the Arrhenius equation. Using the equation I found the activation energy to be 33.8 KJ molˉ¹.
RESULTS.
Acid used = HCl.
Acid used = H2SO4
Acid used = CH3COOH.
CONCLUSION.
The first graph was drawn to show that the rate of the reaction increases as the concentration of the acid increases. The graph shows a line that is slightly curved. This shows that the reaction is second order with respect to the acid. All the three graphs I drew show a line that is slightly curved. So the reaction must be second order with respect to any of these acids.
The table of results can be used to find out if the reaction is second order. In a second order reaction doubling the concentration of the acid would quadruple the rate of the reaction.
Acid used :----> HCl
Concentration = 1M. Rate = 10 × 10ˉ³.
Concentration × 2 = 2M Rate × 4 = 40 × 10ˉ³. Actual rate = 50 × 10ˉ³.
Acid used :----> H2SO4
Concentration = 0.2M. Rate = 0.414 × 10ˉ³.
Concentration × ½ = 0.1M Rate × ¼ = 0.104 × 10ˉ³. Actual rate = 0.150 × 10ˉ³.
Acid used :----> CH3COOH
Concentration = 0.25M. Rate = 0.269 × 10ˉ³.
Concentration × 2 = 0.5M Rate × 2 = 1.076 × 10ˉ³. Actual rate = 1.02 × 10ˉ³.
This shows that as the concentration doubles the rate of the reaction increases by a factor of 4. It also shows that as the concentration is halved the rate of the reaction decreases by a factor of ¼. So now I can say that the order of the reaction is second order with respect to the acids.
A reaction is second order with respect to a reactant if the rate of the reaction is proportional to the concentration of that reactant squared. So when I drew a graph of concentration squared against the rate of the reaction, I got a straight line that went through the origin. This also shows that the reaction is second order with respect to the acids.
EQUATIONS FOR THE REACTION BETWEEN DIFFERENT ACIDS.
Magnesium + Hydrochloric acid ----> Magnesium chloride + Hydrogen.
Chemical equation:
Mg(s) + 2HCl(aq) ----> MgCl2(aq) + H2(g)
Ionic equation:
Mg(s) + 2H (aq) 2Clˉ(aq) ----> Mg² (aq) 2Clˉ(aq) + H2(g)
Clˉ are spectator ions.
Mg(s) + 2H (aq) ----> Mg² (aq) + H2(g)
Magnesium + Sulphuric acid ----> Magnesium sulphate + Hydrogen.
Chemical equation:
Mg(s) + H2SO4(aq) ----> MgSO4(aq) + H2(g)
Ionic equation:
Mg(s) + 2H (aq) SO4²ˉ(aq) ----> Mg² (aq) SO4²ˉ(aq) + H2(g)
SO4²ˉ are spectator ions.
Mg(s) + 2H (aq) ----> Mg² (aq) + H2(g)
Magnesium + Ethanoic acid ----> Magnesium ethanoate + Hydrogen.
Chemical equation:
Mg(s) + CH3COOH ----> CH3COOˉMg² (aq) + H2(g)
Ionic equation:
Mg(s) + CH3COOˉH (aq) ----> CH3COOˉMg² (aq) + H2(g)
CH3COOˉ are spectator ions.
Mg(s) + H (aq) ----> Mg² (aq) + H2(g).
WORKING OUT THE RATE.
Volume of H2 collected = 10cm³.
24dm³ ----> 1 mol of gas
0.01cm³ ----> 1 × 0.01
24
= 0.417 ×10ˉ³mols dmˉ³.
Rate = moles
Time
Acid used = HCl
Rate = k [HCl]²
Where k is the rate constant
k = Rate (mols dmˉ³ sˉ¹)
[HCl]² (mols dmˉ³)²
The units for k = dm³molˉ¹sˉ¹.
Average value of k = 2.661 × 10ˉ
6
= 4.44 × 10ˉ dm³molˉ¹sˉ¹.
Acid used = H2SO4
Rate = k [H2SO4]²
k = Rate (mols dmˉ³ sˉ¹)
[H2SO4]² (mols dmˉ³)²
Units = dm³molˉ¹sˉ¹.
Average value of k = 4.642 × 10ˉ = 7.74 × 10ˉ dm³molˉ¹sˉ¹.
6
Acid used = CH3COOH
Rate = k [CH3COOH]²
k = Rate (mols dmˉ³ sˉ¹)
[CH3COOH]² (mols dmˉ³)²
Units = dm³molˉ¹sˉ¹.
Average value of k = 1.46 × 10ˉ
6
= 2.44 × 10ˉ dm³molˉ¹sˉ¹.
ERRORS.
Resolution = 0.01g.
Average reading = 0.08g.
