1. Calorimetry: :Measuring Heat
Thermochemistry
Calorimetry: :Measuring Heat

2. Calorimetry Calorimetry
How can you measure the change in enthalpy of a reaction?
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3. Calorimetry
Calorimetry is the measurement of the heat flow into or out of a system for chemical and physical processes.
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4. Calorimetry
Calorimetry is the measurement of the heat flow into or out of a system for chemical and physical processes.
In a calorimetry experiment involving an endothermic process, the heat absorbed by the system is equal to the heat released by its surroundings.
In an exothermic process, the heat released by the system is equal to the heat absorbed by its surroundings.
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5. Calorimetry
Calorimetry is the measurement of the heat flow into or out of a system for chemical and physical processes.
The insulated device used to measure the absorption or release of heat in chemical or physical processes is called a calorimeter.
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6. Calorimetry Constant-Pressure Calorimeters
Foam cups can be used as simple calorimeters because they do not let much heat in or out.
Most chemical reactions and physical changes carried out in the laboratory are open to the atmosphere and thus occur at constant pressure.
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7. Calorimetry Constant-Pressure Calorimeters
The enthalpy (H) of a system accounts for the heat flow of the system at constant pressure.
The heat absorbed or released by a reaction at constant pressure is the same as the change in enthalpy, symbolized as ΔH.
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8. Calorimetry Constant-Pressure Calorimeters
The value of ΔH of a reaction can be determined by measuring the heat flow of the reaction at constant pressure.
In this textbook, the terms heat and enthalpy change are used interchangeably.
In other words, q = ΔH.
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9. Calorimetry Constant-Pressure Calorimeters
To measure the enthalpy change for a reaction in aqueous solution in a foam cup calorimeter, dissolve the reacting chemicals (the system) in known volumes of water (the surroundings).
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10. Calorimetry Constant-Pressure Calorimeters
Measure the initial temperature of each solution, and mix the solutions in the foam cup.
After the reaction is complete, measure the final temperature of the mixed solutions.
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11. Calorimetry Constant-Pressure Calorimeters
You can calculate the heat absorbed or released by the surroundings (qsurr) using the formula for the specific heat, the initial and final temperatures, and the heat capacity of water.
qsurr = m  C  ΔT
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12. Calorimetry Constant-Pressure Calorimeters qsurr = m  C  ΔT
m is the mass of the water.
C is the specific heat of water.
ΔT = Tf – Ti
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13. Calorimetry Constant-Pressure Calorimeters
The heat absorbed by the surroundings is equal to, but has the opposite sign of, the heat released by the system.
qsurr = –qsys
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14. qsys = ΔH = –qsurr = –m  C  ΔT
Calorimetry
Constant-Pressure Calorimeters
The enthalpy change for the reaction (ΔH) can be written as follows:
qsys = ΔH = –qsurr = –m  C  ΔT
The sign of ΔH is positive for an endothermic reaction and negative for an exothermic reaction.
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15. Calorimetry Constant-Volume Calorimeters
Calorimetry experiments can also be performed at a constant volume using a device called a bomb calorimeter.
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16. Calorimetry Constant-Volume Calorimeters
In a bomb calorimeter, a sample of a compound is burned in a constant-volume chamber in the presence of oxygen at high pressure.
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17. Calorimetry Constant-Volume Calorimeters
The heat that is released warms the water surrounding the chamber.
By measuring the temperature increase of the water, it is possible to calculate the quantity of heat released during the combustion reaction.
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18. CHEMISTRY & YOU
What type of calorimeter would you use to measure the heat released when a match burns? Describe the experiment and how you would calculate the heat released.
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19. CHEMISTRY & YOU
What type of calorimeter would you use to measure the heat released when a match burns? Describe the experiment and how you would calculate the heat released.
A constant-volume, or bomb, calorimeter would be used to measure the heat released when a match burns. The match would be ignited in the chamber. By measuring the temperature increase in the water and using the equation q = –m  C  ΔT , the heat released, q, can be calculated.
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20. Enthalpy Change in a Calorimetry Experiment
When 25.0 mL of water containing mol HCl at 25.0°C is added to 25.0 mL of water containing mol NaOH at 25.0°C in a foam-cup calorimeter, a reaction occurs. Calculate the enthalpy change (in kJ) during this reaction if the highest temperature observed is 32.0°C. Assume that the densities of the solutions are 1.00 g/mL and the volume of the final solution is equal to the sum of the volumes of the reacting solutions.
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21. Analyze List the knowns and the unknown.1
Use dimensional analysis to determine the mass of the water.
