Concentration's effect on Mg and HCl reaction rate

In this investigation I will be measuring the rate of reaction between magnesium and hydrochloric acid. The rate of a reaction tells us how quickly a chemical reaction happens.

Reaction Rate = change in volume, mass or concentration of substance

Time taken

There are two ways to measure the rate of a reaction, by observing how quickly the reactants are used up or by observing how quickly the products are formed. Measurements of the rate of reaction can be taken in three main ways:

* Measuring the rate of precipitation

* Measuring the volume of gas

* Measuring the change in mass

The results gained from these experiments can be drawn on a graph, which enables the rate of reaction to be worked out.

In a chemical reaction atoms are rearranged. In order for a reaction to occur the molecules must collide by coming together. However not all collisions are effective. This is because in gases and liquids, particles are constantly moving causing millions and millions of collisions every second.

If there were a reaction every time molecules collided all chemical reactions would only take a few seconds. This is why only a small fraction of the collisions between the particles have an effect. When particles collide head on and are fast moving a reaction occurs. This is because if collisions between particles have enough energy a reaction will occur. In gases, liquids and in solution, the particles move at a range of speeds. Some are moving very slowly and others are moving very fast.

To react, particles must collide with enough energy and in the correct orientation for bonds to be broken. This is because for a chemical reaction to take place, some bonds in the reactants must be broken.

In a chemical reaction if the activation energy is low many of the collisions will have enough energy so the reaction will be fast whilst if the activation energy is high fewer collisions will have enough energy so the reaction will be slower. The activation energy for a reaction is the minimum energy needed for a reaction to occur. You can show this on an energy profile for the reaction. For a simple exothermic reaction, the energy profile looks like this:

When magnesium powder and dilute hydrochloric acid are mixed together a reaction occurs. This reaction is exothermic meaning heat is given off.

Magnesium + Hydrochloric Magnesium + Hydrogen + Energy

Acid Chloride

Mg (s) + 2HCl (aq) MgCl2 (aq) + H2 (g) + Energy

The rate of this reaction can be changed by varying the conditions in which the reaction occurs. The factors that affect the rate of reaction are:

* Surface area

* Temperature

* Presence of a catalyst

* Concentration/pressure if gaseous

The surface area of the solid reactants can be changed which has a big effect on the rate of reaction. For example magnesium powder will react faster with hydrochloric acid than magnesium ribbon. This is because although the same mass is used the powder has a larger surface area. This means that more particles are exposed to the acid so there is a greater chance of collisions and the more collisions in a given time the greater the rate of reaction. This is illustrated in the diagram below.

If the same mass of magnesium ribbon was used the reaction would be slower because of the smaller surface area. Only the outside of the ribbon is in contact with the acid particles so to begin with only the outside of the ribbon will react. The magnesium particles inside the ribbon can only react when the outside particles have reacted; this is because they do not come in contact with the acid until this point. Therefore as you increase the surface area you increase the rate of reaction. A simple graph of the affect of surface area on the rate of reaction is shown below.

An increase in temperature will increase the rate of reaction. The kinetic energy of particles is proportional to the temperature. Particles will have more kinetic energy at higher temperatures. This causes them to move faster which means that the collisions will be more frequent between particles in any moment in time. Also they collide more energetically and therefore there is more chance of collisions with energy equal to or greater than the activation energy and so more collisions will cause a reaction. A graph showing the affect of temperature on the rate of reaction is shown below.

Another way to show the effect of temperature on the rate of reaction is to use a Maxwell-Boltzmann curve. A simple curve is shown below.

As you increase the temperature the particles will move faster as they will have more energy, however not all the particles in a substance will move at the same speed. As the temperature has been increased there will be more particles with an energy level equal to or greater than the activation energy and therefore more collisions will bring about a reaction. This affect of the increase in temperature on the number of molecules with energy equal to or greater than the activation energy is shown below.

