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2683 Words Grammar Plagiarism  Writing  Score CHM1311 lab Experiment 3: Enthalpy of Various Reactions
Introduction
A coffee cup calorimeter is an apparatus that is used to measure the quantity of thermal energy gained or lost in a chemical reaction. This experiment utilizes this apparatus, which is made from two styrofoam cups with plastic lids and a thermometer, to measure changes in thermal energy of various reactions. When using this type of apparatus, it is assumed that no heat is transferred between the calorimeter and the surroundings, and that no heat is absorbed or released by the cup. This allows for determination of enthalpy change, which will then allow for the calculation of heat absorbed or released.

The specific heat capacity of a substance is the amount of thermal energy required to heat one gram of that substance by one degree. It is an intensive property, as opposed to the heat capacity, which is an extensive property that depends on the amount of substance present. In this experiment, the heat capacity is determined by measuring the change in temperature of the cold water when a hot metal (that does not react with the water) is placed in it. It is assumed that the specific heat capacity is constant in the temperature range, although in reality it does vary with temperature.

q = mcΔT

The formula above is used to determine the specific heat capacity (c). M represents mass and ΔT represents the change in temperature. The heat lost by the metal when placed in cold water is gained by the water. Thus, the heat gain of water must equal the heat loss of the metal. -q (metal) = q (water)

The specific heat capacity can be used to determine the molar mass of a metal using the equation below. The derivation of this equation was largely based off of an assumption by Dulong and Petit in 1819, that one mole of all metals have roughly the same capacity to absorb heat. They found that the heat capacity per weight of many substances hovered around a certain constant, and using this derived the following equation. cmet x MMmet = 25J/mol◦C

The molar enthalpy can then be determined by the following equation, using the amount of heat released in the reaction.
ΔH = q = -mcΔT n n

This equation can be used for all three parts of the experiment. Procedure
As described in the lab manual (Enthalpy of Various Reactions, Dr. Rashmi Venkateswaran, 2013, Experiment 3, p. 34-38).

Observations/Data/Results

Table 1: Specific Heat Capacity of a Metal (Zinc)

Data Trial 1 Trial 2
Mass of Zinc (g) 10.51 10.77
Mass of Empty Calorimeter (g) 9.18 8.81
Volume of Distilled Water in Calorimeter (mL) 20.00 20.00
Mass of Calorimeter with Water (g) 28.54 27.52
Temperature of Metal in Boiling Water (◦C) 100.00 100.00
Change in Temperate of the Water (◦C) 3.60 4.10
Mass of Water (g) 19.36 18.71
Energy Gained by Water (J) 291.61 320.96
Change in Temperature of Zinc (◦C) -75.0 -73.9
Specific Heat Capacity of Zinc (J/g◦C) 0.36995 0.40327
Actual Molar Mass of Zinc (g/mol) 67.31 61.74
Percent Error for Specific Heat Capacity (%) 95.59 104.20
Percent Error for Molar Mass (%) 102.95 94.43
Observations:
-The zinc metal was shiny and silver. It was broken up into square like pieces, and was very light. It was malleable, and easily breakable into smaller pieces.
-The water in the beaker on the hot plate took some time to boil. -The thermometer was moved straight from the hot beaker to the calorimeter, resulting in a quick temperature change.
-It took about four to four and a half minutes for the final solution to settle at a constant temperature.

Table 1.1: Specific Heat Capacity of a Metal (Zinc); Temperature (ᵒC) vs Time (min); Trial 1 (left) and Trial 2 (Right)
Time (minutes) Temperature (ᵒC) Time (minutes) Temperature (ᵒC)
0:30 21.4 0:30 21.9
1:00 21.4 1:00 22.0
1:30 21.2 1:30 22.0
2:00 21.3 2:00 22.0
2:30 21.4 2:30 22.1
3:00 21.4 3:00 22.1
6:30 25.0 6:00 26.0
6:50 24.9 6:20 26.1
7:10 625.0 6:40 26.1
7:30 25.0 7:00 26.1
7:50 24.9 7:20 26.1
8:10 24.9 7:40 26.0
8:30 24.9 8:00 26.0
8:50 24.9 8:20 25.9
9:10 24.9 8:40 25.9
9:30 24.8 9:00 25.9
9:50 24.9 9:20 25.9
10:10 24.8 9:40 25.9
10:30 24.8
10:50 24.8
11:10 24.8

