Data:
Trial #1: HCl (aq) + NaOH (aq)→ H20 (l) + NaCl (aq)
Trial #2: NaOH (aq) + NH4Cl (aq)→ NH3 (aq) + NaCl (aq) + H20 (l)
Trial #3: HCl (aq) + NH4OH (aq) → NH4Cl (aq) + H20 (l)
Data Analysis:
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Calculating the amount of heat energy: (q) = Cp x m x ∆T
Trial #1: q = (4.18 J/g °C) x (1.03 g/mL x 50mL) (35.22-21.93) = 2860.94 J
Trial #2: q = (4.18 J/g °C) x (1.03 g/mL x 50mL) (23.24-22.34) = 193.74 J
Trial #3: q = (4.18 J/g °C) X (1.03 g/mL X 50mL) (34.38-21.62) = 2746.85 J
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Calculating the Enthalpy Change: Indirect Method
Calculating the net ionic equation: ∆H = ∆H (products) - ∆H (reactants)
Trial #1: HCl (aq) + NaOH (aq)→ H20 (l) + NaCl (aq) → overall
H+ (aq) + OH- (aq)→ H20 (l) → net ionic
∆Hrxn = ∆H H20 (-285.8 kJ/mol) - ∆H OH-(-229.94 kJ/mol) - ∆H H+ (0 kJ/mol)
= -55.9kJ/mol
Trial #2: NaOH (aq) + NH4Cl (aq)→ NH3 (aq) + NaCl (aq) + H20 (l) → overall
OH-(aq) + NH4+ (aq) → NH3 (aq) + H20 (l) → net ionic
∆Hrxn = ∆H H20 (-285.85 kJ/mol) + ∆H NH3 (-80.3 kJ/mol)
- ∆H OH- (-229.94 kJ/mol) - ∆H NH4+ (-132.8 kJ/mol)
= -3.36 kJ/mol
Trial #3: HCl (aq) + NH4OH (aq) → NH4Cl (aq) + H20 (l) → overall
H+ (aq) + NH3 (aq) → NH4+ (aq) → net ionic
∆Hrxn = ∆H NH4+ (-132.8 kJ/mol) – ∆H H+ (0 kJ/mol) - ∆H NH3 (-80.3 kJ/mol)
= -52.5 kJ/mol
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Calculating Enthalpy using Hess’s Law: Indirect method: If a reaction can be written as a sum of several processes, the of the total process equals the sum of the of the various processes. Therefore, subtracting the enthalpy values of trial 1 and 2 should equal trial 3; Hess’s Law can be easily demonstrated in this manner.
(-55.9 kJ/mol) – (-3.36 kJ/mol) = 52.54 (kJ/mol)
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Calculating Percent Error: (accepted value- experimental value) x 100
accepted value
Indirect Method: ( (-52.54) – (-52.5) ) x 100 = .08 % error
52.5
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Calculating Enthalpy Change: Direct Method
Trial #1: (25 mL HCl) x ( 1L ) x (2.0 mol / L) = .05 mol HCl
1000 mL
(2860.94 J) x ( 1 kJ ) / .05 mol = 57.22 kJ/mol HCl
1000 J
Trial #2: (25 mL NaOH) x ( 1L ) x (2.0 mol / L) = .05 mol NaOH
1000 mL
(193.74 J) x ( 1 kJ ) / .05 mol = 3.87 kJ/mol NaOH
1000 J
Trial #3: (25 mL NH3) x ( 1L ) x (2.0 mol / L) = .05 mol NH3
1000 mL
(2746.85 J) X ( 1 kJ ) / .05 mol = 54.94 kJ/mol NH3
1000 J
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Calculating Enthalpy using Hess’s Law: Direct Method: See above for info. On Hess’s Law. Trial 1 – Trial 2 = Trial 3
(-57.22 kJ/mol) - (-3.87 kJ/mol) = -53.35 kJ/mol
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Calculating Percent Error: (accepted value- experimental value) x 100
accepted value
Direct Method: ( (-53.35) - (-54.94) ) x 100 = 2.99 % error
-53.25
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This data supports Hess’s Law by showing the direct correlation in two different methods that the delta H value of a reaction can be calculated using delta H values from the sum of the in trial 1 and trial 2, which makes Hess’s law very easy to see and accept as irrefutable truth. These values which were calculated in different methods were only differentiated due to experimental methods for attaining these values, but these values should be very similar because the two different methods are measuring and calculating the same thing.
Discussion/Conclusion:
The main possible source of error for this laboratory is the loss of heat for the apparatus, it is far from a perfectly closed environment the heat loss could drastically throw off experimentally obtained data. Other sources of error could include human error and equipment malfunction, which could have lead to incorrect data, but the data would have appeared to be correct. However, the percent error calculated during this experiment is relatively low, meaning that high accuracy and precision were obtained as long as the calculated data is correct. Percent error didn’t exceed 3.0% and was as low as .08 %, these figures are extremely low and yield no grounds or reasoning for data rejection. All data and calculations displayed the properties of enthalpy and Hess’s law, and therefore this laboratory experiment is deemed as a success.
This laboratory geared students to understand that all reactions produce some type of heat change, whether it is exothermic or endothermic. This is important information for future laboratory experiments, and to become more familiar with ionic equations and balancing them. These understandings about chemical reactions are essential for further study into various fields of chemistry. Hess’s law also is important information to hold but not essential for advancement in the chemistry field.
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