Prediction:
The background investigation that was performed indicates that the weak acid will have the highest activation energy. This is due to the fact that energy is needed to dissociate the ethanoic acid into ethanoate ions. A lower degree of energy is needed to split the hydrochloric acid into hydrogen and chlorine ions as the strong acid has already fully dissociated. This theory can be supported by Ka values found in the book of data. Therefore, the assumption is that the activation energy for the weak acid will be higher.
Choosing an Appropriate Metal:
In order to examine the kinetics in a reaction an appropriate metal must be chosen. There are two choices of metals in this experiment, Magnesium and Zinc. The metals come in different forms, for example Zinc comes in a granulated and foil forms unlike Magnesium, which is available in a ribbon.
The reactants in the reaction are all in different states. This is called a heterogeneous system. Generally, the smaller the size of reacting particles, the greater the total surface area exposed for the reaction. Therefore, the reaction will take place quicker and the availability of the reactants and the surface area exposed is a factor, which must be kept constant throughout the experiments.
There is a high variability of the surface area exposed of granulated zinc and zinc foil. As a result is not possible to keep a constant amount of surface area in the experiments. Each piece has a different thickness and size and therefore the number of moles in the metal will vary. As the amount of surface area of the metal exposed varies, the amount of acid reacting to the surface of the metal will also vary, producing unreliable results and inconsistencies. This suggests that granulated zinc and zinc foil are unviable options.
Magnesium ribbon will be more suitable for the investigation as it has a constant width, thickness, mass and surface area exposed. As a result there will be the same number of moles per given length. Magnesium ribbon is also better than its contenders (granulated Zinc and Zinc foil) on a practical level, as it is easier to measure and cut more accurately, which will produce more reliable results. Another advantage is that magnesium is higher in the reactivity series in the periodic table than Zinc; therefore the reaction will be quicker.
The disadvantage of using Magnesium ribbon is that the Magnesium has already reacted with Oxygen from the air forming a thin layer of Magnesium oxide on the exposed surface of the ribbon. This can be overcome by rubbing sandpaper on the metal to remove the thin layer of magnesium oxide. This will help ensure that the results of the experiment are reliable and ensure that the acids react with the magnesium and not the formulated oxide layer.
The Acids:
Choice of acids and their varying strengths was based upon the Bronston/Lowry theory. According to the Nuffield Advanced Chemistry Students’ Book p338, acids that ionise completely at moderate dilutions are called strong acids. Consequently they tend to have a large Ka (dissociation constant) value.
Weak acids interact very little with water therefore do not easily dissociate. Ethanoic acid is an example of a weak acid as it dissociates very few hydrogen protons to bases.
Ref: Nuffield Advanced Chemistry Students Book, Longman, p338.
Hydrochloric acid is an example of a strong acid. This is because it completely donates its protons (H+). Dissolving hydrochloric acid in water will in turn formulate an aqueous equilibrium as indicated in the equation below:
HCl (aq) ⇌ H+ (aq) + Cl –(aq) almost completely dissociated
In the reaction above the H+ in the hydrochloric acid acts as the acid and the Cl- acts as the conjugate base. The HCl loses its protons readily; therefore a high concentration of H+ ions is produced when the reaction reaches its homogenous equilibrium.
The dissociation constants in the reaction is equivalent to the equilibrium constant Kc However, the difference between Kc and Ka is that Ka is not influenced by changes in the concentration, but it is influenced by changes in temperature. This enables us to accurately measure the extent to which the acid is dissociated (strength of the acid). Hydrochloric acid is a strong acid and therefore it will have a greater extent of dissociation. Hence, there will be a greater amount of [H+ ] protons in the acid and
[Cl –] ions in the conjugate base. This provides a larger Ka value.
The Ka value is the dissociation constant for the acid.
= [H+ ] [Cl – ]
[HCl (aq)]
= 1.7 x 10-5mol dm-3
The Ka value for hydrochloric acid is not available in the book of data. This is primarily due to the fact that that the acid is very strong, which indicates that it will dissociate into hydrogen and chlorine ions.
