1. Heat Transfer and Calorimetry Dr. Keith Baessler

2. Definitions and Concepts Energy Law of Conservation of Energy System vs Surroundings Types of Systems Heat vs Temperature Endothermic vs Exothermic Specific Heat Heat Capacity Heat Transfer Sign Conventions of q Calorimetry Calorimeters Formula Applications 2

3. 3 Energy is the capacity to do work. Radiant energy comes from the sun and is earth’s primary energy source Thermal energy is the energy associated with the random motion of atoms and molecules Chemical energy is the energy stored within the bonds of chemical substances Nuclear energy is the energy stored within the collection of neutrons and protons in the atom Potential energy is the energy available by virtue of an object’s position

4. 4 First law of thermodynamics – energy can be converted from one form to another, but cannot be created or destroyed. C 3 H 8 + 5O 2 3CO 2 + 4H 2 O Exothermic chemical reaction! Chemical energy lost by combustion = Energy gained by the surroundings system surroundings

5. 5 Heat is the transfer of thermal energy between two bodies that are at different temperatures. Heat versus Temperature Temperature is a measure of the thermal energy. Temperature = Thermal Energy

6. 6 Exothermic process is any process that gives off heat – transfers thermal energy from the system to the surroundings. Endothermic process is any process in which heat has to be supplied to the system from the surroundings. 2H 2 (g) + O 2 (g) 2H 2 O (l) + energy H 2 O (g) H 2 O (l) + energy energy + 2HgO (s) 2Hg (l) + O 2 (g) energy + H 2 O (s) H 2 O (l)

7. 7 The system is the specific part of the universe that is of interest in the study. open mass & energyExchange: closed energy isolated nothing

8. 8 Heat Transfer, q Heat (q) = the transfer of energy which causes the temperature of an object to change units: joules (j), calories (cal) –A calorie is the amount of energy needed to raise the temperature of 1.00 g water by 1°C. 1cal = 4.184 joules Heat spontaneously moves regions of high temperature to regions of lower temperature. A metal spoon at 25°C is placed in boiling water. What happens?

9. 9 The Sign Convention of q Endothermic systems require the surroundings to add energy to the system. q is positive (+) Exothermic reactions release energy to the surroundings. q is negative (-) Energy changes are measured from the point of view of the system

10. 10 The specific heat (s) of a substance is the amount of heat (q) required to raise the temperature of one gram of the substance by one degree Celsius. C = m x s Heat (q) absorbed or released: q = m x s x  t q = C x  t  t = t final - t initial The heat capacity** (C) of a substance is the amount of heat (q) required to raise the temperature of a given quantity (m) of the substance by one degree Celsius. **Heat Capacity (C) is sometimes called the calorimeter constant

11. 11 How much heat is given off when an 869 g iron bar cools from 94 o C to 5 o C? s of Fe = 0.444 J/g o C  t = t final – t initial = 5 o C – 94 o C = -89 o C q = ms  t = 869 g x 0.444 J/g o C x –89 o C= -34,000 J

12. 12 Calorimetry Calorimeter: a closed container used to measure temperature changes in physical and chemical processes. From the temperature changes we can calculate the heat of the reaction, q –q v ; heat measured under constant volume conditions –q p : heat measured under constant pressure conditions q = m x s x  t

13. 13 Two of the most common types of calorimeters are the coffee cup calorimeter and the bomb calorimeter.

14. 14 Coffee Cup Calorimeter The open system allows the pressure to remain constant. Thus we measure q p = ΔH, the enthalpy change. ΔH =change in heat at constant pressure.

15. 15 The 1 st Law of Thermodynamics Keeps Track of Heat Transfer If we monitor the heat transfers (q) of all materials involved, we can predict that their sum will be zero. By monitoring the surroundings, we can predict what is happening to our system. Heat transfers until thermal equilibrium, thus the final temperature is the same for all materials.

16. 16 Example of Using 1 st Law for Heat Transfer A 43.29 g sample of solid is transferred from boiling water (T=99.8°C) to 152 g water at 22.5°C in a Styrofoam coffee cup calorimeter. The T water increased to 24.3°C. Calculate the specific heat of the solid assuming no heat was lost or gained by the calorimeter cup. q sample + q water + q cup = 0 q cup is neglected in problem = 0 Hence q sample = - q water

17. 17 q sample = - q water

18. 18 Chemistry in Action:

19. 19 Chemistry in Action: Fuel Values of Foods and Other Substances 1 cal = 4.184 J 1 Cal = 1000 cal = 4184 J Substance  H combustion (kJ/g) Apple-2 Beef-8 Beer-1.5 Gasoline-34

20. A metal alloy was heated for 2 minutes in a hot water bath at 101.20  C before being transferred into an aluminum calorimeter cup containing water at 23.42 °C. A rise in the temperature of the water in the calorimeter cup was observed. In fact the maximum temperature reached was 28.25 °C. Using the data below, calculate the specific heat of the metal object in cal/g°C. Note that the specific heats of water and aluminum are 1.00 cal/g  C and 0.22 cal/g  C respectively. Wt. Alloy 243.42 g Wt. water in calorimeter cup 213.14 g Wt. Calorimeter cup 45.37 g Initial temp. of water 23.42  C Initial temp. of calorimeter cup 23.42 °C Initial temp. of metal alloy 101.20 °C Final temp. of water 28.25  C 20