Table 2: Enthalpy of Neutralization (HNO3 & HCl acids with NaOH base)
Data HNO3 acid used HCl acid used Trial 1 Trial 2 Trial 1 Trial 2
Volume of Acid (mL) 50.10 50.20 49.00 49.50
Concentration of Acid (mol/L) 1.10 1.10 1.10 1.10
Volume of NaOH (mL) 49.90 50.00 50.50 49.00
Concentration of NaOH (mol/L) 1.00 1.00 1.00 1.00
Mass of empty calorimeter (g) 9.18 9.18 9.18 9.18
Mass of Calorimeter with Solution (g) 110.07 110.78 107.63 107.84
Volume of Final Solution (mL) 100.0 100.20 99.50 98.50
Change in Temperature of Solution (◦C) 6.30 6.80 7.10 6.60
Mass of Final Solution (g) 100.89 101.60 98.45 98.66
Energy Released by Solution (J) -2569.38 -2606.13 -2924.60 -2724.44
Moles of Limiting Reagent (mol) 0.04099 0.05000 0.05050 0.04900
Moles of Water (mol) 0.04099 0.05000 0.05050 0.04900
Heat of Neutralization/Mole of Water (J/mol) -62683.00 -52122.60 -57912.87 -55600.82
Percent Error (%) 109.78 91.28 101.42 97.37
Observations:
-HCl, HNO3, and NaOH are all colourless, clear, and non-viscous solutions, with distinct scents.
-The experiment with the HNO3 lasted longer than the one with the HCl.

Table 2.1: Enthalpy of Neutralization of HNO3;Temperature vs Time; Trial 1 (left) Trial 2 (right)
Time (minutes) Temperature (ᵒC) Time (minutes) Temperature (ᵒC)
0:30 21.5 0:30 21.3
1:00 21.5 1:00 21.3
1:30 21.5 1:30 21.3
2:00 21.6 2:00 21.3
2:30 21.6 2:30 21.3
3:00 21.6 3:00 21.3
4:30 28.1 3:50 28.2
4:50 28.0 4:10 28.2
5:10 27.9 4:30 28.1
5:30 27.9 4:50 28.1
5:50 27.9 5:10 28.0
6:10 27.8 5:30 28.0
6:30 27.8 5:50 28.0
6:50 27.7 6:10 27.9
7:10 27.7 6:30 27.9
7:30 27.7 6:50 27.9
7:50 27.7 7:10 27.8
8:10 27.6 7:30 27.8
8:30 27.6 7:50 27.8
8:50 27.6 8:10 27.8
9:10 27.6 8:20 27.8

Table 2.2: Enthalpy of Neutralization of HCl; Temperature vs Time; Trial 1 (left) Trial 2 (right)
Time (minutes) Temperature (ᵒC) Time (minutes) Temperature (ᵒC)
0:30 22.6 0:30 22.9
1:00 22.7 1:00 22.9
1:30 22.7 1:30 22.9
2:00 22.7 2:00 23.0
2:30 22.7 2:30 23.0
3:00 22.8 3:00 22.9
3:14 29.9 3:15 29.6
3:34 29.8 3:35 29.5
3:54 29.8 3:55 29.5
4:14 29.8 4:15 29.5
4:34 29.7 4:35 29.4
4:54 29.7 4:55 29.4
5:14 29.7 5:15 29.4
5:34 29.6 5:35 29.3
5:54 29.6 5:55 29.3
6:14 29.6 6:15 29.2
6:34 29.6 6:35 29.2
6:54 29.6 6:55 29.2
7:14 29.6 7:15 29.2

Table 3: Enthalpy of Dissolution of an unknown Salt (salt B)

Data Trial 1 Trial 2
Salt Identification B B
Mass of Salt (g) 1.52 1.53
Mass of Empty Calorimeter (g) 8.29 8.42
Volume of Distilled Water (mL) 20.00 20.00
Mass of Calorimeter with Water (g) 27.08 27.73
Change in Temperature of Solution (◦C) -5.40 -5.20
Energy Released by Solution (J) 31.82 30.84
Moles of Salt (mol) 0.015033 0.015132
Enthalpy of Dissolution/Mole of Salt (J/mol) 2116.68 2038.01
Observations:
-The salt was a white-ish powdery substance.
-This part of the experiment required a lot of active mixing.
-This was the longest part, in comparison to the other two parts.
-In this reaction, the temperature dropped, whereas in the other ones, the temperature typically increased.