CH3COOH (aq) ⇌ H+ (aq) + CH3COO–(aq) only partially dissociated
When using both acids we will be able to analyse the comparisons in the different rates of the reaction. This will in turn help determine whether the amount of protons donated in the reaction will influence the amount of collisions made and therefore whether it will increase the rate of the reaction.
Reactions between ionic compounds only work in the solution because the ions are mobile and able to interact. The balanced equation for hydrochloric acid is as follows:
2HCl (aq) + Mg (s) → MgCl2 (aq) + H2 (g)
Ignoring the spectator ions, which are the (Cl-) ions.
2H+ (g)+ Mg(s)→ Mg2+ (aq)+ H2 (g)
The mechanism of the reaction is a description of the sequence of steps, a form of breaking the reaction down into a number of steps. The rate of the reaction depends on the rate of the slow step, which is the rate-determining step. Therefore, in order to increase the rate of the reaction, the speed of the rate-determining step must be increased. This can be achieved if reactants in the rate-determining step are first or second ordered. This will change the amount of the concentration of reactants, which will increase the rate in which products are produced for them to be used in the fast step, increasing the rate of the whole mechanism.
A proposed mechanism for the reaction is shown in the sequence below:
(Rate determining step)
(Slow step)
Mg (s) + H+(aq) → Mg+(s) + H•
(Rate determining step)
(Slow step)
Mg+(s) + H+ → Mg2+ (aq) + H•
(Fast step)
H•(g) + H•(g)→H2 (g)
The mechanism suggests that the reaction consists of two rate-determining steps. Each rate-determining step has high activation energy in the reaction. Therefore it determines the overall rate of the reaction. This suggests that the order is second, as indicated in the indices below:
So Rate= k [H+ (aq)] 2
The magnesium ribbon doesn’t feature in the rate equation because it is a solid. This may occur because the hydrogen ion appears in both rate-determining steps. This is because only liquids, gases and ions have a concentrate. The chlorine doesn’t appear in the rate equation. Any step involving chlorine must come after the rate-determining step. It is reasonable therefore to postulate that the first step involves a proton with magnesium. This is because both the species appear in first order terms in the rate equation. The second step controls the rate of the reaction. It therefore contains the most activation energy. The last step in the mechanism reacts rapidly forming hydrogen gas.
The hydrogen molecule appearing twice in the rate determining step indicates that the reaction is second order in respect to the acid, suggesting the result of the graph will be of second order.
The mechanism above shows that the hydrochloric acid is an aqueous solution. Therefore the concentrated solution contains water as does the equation, which means that the chloride ions and the hydrogen ions in the acid have already dissociated. This may imply that the hydrogen ion in the acid reacts with the solid magnesium ribbon. From the magnesium ion an electron is transferred to produce a hydrogen free radical.
“A free radical is an uncharged attacking group with an odd number of electrons, so they possesses only one of the electron pairs needed for the formation of a new covalent bond.”
Ref: Nuffield advanced chemistry students book, (Longman) p185
This reaction is fairly slow in comparison to the rest; this is due to the fact that the hydrogen ion has to react with the larger magnesium ion. The hydrogen ion then gains another electron from a magnesium ion to form a hydrogen free radical. The magnesium ion has a 2+ charge. This stage of the reaction is called propagation. The rate-determining step has two positively charged reactants, which repel one another.
After this stage of the reaction the hydrogen free radicals end up colliding against each other. This terminates the reaction to produce hydrogen gas, a process known as homolytic fission. This reaction occurs at high speeds due to the fact that the free radical molecules contain a set of unpaired electrons forming an attacking group. These electrons readily react so that the molecule can produce a full shell that will then become inert and stable.
As both molecules contain an odd number of electrons in their outer shell they readily form covalent bonds in which the electrons are shared between the nuclei.
Ethanoic acid is also readily available and is a weak acid; these were strong factors for its choice.