A catalyst is a substance which alters the rate of a reaction without itself being used up or changed chemically during the reaction. Most catalysts reduce the activation energy and so increase the rate of reaction. A catalyst provides an alternative path for a reaction with a lower activation energy. This means that there will be more particles with an energy equal to or greater than the activation energy. Therefore there will be more collisions in which the particles react and so the rate of reaction will have increased. Enzymes also have this affect as they are biological catalysts. Two graphs are shown below one showing the affect of a catalyst on the activation energy and on the rate of reaction.

Many reactions involve catalysts, some examples are shown below.

Reaction

Catalyst

Decomposition of hydrogen peroxide

manganese(IV) oxide, MnO2

Nitration of benzene

concentrated sulphuric acid

Manufacture of ammonia by the Haber Process

iron

Conversion of SO2 into SO3 during the Contact Process to make sulphuric acid

vanadium(V) oxide, V2O5

Hydrogenation of a C=C double bond

nickel

Increasing the concentration of a reagent increases the number of particles in a given volume; this increases the rate of reaction. This is due to the fact that there are more reactant particles in solution and therefore collisions will be more frequent so there is more chance of collisions with energy greater than or equal to the activation energy. Collisions are only effective if they have energy equal to or greater than the activation energy. The activation energy is the minimum amount of energy needed for a reaction to occur. This is because in order for particles to react they must collide with enough energy and in the correct orientation for bonds to be broken. This is because for a chemical reaction to take place, some bonds in the reactants must be broken. This is shown in the diagrams below.

When a reaction first begins there is a high concentration of reactant particles. There are more collisions between the particles and so the rate of reaction is greatest at the beginning of a reaction. As there are more collisions there will be more collisions with an energy equal to or greater than the activation energy so the rate of reaction will be faster at the beginning of a reaction. However as a reaction continues the concentration of reactant particles decreases as they have already reacted. This causes the rate of reaction to decrease. This is because there will be less collisions so there will be fewer collisions with an energy greater than or equal to the activation energy causing the rate of reaction to decrease and eventually be 0 when the reaction stops. The reaction will finish when either of the reactants have run out. This pattern is shown on the graphs on the following page.

The rate of reaction at a particular point can be worked out by drawing a graph of the results. This is done by drawing a tangent at the chosen time as shown in the diagram opposite.

The next step is to draw a right-angled triangle from the tangent as shown in the diagram opposite. The gradient of the tangent can be worked out using the following formula,

Gradient = opposite

Adjacent

You can show the decrease in the rate of reaction at the end of a reaction by drawing a tangent at both the beginning and end of the reaction. This will allow you to compare the rate of reaction at the beginning and end of the reaction.

Variables

The factors that affect the rate of reaction are:

* Surface Area

* Concentration

* Presence of a catalyst

* Temperature

I will be varying the concentration of the hydrochloric acid. The concentrations I will be using are, 0.8M, 1.0M, 1.2M, 1.4M, 1.6M.

Preliminary Work

When carrying out preliminary work I used the highest and lowest concentrations available in order to find a suitable mass of magnesium and volume of hydrochloric acid. This enabled me to see if the reaction was going too fast or too slow to be recorded. This helped me in choosing a suitable mass and volume. Using the highest and lowest concentration also allowed me to make sure that all of the concentrations will react at a reasonable rate.

Method

1. Measure out 25ml of 0.8M hydrochloric acid using a 25ml measuring cylinder. Pour the acid into a 100ml beaker.

2. Place the boat on the balance and tare the balance. Add 0.2g of magnesium powder using a spatula.

3. Place the beaker with acid and the boat with magnesium on the balance and tare it.

4. Pour the magnesium powder into the beaker of hydrochloric acid and put the empty boat back on the balance. Start the stopwatch at the same time as the magnesium is added.