Table 3.1: Enthalpy of Dissolution of a salt (salt B); Temperature (ᵒC) vs Time (min); Trial 1 (left) trial 2 (right)
Time (minutes) Temperature (ᵒC) Time (minutes) Temperature (ᵒC)
0:30 21.5 0:30 21.4
1:00 21.6 1:00 21.4
1:30 21.6 1:30 21.4
2:00 21.6 2:00 21.4
2:30 21.6 2:30 21.4
3:00 21.6 3:00 21.4
4:10 16.8 3:50 19.1
4:30 16.1 4:10 18.2
4:50 16.3 4:30 16.5
5:10 16.2 4:50 16.0
5:30 16.0 5:10 15.9
5:50 16.1 5:30 16.0
6:10 16.2 5:50 16.1
6:30 16.3 6:10 16.2
6:50 16.4 6:30 16.4
7:10 16.5 6:50 16.5
7:30 16.6 7:10 16.6
7:50 16.6 7:30 16.7
8:10 16.7 7:50 16.8
8:30 16.8 8:10 16.9
8:50 16.9 8:30 17.0
9:10 17.2 8:50 17.1
9:30 17.4 9:10 17.1
9:50 17.6 9:30 17.3
10:10 17.8 9:50 17.4
10:30 17.8 10:30 17.4
10:50 17.9 10:50 17.4
11:10 18.0 11:10 17.4
11:30 18.0 11:30 17.4
11:50 18.1 11:50 17.4
12:10 18.1
12:30 18.1

Calculations
Part 1: Specific Heat Capacity of a Metal
1. Change in temperature of the water:
ΔwaterT = Tf - Ti = 25.0◦C – 21.4◦C = 3.60◦C

2. Energy gained by the water: mwater = 28.54g – 9.18g q = mcΔT = 19.36g = (19.36g)(4.184J/g◦C)(3.60◦C) = 291.61J

3. Change in temperature of zinc:
ΔmetalT = Tf - Ti = 25.0◦C – 100.0◦C = -75.0◦C

4. Specific heat capacity of zinc: qwater = - qmetal qwater = -cmet x mmet x ΔTmet
291.61J = (-cmet)(10.51g)(-75◦C) cmet = 0.36995 J/g◦C

5. Molar mass of zinc:
MMmet = 24.9 J/g◦C 0.36995J/g◦C = 67.31g/mol

6. Percent error:
% error Mm = actual % error c = actual theoretical x 100 theoretical x 100 = 67.31g/mol = 0.36995 J/g◦C 65.38g/mol x 100 0.387 J/g◦C x 100 = 102.95% = 95.59%

Part 2: Enthalpy of Neutralization- HNO3- Trial 1
7. Change in temperature of the solution:
ΔsolnT = Tf - Ti = 27.9◦C – 21.6◦C = 6.30◦C

8. Volume of the solution:
VT = VHNO3 + VNaOH = 50.10mL + 49.90mL = 100.00mL

9. Mass of the final solution:
D = m measured mass = 110.07g – 9.18g V = 100.89g msoln = (1.00g/mL)(100.00mL) = 100.00g
The values of the theoretical mass and the measured mass differ by merely 0.98g, however it can be said that the value of the measured mass is more accurate in this experiment because it is not based off of an assumption (the density is assumed in the theoretical mass).

10. Energy released: qsoln = - mcΔT = - (100.89g)(4.184J/g◦C)(6.3◦C) = -2569.38J

11. Number of moles of the limiting reagent: nHNO3 = C x V nNaOH = C x V = (1.10M)(0.05010L) = (1.00M)(0.04990L) = 0.05511mol = 0.04099mol
NaOH (OH-) is the limiting reagent.

12. Number of moles of water formed: nH20 = 0.04099mol OH- x 1 mol H20 1 mol OH- = 0.04099mol

13. Heat of Neutralization per mole of water:
ΔNH◦ = qsoln nH20 = -2569.38J 0.04099mol = -62683J/mol

14. Compare the heat of neutralization per mole of the 2 acids: ΔNH◦ HNO3 Trial 1 = -62683.00J/mol ΔNH◦ HCl Trial 1 = -57912.87J/mol
The heat of neutralization/mole when using HNO3 was much greater in value then compared with HCl.

