Energy

1. Energy A Give and Take

2. 10.1 The Nature of Energy • Energy: the ability to do work or produce heat • Potential energy (store energy): energy due to position or composition

3. Kinetic energy (motion energy): energy due to motion of the object and depends on the mass of the object and its velocity • KE = ½ (mv2) • Law of conservation of energy: that energy can be converted from one from to another but can be neither created or destroyed. • Energy of the universe is constant

4. The nature of energy • Work: force acting over a distance • w = F x d • State function: property of the system that changes independently of its pathway

5. Temperature and Heat • Temperature: is a measure of the random motions of the components of a substance • E.g H2O molecules move rapidly in hot water than in cold water • Heat: a flow of energy due to a temperature difference • Tfinal = average temp from mixing (hot & cold temp)

6. Exothermic and Endothermic Process • System – everything we focus on in experiment • Surroundings– everything other the system • exothermic (energy flows out of system to surrounding (via heat) • endothermic ( energy flows into system from surrounding (via heat)

7. Examples • Identify whether these process are exothermic or endothermic • Your hand gets cold when you touch ice • The ice melts when you touch it • Propane is burning in a propane torch • Two chemicals mixing in a beaker give off heat

8. Thermodynamics • Is the study of energy. • First law of thermodynamics: the energy of the universe is constant • Internal energy – energy of the system • ∆E = q + w • ∆ => change in the function • q => represents heat • w => represents work

9. 10.5 Measuring Energy changes • calorie: the amount of energy (heat) required to raise the temperature of one gram of water by 1oC • 1Calorie = 1000 calories • Joule (J) – SI unit • 1 calories = 4.184 joules

10. Converting Calories to Joules • Express 60.1 cal of energy in units of Joules • How many calories of energy corresponds to 28.4 J?

11. Calculating Internal Energy • Calculate ΔE for q = 34 J, w = -22 J • ΔE = q + w • ΔE = 34 J + (-22 J) = 12 J Is this exothermic or endothermic? ΔE > 0, therefore it is endothermic

12. Specific heat • The amount of energy required to change the temperature of one gram of a substance by 1oC • Denoted as s • Heat required = specific heat x mass x change in temp • Q = s x m x ∆T

13. Calculating Energy Requirements • Determine the amount of energy (heat) in joules required to raise the temperature of 7.40 g water from 29.0oC to 46.0 oC Energy required or Q = s x m x ΔT swater = 4.184 J/goC Q = 4.184 J x 7.40 g x 17°C g °C = 526 J

14. A 5.63 g sample of solid gold is heated from 21oC to 32oC. How much energy in Joules and calories is required? • Q= s x m x ΔT sgold = 0.13 J/g oC Q= 0.13 J x 5.63 g x 11°C g °C = 8.1 J

15. A 55.0 g aluminum block initially at 27.5°C absorbs 725 J of heat. What is the final temperature? • Q= s x m x ΔT • ΔT= Tfinal – Tinitial saluminum = 0.89 J/g oC 725 J = 0.89 J x 55.0 gx Tf -27.5°C g°C 725 J = Tf – 27.5 °C 48.95J/°C

16. 725 J = Tf – 27.5 °C 48.95J/°C 14.8 °C = Tf – 27.5 °C Tf = 42.3 °C

17. A sample of gold requires 3.1 J of energy to change its temperature from 19oC to 27oC. What is the mass of this sample of gold? Q = s x m x ΔT sgold = 0.13 J/g oC

18. 3.1 J = 0.13 J x m x 8 °C g °C 3.1 J= m 8 °C x 0.13 J/g°C m = 2.98 g or 3.0 g