5. Record the mass loss every 15 seconds and record results in a table.

Results

Mass of magnesium = 0.2g

Concentration of acid = 0.8M

Volume of acid = 25ml

Time (s)

Mass loss (g)

0

0.00

15

0.08

30

0.14

45

0.17

60

0.21

75

0.23

90

0.25

105

0.27

120

0.29

135

0.30

150

0.31

165

0.33

180

0.33

I have used magnesium powder not magnesium ribbon because the magnesium powder has a greater surface area which will mean that less magnesium will need to be used in order for it to react at a reasonable rate.

After looking at the results for the 0.8M acid with 0.2g of magnesium and 25ml of acid I have found that the reaction was too fast too record. I reduced the mass to 0.1g to see if this was a suitable mass. This should reduce the rate of reaction as there are less reactant particles so less collisions.

Method

1. Measure out 25ml of 0.8M hydrochloric acid using a 25ml measuring cylinder. Pour the acid into a 100ml beaker.

2. Place the boat on the balance and tare the balance. Add 0.1g of magnesium powder using a spatula.

3. Place the beaker with acid and the boat with magnesium on the balance and tare it.

4. Pour the magnesium powder into the beaker of hydrochloric acid and put the empty boat back on the balance. Start the stopwatch at the same time as the magnesium is added.

5. Record the mass loss every 15 seconds and record the results in a table.

Results

Mass of magnesium = 0.1g

Concentration of acid = 0.8M

Volume of acid = 25ml

Time (s)

Mass loss (g)

0

0.00

15

0.07

30

0.10

45

0.11

60

0.14

75

0.15

90

0.16

105

0.16

120

0.17

135

0.17

150

0.17

165

0.17

180

0.17

The results of this experiment show that the reaction is still quite fast. To solve this problem I have reduced the volume of acid to 15ml as this will reduce the rate of reaction as there are less reactant particles so less collisions. I also decided to record the mass loss every 5 seconds not 15 seconds and this will make my results more accurate and will allow me to plot a better graph.

Method

1. Measure out 15ml of 0.8M hydrochloric acid using a 25ml measuring cylinder. Pour the acid into a 100ml beaker.

2. Place the boat on the balance and tare the balance. Add 0.1g of magnesium powder using a spatula.

3. Place the beaker with acid and the boat with magnesium on the balance and tare it.

4. Pour the magnesium powder into the beaker of hydrochloric acid and put the empty boat back on the balance. Start the stopwatch at the same time as the magnesium is added.

5. Record the mass loss every 5 seconds and record the results in a table.

Results

Mass of magnesium = 0.1g

Concentration of acid = 0.8M

Volume of acid = 15ml

Time (s)

Mass loss (g)

0

0.00

5

0.05

10

0.10

15

0.10

20

0.11

25

0.12

30

0.13

35

0.14

40

0.15

45

0.16

50

0.17

55

0.17

60

0.18

65

0.18

70

0.20

75

0.20

80

0.21

85

0.21

90

0.22

95

0.23

100

0.23

105

0.23

110

0.23

115

0.23

120

0.23

This reaction seemed to work well so I have repeated the experiment using acid of concentration 1.6M. I have done this in order to check that the mass and volumes chosen are suitable for all the concentrations. I need to ensure that the reaction will not be too fast or too slow to record.

Results

Mass of magnesium = 0.1g

Concentration of acid = 1.6M

Volume of acid = 15ml

Time (s)

Mass loss (g)

0

0.00

5

0.19

10

0.20

15

0.23

20

0.24

25

0.26

30

0.27

35

0.28

40

0.29

45

0.30

50

0.30

55

0.30

60

0.30

65

0.30

70

0.30

75

0.30

These results show that the volume of acid and mass of magnesium are suitable as both the highest and lowest concentrations will react and neither reacts too quickly or too slowly. I have decided to use 0.1g of magnesium powder and 15ml of hydrochloric acid. I will be recording the mass loss every 5 seconds.

Fair Test

The test will be fair because I will only be changing the concentration of hydrochloric acid.