15. Percent error:
% error = actual theoretical x 100 = -62683J/mol -57100J/mol x 100% = 109.78%

Part 3: Enthalpy of Dissolution of a Salt (Salt B)
16. Change in temperature of the solution:
ΔsolnT = Tf - Ti = 16.2◦C -21.6◦C = -5.40◦C

17. Energy released by the solution: qsoln = - mcΔT = -(1.52g)(3.877J/g◦C)(-5.40◦C) = 31.82J

18. Enthalpy of dissolution per mole of salt: nsalt = 1.52g ΔsH◦ = 31.82J 101.11g/mol 0.015033mol = 0.015033mol = 2116.68J/mol

Discussion

The purpose of this experiment was to determine the enthalpy of various reactions, namely three, using a coffee cup calorimeter. In the first part of the experiment, metal zinc was heated to 100ᵒC and placed in cold water which it did not react with. The heat was transferred to the water resulting in an increase in the temperature of water, hence it was an exothermic reaction. In analyzing the graph of this reaction (Figure 1 or Figure 2), it can be seen that the temperature of the water rises from the initial temperature greatly, but as time passes it decreases slightly. This can be attributed to the material of the cheap lids. While mixing, heat would escape from the poorly structured and attached lid. Or from the hole on top of calorimeter’s lid for the thermometer probe. Apart from these experimental errors, the value for the specific heat capacity of the metal proves to be reasonable.In calculating the molar mass of zinc, an approximation was used, which inherently paves the way for errors. The actual value obtained in trial 1 was closer tot he theoretical value than that of trial 2, however both values were quite similar. The slight differences, however, can be attributed to the different masses, and that not all metals have the same capacity to absorb heat. In the second part of the experiment, the enthalpy of neutralization of HNO3 and HCl acids with the base of NaOH were measured. What was to be fixed for a repeat of this experiment would be in the process of mixing. Typically, in this part of the experiment the values spread from lower to higher, because of difficulties in mixing. This was not the case, as shown by the results above, and Figures 3, 4, and 5. Rather, the temperatures slowly decrease. This can be justified by the same reasoning used above, that the material of the lids of the calorimeters was poorly structured, resulting in heat loss. And that the hole of the thermometer probe allowed for gradual heat loss. This was one similarity between the HCl and HNO3 reactions. The reactions differed in the heats of neutralizations, as it was a much larger value for the HNO3 acid. This is because when sodium hydroxide reacted with nitric acid, this solution gained more heat from the base, resulting in a larger final temperature compared to when sodium hydroxide reacted with hydrochloric acid. The theoretical heat of neutralization per mole of water with a strong acid and base, however, is always -57.1kJ/mol. This is because strong acids, bases, and the salt they form are all completely ionized in dilute aqueous solutions. In the third and final part of the experiment, the enthalpy of dissolution for an unknown salt was determined. This experiment was the most time consuming, as the temperature of the salt mixed with water varied for a long time. Figures 6 and 7 demonstrate that after the time of mixing, the temperature dropped dramatically (mixing techniques), and then as time passed gradually increased until it stabilized. This gradual increase can be explained by comparing it to a cup of hot coffee; if the cup is being mixed and has some exposure to air, the temperature will try to equilibrate to room temperature and thus increase. The later stability of the graph represents the time when the salt is dissolved. The unknown salt is Potassium Nitrate (KNO3), as deduced by its molar mass of 101.11g/mol. The value for the enthalpy of dissolution of KNO3 was a positive value indicating the lattice energy was greater than the hydration energy, and thus it was an endothermic reaction. The literature value for the enthalpy of dissolution of KNO3 was not found.

Conclusion In conclusion, the specific heat capacity of zinc was measured to be 0.36995J/g◦C with a percent error of 91.59%, and a molar mass of 67.31g/mol and a percent error of 102.95%. The neutralization per mole of water with HNO3 used was -62683.00J/mol with a percent error of 109.78% and with HCl was -57912.87J/mol with a percent error of 101.42%. The enthalpy of dissolution per mole of potassium iodide was 2116.68J/mol.

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