The balanced equation for the weak acid (ethanoic acid) and the magnesium ribbon is shown below:
2CH3COOH (aq) + Mg (s) →(CH3COO-)2 Mg2+ (aq) +H2 (g)
Ignoring the spectator ions
2H+(aq) + Mg→ Mg 2+ (aq) + H2 (g)
The mechanism in this reaction will be the same as the one for hydrochloric acid, as indicated above. This is because the spectator ions (Cl-) and CH3COO- are omitted in the ionic equation, therefore they can be ignored.
Preliminary Experiments:
The aim of the preliminary experiment is to formulate an accurate method in conducting the final experiment. Three experiments will be conducted which will help ensure the validity and reliability of the results. In this experiment the variables, which may contribute to inaccuracies in the final results, will be explored. The aim of the preliminary experiments is to determine the following:
- Whether the oxide layer on the magnesium makes a difference in terms of the rate of the reaction.
- Suitable range of concentrations to use.
- Suitable temperature range to use.
- The type of metal which is going to be used
- If the experiment is exothermic.
The Volume of Acid:
In the preliminary experiment we are going to use a 10.0cm3 of acid. This is because it is a standard measurement, which can be accurately and easily measured using a pipette. The volume of acid is in excess in comparison to the length of magnesium ribbon. Therefore, it will ensure that the reaction will be completed methodically and successfully. As the reaction is exothermic the excess volume of acid will act as a temperature regulator absorbing any heat produced.
Apparatus for the Preliminary Experiment:
The apparatus chosen for the preliminary experiment was an important consideration because it would ultimately determine the selection procedure for the apparatus, which would be used in the final experiment.
When conducting the first preliminary experiment it was concluded that using a beaker gave far too much surface area making the reaction quicker. In the second experiment it was then decided to place the acid in a boiling tube that has a diameter of 2.0 cm. The main advantage in using the boiling tube is that it provides an increase in depth ensuring that the reaction takes place.
When the physical apparatus had been selected, a suitable acid and concentration of that acid was needed, both strong and weak. The three contenders for the choice of strong acid were Sulphuric, Hydrochloric and Nitric acid. A 1 cm length of magnesium ribbon was placed into 10cm of acid in a 1.00M solution, at room temperature. Each experiment was timed in seconds for the metal to be displaced.
The results of the experiment are shown in the table below:
The results table above shows that at room temperature the nitric acid is very slow. This acid will not be viable as it is too time consuming. The sulphuric acid has the quickest time; this implies sulphuric acid has the highest rate of dissociation. In contrast to nitric acid, sulphuric acid reacts too quickly with the magnesium ribbon, which means it cannot be measured accurately. It can be finally observed that hydrochloric acid is the most suitable for the experiment as it completely reacts with the magnesium in the most suitable time.
Ethanoic acid is the only weak acid available; therefore it will be used to draw comparisons to investigate the difference between strong and weak acids.
The experiment was initially conducted with a 2.00M solution of hydrochloric acid at room temperature and at 60.0oC. At 60.0 oC the reading was 18.22 seconds. The intense speed of the reaction proved a disadvantage as it made the experiment hard to control. It was then decided against using this concentration. This could have been overcome by using a higher volume of acid e.g. 20.00cm3. This is an awkward measurement to use with a standard sized 10.00 cm3 pipette, as it would have to be measured twice and placed into the test tube. This will lead to more inaccuracies in the experiment. It was finally decided to use a concentration of 1.00 molar.
Choosing the Length of Magnesium Ribbon:
To choose the appropriate length of magnesium ribbon several different lengths were to be tested in a reaction with 10.00cm3 of 1.00M acids at room temperature (21.0 oC). The following results were obtained:
The results above show some anomalies, for example the 1.50 cm reacted in a longer time than with the 2.0cm. This is inconsistent with the background research that was conducted because it suggested that a longer length of magnesium ribbon would result in a slower rate of reaction.
The table below shows the time it took to displace the various lengths of magnesium at 1.00M solution of hydrochloric acid at 60 oC.
It is clear from the table above that the timings produced with the 1.00M of solution are manageable to use.
