Energy and Phases

1. Energy and Phases

2. Potential Energy - stored energy (stored in bonds, height) • Kinetic Energy - energy of motion, associated with heat

3. Forms of Energy • Light - light waves, electromagnetic radiation • Electrical • Chemical • Heat • Mechanical - moving parts, machines • Atomic/Nuclear - changes in mass of atom

4. Conservation of Energy • Energy can be converted from one form to another but never destroyed. • The total amount of energy is always constant.

5. Exothermic Reactions • Energy (heat) exits • Releases energy (heat) when new products are formed • Potential Energy of the reactants is greater than the potential energy of the products • Surroundings feel warm because heat was released • Heat is a product Ex: AB  A + B + heat

6. Endothermic Reactions • Endothermic – energy (heat) goes in • Reaction absorbs heat from the environment • Potential energy of the reactants is less than the potential energy of the products • Surroundings feel cold because heat was absorbed from the surrounding • Heat is a reactant • Ex: A + B + heat  AB

7. Activation Energy • The energy needed to break the bonds (to get the reaction started) • Energy to get over the “hill”, difference between your starting point and the top of the hill • All reactions need activation energy

8. Heat of Reaction (H) • Heat absorbed/released by a reaction • H = Hp - HR • If H is negative, Exothermic • If H is positive, Endothermic

9. Table I • Is the reaction, C(s) + O2(g) → CO2(g), exothermic or endothermic? • What is the value of ΔH for the reaction, CO2(g) → C(s) + O2(g)

10. Catalyst • Lowers the activation energy • Speeds up the reaction

11. Heat and Temperature • Heat and temperature are not the same • Heat – measure of the energy transferred from one substance to another • Temperature – measure of the average kinetic energy of a substance’s particles • The faster the particles move (more KE), the higher the temperature • Heat flows from high to lower until an equilibrium is established (Hot  Cold) • Calorimeter – measures the heat given off by a reaction

12. Heat/Temp Examples • If two systems at different temperatures have contact with each other, heat will flow from the system at a. 20oC to a system at 303K b. 30oC to a system at 313K c. 40oC to a system at 293K d. 50oC to a system at 333K 2. Which is not a form of energy? a. Light b. Temperature c. Heat d. Motion 3. Which has the most kinetic energy? a. 10.0 g of H2O at 70oC b. 10.0 g of H2O at 5oC c. 25.0 g of H2O at 60oC d. 25.0 g of H2O at 10oC

13. Solids • Definite shape • Definite volume • Very high attractive forces between molecules • Neighboring particles are very close together • Crystalline structure

14. Solids • Kinetic Energy – solids have KE • Particles are constantly vibrating (around their positions in the crystal) • Positions do not change in relation to the other particles in the crystal • At absolute zero (O K) all movement stops (theoretically)

15. Melting Point • As energy is added to the solid, KE of the particles increases until they have sufficient energy to overcome the forces holding them in the crystal, the substance begins to melt • Temperature where the solid and liquid phase exist in equilibrium

16. Heat of Fusion • Heat needed to melt • Amount of heat needed to convert a unit mass of a substance from a solid to liquid at a constant temperature • q = mHf q = heat (J) m = mass (g) Hf = heat of fusion (J/g) Hf for water = 334 J/g

17. Hf Examples • How much energy is needed to change 75g of ice at OoC to water at the same temperature? • 11,000J of heat are released as a sample of water at OoC freezes. Calculate the mass of the sample.

18. Melting/Freezing • Melting is an endothermic process, the H2O absorbs 334J/g • Freezing is the reverse process, so it is exothermic, the H2O releases 334J/g

19. Sublimation • Solid phase  gas phase (skip liquid) • Endothermic Example: dry ice CO2(s)  CO2(g) Iodine I2(s)  I2(g) Naphthalene (moth balls)

20. Liquids • Definite Volume • Take the shape of the container • High attractive forces between molecules (but not as high as those found in solids)

21. Evaporation (Vaporization, Boiling) • Change from liquid to gas, endothermic • As temperature increase, rate of evaporation increases • Increased temperature, more KE, easier to overcome intermolecular forces and to break them • Condensation - change from gas to liquid, exothermic

22. Vapor Pressure • Pressure of a gas on a liquid • In a closed system (sealed container) the vapor evaporating from the liquid exerts pressure on the liquid • Vapor Pressure increases as the temperature of the liquid increases • It has specific values for each substance at any given temperature • Table H

23. Boiling Point • Temperature where vapor pressure = atmospheric pressure • Water boils at 100oC at 101.3kPa (1atm) • Pressures below 101.3kPa (high elevations), water boils below 100oC • Pressures greater than 101.3kPa (below sea level), water boils above 100oC

25. Table H Examples • What is the vapor pressure of water at 105oC? • If the pressure is 30kPa, what temperature will water boil at? • What pressure is needed for ethanol to boil at 50oC? • Which liquid on Table H has the strongest intermolecular forces?

26. Heat of Vaporization • Heat to vaporize • Amount of heat needed to convert a unit mass of a substance from liquid to vapor at a constant temperature • Does not increase KE, so temperature does not change

27. Heat of Vaporization • q = mHv q = heat (J) m = mass (g) Hv = heat of vaporization Hv for water = 2260 J/g Examples: • How much heat must be supplied to evaporate 50.g of H2O at 100oC? • 12,750 J of heat are used to boil a sample of water. Calculate the mass of the sample.

28. Boiling and Condensation • Boiling is an endothermic process, the H2O absorbs 2260J/g • Condensation is the reverse process, so it is exothermic, the H2O releases 2260J/g

29. Deposition • Reverse of sublimation • Direct change from the gas to the solid phase, skip liquid • Exothermic • Example: Frost

30. Phase Diagrams / Heating and Cooling Curves • When a sample of matter is heated its temperature usually increases.  Increase in KE causes an increase in temp. • Sometimes matter can gain or lose heat without changing temperature.  What is going on at this point?

31. Points to remember: • When heat is used to increase the speed of particles, the temperature increases – at this point the KE is changing • When heat does not cause a change in temperature, it is being used to change phases (phase changes occur at the flat parts of the graph) – at this point PE is changing • Melting Point - point when the solid begins to melt, both the solid and liquid phases are present • Boiling Point - point when the liquid begins to boil, both the liquid and gas phases are present

32. Rate of Heating • Total amount of heat absorbed = time x rate of heating q= t x rate Example: A solid is heated at a constant rate of 150 J/min for 3 minutes. How much heat is absorbed?

33. Heating Curve

34. Cooling Curve

35. Change in Temperature Calcs • q = mcT • q = heat (Joules) • m = mass • C = specific heat capacity (for water; c = 4.18 J/goC) • T = change in temperature

36. Examples • How many joules does it take to raise the temperature of a 5.0g sample of water by 20.oC? • A 1000. gram mass of water in a calorimeter has its temperature raised 5.0oC. How much heat energy was transferred to the water?