I will use:

* The same volume of acid

* The same mass of magnesium

* The same temperature of acid (room temperature)

* The same surface area of magnesium

* No catalyst

* The same sized beaker

Prediction

I predict that as the concentration increases the rate of reaction will increase. This is because the more concentrated the acid is the more particles there are in a particular volume. As there is a greater number of particles there will be more frequent collisions so there is more likely to be a collision between particles that have energy equal to or greater than the activation energy. A graph showing the activation energy is shown below.

In order for the particles to react they must have energy equal to or greater than the activation energy. They must collide with enough energy and in the correct orientation in order for a reaction to occur. A higher concentration of acid will cause the reaction to be faster because there will be more collisions that cause a reaction. This is shown in the diagrams below.

As there are more particles there will be more productive collisions.

Safety

When carrying out the experiment I need to wear safety spectacles and make sure that no equipment is near the edge of the table. Hydrochloric acid is corrosive and causes burns so I must wear eye protection. Hydrochloric acid is also dangerous with magnesium so I need to make sure I only use low concentrations of acid and a small volume of acid. Powdered magnesium is dangerous with acid because of the large surface area so I will use a small mass of magnesium. The gas produced in the reaction between magnesium and hydrochloric acid is hydrogen, which is a highly flammable gas. Therefore I must only use small quantities of magnesium and hydrochloric acid so ensure that less hydrogen is released.

Apparatus

1 x boat

1 x spatula

1 x balance

1 x 25ml measuring cylinder

1 x beaker (100ml)

1 x stop watch

Magnesium Powder

Hydrochloric Acid of concentration 0.8M, 1.0M, 1.2M, 1.4M, 1.6M

Method

1. Set up apparatus as shown in the diagram opposite

2. Measure out 15ml of 0.8M hydrochloric acid using a 25ml measuring cylinder. Pour the acid into a 100ml beaker.

3. Place the boat on the balance and tare the balance. Add 0.1g of magnesium powder using a spatula.

4. Place the beaker with acid and the boat with magnesium

on the balance and tare it.

5. Pour the magnesium powder into the beaker of hydrochloric acid and put the empty boat back on the balance. Start the stopwatch at the same time as the magnesium is added.

6. Record the mass loss every 5 seconds and record the results in a table.

7. Repeat steps 1 to 6 with concentrations 1.0M, 1.2M, 1.4M and 1.6M

8. Repeat steps 1 to 7 twice more to make results more reliable and find the average mass loss at each time for each concentration.

Mass loss (g)

Time (s)

Try 1

Try 2

Try 3

Average

0

0.00

0.00

0.00

0.00

5

0.05

0.04

0.03

0.04

10

0.06

0.06

0.06

0.06

15

0.07

0.07

0.07

0.07

20

0.08

0.08

0.08

0.08

25

0.09

0.09

0.10

0.09

30

0.10

0.10

0.10

0.10

35

0.11

0.10

0.11

0.11

40

0.12

0.11

0.12

0.12

45

0.12

0.11

0.13

0.12

50

0.13

0.12

0.14

0.13

55

0.13

0.12

0.14

0.13

60

0.14

0.13

0.14

0.14

65

0.14

0.13

0.15

0.14

70

0.15

0.13

0.16

0.15

75

0.15

0.13

0.16

0.15

80

0.15

0.13

0.16

0.15

85

0.15

0.14

0.17

0.15

90

0.15

0.14

0.18

0.16

95

0.15

0.14

0.18

0.16

100

0.15

0.15

0.18

0.16

105

0.15

0.15

0.18

0.16

110

0.15

0.15

0.18

0.16

115

0.15

0.15

0.18

0.16

120

0.15

0.15

0.18

0.16

125

0.15

0.15

0.18

0.16

130

0.15

0.15

0.18

0.16

135

0.15

0.15

0.18

0.16

140

0.15

0.15

0.18

0.16

Results for 0.8M

Mass loss (g)