You must also calculate ΔT.
Use ΔH = –qsurr = –m  C  ΔT to solve for ΔH.
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22. Analyze List the knowns and the unknown.1
KNOWNS
UNKNOWN
Cwater = 4.18 J/(g·°C)
Vfinal = VHCl + VNaOH
= 25.0 mL mL = 50.0 mL
Ti = 25.0°C
Tf = 32.0°C
densitysolution = 1.00 g/mL
ΔH = ? kJ
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23. Calculate Solve for the unknown.2
First calculate the total mass of the water.
mwater = 50.0 mL  = 50.0 g
1 mL
1.00 g
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24. Calculate Solve for the unknown.2
Now calculate ΔT.
ΔT = Tf – Ti = 32.0°C – 25.0°C = 7.0°C
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25. Calculate Solve for the unknown.2
Use the values for mwater, Cwater, and ΔT to calculate ΔH.
ΔH = –qsurr = –mwater  Cwater  ΔT
= –(50.0 g)(4.18 J/(g·oC))(7.0°C)
= –1500 J = –1.5 kJ
Use the relationship 1 kJ = 1000 J to convert your answer from J to kJ.
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26. Evaluate Does the result make sense?
Sample Problem 17.3
Evaluate Does the result make sense? 3
The temperature of the solution increases, which means that the reaction is exothermic, and thus the sign of ΔH should be negative.
About 4 J of heat raises the temperature of 1 g of water 1°C, so 200 J of heat is required to raise 50 g of water 1°C. Raising the temperature of 50 g of water 7°C requires about 1400 J, or 1.4 kJ.
This estimated answer is very close to the calculated value of ΔH.
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27. The initial temperature of the water in a constant-pressure calorimeter is 24°C. A reaction takes place in the calorimeter, and the temperature rises to 87°C. The calorimeter contains 367 g of water, which has a specific heat of 4.18 J/(g·°C). Calculate the enthalpy change during this reaction.
Move solution up so it doesn’t run off the slide. Also, the text wraps differently in this slide versus the previous one.
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28. The initial temperature of the water in a constant-pressure calorimeter is 24°C. A reaction takes place in the calorimeter, and the temperature rises to 87°C. The calorimeter contains 367 g of water, which has a specific heat of 4.18 J/(g·°C). Calculate the enthalpy change during this reaction.
ΔH = –m  C  ΔT
= –367 g  4.18 J/(g·°C)  (87°C – 24°C)
= –97000 J = –97 kJ
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29. Thermochemical Equations
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30. Thermochemical Equations
In a chemical equation, the enthalpy change for the reaction can be written as either a reactant or a product.
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31. Thermochemical Equations
In the equation describing the exothermic reaction of calcium oxide and water, the enthalpy change can be considered a product.
CaO(s) + H2O(l) → Ca(OH)2(s) kJ
Calcium oxide is one of the components of cement.
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32. Thermochemical Equations
A chemical equation that includes the enthalpy change is called a thermochemical equation.
CaO(s) + H2O(l) → Ca(OH)2(s) kJ
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33. Thermochemical Equations
Heats of Reaction
The heat of reaction is the enthalpy change for the chemical equation exactly as it is written.
Heats of reaction are reported as ΔH.
The physical state of the reactants and products must also be given.
The standard conditions are that the reaction is carried out at kPa (1 atm) and 25°C.
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34. Thermochemical Equations
Heats of Reaction
Each mole of calcium oxide and water that reacts to form calcium hydroxide produces 65.2 kJ of heat.
CaO(s) + H2O(l) → Ca(OH)2(s)
ΔH = –65.2 kJ
In exothermic processes, the chemical potential energy of the reactants is higher than the chemical potential energy of the products.
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35. Thermochemical Equations
Heats of Reaction
Baking soda (sodium bicarbonate) decomposes when it is heated. This process is endothermic.
2NaHCO3(s) + 85 kJ → Na2CO3(s) + H2O(l) + CO2(g)
The carbon dioxide released in the reaction causes muffins to rise while baking.
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36. Thermochemical Equations
Heats of Reaction
2NaHCO3(s) + 85 kJ → Na2CO3(s) + H2O(l) + CO2(g)
Remember that ΔH is positive for endothermic reactions. Therefore, you can write the reaction as follows:
2NaHCO3(s) → Na2CO3(s) + H2O(l) + CO2(g)
ΔH = 85 kJ
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37. Thermochemical Equations
Heats of Reaction
The amount of heat released or absorbed during a reaction depends on the number of moles of the reactant involved.
The decomposition of 2 mol of sodium bicarbonate requires 85 kJ of heat.