A 2.00cm strip of magnesium ribbon will be used in the final experiment, as it is easier to measure out and to cut practically. The experiment shall have set temperatures that will range from 20.0 oC to 60.0 oC in ten-degree intervals. The highest temperature will be 60.0 oC because as observed in the background research, any temperature above this would lead to the acid evaporating. Each temperature should be repeated twice in order to rule out any anomalous results. This will then improve the reliability of the experiment.
Quantities Used in the Final Experiment:
The following quantities will be used in the final experiment:
10cm of magnesium ribbon weighed = 0.07g
Therefore 2.0 cm weighs 0.014/5 = 0.014g
Number of moles of magnesium strip = mass/Ar
= 0.014/24
= 5.8 x 10-04 moles
It is important to note that the volume of acid used in the experiment is an important factor, which must be kept constant. This is because in theory if more magnesium ribbon was added to excess in the system, eventually there will be a blockage point in which the magnesium 2+ ions will find it difficult to get away from the active site. These will in turn slower the rate of the reaction. This is called crowding effects.
Using 10 cm3 of 2 molar acids
Number of moles = (volume/ 1000) x concentration
= 10/1000x 2
= 0.02 moles of acid
Ratio of acid to Mg = 0.02: 5.8 x 10-4 = 34: 1
It is also clear that the number of moles of acid is greater than the number of moles of magnesium. This indicates that the volume of acid used is in excess because it is five times larger. This will add an advantage to the experiment, as it will ensure that the reaction on hand is completed. Also the excess acid will absorb any heat evolved, as it is an exothermic reaction because bonds are being broken.
The volume of hydrogen produced = number of moles x 24000cm3
= 8.3 X 10-4 moles x 24000
= 17 cm3 of hydrogen gas evolved.
The Variables:
The dependent variable in the reaction is the temperature varying from the first experiment. This will determine the activation energy of the reaction. Another factor to consider is that the magnesium ribbon has already reacted with the air, forming a layer of magnesium oxide. Before each experiment the magnesium ribbon was rubbed ten times in order to remove the magnesium oxide layer and leave the pure metal. The number of strokes will be kept constant as too many can decrease the mass of the magnesium, therefore decreasing the number of moles which may decrease the rate of the reaction. Also the same concentration of acid is needed for the first experiment.
The water, which will start at room temperature, will absorb some of the excess heat of the exothermic reaction to keep the temperature the same.
Apparatus:
The following diagram shows the apparatus, which was used for all the experiments:
Final Practical Procedure:
When conducting the final experiment it was important to note that any change to the compiled method (compiled from the preliminary experiment) would result in inaccurate results. The final experiment was conducted using the following method.
- Arrange the apparatus shown in the diagram.
- Cut a 10.0 cm strip of magnesium ribbon
- Stroke x10 with sandpaper over the magnesium ribbon
- Measure out 2.0cm and mark with a pencil. Then cut the magnesium ribbon
-
Fill the beaker 200.0 cm3 with tap water halfway
-
Place 10.0 cm3 using a pipette of the required acid i.e. HCl or ethanoic acid in a boiling tube
- Place the boiling tube containing the acid in the beaker of water.
- Place a thermometer in the acid
- Heat the system using the Bunsen burner until the required temperature and stop when the temperature is accessed
- Remove the Bunsen burner to avoid over temperature rising above the required temperature.
- Place the magnesium strip immediately after removing the heat
- Immediately start the stop clock. And stop when the magnesium disappears via effervescence.
- Repeat the process three times for every temperature required
- Place results in a table.
The second experiment on hand focuses on the order of the reaction as the concentration of the acid is varied.
The order of the reaction with respect to the reactant can be found and analysed using an initial rate method. This will be done by analysing the effect of concentration of the reactant on the rate of reaction. In this experiment the reactant will be is hydrochloric acid. The method used will be similar to the first experiment however the temperature will be kept constant.
Rate = k [X]a
If Ln is removed from each side of the equation, Ln Rate = a Ln X + Ln constant.
A graph with Ln Rate on the vertical axis and Ln [H+] on the horizontal axis will be produced, where the gradient will be the order of the reaction.