Time (s)

Try 1

Try 2

Try 3

Average

0

0.00

0.00

0.00

0.00

5

0.05

0.05

0.06

0.05

10

0.09

0.07

0.08

0.08

15

0.10

0.10

0.09

0.10

20

0.11

0.10

0.11

0.11

25

0.13

0.10

0.13

0.12

30

0.14

0.12

0.14

0.13

35

0.14

0.12

0.14

0.13

40

0.15

0.14

0.14

0.14

45

0.15

0.14

0.14

0.14

50

0.16

0.14

0.15

0.15

55

0.16

0.15

0.15

0.15

60

0.16

0.16

0.16

0.16

65

0.16

0.16

0.16

0.16

70

0.17

0.16

0.16

0.16

75

0.17

0.17

0.17

0.17

80

0.17

0.17

0.17

0.17

85

0.17

0.17

0.17

0.17

90

0.17

0.17

0.17

0.17

95

0.17

0.17

0.17

0.17

100

0.17

0.17

0.17

0.17

105

0.17

0.17

0.17

0.17

110

0.17

0.17

0.17

0.17

115

0.17

0.17

0.17

0.17

120

0.17

0.17

0.17

0.17

125

0.17

0.17

0.17

0.17

130

0.17

0.17

0.17

0.17

135

0.17

0.17

0.17

0.17

140

0.17

0.17

0.17

0.17

Results for 1.0M

Mass loss (g)

Time (s)

Try 1

Try 2

Try 3

Average

0

0.00

0.00

0.00

0.00

5

0.08

0.08

0.05

0.07

10

0.10

0.10

0.09

0.10

15

0.12

0.12

0.11

0.12

20

0.12

0.14

0.13

0.13

25

0.13

0.14

0.13

0.13

30

0.13

0.15

0.14

0.14

35

0.14

0.15

0.15

0.15

40

0.15

0.15

0.15

0.15

45

0.15

0.16

0.15

0.15

50

0.16

0.16

0.16

0.16

55

0.16

0.16

0.16

0.16

60

0.16

0.17

0.16

0.16

65

0.17

0.17

0.17

0.17

70

0.17

0.17

0.17

0.17

75

0.17

0.17

0.17

0.17

80

0.17

0.18

0.17

0.17

85

0.17

0.18

0.17

0.17

90

0.17

0.18

0.17

0.17

95

0.17

0.18

0.17

0.17

100

0.17

0.18

0.17

0.17

105

0.17

0.18

0.17

0.17

110

0.17

0.18

0.17

0.17

115

0.17

0.18

0.17

0.17

120

0.17

0.18

0.17

0.17

125

0.17

0.18

0.17

0.17

130

0.17

0.18

0.17

0.17

135

0.17

0.18

0.17

0.17

140

0.17

0.18

0.17

0.17

Results for 1.2M

Mass loss (g)

Time (s)

Try 1

Try 2

Try 3

Average

0

0.00

0.00

0.00

0.00

5

0.11

0.08

0.09

0.09

10

0.12

0.10

0.11

0.11

15

0.13

0.11

0.14

0.13

20

0.14

0.12

0.15

0.14

25

0.15

0.13

0.15

0.14

30

0.15

0.13

0.16

0.15

35

0.15

0.14

0.16

0.15

40

0.16

0.15

0.16

0.16

45

0.16

0.15

0.17

0.16

50

0.17

0.16

0.17

0.17

55

0.17

0.16

0.17

0.17

60

0.18

0.17

0.18

0.18

65

0.18

0.17

0.18

0.18

70

0.18

0.17

0.18

0.18

75

0.18

0.17

0.18

0.18

80

0.18

0.17

0.18

0.18

85

0.18

0.17

0.18

0.18

90

0.18

0.17

0.18

0.18

95

0.18

0.17

0.18

0.18

100

0.18

0.17

0.18

0.18

105

0.18

0.17

0.18

0.18

110

0.18

0.17

0.18

0.18

115

0.18

0.17

0.18

0.18

120

0.18

0.17

0.18

0.18

125

0.18

0.17

0.18

0.18

130

0.18

0.17

0.18

0.18

135

0.18

0.17

0.18

0.18

140

0.18

0.17

0.18

0.18

Results for 1.4M

Mass loss (g)

