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 (nucleus) of an 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. Reaction coordinate Exothermic Potential Energy Diagram Potential Energy of the Reactants = Potential Energy of the Products = Activation Energy = Heat of Reaction = 700kJ Potential Energy 500kJ 300kJ

7. Endothermic Reactions • 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

8. Endothermic Potential Energy Diagram • Potential Energy of the Reactants = • Potential Energy of the Products = • Activation Energy = • Heat of Reaction =

9. 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

10. Heat of Reaction (H) • Heat absorbed/released by a reaction • H = HP - HR • If H is negative, Exothermic • If H is positive, Endothermic • The heat of reaction for selected reactions is given on Reference Table I

11. 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)

12. Catalyst • Speeds up the reaction • Lowers the activation energy • Provides an alternate pathway for the reaction (still start and end in the same spot)

13. Heat and Temperature • Heat and temperature are not the same • Heat – type of energy • Temperature – measures 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 low until an equilibrium is established (Hot  Cold)

14. 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

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

16. 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)

17. Liquids • Definite Volume • Take the shape of the container • High attractive forces between molecules (but not as high as those found in solids) • Particles move freely • Particles have more kinetic energy (compared to the solid phase)

18. Gases • No Definite Volume • No Definite Shape • They take the shape of the container and volume of their container • Very spread out • Weak attractive forces • Particles move rapidly • Particles have very high kinetic energy

19. Melting • Solid to liquid phase change • Also called fusion • Endothermic

20. Freezing • Liquid to solid phase change • Also called solidification or crystallization • Exothermic

21. Boiling • Liquid to gas phase change • Also called vaporization or evaporation • Endothermic

22. Condensation • Gas to liquid phase change • Endothermic

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

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

25. Melting Point • Temperature where the solid and liquid phase exist in equilibrium • Equivalent to the freezing 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 and the particles spread apart • Melting

26. Boiling Point • Temperature where the liquid and gas phase exist in equilibrium • Equivalent to the point of condensation • As energy is added to the liquid, KE of the particles increases until they have sufficient energy to overcome the intermolecular forces holding them in place, the particles enter the gas phase and spread out

27. 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

28. Vapor Pressure / Boiling Point • A liquid will boil when its 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

30. 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?

31. Heat of Fusion • Heat needed to melt • Amount of heat needed to convert 1.0 gram 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 (Table B)

32. 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.

33. 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

34. Heat of Vaporization • Heat to vaporize / boil / evaporate • Amount of heat needed to convert 1.0 g of a substance from liquid to vapor at a constant temperature

35. Heat of Vaporization • q = mHv q = heat (J) m = mass (g) Hv = heat of vaporization Hv for water = 2260 J/g (Table B) 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.

36. 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

37. Change in Temperature Calcs • q = mcT • q = heat (Joules) • m = mass • C = specific heat capacity • Specific heat capacity = the amount of energy needed to raise the temperature of one gram by one degree Celsius • (for water; c = 4.18 J/goC) *given on Table B • T = change in temperature

38. Examples • How many joules does it take to raise the temperature of a 5.0g sample of water from 20.oC to 65oC? • A 1000. gram mass of water in a calorimeter has its temperature raised 5.0oC. How much heat energy was transferred to the water? • 17,500 J of heat are absorbed as a sample of water is heated from 12oC to 35oC. Find the mass of the sample.

39. 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?

40. 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?

41. 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

42. Heating Curve

43. Heating Curves • Diagonals • Only one phase is present • Temperature is increasing • Kinetic energy is increasing • Potential energy remains constant • Horizontals • Phase change is occurring • 2 phases are present • Temperature is constant • Kinetic energy remains the same • Potential energy increases

44. Cooling Curve

45. Cooling Curve • Diagonals • Only one phase is present • Temperature is decreasing • Kinetic energy is decreasing • Potential energy remains constant • Horizontals • Phase change is occurring • 2 phases are present • Temperature is constant • Kinetic energy remains the same • Potential energy decreasing