1. Thermodynamics

2. The ability to do work I. The Nature of Energy Potential Energy
Kinetic Energy
Radiant Energy

3. 2. Thermal (heat) - energy caused by the random internal motion of
Kinetic Energy
Energy of motion
Types
1. Mechanical - moving
parts of a machine
2. Thermal (heat) - energy
caused by the random
internal motion of
particles of matter

4. Potential Energy
Stored energy
Types
1. Electrical - ex. battery
2. Gravitational - ex. water
behind dam - used for
electricity
3. Chemical - ex. chemical
bonds in food

5. Law of Conservation of Energy
In any process, energy is neither created nor destroyed.
ex. 1 hitting a baseball transfers kinetic energy from the bat to the ball
ex. 2 igniting a match changes chemical energy into heat and light

6. Laws of Thermodynamics
0th law- Defines temperature
If two systems are in thermal equilibrium respectively with a third system, they must be in thermal equilibrium with each other.
1st law- Energy cannot be created or destroyed.
When energy passes, as work, as heat, or with matter, into or out from a system, its internal energy changes in accord with the law of conservation of energy.
2nd law- Systems have a tendency to move toward disorder
 In a natural thermodynamic process, the sum of the entropies of the participating thermodynamic system increases.
3rd law- All molecular motion stops at 0 K
 The entropy of a system approaches a constant value as the temperature approaches absolute zero.

7. TEMPERATURE VS. HEAT Temperature Heat
A measure of the average kinetic energy of molecules in motion
Remember:
Kelvin = Celsius
Heat
Total amount of energy that flows between matter
Flows from matter of higher temperature to matter of lower temperature
“Hot” molecules quickly move into areas of slower moving “cold” molecules

8. Temperature and the Temperature Scale
Chapter 10
Visual Concepts
Temperature and the Temperature Scale

9. Heat -particles are always moving
The transfer of kinetic energy from a hotter object to a colder object.
Symbol is q
-particles are always moving
-when you heat water  molecules move faster

10. EXOTHERMIC REACTIONS Chemical reactions that release thermal energy
Feels hot – temperature rises
Examples: condensation, freezing

11. ENDOTHERMIC REACTIONS
Chemical reactions that absorb thermal energy
Feels cold – temperature drops
Examples: boiling, evaporation, melting

12. Kelvin Scale (K) K = °C + 273 Measuring Heat
0 K - point at which there is no molecular motion (absolute zero)
All Kelvin temperatures are positive
K = °C + 273
What is the boiling point of water in Kelvin?
What is the freezing point of water in Kelvin?

13. Energy stored in food = Calorie (Cal) 1 Cal = 1000 cal = 1 kcal
Measuring Heat
Calorie = amount of heat needed to raise the temperature of 1 gram of water by 1 degree Celsius
(1 cal = 1 g x 1 C°)
Energy stored in food = Calorie (Cal)
1 Cal = 1000 cal = 1 kcal
Joule (J) : 1 cal = J

14. Specific Heat Specific Heat
Amount of heat required to raise 1 gram of a substance 1°C.
physical property
Liquid water J/g°C
Fe J/g°C

15. Metals have low heat capacity
Specific Heat
Water has high heat capacity:
Absorbs a large quantity of heat with only a small increase in temperature
Gives up a large quantity of heat with only a small decrease in temperature
Metals have low heat capacity
Small amount of heat  large temperature change

16. HEAT q = cp• m • ΔT Where q = heat released (-) or heat absorbed (+)
cp = specific heat
m = mass
ΔT (means change in Temperature)
= Final Temp – Initial Temp

17. HEAT q = cp• m • ΔT Where q = Joules (J) or calories (cal)
cp = J/g•K or J/g•˚C or cal/g •˚C
m = grams
ΔT = K or °C

18. SAMPLE PROBLEM A
If 75 g of iron (cp = J/g•K) is heated from 274 K to 314 K, what is the amount of heat that is absorbed?

19. SAMPLE PROBLEM B
A 4.0 g sample of glass was heated from 274 K to 314 K and was found to have absorbed 32 J of energy as heat. What is the specific heat of this type of glass?

20. Calorimetry
Calorimeter - an insulated device to measure temperature changes
Can determine enthalpy changes (heat) of reactions

21. Calorimetry
Exothermic process  releases heat  temperature of surroundings increases
Endothermic process  absorbs heat  temperature of surroundings decreases

22. Transferred Surroundings Put hot iron ring into cool water
Calorimeter
qrxn = - qsur
Transferred Surroundings
Put hot iron ring into
cool water
Leave until the temp
Is the same for both
How does heat lost by
the iron compare to
heat gained by the
water?
Heat lost 
Iron
Heat gained
Water

23. THERMOCHEMISTRY- Study of the changes in heat in chemical reactions
All chemical reactions involve changes in energy.
Heat energy is either absorbed or released.

24. Enthalpy Enthalpy (H)-heat content of a substance
Depends on temperature, physical state, and composition
Enthalpy Change - the amount of heat absorbed or released during a chemical reaction; ΔH
ΔH = Hproducts – Hreactants

25. Bond Energy and Enthalpy
H2 + O2  H2O
Energy associated with chemical bonds is referred to as bond energy.
Energy is released when bonds are FORMED

26. Bond Energy and Enthalpy
2H2 + O2  2H2O
Energy is released when bonds are FORMED
H H
H O H
H H
H O
O O H

27. Reactions that release heat to their surroundings Combustion
Exothermic Reactions
Reactions that release heat to their surroundings
Combustion
C3H8(g) + 5 O2(g)  3 CO2(g) + 4 H2O(g) + heat
- Heat is produced because the energy released as new bonds are formed (products) > energy required to break the old bonds (reactants)

28. Al(s) +Fe2O3(s) Al2O3(s) +Fe(l) +heat
Exothermic Reaction
Products have
lower potential
energy than reactants.
energy released
Al(s) +Fe2O3(s) Al2O3(s) +Fe(l) +heat

29. If heat is released,
Hproducts < Hreactants
DH is negative (exothermic)

30. Endothermic Reactions
Reactions that absorb heat from surroundings
Ammonium nitrate in water
NH4NO3(s) + heat  NH4+(aq) + NO3-(aq)
Energy released as new bonds are formed (products) < energy required to break bonds (reactants)
This energy must be supplied by surroundings and is stored in the chemical bonds of the products

31. NH4NO3(s) + heat  NH4+(aq) + NO3-(aq)
Endothermic Reaction
Reactants have
lower potential energy than products
energy absorbed
NH4NO3(s) + heat  NH4+(aq) + NO3-(aq)

32. If heat is absorbed,
Hproducts > Hreactants
DH is positive (endothermic)

33. Calculating Heat of Reaction
How much heat will be released when 136 g of hydrogen peroxide decomposes?
2 H2O2(l)  2 H2O(l) + O2(g) ΔH° = -190 kJ
-190 kJ
2 mol H2O2
Ratio of coefficients

34. N2(g) + 2 O2(g)  2 NO2(g) ΔH°= +68.0 kJ
2. How much heat will be released when 184 g of NO2 is formed at STP?
N2(g) + 2 O2(g)  2 NO2(g) ΔH°= kJ

35. 4Al(s) + 3O2(g)  2Al2O3(s) ΔH = -3352 kJ
3. What mass of oxygen is needed to produce 1500 kJ of energy at STP?
4Al(s) + 3O2(g)  2Al2O3(s) ΔH = kJ