Calorimetry. Chapter 5. Week 2 Unit 1. Calorimetry. Since we cannot know the of the reactants and products, we measure H through, the of.

Transcription

1 Chapter 5 Week 2 Unit 1 Calorimetry John D. Bookstaver St. Charles Community College Cottleville, MO Calorimetry Since we cannot know the of the reactants and products, we measure H through, the of. 1

2 Heat Capacity (C) and Specific Heat (Cs) We define specific heat capacity as the Units = J/g ⁰C Heat capacity is the amount of Units = J/ ⁰C Note: Cm molar heat capacity Units = J/mole ⁰C Specific Heat Values The amount of energy required to raise the temperature of 1g substance by 1 K (or 1 C) is its specific heat capacity. 2

3 Heat Capacity and Specific Heat Specific heat, then, is Specific heat = heat transferred mass temperature change Cs = q = What Are The Variables/Units? Energy (heat) required, q = Cs m ΔT q = (heat) released or absorbed ( ) Cs = specific heat capacity ( or ) m = mass ( ) ΔT = change in temperature ( or ) Copyright Cengage Learning. All rights reserved 3

4 Example 1: What is the heat capacity of g of water? Example 2: A sample of pure iron requires cal of energy to raise its temperature from 23ºC to ºC. What is the mass of the sample? (The specific heat capacity of iron is J/gºC.) Copyright Cengage Learning. All rights reserved 4

5 Example 3: How much heat is needed to warm g of water (about 1 cup) from C to 98 C? (b) What is the heat capacity of water? Example 4: a. Large beds of rocks are used in some solar-heated homes to store heat. Assume that the specific heat of the rocks is J/g K. Calculate the quantity of heat absorbed by kg of rocks if their temperature increases by 12.0 C. b. What temperature change would these rocks undergo if they emitted kj of heat? 5

6 Example 5: Calculate the mass of water that can be heated from 20⁰C to ⁰C if 10,000 BTU are applied to the water. Example 6: g of water at 80⁰C is mixed with 50g of cold water at ⁰C. Calculate the final temperature when the water is mixed. Assume no heat loss to the surroundings. 6

7 Example 7: You have a glass container with g of water at 10 C. You add a g iron ball at C to the water. Calculate the final temperature of the water. Copyright Cengage Learning. All rights reserved 13 Constant Pressure Calorimetry By carrying out a reaction in aqueous solution in a simple calorimeter such as this one, one can indirectly measure the change for the system by measuring the heat change for the in the calorimeter. 7

8 Coffee Cup Calorimetry Two aqueous solutions, each containing a reactant is mixed in a coffee cup calorimeter. The reactants are the system and the aqueous solution (water) and calorimeter are the surroundings. In an process heat would be transferred from the reaction to the water (since the temperature change of the solution is measured, the temperature rises). In an process heat would be transferred from the water to the reaction (temperature decreases). Coffee Cup Calorimetry By measuring the temperature the flow of between and can be monitored. As heat is by the solution ( process), heat is lost by the reaction (system) and vice versa. 8

9 Calculation q reaction Most reactions are exothermic: Exothermic = = q reaction If the reaction was Endothermic = Example 8: Coffee Cup Calorimetry g of coffee in a cup is initially at 80 ⁰C. It loses kj of heat. What is the final temperature of the coffee? Assume coffee has same specific heat as water. 9

10 Example 9: When a student mixes ml of 1.0 M HCl and 50 ml of M NaOH in a coffee-cup calorimeter, the temperature of the resultant solution increases from C to 27.5 C. Calculate the enthalpy change for the reaction in kj/mol HCl. Assume no loss of heat Assume the density of the solution is 1.0 g/ml, and that its specific heat is 4.18 J/g K Example 10: When a g sample of solid NaOH dissolves in 100g of water in a coffee cup calorimeter, the temperature rises from 23.6 ⁰C to ⁰C. Calculate H (in kj/mol NaOH) for the solution process. Assume the density of the solution is 1.0 g/ml, and that its specific heat is 4.18 J/g K 10

11 Bomb Calorimetry Reactions can be carried out in a sealed such as this one. The heat absorbed (or released) by the is a very good approximation of the change in for the. Bomb Calorimetry The bomb calorimeter is a stainless steel reaction vessel used to accurately measure the of hydrocarbon fuels by reaction with excess oxygen gas. As combustion occurs heat is and the in temperature of the surrounding water is. The manufacturer supplies the heat capacity for the entire calorimeter, including the water it holds. Pure oxygen gas is pumped into the bomb, which is then sealed tightly. The reaction is started by an electric current. 11

12 Bomb Calorimetry Because the in the bomb calorimeter is, what is measured is really the change in, E, not H. for most reactions, the To calculate the heat of combustion from the measured temperature, we use the equation: Bomb Calorimetry q rxn = 12

13 Example 11 A g sample of octane (C 8 H 18 ) was burned in a bomb calorimeter whose total heat capacity is kj/⁰c. The temperature of the calorimeter and its contents rises from ⁰C to 28.78⁰C. Calculate the heat of combustion of octane per gram and per mole. Example 12: ml of 2.0M H 2 SO 4 is mixed with 40.0mL of M NaOH in a calorimeter and the resulting heat of neutralization causes the solution to change in temperature from 20⁰C to ⁰C. Calculate the heat of neutralization per mole of H 2 SO 4 neutralized. 13

14 Example 13: A g sample of lactic acid (HC 3 H 5 O 3 ) is burned in a calorimeter whose heat capacity is 4.812kJ/⁰C. The temperature increases from 23.10⁰C to ⁰C. Calculate the heat of combustion of lactic acid, per mole and per gram. Calorimetry 5.49, 5.51, 5.53, 5.55, 5.57, 14

15 References: Brown, T.L., Bursten, B.E., LeMay, H.E., Murphy, C.J. and Woodward, P.M. (2012). Chemistry The Central Science. (12 th ed.). United States Of America: Pearson Prentice Hall Brown, T.L., Bursten, B.E., LeMay, H.E. and Murphy, C.J. (2009). Chemistry The Central Science. (11 th ed.). Upper Saddle River, NJ: Pearson Prentice Hall. Zumdahl, S.S. and Decoste, D.J. (2010). Introductory Chemistry: A Foundation. (7 th ed.). Belmont, CA: Brooks/Cole Cengage Learning. 15