The activation energy (minimum energy required for a chemical reaction to occur) for the reaction between Magnesium and Hydrochloic acid will be measured, through obtaining values for the rate at different temperatures. From this a graph of ln(k) against 1/T (Kelvin) will be plotted and the gradient will be calculated, which with the Arhenius equation ( explained in the analysis section), the activation energy will be calculated.
The reaction rate for a reactant or product in a particular reaction is defined as the fraction of the chemical that is formed, or removed in moles per unit time per unit volume. The first method that could be used is known as ‘following a reaction’, and some of the different techniques that could be used include:
-
Titration: A mixture is made up and samples are taken out from it, by using a pipette. Then the reaction must be quenched (slowing it rapidly at a measured time from the start of the reaction), by using for example ice, or through the removal of a catalyst. The samples can then be titrated depending on the reaction mixture. Where in this particular case sodium hydroxide could be used.
-
Dilatometry: This method is only really usable when two liquids react together, where the volume of the whole mixture changes slightly, and is measured using a dilatometer. However, this method is not suitable for this investigation as it involves a metal and an acid and the volume is unlikely to change by a measurable amount.
-
Gaseous product: When a metal and an acid react, a salt and gas is given off. In this experiment hydrogen gas is produced (as shown in the equations), and so this could be collected, and its volume or mass could be measured.
Having considered all the possible methods, I have come to a decision that the collection of the hydrogen gas that evolves will be the most appropriate and adequate method to measure the rate of the reaction, and ultimately obtain the order for the reaction as well as an overall rate equation.
The other method is known as ‘initial rates’ where in this particular reaction a large excess of acid would be used so that the quantity of acid used by the time the reaction is finished has very little effect on the overall acid concentration. We can therefore assume that the acid concentration would remain constant. The starting concentration can then be used with a particular time interval associated with the reaction; this could be the time for a certain volume of hydrogen to be produced or the time for the magnesium to all react. If this experiment is repeated over a broad range of concentrations and in each case the time is recorded for the same event to occur, then a whole report of the rate of reaction can be constructed.
With such a reaction the independent variables that could be controlled and varied include the temperature at which the reaction takes place, pressure, use of a catalyst, the surface area of the magnesium ribbon used, the concentration of the acid used, the type of acid used such as mono/di/tri protic, and strong/weak acids.
The dependent variables that could be investigated and measured include, the time taken for a specific volume (e.g. 20cm³) of hydrogen gas to evolve, or for the Magnesium ribbon to react; the change in mass; the conductivity of the solution ( as H+ decrease); and the pH of the solution.
When a variable is changed, all he others must remain constant in order to ensure a fair test, and that I can be certain that only the changing variable is affecting the reaction. For example, if when the temperature is changed, I must ensure that other factors such as surface area, temperature, concentration and type of acid must remain constant and the same is repeated for each experiment.
Risk assessment
To ensure that my experiment will be carried out safely, safety goggles will be on at all stages of the experiment, and protective clothing must be worn at all time whilst conducting the investigation. The practical will be carried out in plenty of space, clear of bags and coats. It is also imperative to consider the safety of other pupils that may be close to the practical. Therefore no one who is not involved in the practical shall be allowed in close vicinity of the experiment. 2M Hydrochloric acid is a corrosive, irritant substance, and so care will be taken to avoid contact with skin and preventing spillages. Care must be taken with all other acids, and when making different concentrations. Although only small volumes of Hydrogen gas will be produced, my experiment will be positioned away from open flames, due to the flammability of Hydrogen and the potential danger this can cause. All experiments were checked by a teacher before they were carried out, to ensure they were safe, and a Risk Assessment form was handed in, and a copy is also attached at the end of this project.
Predictions
In the reaction that I am studying there are only two reactants, Mg (s) and HCl (aq), and as the magnesium is a solid it does not have a concentration so it can be omitted. So the general rate equation above can be rewritten as:
Rate = k [HCl (aq)]
Although, the order of a reaction cannot be calculated from a balanced chemical equation, results from my second preliminary experiment, show that there is some form of order, and a zero order is unlikely for hydrochloric acid.
Phosphoric acid is not a particularly strong acid and is weaker than sulphuric acid and hydrochloric acid. Each successive dissociation step occurs with decreasing ease. Thus, the ion H2PO4¯ is a very weak acid and HPO42 ¯ is an extremely weak acid. I therefore predict that the rate, at which hydrogen gas is produced, will be slower than hydrochloric and sulphuric acid.
For the second part of this investigation, my aim to obtain a value for the activation energy for the reaction between Magnesium and Hydrochloric acid. I will use the Arrhenius equation:
ln k =constant – (Ea / R) (1/T)
Where k is the rate constant, Ea is the activation energy (J mol-1), R is the molar gas constant (8.31 JK-1 mol-1), and T is the absolute temperature (Kelvin). The equation can be compared to the equation of a straight line y = mx + c where y = ln k, and x = 1/T. In a graph of such a line the gradient would be equal to the negative of the activation energy of the reaction divided by the value of R (molar gas constant), and so from this the activation energy can be found. Due to possible errors in this experiment, it is unlikely that a direct straight line will be created, and so a line of best fit will be drawn. The activation energy cannot be predicted through chemical theory alone, however I predict a reasonably linear graph when ln k (y axis) and 1/T (x axis) are plotted, and the gradient of the graph should be negative. I came across a website which had gathered various results from previous students for the activation energy of this reaction, and a mean value of
23 kg mol-1 was given, and I will use this value as guideline of what to expect, and compare my value to it.
The analysis will be sectioned into two separate parts, concerning the two different variables investigated, and then these will later be brought together in the conclusion.
Following the reaction
In all of the experiments in this section, 10 cm³ of 1M acid was used as this was found to be the most adequate volume and quantity to obtain as reliable and accurate data given the laboratory equipment provided. As the reaction progressed, the concentration of the acid decreased, as the acid reacted and was used up. Magnesium was in excess, so we can therefore assume that the reaction will stop when all of the acid had reacted. The volume of hydrogen at t final, which is V final, will depend on the initial concentration of the acid.
Example: If [HCL] initial is high, V final will be large
If [HCL] initial is low, V final will be small
At anytime, the concentration of acid remaining [HCL]t, will be proportional to V final-Vt.
Measuring the Activation Energy by varying the temperature
In this investigation my aim was to calculate a value for the activation energy between Magnesium and the monoprotic acid Hydrochloric acid. This requires the use of the Arrhenius equation: ln k =constant – (Ea / R) (1/T). I will therefore plot a graph with ln rate on the y-axis, and 1 / temperature (per Kelvin), and should hopefully obtain a fairly straight line with a negative gradient, and it is this gradient that will have to be calculated.
From the graph on the right, we can calculate the activation energy. Given that the gradient of the line is equal to –Ea/R, where is the molar gas constant 8.31 Jk-1mol-1.
Ea = -gradient x 8.31
The gradient can be calculated from choosing two points (preferable furthest apart to reduce error), and then it’s the change in y/change in x. From the graph the value for the gradient = -2362
Ea = -(-2362) x 8.31 = 19628.22 J
= 19.63 KJ (2.d.p) - this is the minimum energy required for the chemical reaction to take place. This is just an approximate value, as errors are introduced due to the equipment that was used, such as the thermometer and stop watch, and human error in determining when the solid magnesium strip had fully dissolved. This value is in the range of the values obtained by reordered by other students, and is close to the mean value of 23 KJmol-1.
Conclusion
In this final part of the report I will bring together all of the data that I have analysed, and see whether I have achieved my aims outlined in my plan. Firstly, the orders and the rate equations for the different types of acids, hydrochloric, sulphuric, and phosphoric acid at room temperature were obtained. The hydrochloric acid reaction gave a first order reaction, where the concentration vs. time graph showed exponential decay, the half life seemed fairly constant, as it always took approximately (given the possible errors) the same amount of time for the concentration to halve.. For the diprotic, Sulphuric acid I found that the order was likely to be second due to the half life not being constant, it increased as the reaction went on. As for the phosphoric acid, it was unclear whether it was first or second order. However, having drawn first order plots it seemed that I as more likely to be a first order reaction.