1. Phase Changes and Heat

2. Phase Changes

3. THE NATURE OF ENERGY Energy - the ability to do work or produce heat
2 types: potential and kinetic
Potential – stored energy
Kinetic – energy of motion

4. Phase Change Terminology
Solid (s)– volume and shape are constant
Has low KE, high PE
Liquid (l)– volume is constant, but shape is not defined
As temperature increases, KE increases and PE decreases
Gas (g)– no defined volume or shape
Has high KE, low PE

5. Heating Curve

6. Another View of Phase Change Diagram

7. Cooling Curve High Kinetic Energy Gas Condensation Boiling Temp (C)
Liquid
Temp (C)
Freezing
Melting
Solid
Low Kinetic Energy
Time (min)

8. In Words…
Melting point – temperature at which a solid changes to a liquid
Freezing point – temperature at which a liquid changes to a solid
Boiling point – temperature at which a liquid changes to a gas
Condensation point – temperature at which a gas changes to a liquid
Sublimation – temperature at which a solid changes to a gas
Think of dry ice

9. Phase Changes and PE
Heat energy during the melting phase is being used to overcome attractive forces between molecules
Solid phase: molecules stay in place and vibrate
Liquid phase: molecules can flow past each other
With a phase change, there is a change in potential energy (energy of position of molecules next to each other)

10. Phase Change and PE
Melting point and freezing point are at the same temperature
Melting: PE increases
Freezing: PE decreases
Boiling point and condensation point are the same temperature
Boiling: PE increases
Condensation: PE decreases

11. Practice
Identify each phase and energy type (KE/PE) for sections a, c, and e: e d
Temp. c b a
Time

12. Heat

13. TEMPERATURE AND HEAT Temperature Heat
Measure of the average kinetic energy of the particles in a substance
Heat
The transfer of energy from one substance to another due to temperature differences
Always from high E to low E

14. We can measure heat… q = m x ∆T x c q = heat m = mass
∆T = change in temperature
Tfinal – T initial
c = specific heat capacity (specific heat)
The quantity of heat required to raise temp. of an object by 1oC

15. Using the formula We need 3 out of the 4 so we can solve for the 4th
q = m x ∆T x c
If have q, m, and c – what solving for?
If have m, c, and ∆T – what solving for?

16. Value of “q” What is “q” – HEAT q can be positive or negative
If q is positive – energy is added
Therefore reaction is endothermic
If q is negative – energy is removed/leaving
Therefore reaction is exothermic

17. You said what??
Exothermic process: a change (ie. a chemical reaction) that releases heat.
Ie: Burning fossil fuels
Think “exit”
Endothermic process: a change (ie. a chemical reaction) that absorbs heat.
Photosynthesis is an endothermic reaction (requires energy input from sun)

18. Practice Problem – #3 on Worksheet #2 in your packet
How much heat is required to raise the temperature of 854 g of water from 23.5oC to 85.0oC? (specific heat of water is 4.184J/goC)

19. Heat Transfer

20. Heat Transfer Heat will transfer from one object to something else
This is a transfer of energy from a place of high E to a place of low E
The transfer of E with change the temperatures

21. Heat Transfer in Water
reaction
Exothermic reaction, heat given off & temperature of water rises
reaction
Endothermic reaction, heat taken in & temperature of water drops

22. But what about temperature?
Exothermic reaction:
-temperature of water will increase
-temperature of object will decrease
UNTIL temperature of water and object are equal!!
Therefore Tfinal for water and object are the same
reaction

23. Endothermic reaction: -temperature of water will decrease
-temperature of object will increase
UNTIL temperature of water and object are equal!!
Therefore Tfinal for water and object are the same
reaction

24. Therefore… - q lost = q gained We can say that: q lost = q gained
But wait… one is going to be negative
Therefore we must:
- q lost = q gained

25. But what does q equal? -(m x ∆T x c) = m x ∆T x c - q lost = q gained
q = m x ∆T x c
Therefore we can say that:
- q lost = q gained
-(m x ∆T x c) = m x ∆T x c

26. Heat of Fusion

27. Heat of Fusion/Vaporization
How much heat is needed to change a solid to a liquid, or a liquid to a gas?
The amount of heat needed depends on:
How much substance you have (mass)
The substance itself – but we cannot use C because we are talking about phase changes…
So what do we use???

28. Formulas
Enthalpy change (ΔH)- Heat released or absorbed during a chemical reaction
q = m Hfus
q = m Hvap

29. Heat of Fusion
How much heat is needed to melt 25.0 g of ice? (Hfus ice is 334 J/g)
q = m Hfus
q = (25.0 g)(334 J/g)
q =

30. Heat of Vaporization
How much heat is needed to evaporate 25.0 g of water?
q = m Hvap
q = (25.0 g)(2260 J/g)
q =

31. Heat of Solution

32. Heat of Solution Heat of solution
The heat produced by a chemical reaction, or the heat required for a chemical reaction to occur
Heat is absorbed from the atmosphere or released into the atmosphere

33. Heat of Solution Endothermic reaction Exothermic reaction
Requires heat energy from the environment to get reaction to run
Heat is transferred from the environment to the reaction
Is a positive value
Exothermic reaction
Produces heat
Heat is transferred from the reaction to the environment
Is a negative number

34. Boiling Point Does water always boil at 100ºC? Boiling point: NO.
Boiling point depends on atmospheric pressure
Boiling point:
Is the temperature at which vapor pressure of the liquid is the same as the atmospheric pressure
Pressure of the bubbles = air pressure in the room