Energy

1. Energy Chapter 10

2. 10.1 The Nature of Energy • Energy – the ability to do work or produce heat • Potential Energy – due to position or composition • Kinetic Energy – due to motion • Depends on the mass of the object (m) and its velocity (v) • KE = ½ mv2

3. The Law of Conservation of Energy • Energy can be converted from one form to another but can be neither created nor destroyed • The energy in the universe is constant

4. Work • Work = Force x distance • W = Fd • Frictional Heating – 2 surfaces in contact with each other • Depends on surface and force pushing the surfaces together

5. State Function • The property of the system that changes independently of its pathway • The pathway is how you get there • Example • If you travel from Chicago to Denver what are state functions? • The route you take to get there is your pathway, so it is not a state function • Change in elevation doesn’t depend on how you get there so it is a state function

6. 10.2 Temperature and Heat • Temperature – Measure of the random motion of the components of a substance

7. Heat – The flow of energy due to a difference in temperature

8. 10.3 Exothermic and Endothermic Processes • System – part of universe we are looking at • Surroundings – everything else • Exothermic – energy flows out of a system • Endothermic – energy flows into a system

9. Where does energy as heat come from in exothermic reactions? • It depends on the potential energy between the products and reactants

10. 10.4 Thermodynamics • Law of Conservation of Energy (a.k.a. The First Law of Thermodynamics) • Energy can neither be created nor destroyed under normal conditions • The energy of the universe is constant

11. E = internal energy • E is the sum of the kinetic energy and the potential energy • Can be changed by the flow of work, heat, or both • ∆ = change in; called “delta” • w = work • q = heat • ∆E = q + w • Change in internal energy equals heat plus work

12. Thermodynamic quantities are made up of a number that shows magnitude and a sign that shows whether energy is flowing into the system (endothermic = + ) or out of the system (exothermic = - )

13. 10.5 Measuring Energy Changes • calorie = amount of energy required to raise the temperature of 1 gram of water by one degree Celsius • 1000 calories (1 kilocalorie) is what we refer to as a “Calorie” with a capital C • 1 calorie = 4.184 joules • 1 cal = 4.184 J • To go from calories to joules multiply by 4.184 • To go from joules to calories divide by 4.184

14. And now for a problem! • How much heat, in joules, is required to raise the temperature of 7.40 g water from 29.0 °C to 46.0 °C? • We know we need 4.184 J of energy raise 1 g of water 1 °C • We have 7.40 g of water so it will take 7.4 g x 4.184 J to raise it 1 °C • We also need to raise the temperature 17 °C so 17.0 °C x 7.4 g x 4.184 J/ g x °C • So we need 526 J of energy • Now try this • Calculate the joules of energy required to heat 454 g of water from 5.4 °C to 98.6 °C?

15. So we know that the amount of energy we need to raise the temperature of a substance depends on the amount of substance and the change in temperature • But the substance also plays a big part • Specific Heat Capacity = the amount of energy needed to raise the temperature of 1 g of a substance 1 °C

16. Specific Heats • Liquid water = 4.184 J • Aluminum = 0.89 J • Gold = 0.13 J • This explains why certain things heat up faster than others • The pot heats up faster than the water in it • The water in the pool is colder that the cement around it

17. Now for another equation • The amount of energy required = the specific heat x mass x change in temperature • Q = m x Cp x ∆T • Try this sample • A 1.6 g sample of metal that looks like gold requires 5.8 J of energy to change its temperature from 23 °C to 41 °C. Is the metal gold? (Hint – you are finding what s is and comparing to what you know about gold’s specific heat) • Answer = No; Gold’s s = 0.13 J/ g °C but this substance has an s = 0.20 J / g °C

18. 10.6 Thermochemistry (Enthalpy) • Enthalpy (symbol = H) is the same as the flow of heat • ∆Hp = heat • P tells us it occurred under constant pressure • ∆ means “change in” • So the enthalpy for a reaction at constant pressure is the same as heat

19. Calorimetry • Calorimeter = device used to determine the heat associated with a chemical reaction • Reaction is run in calorimeter and temperature change is observed • We can use calorimeter to find ∆H • Once we know ∆H for some reactions we can use those to calculate ∆H for other reactions

20. 10.7 Hess’s Law • The change in enthalpy for a given process is independent of the pathway for the process (this means it is a state function) • Hess’s Law states that the change in enthalpy from reactants to products in a reaction is the same whether it takes place in one step or a series of steps • N2 + 2O2→ 2NO2 ∆H = 68 kJ or • N2 + O2→ 2NO ∆H = 180 kJ • 2NO + O2 → 2NO2 ∆H = -112 kJ • So 180 kJ + (-112 kJ) = ∆H = 68 kJ

21. Characteristics of Enthalpy Changes • If a reaction is reversed, ∆H is reversed • Xe + 2F2 → XeF4 ∆H = -251 kJ • XeF4 → Xe + 2F2 ∆H = +251 kJ • Magnitude of ∆H is proportional to quantities of reactants and products • Xe + 2F2 → XeF4 ∆H = -251 kJ • 2(Xe + 2F2 → XeF4) ∆H = -502 kJ

22. 10.8 Quality versus Quantity of Energy • One of the most important characteristics is that it is conserved • Eventually all energy will take the form of heat and spread evenly throughout the universe and everything will be the same temperature • This means work won’t be able to be done and universe will be dead; called “heat death” • We care more about what kind of energy (quality) than the amount of energy (quantity)

23. 10.9 Energy and Our World • Fossil Fuels formed by decaying products of plants • Petroleum • Natural Gas • Coal • Greenhouse Effect – Visible light travels through atmosphere, converted to infrared radiation (heat) which is absorbed by certain molecules, H20 and CO2 mainly, which radiate it back to earth

24. 10.10 Energy as a Driving Force • Energy Spread – in any given process, concentrated energy is dispersed widely • Happens with every exothermic reaction • When gas is burned, energy stored is dispersed into surrounding air • Matter Spread – molecules of a substance are spread out and occupy a larger volume • Salt dissolves in water due to matter spread • These 2 processes are important driving forces that cause events to occur

25. Entropy • Invented function that keeps track of disorder • Entropy (S) is a measure of disorder or randomness • So a cube of ice has a a lower S value than steam • Energy spread and Matter spread lead to greater entropy • The entropy in the universe is always increasing