Planning A:

Refer to lab handout entitled, Heat of Reaction for the Formation of Magnesium Oxide.

Planning B:

Refer to lab handout entitled, of Reaction for the Formation of Magnesium Oxide.

Data Collection:

Quantitative Table I:

CompoundTrialMass±0.001gVolume of HCl±0.5mLTemperature of HCl±0.5°CTime (seconds)Temperature of solution±0.5°C





Qualitative Table I:

  • MgO is small, white granular pieces, a fine white powder.
  • HCl is clear and has a slight odour. Very sour scent.
  • Fizzing
  • Steaming. A gas is beginning to form, only a little bit.
  • Feel the solution getting hotter while conducting experiment. Very exothermic.
  • The MgO has completely dissolved in the HCl.
  • The new solution is clear and odourless despite the previous odour of the HCl.
  • Mg is small grey chunks like little shiny stones
  • HCl is clear and has a slight odour. Very sour scent.
  • Fizzing
  • Steaming. A lot of gas is forming.
  • Feel the solution getting a lot hotter while conducting the experiment. Extremely exothermic.
  • Many bubbles are forming.
  • Bad odour is being conducted as the Mg and HCl mix.
  • The Mg has completely dissolved in the HCl.
  • The new solution is clear.
  • There is a very strong and bad odour. Almost smells like rotten eggs but not as strong.

Data Analysis:

Therefore, the average temperature high for the MgO and HCl solution in trial 1 was 30.0⁰C.

Therefore, the average temperature high for the MgO and HCl solution in trial 1 was 46.0⁰C.

Therefore, the average temperature high for the MgO and HCl solution in trial 2 was 29.0⁰C.

Therefore, the average temperature high for the Mg and HCl solution in trial 2 was 44.5⁰C.

Data Processing:


This investigation was conducted in order to determine the enthalpy of formation for magnesium oxide by manipulation of the three equations given. Through experimentation, it was found that the enthalpy of change for the combustion of magnesium is -593.3KJ/mol and that the thermochemical equation (target equation) for the combustion of magnesium is (see right).

It is evident that the two equations which were used in this experiment were exothermic since the enthalpy of change that resulted was a negative value, therefore the experiment was successful. Furthermore, the results are quite accurate as the value reached through the experiment, -24410.6 KJ/Kg, is very close to the theoretical value, -25020KJ/Kg, which is seen through the low percent error of only 2.4%.

Since the percent error is a very small value it is evident that the experiment was a success, however, this does not mean that it was perfect. The Sources of error, in this case, would have been quite minimal resulting in a small change that led to a value slightly lower than the expected value. One error that may have caused a lower enthalpy of change value than expected could have been that heat escaped from the calorimeter used during the experiment.

There were two holes on the lid of the calorimeter and one was being used for the thermometer, however the second, although very small, was left open. These holes could have let heat escape as the reaction was taking place which would have lowered the final temperature value. What’s more, is that the lid was not as tight as it could have been since it simply snapped onto the container being used as a calorimeter and was not airtight, which also could have let some heat escape.

Both of these conditions would have lead to a lower final temperature value. Consequently, the heat value, or Q, would have been lower which also would have lead to a lower enthalpy value, like the one that was found. However, due to the small sizes of the holes and the security of the lid, it is unlikely that a large amount of heat would have escaped which is why only a minimal change would occur, much like in the case of this experiment. To prevent even the slightest anomalies, in the future any holes on the calorimeter can be covered by tape or another item that could block the passage.

The top of the calorimeter could also be covered with aluminum, this would not only cover the holes but would secure the space under the lid so any heat that may escape would stay within the area due to the aluminum. Aluminum could also be tucked in the space between the lid and the calorimeter to once again lock the heat in. This way, the calorimeter will be more effective and maintain all the heat of the reaction resulting in values that are completely accurate and decreasing even the slightest errors.

Another discrepancy that may have occurred during the experiment was that the magnesium strip may have reacted with the oxygen in the air before it was poured into the calorimeter. This error would only be specific to the magnesium as the magnesium oxide has already reacted with oxygen and no further reaction would occur.

The procedure for the experiment does state that the HCl is to be measured first and then the magnesium, the importance of this step is not emphasized and so in a group of two, like the one for this experiment, the step would be broken into two parts as one partner handles the magnesium and the other handles the HCl.  So as the magnesium was being carried from the measuring area to the workstation or while it was sitting on the countertop or being poured in, it could have reacted with the oxygen in the atmosphere and combusted. Consequently, this would have led to a decrease in the mass of the magnesium, one that would have been unknown at the time.

This would have led to a lower temperature value for the reaction and when the calculation for the heat (Q) was done the lower temperature value would have led to a lower value for the heat of reaction since it is calculated using Q=mc t. Once again, the change that would have occurred would have been minimal as it is difficult for large amounts of magnesium to react with oxygen in such a short amount of time without the use of a catalyst. However, this would explain the small error in this experiment as the discrepancies were not that high.

In the future the procedure should emphasize the importance of measuring and pouring the magnesium after the HCl has been measured and poured into the calorimeter, this would prevent other reactions from occurring. The electronic balance, or any balance, should be placed at the work station and as soon as the right amount of magnesium has been measured out it should be poured into the calorimeter immediately. The probability of a reaction occurring would definitely decrease as the calorimeter and balance would be in close proximity to one another.

The final error that may have occurred during the experiment was the loss of heat during the pouring stage. It is evident that the reactions all began immediately, seeing as they were very short, between 90-120 seconds. This means that heat was being produced immediately and the lid was not on the calorimeter to keep the heat from escaping. Once again, the loss of heat would have resulted in a lower enthalpy value.

In the future, to prevent this, the partner that is not pouring the magnesium should hold the lid close to the calorimeter and only open it at a small angle so that there is just enough space for the other partner to pour in the magnesium or magnesium oxide. Once it is all poured in the lid should be snapped down immediately. Although this does not completely stop heat from escaping it certainly decreases the amount that can escape.

Therefore the discrepancies in the experiment, however small, could have led to a lower value than expected which resulted in the low percent error.

author avatar
William Anderson (Schoolworkhelper Editorial Team)
William completed his Bachelor of Science and Master of Arts in 2013. He current serves as a lecturer, tutor and freelance writer. In his spare time, he enjoys reading, walking his dog and parasailing. Article last reviewed: 2022 | St. Rosemary Institution © 2010-2024 | Creative Commons 4.0

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