Energy and Rates Analysis of Chemical Reactions

Evidence

Enthalpy test:

Rate test:

Analysis

Q: How does the molar enthalpy of reaction of magnesium vary with different acids, namely, hydrochloric, sulfuric, and acetic acids?

A: The three molar enthalpies were all relatively close.  The difference between the enthalpies is most likely due to human error, or other complications during the experiment.  Such as the inaccuracy of the scale, thermometer, calorimeter or any other apparatus used.  A major source of error could have been using the specific heat capacity for water for each acid.  Ideally the three molar enthalpies should be the same, because the molar enthalpy for magnesium is dependant upon magnesium, not the other reactant.

(In Ethanoic acid: 2.8E2 kJ/mol, In Sulfuric acid: 3.5E2 kJ/mol, In Hydrochloric acid: 3.5E2 kJ/Mol)

Q: How does the rate of reaction of magnesium vary with these acids?

A: The rate of reaction varies noticeably in the three acids.  The rate was fastest in Sulfuric acid, at 5.29E-5 mol/s, second fastest in Hydrochloric acid at 1.87E-5 mol/s, and slowest in Ethanoic acid at 7.61E-6 mol/s.  Since the surface area of the magnesium and the temperature of the acid was essentially the same in every sample, the only reason that the rates would vary this much is due to their different chemical properties.

Evaluation

        As predicted before the experiment, there was a lot of fizzing and bubbling during the reaction.  Also, as predicted, the system increased in temperature in all segments of the experiment, due to the exothermic nature of the reaction.

        Contrary to the hypothesis, the Carboxylic acid had the slowest rate of reaction with magnesium.  This could be due to the slight polarity of Ethanoic acid.  Since magnesium is a single atom molecule, it is not polar; therefore it would dissolve more readily in a non-polar solution.  This would account for its higher reaction rate with less polar hydrochloric acid and Sulfuric acid.

        As hypothesized, the molar enthalpy for magnesium did not change a lot, and the differences are so minute that they can be accounted for by error.  The molar enthalpy change should be the same for each segment because the acid is only being used to see how much heat is given off by magnesium.  If the specific heat capacities were more accurate, the enthalpies would be closer to each other.  Another source of error could be how well the thermometers were being read; digital thermometers would have given a more accurate reading.  The calorimeters were not 100% closed systems.  There could have been minor energy loss through the hole in the top, or the sides.

Synthesis

        Before the investigation was carried out, it was predicted that Ethanoic acid would react faster than Hydrochloric acid and Sulfuric acid with the magnesium ribbon.  After carrying out the lab, this was proven wrong.  The carboxylic acid was actually the slowest of the three.  As previously discussed, this could be due to the fact that it is more polar than Hydrochloric and Sulfuric acid.  This could also be due to the more simplistic structures of Hydrochloric acid and Sulfuric acid, therefore there would be less intermediates and a more straightforward reaction.

        The rate in the experiment was measured in respect to consumption of magnesium.  This seems to be the easiest and most straightforward method of measuring the rate of the reaction.  Another way that the rate could have been measured is in respect to production of hydrogen gas.  The investigation would have been carried out the exact same, but this time after the reaction had finished, the timer would be stopped, and the remaining solution would be measured in terms of mass.  This mass would be subtracted from the sum of the original mass of the aqueous acid and the mass of the magnesium.  This difference would b the mass of the hydrogen gas produced.  This number would then be divided by 2.02, the molar mass of hydrogen gas.  The quotient would be the number of moles of hydrogen gas produced.  By dividing this by the time the reaction took to finish, you would obtain the rate of production for hydrogen gas in mol/s.

        The more heat that is added to a system, the faster the molecules move, therefore there would be more molecules with the necessary amount of energy to create an activated complex or have a successful collision.  Exothermic reactions produce heat energy, therefore the solution of acid and Magnesium would get warmer the more it reacted.  Since the solution is getting hotter, the molecules are increasing in energy levels, and therefore the rate of the reaction is increasing.  A way to control the exothermic nature of this reaction would be cooling it down as the reaction goes on.  Adding more Freon or ice to the surrounding atmosphere to absorb the heat given off by the reaction.  Not enough to cool down the solution, just enough to keep it at a stable temperature.

        Something like this would be effective in providing heat in a hand warmer, but for multiple reasons, unsafe.  In order to generate enough heat to keep somebody’s hands warm in the winter, the acid would have to be pretty concentrated.  This concentrated acid would be a hazard to the person if it ever leaked, caustic acid can cause irritation to skin and in some cases severe burns.  Also there is a considerable amount of hydrogen gas produced in this reaction, which needs to go somewhere.  The package with the acid and Magnesium in it would have to have ventilation for the gas to escape.  This would be a place for the acid to leak out also.  Acid and magnesium is not the ideal exothermic reaction to provide heat for hand warmers due to the risk of acid burns.

Calculations