Investigating the kinetics involved in the reaction of metals with acids.

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Background

In this coursework, I will be investigating the kinetics involved in the reaction of metals with acids. I will be concentrating on two aspects of the reaction, the order and activation energy of the reaction. I will be studying whether the order of the reaction changes with the type of acid used, and also the effect the type of acid used has on the activation energy.

        The order of a reaction shows how the reaction rate is affected by the concentrations of the reactants. Considering the equation shown below.

A  +  B  +  C   D

If the order is zero with respect to reactant A, the rate is unaffected by changes in concentration of A. If the order is one with respect to reactant B, the rate is doubled by doubling the concentration of reactant B. Finally, if order is two with respect to a reactant C, the rate is quadrupled by doubling the concentration of C. Therefore, the overall order of the reaction is given by adding the orders.

Rate = k[A]0[B]1[C]2

= 0  + 1  +  2

= 3

The rate constant of the reaction is k.

        The reaction rate can only be calculated from experimental values. Therefore, in this investigation I will be plotting a concentration/time graph to show the overall order of the acid-metal reaction. The shape of the graph will indicate the order by measuring the half-life of the reactant. The shapes of graphs for zero, one and second order reactions are shown below.

Then plotting the rate against the concentration will confirm the order of the reaction. Rate against concentration will be a horizontal line for zero order, directly proportional line for first order and rate increasing at an increasing rate against concentration for second order. (See blue lines)

        The activation energy of a reaction is the minimum energy required for the reaction to occur. In a gas or a solution, particles are in constant motion and they collide both with each other and any solid particles. A reaction can only take place if the energy of the collisions exceeds the activation energy of the reaction. In this investigation the activation energy will be supplied by heating the acid and metal. Reactions with small activation energy often take place rapidly whereas ones with high activation energies will stop the reactants from taking part in the reaction.

 

The Arrhenius equation can be used to work out the activation energy. The equation is

                                           In k = constant – EA (1/T)

                                                                         R

In the above equation, k is the rate constant of the reaction, R is the gas constant 8.31J K-1 mol-1, EA is the activation energy in J mol-1 and T is the temperature in kelvins.

The Arrhenius equation is written in the same form as the equation to a straight line, y = mx + c. After a graph of In rate and 1/T is obtained, the gradient of that graph multiplied by the gas constant will determine the value of the activation energy.

Aim

The aim of this investigation is to find out how the type of acid, strong and weak, affects the activation energy and order of a metal acid reaction.

        The strong and weak acid I have chosen for the reactions are hydrochloric acid and ethanoic acid respectively. A strong acid is one that dissociates into H+ ions very easily in water unlike a weak acid, which does the opposite. The word, chemical and ionic equations are shown below.

Magnesium(s) + Hydrochloric acid(aq)  Magnesium Chloride(aq) + Hydrogen(g)

Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)

 Mg(s) + 2H+(aq)   Mg2+(aq) +  H2(g)

Magnesium(s) + Ethanoic acid(aq)  Magnesium Ethanoate(aq) + Hydrogen(g)

Mg(s) + CH3COOH(aq)  Mg(CH3COO)2(aq) + H2(g)

 Mg(s) + 2H+(aq)   Mg2+(aq) +  H2(g)

Preliminary Experiment

        This experiment was used to gain an understanding of what concentrations of acid should be used and different ways in which the metal could be added to the acid. Magnesium was the metal used for this experiment as it is close to magnesium in the reactivity series and therefore, reacts similarly. The concentrations of hydrochloric and ethanoic acid used were 1M – 4M, and the two temperatures used for activation energy were 64oC and 30oC. The method for the order experiment is as follows. Firstly connect a conical flask, containing 20cm3 of acid, to gas syringe. Drop 2.5cm of magnesium ribbon into the flask and then immediately seal the flask with a bung. Start timing until 10cm3 of hydrogen has been collected in the gas syringe.

Repeat the procedure of adding the zinc ribbon to the acid like in the previous experiment, except this time carry out the experiment at two different temperatures.

The results for the order and activation energy experiment are shown below.

From the results for the order reaction, it can be seen that using 4M concentration of hydrochloric acid made to reaction occur in two seconds. The reaction occurred in a too fast time for it to be measured accurately. Therefore, when magnesium is used, which is more reactive than zinc, the time will be even quicker. This will mean that it will be very difficult to measure an accurate time for 4M concentrations. Hence I will not be using this concentration for the actual experiment. As the 1M-concentration reaction did not take a long time to occur for both acids, it will be possible to use 0.5M concentrations. Adding this concentration to the actual experiment will make the order graph more accurate as there are more results to plot.

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        Looking at the activation energy results I think it will be possible to lower the temperature more to gain extra results, which will result in a better indication of what the activation energy is. I also think that 60oC is a suitable starting temperature as it is not close to the boiling point of either acids.

Apparatus

  • Hydrochloric Acid - 0.5M, 1.0M, 1.5M, 2.0M, 2.5M & 3M.
  • Ethanoic Acid - 0.5M, 1.0M, 1.5M, 2.0M, 2.5M & 3M.
  • Magnesium Ribbon – 1 metre.
  • Gas syringe – 50cm3.
  • Conical Flask – 100cm3.
  • Stop ...

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