Therefore, the decomposition of 4 mol of the same substance would require twice as much heat, or 170 kJ.
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38. Thermochemical Equations
Heats of Reaction
To see why the physical state of the reactants and products must be stated, compare the following two equations.
difference = kJ
H2O(l) → H2(g) + O2(g) ΔH = kJ 1 2
H2O(g) → H2(g) + O2(g) ΔH = kJ
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39. Thermochemical Equations
Heats of Reaction
To see why the physical state of the reactants and products must be stated, compare the following two equations.
difference = kJ
H2O(l) → H2(g) + O2(g) ΔH = kJ 1 2
H2O(g) → H2(g) + O2(g) ΔH = kJ
The vaporization of 1 mol of liquid water to water vapor at 25°C requires 44.0 kJ of heat.
H2O(l) → H2O(g) ΔH = 44.0 kJ
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40. Using the Heat of Reaction to Calculate Enthalpy Change
2NaHCO3(s) + 85 kJ → Na2CO3(s) + H2O(l) CO2(g)
Calculate the amount of heat (in kJ) required to decompose 2.24 mol NaHCO3(s).
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41. Analyze List the knowns and the unknown.1
Use the thermochemical equation to write a conversion factor relating kJ of heat and moles of NaHCO3. Then use the conversion factor to determine ΔH for 2.24 mol NaHCO3.
KNOWNS
amount of NaHCO3(s) that decomposes = 2.24 mol
ΔH = 85 kJ for 2 mol NaHCO3
UNKNOWN
ΔH = ? kJ for 2.24 mol NaHCO3
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42. Calculate Solve for the unknown.2
Write the conversion factor relating kJ of heat and moles of NaHCO3.
85 kJ
2 mol NaHCO3(s)
The thermochemical equation indicates that 85 kJ are needed to decompose 2 mol NaHCO3(s).
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43. Calculate Solve for the unknown.2
Using dimensional analysis, solve for ΔH.
85 kJ
2 mol NaHCO3(s)
ΔH = 2.24 mol NaHCO3(s) 
= 95 kJ
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44. Evaluate Does the result make sense?3
The 85 kJ in the thermochemical equation refers to the decomposition of 2 mol NaHCO3(s).
Therefore, the decomposition of 2.24 mol should absorb more heat than 85 kJ.
The answer of 95 kJ is consistent with this estimate.
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45. Thermochemical Equations
Heats of Combustion
The heat of combustion is the heat of reaction for the complete burning of one mole of a substance.
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46. Thermochemical Equations
Heats of Combustion
Small amounts of natural gas within crude oil are burned off at oil refineries.
CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) kJ
This is an exothermic reaction.
Burning 1 mol of methane releases 890 kJ of heat.
The heat of combustion (ΔH) for this reaction is –890 kJ per mole of methane burned.
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47. Heats of Combustion at 25°C
Substance
Formula
ΔH (kJ/mol)
Hydrogen
H2(g)
–286
Carbon
C(s, graphite)
–394
Methane
CH4(g)
–890
Acetylene
C2H2(g)
–1300
Ethanol
C2H6O(l)
–1368
Propane
C3H8(g)
–2220
Glucose
C6H12O6(s)
–2808
Octane
C8H18(l)
–5471
Sucrose
C12H22O11(s)
–5645
Like other heats of reaction, heats of combustion are reported as the enthalpy changes when the reactions are carried out at kPa and 25°C.
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48. Which of the following thermochemical equations represents an endothermic reaction?
A. Cgraphite(s) + 2 kJ Cdiamond(s)
B. 2H2(g) + O2(g) H2O kJ
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49. Which of the following thermochemical equations represents an endothermic reaction?
A. Cgraphite(s) + 2 kJ Cdiamond(s)
B. 2H2(g) + O2(g) H2O kJ
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50. Key Concepts & Key Equation
The value of ΔH of a reaction can be determined by measuring the heat flow of the reaction at a constant pressure.
In a chemical equation, the enthalpy change for the reaction can be written as either a reaction or a product.
qsys = ΔH = –qsurr = –m  C  ΔT
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51. Glossary Terms
calorimetry: the precise measurement of heat flow out of a system for chemical and physical processes
calorimeter: an insulated device used to measure the absorption or release of heat in chemical or physical processes
enthalpy (H): the heat content of a system at constant pressure
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52. Glossary Terms
thermochemical equation: a chemical equation that includes the enthalpy change
heat of reaction: the enthalpy change for a chemical equation exactly as it is written
heat of combustion: the heat of reaction for the complete burning of one mole of a substance
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