Time (s)

Try 1

Try 2

Try 3

Average

0

0.00

0.00

0.00

0.00

5

0.12

0.11

0.13

0.12

10

0.13

0.12

0.14

0.13

15

0.15

0.13

0.15

0.14

20

0.16

0.14

0.16

0.15

25

0.16

0.14

0.16

0.15

30

0.17

0.15

0.16

0.16

35

0.17

0.15

0.17

0.16

40

0.18

0.15

0.17

0.17

45

0.18

0.16

0.18

0.17

50

0.18

0.17

0.18

0.18

55

0.18

0.17

0.18

0.18

60

0.18

0.17

0.19

0.18

65

0.18

0.17

0.19

0.18

70

0.18

0.17

0.19

0.18

75

0.18

0.17

0.19

0.18

80

0.18

0.17

0.19

0.18

85

0.18

0.17

0.19

0.18

90

0.18

0.17

0.19

0.18

95

0.18

0.17

0.19

0.18

100

0.18

0.17

0.19

0.18

105

0.18

0.17

0.19

0.18

110

0.18

0.17

0.19

0.18

115

0.18

0.17

0.19

0.18

120

0.18

0.17

0.19

0.18

125

0.18

0.17

0.19

0.18

130

0.18

0.17

0.19

0.18

135

0.18

0.17

0.19

0.18

140

0.18

0.17

0.19

0.18

Results for 1.6M

Conclusion

The results tables show that when using the 0.8M acid after 10 seconds the average mass loss was 0.06g.

After 100 seconds there was an average mass loss of 0.16g.

When using the 1.0M acid after 10 seconds the average mass loss was 0.08g.

After 100 seconds there was an average mass loss of 0.17g.

When using the 1.2M acid after 10 seconds the average mass loss was 0.10g.

After 100 seconds there was an average mass loss of 0.17g.

When using the 1.4M acid after 10 seconds the average mass loss was 0.11g.

After 100 seconds there was an average mass loss of 0.18g.

When using the 1.6M acid after 10 seconds the average mass loss was 0.13g.

After 100 seconds there was an average mass loss of 0.18g.

I have drawn a graph of the results showing the mass loss against time for 0.8M, 1.0M, 1.2M, 1.4M, 1.6M acid. I have plotted the points and drawn a line of best fit for each concentration. This has allowed me to work out the rate of reaction at certain times by finding the gradient of the line at this point. The equation I have used for this is:

Gradient = opposite

Adjacent

I have found that when using the 0.8M acid at 5 seconds the rate of reaction was 0.0054gs-1.

When using 1.0M acid at 5 seconds the rate of reaction was 0.0067gs-1.

When using 1.2M acid at 5 seconds the rate of reaction was 0.0085gs-1.

When using 1.4M acid at 5 seconds the rate of reaction was 0.0092gs-1.

When using 1.6M acid at 5 seconds the rate of reaction was 0.0100gs-1.

My results show that the rate of reaction increases as the concentration increases. The higher the concentration the faster the mass loss in grams. My graph shows this because it took 28 seconds for 0.10g to be lost when using hydrochloric acid of concentration 0.8M.

It took 16.5 seconds for 0.10g to be lost when using hydrochloric acid of concentration 1.0 M.

It took 9.5 seconds for 0.10g to be lost when using hydrochloric acid of concentration 1.2M.

It took 6.5 seconds for 0.10g to be lost when using hydrochloric acid of concentration 1.4M.

It took 2.5 seconds for 0.10g to be lost when using hydrochloric acid of concentration 1.6M.

