The Flow of Energy

1. The Flow of Energy

2. Think about this question for 15 seconds… What does a thermometer measure? Discuss with your neighbor what your answer is Share what your neighbor said

3. Suppose 2 identical candles are used to heat 2 samples of water. One sample is a cup of water, the other is 10 gallons in a large bucket. 1. How does the change in temperature of the samples compare? 2. How will the amount of heat received by each container compare?

4. Thermochemistry: the study of heat changes that occur during chemical reactions Energy: the ability to do work or supply heat energy is not considered to be matter- no mass and does not take up space

5. Potential Energy: stored energy, or energy of position -gas in a car has potential energy- it is stored, so that when you start the car, the car can use the energy of breaking bonds to make the car move Kinetic Energy: energy of motion the faster you move, the more kinetic energy you have - the potential energy of the gas is converted to kinetic energy when the car moves

6. Law of Conservation of Energy: Energy cannot be created or destroyed, but it can be transferred to another object, or changed from one form to another Heat – (q) the flow of energy from one object to another due to temperature differences Transferred from warmer object to cooler object The transfer occurs until an equilibrium is reached

7. When chemical reactions occur, there is either a release of heat or an absorption of heat System: what is being studied Surroundings: everything around the system in nature Universe: the system and the surroundings together It is the system that will gain energy from or lose energy to the surroundings

8. 3 Types of Systems: • Open- free exchange of matter and energy with surroundings. - a saucepan of soup on a stove 2. Closed- no exchange of matter, but an exchange of energy with surroundings. - placing a lid on the saucepan so the soup does not boil out 3. Isolated- no exchange of matter or energy between system and surroundings - placing soup in a thermos

9. Determine the system and the surroundings in the following situations: • You are sitting around a campfire. 2. You place 2 solutions in a test tube. 3. A piece of hot copper is placed in a beaker of water.

10. Endothermic Process: when the system absorbs heat from the surroundings. q is a positive value Example- when you sit around a campfire (surroundings) and you (system) start to become warm Exothermic Process: system loses heat to the surroundings q is negative Example- After exercising, you perspire, you (system) is giving off heat to the surroundings

11. An ice cube is placed into a glass of room temperature lemonade in order to cool the temperature down. 1. Is the heat of the ice cube gained or lost? 2. Is this endothermic or exothermic? 3. Is q positive or negative?

12. Units of Energy: joule (J)- SI unit for heat or energy may see kJ where 1kJ = 1000 J calorie(cal)- amount of energy needed to raise the temperature of 1g of pure water 1o C Calorie(Cal)- the nutritional unit 1Cal = 1kcal = 1000cal 1J = 0.2390 cal 4.184 J = 1 cal

13. Convert: • 1656.70 J to cal • 483.12 cal to J • 0.56721 Cal to J • 395.951 cal • 2021.4 J • 2373.3 J

14. Heat Capacity: the amount of heat needed to change the temperature of a system 1o C Specific Heat Capacity: the amount of heat needed to raise the temperature of 1g of a substance by 1o C The higher the specific heat capacity, the more energy needed to raise the temperature Metals tend to have a low specific heat because they heat up quickly Water has a very high specific heat capacity- it takes a lot of energy to raise the temperature

15. Calculating the Specific Heat Capacity: q = heat (J) m = mass (g) DT = temperature change Tfinal - Tinitial (o C) The formula can be solved algebraically for an unknown Do Problems 1 – 3 pg 299

16. Measuring Heat Changes • Calorimetry- the measurement of heat change in a chemical reaction The heat released by the system is equal to the heat absorbed by the surroundings For a system at constant pressure- the heat is the same as a property called Enthalpy (D H) So… q = DH = C• m • DT

17. When D H: -is positive-endothermic reaction- heat is absorbed into the system -is negative- exothermic reaction- heat is released from the system When the reaction is in an aqueous environment, the mass, D T, and C of the water is used because the energy released or absorbed by the system is the same as the energy released or absorbed by the water

18. Thermochemical Equations: -A balanced equation that includes the heat changes -Done at Standard Conditions: physical states of the substances given at 25o C and pressure at 101.3kPa or 1 atmosphere -From the balanced equation, you can determine the heat absorbed or released for a certain amount of substance

19. Example: Baking soda is used in cooking to make cakes rise- 2NaHCO3(s) Na2CO3(s) + H2O(g) + CO2(g) DH = 129 kJ If you start with 2.24 moles of NaHCO3, how much heat is needed to decompose? 2.24molesNaHCO3 x 129 kJ 2 mol NaHCO3 = 144 kJ

20. Heat of Combustion: DH for the complete combustion of 1 mole of a substance Values are on Table 11.4 on page 305

21. Heat in Changes of State • Heat of Fusion- DHfus- endothermic, melting from a solid to a liquid • Heat of Solidification- DHsolid- exothermic, liquid forms a solid DHfus= - DHsolid The heat absorbed as a substance melts is equal to the heat released as it freezes

22. Heat of Vaporization- DHvap – endothermic, heat absorbed as a substance vaporized from a liquid to a vapor • Heat of Condensation-DHcond – exothermic, heat released as a substance condenses from a vapor to liquid DHvap = -DHcond

23. Heat of Solution- DHsoln – can be endothermic or exothermic, heat absorbed or released when a solute is dissolved into a solvent