What Is The Determination Of The Neutralization Of A Coffee Cup Calorimeter
In this experiment, a coffee cup calorimeter was used to measure the temperature readings of a neutralization reaction and magnesium oxide, MgO. Calorimetry is used to measure amounts of heat transferred to or from a substance.2 The difference of temperatures was used to calculate the heat energy given off by each sub-reaction. These values were solved by using Hess’s Law which determined the overall enthalpy changes of the neutralization reaction and MgO formation. Hess's Law states that the heat evolved or absorbed in a chemical process is the same whether the process takes place in one or in several steps.1 If this experiment were to be completed again, less human error could be avoided by inserting a step that states to thoroughly rinse
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Temperature Changes in Reactions 1-3
Condition Reaction 1 Reaction 2 Reaction 3
Tmax (°C) 32.5 21.3 33.6
Tinitial (°C) 21.9 20.9 21.0
ΔT (°C) 10.6 0.4 12.6
Heat Energy Absorbed by Surroundings in Reactions 1-3
Reaction q(kJ)
1 4.5680 kJ
2 0.1724 kJ
3 5.4300 kJ
The products and reactants in reactions 1 and 2 were manipulated. In reaction 2, the equation was flipped so NH4+ was a product and NH3 was a reactant used later in reaction 3. Since the second reaction was flipped, the sign must also change from a negative to a positive. The spectator ions were canceled out to yield only the final equations for the reactions, specifically reaction 3.
The ΔHrxn was calculated by first determining the limiting reactant. The volume was multiplied by the molarity. This gives the value in moles. All the chemical reactions had the same limiting reactant value so therefore they were all equally as limiting. The q values calculated for each reaction were divided by .1 moles. This value was determined by multiplying .05 L solution by the 2 M. This gives the value in kJ/mol.
Measured Enthalpy Changes for Reactions 1-3
Reaction ΔHrxn(kJ/mol)
1 45.68 kJ
2 1.724 kJ
3 54.23 …show more content…
Equation 4 Measured Molar Enthalpy Change
Figure 4. Combination of 100 mL 6 M HCl and approximately 1.0 g MgO Reaction
Figure 5. Combination of 120 mL 2 M HCl and approximately 0.5 g Mg Reaction The given theoretical value was 602 kJ/mol and the measured ΔH was -552.05 kJ/mol. The calculated percent error was 8.30% which is very good, considering less than fifteen percent is a reasonable percent error value. Therefore, the theoretical and measured values were very close suggesting that the summation of the ΔH of the sub-reactions were very similar to the ΔHf value.
The difference between ΔHrxn and ΔHf is that ΔHrxn is the differences between ΔH which give the amount of heat that is released or absorbed and the AHf is the change in enthalpy when products are formed.
The two parts, B and C, had all exothermic reactions. The values both theoretical and measured were very similar, therefore our percent errors were relatively low. This supports the overall main focus of this experiment, Hess’s Law. The coffee-cup calorimeter successfully showed the overall enthalpy changes when the heat was being released from the system to the surroundings in both reactions: neutralization reaction and formation of
Temperature Changes in Reactions 1-3
Condition Reaction 1 Reaction 2 Reaction 3
Tmax (°C) 32.5 21.3 33.6
Tinitial (°C) 21.9 20.9 21.0
ΔT (°C) 10.6 0.4 12.6
Heat Energy Absorbed by Surroundings in Reactions 1-3
Reaction q(kJ)
1 4.5680 kJ
2 0.1724 kJ
3 5.4300 kJ
The products and reactants in reactions 1 and 2 were manipulated. In reaction 2, the equation was flipped so NH4+ was a product and NH3 was a reactant used later in reaction 3. Since the second reaction was flipped, the sign must also change from a negative to a positive. The spectator ions were canceled out to yield only the final equations for the reactions, specifically reaction 3.
The ΔHrxn was calculated by first determining the limiting reactant. The volume was multiplied by the molarity. This gives the value in moles. All the chemical reactions had the same limiting reactant value so therefore they were all equally as limiting. The q values calculated for each reaction were divided by .1 moles. This value was determined by multiplying .05 L solution by the 2 M. This gives the value in kJ/mol.
Measured Enthalpy Changes for Reactions 1-3
Reaction ΔHrxn(kJ/mol)
1 45.68 kJ
2 1.724 kJ
3 54.23 …show more content…
Equation 4 Measured Molar Enthalpy Change
Figure 4. Combination of 100 mL 6 M HCl and approximately 1.0 g MgO Reaction
Figure 5. Combination of 120 mL 2 M HCl and approximately 0.5 g Mg Reaction The given theoretical value was 602 kJ/mol and the measured ΔH was -552.05 kJ/mol. The calculated percent error was 8.30% which is very good, considering less than fifteen percent is a reasonable percent error value. Therefore, the theoretical and measured values were very close suggesting that the summation of the ΔH of the sub-reactions were very similar to the ΔHf value.
The difference between ΔHrxn and ΔHf is that ΔHrxn is the differences between ΔH which give the amount of heat that is released or absorbed and the AHf is the change in enthalpy when products are formed.
The two parts, B and C, had all exothermic reactions. The values both theoretical and measured were very similar, therefore our percent errors were relatively low. This supports the overall main focus of this experiment, Hess’s Law. The coffee-cup calorimeter successfully showed the overall enthalpy changes when the heat was being released from the system to the surroundings in both reactions: neutralization reaction and formation of