Percentage uncertainty = 0.01 × 100
0.08
= 12.5%
Resolution = 0.1cm³
Average reading = 10cm³.
Percentage uncertainty = 0.1 × 100
10
= 1%
Resolution = 1cm³
Average reading = 50 + 50 + 50 + 10 + 10
5
= 170 = 34 cm³
5
Percentage uncertainty = 1 × 100
34
= 2.94%
Resolution = 0.1 cm
Average reading = 1 cm
Percentage uncertainty = 0.1 × 100
1
= 10%
Resolution = 1 cm³
Average reading = 50 + 50 + 50 + 90 + 90
5
= 330 = 66 cm³
5
Percentage uncertainty = 1 × 100
66
= 1.52%
Resolution = 1 cm³
Average reading = 10 cm³
Percentage uncertainty = 1 × 100
10
= 10%
Resolution = 1ºC
Average reading = 27 + 40 + 60 + 80 + 0
5
= 207 = 41.4º
5
Percentage uncertainty = 1 × 100
41.4
= 2.42%
Resolution = 0.01 s
Average reading =
20+98.3+413+2201.7+2903.3+6416.7+13.7+36.7+276+616.7+2415+6675+110+343.3+980+3041.7+3713.3+7233+325+98.3+26.7+25+14.3 .
23
= 37996.7 = 1652
23
Percentage uncertainty = 0.01 × 100
1652
= 0.0006%
Total percentage uncertainty = 12.5% + 1% + 2.94% + 10% + 1.52% + 10% + 2.42% + 0.0006%
= 40.38%
After calculating the percentage uncertainties for the apparatus I used I found the electronic balance to have the biggest percentage uncertainty. The percentage uncertainty for the electronic balance was 12.5%. This shows that this is a very large error and therefore the mass of the magnesium strip of 10 cm cannot be accurate. So the moles of magnesium calculated by using this mass would also be wrong and this would mean the volume of HCl needed to react with a magnesium strip of 1cm would also be wrong because I used the initial mass of the magnesium to work this out. But since I used an excess of the acids this error would not make a big difference to the final results I got.
The other apparatus which had big percentage uncertainties was the 15cm ruler and the gas syringe. The ruler has a percentage uncertainty of 10 % and the gas syringe had the same value. These percentage uncertainties would affect the accuracy of the results I got. I had to use the ruler to cut out 1cm of magnesium. But since the percentage uncertainty for this apparatus is so large then it mite be that when I measure out 1cm of magnesium it could have been more than 1cm in some of the reactions and it could have been less than 1cm for the other reactions. This shows that the length of magnesium I used was not consistent through out the whole of this experiment.
The thermometer I used had a percentage uncertainty of 2.42%. Although this is not a very large number I can still make this percentage error to be very small by using a thermometer which has a resolution of 0.1ºC if I do this experiment again.
Altogether the percentage uncertainties of the apparatus I used worked out to be 440.38%. This shows that the apparatus I used were not very accurate in the results they gave. This in turn would have affected the accuracy of the final results I got.
The main difficulties I found in doing this experiment was to put the bung into the test tube as soon as I added the magnesium and to also start the stop clock as soon as I added the magnesium. I did try my best to put the bung in as soon as I added the magnesium but still some of the gas would still escape before I could put the bung on. I also found it very hard to start the stop clock immediately after I added the magnesium. Some time did pass before I could start the stop clock because I had to put the bung in first. To overcome this error I found the average time taken for the reaction.
All the concentrations I used in this experiment where made from 2M acid. This means the accuracy of the molarity and the amounts of water added would affect the accuracy. But since the percentage uncertainties for the apparatus I used to make these dilutions were very small, I don’t think this would give me a percentage error in the overall results of the experiment.
CONCLUSION.
After doing this experiment I can conclude that the rate of the reaction increases as the concentration of the acids increase. The overall order of the reaction between magnesium and HCl is second order with respect to the acid.
Rate = k [HCl]²
k = 4.44 × 10ˉ dm³molˉ¹sˉ¹.
[Mg] is not added in the rate equation because the magnesium is a solid and its concentration does not vary even though its mass is decreasing.
The overall order of the reaction between Mg and H2SO4 is also second order with respect to the acid.
Rate = k [H2SO4]
k = 7.74 × 10ˉ dm³molˉ¹sˉ¹.
The overall order of the reaction between Mg and CH3COOH is second order with respect to the acid.
Rate = k [CH3COOH]
k = 2.44 × 10ˉ dm³molˉ¹sˉ¹.
BIBLIOGRAPHY.
BOOKS:
- Nuffield advanced chemistry students’ book.
- A –level study guide, chemistry – Philip Barratt & Michael Cox.
- Study and revise As and A2 chemistry – Whsmith.
- Chemistry in context – Longman.
WEBSITES:
