An Experiment to Investigate the Kinetics of the Reaction of Magnesium with Acid
Jason Kotiah 7/9/2000
An Experiment to Investigate the
Kinetics of the Reaction of Magnesium with Acid
Aim
The aim is to find out the kinetics and determine the rate of reaction of magnesium with hydrochloric acid. To do this we have to measure the amount of hydrogen gas being produced within a given time and from this, the order of magnesium and HCl(aq) can be worked out. Also the activation energy can be worked out, by increasing the temperature of the HCl(aq) and from this, measure the amount of hydrogen produced within certain time intervals.
Background Theory
The collision theory is a model, which has been used to account for the dependence of rate of reaction on temperature. The collision theory involves reacting particles such as HCl(aq) moving towards each other such as magnesium and colliding in such a way that bonds are broken and new bonds formed. In many reactions between gases, it is the actual collision, which controls the rate of the reaction. For a collision which actually results in a reaction, the kinetic energy possessed by the colliding particles of HCl(aq) and magnesium, must be more than a certain minimum energy, E e.g. a mole of colliding particles at a temperature T, the rate of reaction can be found from:
ln(rate) = ln(collision rate) - EA/R (1/T)
where R is the gas constant, 8.314J K-1 mol-1 and EA is the activation energy. To find out the order of a reaction such as HCl(aq) (A) and magnesium (B), experimental results can be analysed to show that:
Rate = K[A]m [B]n
This is a rate equation. The power to which the concentration of A is raised in the experimental rate equation is called the order of a reaction with respect to A. The order with respect to A is m and the order with respect to B is n. So therefore the overall order is the sum of the individual orders is m + n in this case. Once these orders are known, the rate may be predicted at any concentrations of A and B.
The activation energy (Ea) of a reaction between magnesium and HCl(aq) is the minimum energy required by the magnesium and the HCl(aq) particles for the reaction to occur. The Arrhenius equation predicts that the rate constant, k, is related to temperature by:
lnk = ln A-Ea/R(1/T)
The second term involves the activation energy Ea, the gas constant R and the thermodynamic temperature T (in Kelvin). Each reaction has a particular value of A and Ea. Although the Arrhenius equation seems complicated, it is actually the equation of a straight line (y = mx + c), where ...
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The activation energy (Ea) of a reaction between magnesium and HCl(aq) is the minimum energy required by the magnesium and the HCl(aq) particles for the reaction to occur. The Arrhenius equation predicts that the rate constant, k, is related to temperature by:
lnk = ln A-Ea/R(1/T)
The second term involves the activation energy Ea, the gas constant R and the thermodynamic temperature T (in Kelvin). Each reaction has a particular value of A and Ea. Although the Arrhenius equation seems complicated, it is actually the equation of a straight line (y = mx + c), where y equals ln rate, c equals the constant (-8.31Kjmol). As a result the gradient of the line graph is equal to -Ea/R therefore to find he activation energy, in joules, you multiply the gradient by the constant, R which is equal to
8.314J K-1 mol-1. So, to find the EA of this reaction it is necessary to find the rate of reaction at varying temperatures.
Rate of reaction is proportional to 1/time taken to reach a fixed quantity of product. So, we can find this by timing how long it takes for the reaction to reach a fixed point. Instead, we can use several time intervals (points) and take the ln rate (1/time for reaction) and record temperatures in Kelvin (Celsius +273) and use this to find 1/T to draw the graph. Then, using the equation EA = -gradient x R, I can find EA. This graph below shows how the energies of particles are distributed. It shows the number of particles in each small range of kinetic energy. The fraction of particles with energy greater than EA is equal to the shaded area divided by the total area:
Experiments (orders) such as this one show that the rates of most reactions can be related to the concentration of individual reactants by an equation of the form:
Rate = k(X)n
This expression, in which X is the reactant (HCl(aq)) and n is usually 0,1 or 2, it is known as rate equation or rate law. The value of n gives the order of the reaction. When n = 0, the reaction is said to be zero order with respect to HCl(aq):
Rate = k[HCl(aq)]0
But, since [HCl(aq)]0 = 1,
Reaction rate = k
When n = 1, the reaction rate is proportional to [HCl(aq)]1 and the reaction is said to be first order with respect to HCl(aq).