The second experiment will make use of the same acids as the previous experiment, where ethanoic acid is used for the weak acid and hydrochloric for the strong acid respectively. In this experiment with a constant temperature and a variability of the concentration of hydrochloric acid, it is likely that the hydrochloric acid will produce faster timings because it is a stronger acid.
The rate equation and procedure will be kept the same as the previous experiment. Also the errors are the same in the experiment. The temperature of the experiment must be kept constant throughout. Placing the system into a water bath to maintain room temperature will do this. Also using excess in terms of volume of acid will help maintain a constant temperature and concentration, for the same reasons as the fist experiment.
In conducting the second experiment, firstly the rate of the reaction will be calculated using the above method. Secondly, the experiment will be repeated for each concentration used in order to produce more results so an average can be taken. From the compiled results a graph will be constructed with a line of best fit. The graph will be used to determine the gradient of the line by Ln rate / Ln concentration. This will then in turn define the order of the experiment with respect to the acid used.
From the order of the reaction it is possible to predict a likely mechanism for the reaction. It is likely that order will be of second order and therefore the graph will be a curve. The curve will be relatively deep and the half-life won’t be constant but will increase as the reaction proceeds.
Rate=k [H+(aq)]a
Ln rate=a Ln [H+ ] + Ln k
Y= mx + c
The indices ‘a’ represent the order of the reaction, which is likely to be second. If this prediction is correct then it suggests that the reaction consists of two rate-determining steps, which are slow steps. Each rate-determining step has highest activation energy in the reaction. Therefore it determines the overall rate of the reaction.
Prediction:
As the hydrochloric acid has pre-dissociated previously, increasing the concentration of solute will not have an overall effect on the actual order of the reaction.
Section 2: Implementation
The following shows the results obtained from both experiments compiled into a tabulated format so they can be analysed:
Investigation 1 (10.00cm 3 ) 1.00M hydrochloric acid (strong acid) + magnesium ribbon:
Investigation 1 (10.00cm3) 1.00M Ethanoic acid (weak acid) + magnesium ribbon:
Investigation 2 (10.00cm3) of Hydrochloric acid in various concentrations:
Section 3: Evaluation & Conclusion
Calculations:
The graphs produced earlier help in determining the activation energy of the reaction. This can be done by using the following equation and applying it to the graphs on produced:
Ln (rate of reaction)= K – EA (1/T)
R
As with previous graphs axis, the (1/T) in Kelvin has been plotted on the Y-axis and the k (Ln rate) has been plotted on the X-axis. Ln K will then determine the gradient of the graph / (1/T). This is equivalent to – EA
In order to determine the activation energy the gradient found from the graphs must then be multiplied by the gas constant R that is 8.31. Therefore for the first investigation using hydrochloric acid and magnesium ribbon the activation energy can be worked out by
Experiment 1 (a) (strong acid):
The results for the reaction between magnesium ribbon and the strong acid (hydrochloric acid) was hand drawn previously. Following the procedure outlined above to find the calculation, it was then possible to work out gradient and therefore the activation energy of the strong acid. The calculation is outlined below to find out the activation energy:
Gradient = change in y / change in x
= -0.800/ 2.65 X 10-4
= - 3018. 87 (2dp)
The gradient of the graph is a negative number this represents that the reaction is an exothermic one as bonds are being broken.
Experiment 1 a (hydrochloric acid) activation energy calculation
The following shows the calculation to calculate the activation energy for the reaction between hydrochloric acid and magnesium ribbon:
Activation energy =gradient x R (gas constant)
= -3018.87 x 8.31
=- 25086.81(2dp) J mol-1
= -25086.81/ 1000
= +25.08 kJ mol-1
Experiment 1b (ethanoic acid) activation energy calculation
For the reaction between ethanoic acid and magnesium ribbon, the activation energy can be found by using the same method as above
Activation energy =gradient x R (gas constant)
=--(-4980) x 8.31
=+41383.80 J mol-1
= +41383.80 /1000
= +41.3838
= +41.40 kJ mol-1
It is now possible to compare and contrast the data from the activation energies calculated above to reach some conclusions about the experiments.