This clearly shows that the higher the concentration the faster the rate of reaction. I have used the graphs drawn to work out the average rate of reaction of each concentration. I have worked this out using the following formula:

Average rate of = total mass loss

Reaction time

I have taken the time as the time it took each reaction to reach the total mass loss. The average rate of reaction of 0.8M acid was 0.0017gs-1

The average rate for 1.0M acid was 0.0021gs-1,

The average rate for 1.2M acid was 0.0024gs-1,

The average rate for 1.4M acid was 0.0026gs-1

The average rate for 1.6M acid was 0.0030 gs-1.

The graphs that I have drawn are curves and clearly show that the rate of reaction is faster at the start of the reaction than at the end. The change in the gradient of the curve shows this. I have worked out the rate of reaction at 5 seconds for the 0.8M acid and found the rate to be 0.0054gs-1. The rate of reaction at 80 seconds for the 0.8M acid was 0.0007gs-1. This shows the decrease in rate as the reaction proceeds. This is because when a reaction first begins there is a high concentration of reactant particles. There are more collisions between the particles and so the rate of reaction is greatest at the beginning of a reaction. As there are more collisions there will be more collisions with energy equal to or greater than the activation energy so the rate of reaction will be faster at the beginning of a reaction. However as a reaction continues the concentration of reactant particles decreases as they have already reacted. This causes the rate of reaction to decrease. This is because there will be less collisions so there will be fewer collisions with an energy greater than or equal to the activation energy causing the rate of reaction to decrease and eventually be 0 when the reaction stops. The reaction will finish when either of the reactants have run out. This pattern is shown on the graphs below.

The rate of reaction increases as concentration increases because a higher concentration contains more particles in a given volume. This is shown in the diagram below.

This is due to the fact that there are more reactant particles in solution and therefore collisions will be more frequent. This means that there is more chance of collisions with energy greater than or equal to the activation energy. Collisions are only effective if they have energy equal to or greater than the activation energy. The activation energy is the minimum amount of energy needed for a reaction to occur. This is because in order for particles to react they must collide with enough energy and in the correct orientation for bonds to be broken. This is because for a chemical reaction to take place, some bonds in the reactants must be broken.

The equation for this reaction is:

Magnesium + Hydrochloric Magnesium + Hydrogen + Energy

Acid Chloride

Mg(s) + 2HCl(aq) MgCl2(aq) + H2 (g) + energy

My conclusion agrees with my prediction because I predicted that as the concentration is increased the rate of reaction also increased so the waste gas, hydrogen will be given off faster at a high concentration than at a low concentration. This is what my results proved to be correct meaning that my prediction was correct.

Evaluation

I used the same conditions for each individual experiment. I used the same volume of hydrochloric acid in each experiment. I also kept the mass of magnesium the same and used the same sized beaker in each experiment. I carried out each experiment 3 times in order to gain more reliable results. However I do not think that my results were that accurate. This is because I used magnesium powder, which meant that the surface area was not constant. Although the mass was the same I did not check that the surface area was exactly the same as this would have been extremely difficult to do. Another problem was keeping the temperature constant. The solution was not heated however the temperature of the room would not have been constant. This is because he experiment was carried out over a number of different days meaning that the temperature would not be exactly the same. Another variation in the temperature was caused by the fact that the reaction is exothermic. The magnesium was weighed out in a boat however it was difficult to get all the magnesium out of the boat and into the beaker of hydrochloric acid. This would have caused a slight variation in mass of magnesium in each experiment, which would have affected the rate of reaction. Also the balance used only recorded the change in mass loss to 2 decimal places. This caused the value on the balance to fluctuate greatly, which made it difficult to record the mass loss.