When n = 2, the reaction rate is proportional to [HCl(aq)]2 and the reaction is said to be second order with respect to HCl(aq).
This graph shows the variation of reaction rate with concentration for reactions, which are zero, first and second order.
Overall Preliminary Work
The basic principle of this experiment was to add magnesium to HCl(aq) and then measure the rate of the reaction at different time intervals or finding out how long it took for the reaction to finish. To start off the preliminary stages of this experiment, I had to find out what form of magnesium to use: - ribbon, filings or powder. These types of magnesium have different surface areas and therefore some might react with HCl(aq) quicker or produce the same amount of hydrogen in a quicker time. To find out what type of magnesium I was going to use, I reacted 0.3grams of each type of magnesium with 10cm3 of 2M HCl(aq).
Type of Magnesium
Time taken to react/sec
Ribbon
29
Filings
4
Powder
24
After reacting these various types of magnesium I decided to use filings because:
* It reacted the quickest overall therefore I only have to reduce the concentration of HCl(aq), so the reaction between magnesium and HCl(aq) can go on for a longer time span, until all the reactants have been used up.
* Also, the time taken for the reaction to finish was reasonable-therefore the reaction is slow enough to calculate rates with higher temperatures (for activation energy calculations) but fast enough to allow enough time for other reactions which needed to be done.
I found that out, that using ribbon was more time consuming to use due to measuring it out then cutting it into your specified length. Using magnesium filings as a much easier process as you only had to weigh your specified amount on the digital scales, but its probably not as accurate as you may not be able to attain the same weight (grams) of magnesium after each experiment, (i.e. + or - 0.1g). The main factors that affect these experiments are: Magnesium surface area, HCl(aq) volume (below excess), HCl(aq) concentrations, magnesium mass/length and temperature.
As for my preliminary experiments, I tried using various amounts of magnesium with 10cm3 of 2M HCl(aq):-
0.6g-the hydrogen gas produced went off the 100cm-gas syringe scale
0.1g-the amount of hydrogen produced was too slow
0.3g-the amount of hydrogen produced did not exceed the scale and happened within a 60-second time limit.
I added 0.6g of Magnesium filings to 2M Hydrochloric acid (the highest concentration available in all experiments) and measured how much Hydrogen gas was produced and how long it took to produce it. The apparatus is shown below:
Within seconds of the start of the reaction, the amount of Hydrogen gas produced by the reaction went off the 100cm gas syringe scale. As a result I reacted smaller amounts of Magnesium with 2M HCl(aq).
If I were to use 0.1g of Magnesium for the activation energy experiments, then I would have to react this amount of Magnesium at much higher temperatures then room temperature. This would cause problems as the Magnesium would react so quickly that the rate would be near infinity (rate = time). Even with lower concentrations than 2M Hydrochloric acid the problem of not getting results of a large enough time range still remains. Also the amount of Hydrogen gas produced was very small, so to get an end point for all of the reactions (e.g. how long it took to collect 10cm or 20cm) would be impossible. As a result, I chose 0.3g of Magnesium filings. This amount of Magnesium would be used to find both the activation energy of the reaction between Hydrochloric acid and Magnesium and also the order of HCl(aq) in this reaction.
The Order of Hydrochloric acid in this reaction with Magnesium
Aim: to find the order of hydrochloric acid with respect to its reaction with Magnesium. It is not possible to find the order of Magnesium, because Magnesium is a solid and therefore it's concentration does not change during the reaction.
Key variables: concentration of HCl(aq)
Control Variables: temperature, amount of magnesium and volume of acid
Preliminary work: to calculate the order of HCl(aq), the concentrations needs to be varied. The maximum concentration for Hydrochloric acid for this investigation was 2M and the lowest concentration I decided to use was 0.1M. Lower concentrations of acid reacting with Magnesium would take a lot of time.
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