The activation energy for the weak acid was stronger than that of the hydrochloric acid. The ethanoic acid (weak acid) partially dissociated, so there were fewer hydrogen molecules to react with the magnesium ribbon.
Le Chatelier’s principle can bring balance to the argument in that
“ The position of the equilibrium shifts in direction which opposes the change in the conditions”.
Ref: Nuffield Advanced Chemistry Students’ Book, p167
Therefore the ethanoic acid will dissociate hydrogen ions, which are used up in the reaction with the magnesium ribbon. More hydrogen ions are made to compensate. This would require energy for an optimum amount of dissociation, producing higher activation energy.
In the system, (hydrochloric acid (strong acid) and magnesium) the concentration of hydrogen ions decreased. This is due to the fact that the hydrogen ions fully dissociate in the solution. Therefore lower activation energy is required. Therefore there is less energy needed to break the hydrogen chlorine bonds. In a weak acid the hydrogen ion is used in the reaction with the magnesium ribbon. Therefore more hydrogen ions are produced to compensate.
Experiment 2 (order of reaction):
The order of the reaction is determined by the gradient of the graph, which is 2.246 Shown in the graph above. This indicates that there are two slow steps (rate determining steps) in the mechanism for hydrochloric acid and magnesium ribbon.
Therefore the rate equation for the reaction is
Rate = [H+] 2
The main disadvantage of this result is that the order of the reaction should have been of 2.00. This would have given a perfect second order. However, the results obtained did not achieve this. The result obtained in the reaction was 2.246- 2.00 = 0.246.
This was due to the temperature fluctuation in the system, which could have been prevented if temperatures were observed more frequently. However the water bath and the larger volume of acid didn’t completely absorb the excess heat, this lead to inaccuracies.
This may be improved by using a larger volume of acid and a larger water bath to maintain the temperature.
Errors in the Experiment:
As with all experiments, limitations can influence the validity of the results. It was important therefore to examine any limitations and determine ways in which they could be overcome.
When conducting the experiment it was important to ensure that a constant length and thickness of magnesium ribbon was used. This proved to be a key factor in the validity of the results obtained from the experiment. It is important to understand the consequences of failing to keep a constant length of magnesium ribbon in the experiment i.e. smaller than 2.0cm. This would in turn have decreased the rate of the reaction, as there are fewer atoms available on the solid to react with. In order to overcome this problem, a 10.0 cm strip of magnesium was cut prior using a ruler and mark off 2.0 cm with a pen. The Magnesium was measured with a ruler to 1 mm in 20 mm.
Error = 1/20 x 100 = 5.0%
As discussed earlier, the magnesium ribbon would have already reacted with the air to form a layer of magnesium oxide, slowing the rate of the reaction (as the acid will have to react with the oxide layer before the actual element itself). The layer of oxide may vary in thickness, increasing or decreasing the rate of the reaction according to the thickness of each piece of magnesium. This problem was overcome by using sandpaper to stroke the ribbon an equal number of times so that its mass did not vary significantly.
The boiling tubes used in the experiment were cleaned using deionised water to remove any contaminations. This ensured consistency and more accurate results throughout all experiments so that fair conclusions could be drawn.
The volume of acid used in each experiment was measured accurately to avoid any potential errors in the results. It is important to understand that for example if a smaller amount of acid was used, there would be less molecules of the acid to react with the magnesium. This will in turn slow down the rate of reaction, as there would be fewer molecules per second colliding with the active sites on the magnesium ribbon resulting in less activated complexes formed per second. Therefore fewer products formed, slowing down the rate of reaction.
This may have influenced the validity of the experiment at any given temperature on the comparisons made on the hydrochloric and ethanoic (acetic) acid. Therefore the exact effect cannot be calculated.
- Measuring the volume of acid:
The volume of acid was measured accurately using a 10.0 cm3 graduated pipette. This form of measurement is more precise then using a measuring cylinder. The pipettes are accurate to 0.05 cm3. Therefore, the pipette used in the experiment had an error of +/- 0.05 cm in every 10.0 cm3
The percentage error in measuring the volume of acid was:
0.05 x 100 = 0.50%
10.0
The volume of acid used in the experiment will be in excess in comparison to the amount of solid magnesium used. As a direct consequence the concentration of the acid and temperature was maintained in the experiment. This ensured the reaction was completed with no disturbance to the final calculations on the activation energy.