I do not think that my results are that accurate because even allowing for experimental error when I repeated the results some of the readings were not the same. This is shown on my graphs as I have some points on my graphs that do not quite fit which shows that not all my readings were accurate. I think that this is mainly due to the fact that the balance was only to 2 decimal places as it meant that the same reading appeared many times before the value rose. Also I rounded the averages to 2 decimal places because the readings I took were only to 2 decimal places. This has caused some of the reactions to have the same total mass loss, however this is not accurate as if the results were all to 3 decimal places the total mass loss would alter slightly. Recording the results to 3 decimal places would have improved the shape of my graphs and there would be fewer points that did not fit. Although the results of each experiment are not exactly the same none were sufficiently different to be considered anomalous. Had there been any anomalous results I would have left them out of the average to try and maintain the accuracy of the results.

The experiments for each of the different concentrations did not have the same total mass loss however I have carried out a calculation as shown below which shows that the mass of hydrogen should be the same in each experiment.

Mg(s) + 2HCl(aq) MgCl2(aq) + H2 (g)

Number of Moles of Mg = mass

Mr

= 0.1

24

= 0.00417 moles

Mg:H2

1:1

0.00417:0.00417

Number of moles of H2 = 0.00417 moles

Mass of H2 = number of moles x Ar

= 0.00417 x 2

= 0.00834g

This shows that during the reaction 0.00834g of hydrogen should be formed. To ensure that this mass is constant for all the concentrations used I have done another calculation as shown below.

Mg:HCl

1:2

0.00417:0.00834

This tells me that if there are at least 0.00834 moles of hydrochloric acid then 0.0041g of hydrogen will be formed.

Number of moles of 0.8M HCl = Concentration x volume

1000

= 0.8 x 15

1000

= 0.012moles

This calculation shows that when using the lowest concentration of HCl (0.8M) there are a sufficient number of moles in order for 0.00834g of hydrogen to be formed. This indicates that for each concentration the total mass loss should be 0.00834g. However the lowest total mass loss recorded was 0.16g which is much too large. This shows inaccuracies in the method. I think that there was a greater mass loss than expected because during the reaction there would have been spray causing some of the solution to be lost. This would have caused a greater mass loss and would also explain why each concentration did not have the same total mass loss. When using the 1.6M acid the reaction was more vigorous than that of the 0.8M acid. Therefore more acid would have been lost due to spray with the 1.6M acid causing a greater total mass loss.

If I were to repeat this investigation I would not use the same equipment, as there were many inaccuracies. To try and stop the variation of surface area magnesium ribbon could be used. However this would mean than more magnesium would be needed in order for a reaction to take place because magnesium ribbon has a smaller surface area than magnesium powder. Also to stop the variation in temperature a water bath could be used. This would allow me to carry out the experiment at a set temperature. I would also use a balance that records the change in mass to 3 decimal places. I think that this is necessary because the change in mass was not that large due to the fact that the gas given off was hydrogen.

Hydrogen is very light causing the mass loss to be small. Using a balance to 3 decimal places would record the loss in mass more accurately. I could also have hooked the balance up to a computer causing the results to be recorded electronically. This would make the results much more accurate. I used a measuring cylinder to measure the 15ml of hydrochloric acid, this was because there was not a 15ml pipette available. However using a pipette and pipette filler would have been a more accurate way of measuring the acid.

I think that I have sufficient evidence so support my conclusions however the results could be made more accurate by using a larger range of concentrations or recording the change in mass at a greater number of intervals. This would make the line of best fit more accurate. I could also try to take more readings at the start of the reaction, as this is when the rate of reaction is fastest. This would also make the graph easier to plot. I could also repeat each experiment a greater number of times in order to make my results more reliable. I could extend the experiment by changing the product collected. This could be done by using marble chips instead of magnesium causing carbon dioxide to be collected. This would also make the change in mass easier to record because carbon dioxide gas has a Mr value of 44 whereas hydrogen has a Mr value of 2. I could also investigate other factors that affect the rate of a reaction such as surface area and mass of magnesium, volume of acid and temperature to see if I get similar results.