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 Light - light waves, electromagnetic radiation Electrical Electrical Chemical Chemical Heat Heat Mechanical - moving parts, machines Mechanical - moving parts, machines Atomic/Nuclear - changes in mass of atom 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 - H R If  H is negative, Exothermic If  H is positive, Endothermic

9. Table I 1.Is the reaction, C(s) + O 2 (g) → CO 2 (g), exothermic or endothermic? 2.What is the value of ΔH for the reaction, CO 2 (g) → C(s) + O 2 (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 1.If two systems at different temperatures have contact with each other, heat will flow from the system at a. 20 o C to a system at 303K b. 30 o C to a system at 313K c. 40 o C to a system at 293K d. 50 o C to a system at 333K 2. Which is not a form of energy? a. Lightb. Temperature c. Heatd. Motion 3. Which has the most kinetic energy? a. 10.0 g of H 2 O at 70 o Cb. 10.0 g of H 2 O at 5 o C c. 25.0 g of H 2 O at 60 o Cd. 25.0 g of H 2 O at 10 o C

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

14. 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 = mH f q = heat (J) m = mass (g) H f = heat of fusion (J/g) H f for water = 334 J/g

17. H f Examples 1.How much energy is needed to change 75g of ice at O o C to water at the same temperature? 2.11,000J of heat are released as a sample of water at O o C freezes. Calculate the mass of the sample.

18. Melting/Freezing Melting is an endothermic process, the H 2 O absorbs 334J/g Freezing is the reverse process, so it is exothermic, the H 2 O releases 334J/g

19. Sublimation Solid phase  gas phase (skip liquid) Endothermic Example: dry ice CO 2 (s)  CO 2 (g) IodineI 2 (s)  I 2 (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 100 o C at 101.3kPa (1atm) –Pressures below 101.3kPa (high elevations), water boils below 100 o C –Pressures greater than 101.3kPa (below sea level), water boils above 100 o C

25. Table H Examples 1.What is the vapor pressure of water at 105 o C? 2.If the pressure is 30kPa, what temperature will water boil at? 3.What pressure is needed for ethanol to boil at 50 o C? 4.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 = mH v q = heat (J) m = mass (g) H v = heat of vaporization H v for water = 2260 J/g Examples: 1.How much heat must be supplied to evaporate 50.g of H 2 O at 100 o C? 2.12,750 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 H 2 O absorbs 2260J/g Condensation is the reverse process, so it is exothermic, the H 2 O 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/g o C)  T = change in temperature

36. Examples 1.How many joules does it take to raise the temperature of a 5.0g sample of water by 20. o C? 2.A 1000. gram mass of water in a calorimeter has its temperature raised 5.0 o C. How much heat energy was transferred to the water?