The concentrations of acid are made to the nearest 0.01 mol dm -3.
Error for 1.00 M acid = 0.01 / 1.00 x 100 = 1.00%
The temperature changes were observed by placing a thermometer in the system (a thermometer consists of graduations of 1.0 oC). The effect of temperature in the experiment was an important element, which must be considered. This is because it may have influenced the rate in which an activated complex was reached. If at room temperature (25 oC), the reading is more than 24.5 oC but less then 25.5 oC, an error of +/- 0.5 oC will be present. Therefore the percentage error for the reaction is:
0.5 x 100 => Error = 2.0%
25
The stopwatches used in the experiment have an accuracy of 0.01 seconds. The time in which the magnesium ribbon disappeared via effervescence was determined visually. The error in judging when the magnesium had disappeared was larger that the error from the stop clock. As a result this error can be ignored.
Another factor, which had to be kept constant, is the amount of surface area showing on the magnesium. If the boiling tube had slanted then there would have been more surface area to volume ratio. Therefore, there would be more reaction surface on the magnesium, increasing the rate of the reaction. This was overcome by holding the boiling tubes upright.
Total 8.50%
The total percentage error = 8.50% approximately
Percentage errors for the activation energy
The activation energy can be compared and contrasted to find out the minimum value and the maximum value for both values
The activation energy
In order to find the percentage error of each calculation the following calculation must be done;
Hydrochloric acid
25.08/100 x 8.50= 2.1318 kJ mol-1
Ethanoic acid
41.40/100x 8.5 = 3.519 kJ mol-1
The error calculation for the order in investigation 2
Hydrochloric acid
Order = 2.24
2.24/100 =0.0224 x 8.5%
= 0.1904
= 0.2
Sources of Error:
The error bars on the graph 1.1b (weak acid) and graph 2 (order) consist of error bars. These bar show the potential errors of an 8.5% probability in the results produced.
This can give a better indication of how far the activation energy has been affected due to potential experimental error.
The investigation consisted of sources of error in terms of experimental measures. This has an effect on the following:
- The magnesium oxide on the magnesium ribbon was removed by stroking the magnesium using sandpaper. Although the ribbon was stroked the same number of times the mass of the magnesium may have altered in terms of mass on each piece, influencing the rate of the reactions taking place. For example if the oxide layer were thinner on a piece of magnesium, it would have required fewer strokes. The overall effect of this occurring is that the magnesium strip would disappear slower then in other reactions producing inconsistent results.
- The amount of magnesium ribbon varied in each reaction in terms of how much the amount of surface area to volume ratio was available for the acid to react with.
- Another source of error was that it was hard to distinguish the precise timing in which the magnesium disappeared. This was because there where always excess bubbles consisting of hydrogen gas, distorting the observation and recordings of when the magnesium ribbon actually disappeared i.e. when the reaction took place
- The reflex action of starting the stop clock and placing the magnesium ribbon into the appropriate acid. This can influence the rate of reaction and therefore the activation energy calculated.
Possible improvements for the investigation
As with all experiments there is scope for improvement. The following shows the ways, which the experiment could be improved if it was conducted again:
- A wider range of temperatures could have been used in the experiment. This would have produced a more accurate line of best fit causing a more reliable gradient produced. This will then lead to more reliable calculation of the activation energy.
- Another source of error, which could have been improved on, is the fact that the reflex time would have been different in each experiment. This could have been overcome by using one person to time the experiment while another observes the reactant disappearing.
- Another improvement, which can be made to the investigation, is to weigh the magnesium ribbon after cutting it into 2.00cm strips. A consistent mass will mean consistent and reliable results. This will help ensure the validity of the experiment
- Another improvement for the activation energy investigation is to place a extra thermometer in the water bath as well to maintain accuracy in terms of temperature.
