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Ch and 11 Thermodynamics
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2. Energy Definition: the capacity to do work, or to produce heat;
SI (metric) unit is Joule
Non-SI unit is calorie
1 cal = 4.18 J
1 calorie = the amount of energy needed to raise 1 gram of water 1° C
Heat - energy that is transferred from a warm object to a cooler object; represented by “q”
Temperature - a measure of the average kinetic energy of an object

3. Specific heat/Change of Temp
Quantity of heat needed to raise 1 gram
of a substance 1° C; unit is J/g°C
FORMULA: q= c x m x ΔT
q= heat (J)
c = specific heat (J/g°C)
m = mass (g)
T = change in temp. (°C)

4. Specific Heat Practice Problem #1
Calculate the amount of heat in joules needed to warm 250. g of water from 25.0°C to 95.0°C. (c=4.184 J/°C g)

5. Practice Problem #2
How much heat is lost when grams of Al is cooled from °C to 62.0 °C? The specific heat of Al is J/g°C

6. Change of state and heat
heat of fusion- amount of heat needed to melt 1 gram of a substance at its melting point
Hf copper = 205 J/g
Hf water = 80 cal/g= 334 J/g
q = mHf
heat of vaporization- amount of heat needed to boil 1 gram of a substance at its boiling point
Hv water = 540 cal/g = 2260 J/g
q= mHv

7. Heating/Cooling Curves
Change of Temperature
Change of Phase
Know which formula to use when!

8. Change of state Practice Problems
Calculate the amount of heat, in Joules, needed to melt 70.0g of copper at its melting point.
Calculate the heat required, in calories, to change 250g of water at 100°C to steam at 100°C.
Calculate the amount of heat needed to change 20g of ice at -10.0°C to water at 80.0°C.

9. Enthalpy (H) a measure of heat content of a system
H = change in heat content that accompanies a process
Hrxn = Hfinal - Hinitial
Hrxn = Hproducts - Hreactants
** ΔHrxn can also be written as ΔHf, for heat of formation**
**Chemical systems in the world tend to achieve the lowest possible energy. Would this occur in an exothermic or an endothermic reaction?

10. Exothermic reactions
chemicals react and give off heat (feel hot); H is negative; products are more stable
4Fe + 3O2  2Fe2O kJ
reactants
products

11. Endothermic reactions
chemicals need to absorb energy in order for the reaction to take place (feel cool); H is positive; reactants are more stable
27 kJ + NH4NO3  NH4+ + NO3-
products
reactants

12. Practice problems #1
CO (g) + NO (g)  CO2 (g) + N2 (g)

13. Practice #2
CH4 (g)+ O2 (g)  CO2 (g)+ H2O (g)

14. Practice #3
N2 (g) + O2 (g)  NO2 (g)

15. Entropy (S) Measure of disorder or chaos in a system
Law of disorder – states that things move spontaneously in the direction of maximum chaos or disorder
ΔSsystem = SP - SR
If ΔS is +, there is an increase in entropy.
If ΔS is -, there is a decrease in entropy

16. Rules for disorder 1. Entropy increases as particles move apart.
Gas > Liquid > Solid
I2 (s)  I2 (g)
2. Entropy increases when you divide a substance into parts (when the total number of products > the total number of reactants)
2H2O  2H2 + O2
3. Entropy increases as temp increases b/c particles move faster. (unit for entropy is J/K mole)
4. Entropy increases when you dissolve a solid into a liquid. Entropy decreases when you dissolve a gas into a liquid.

17. Entropy Examples Water (liquid) - water (solid) KCl(s)  KCl (l)
C(s) + O2 (g)  CO2 (g)

18. Gibbs’ Free Energy (G) Energy available to do work
Relates enthalpy (H) and entropy (S), using the equation:
ΔG = ΔH – TΔS
ΔG = GP - GR

19. Still confused? Try these online notes

20. Spontaneous Reactions
If ΔG is negative, the reaction will occur spontaneously.
If ΔG is positive, the reaction will NOT occur spontaneously.

21. Quick Spontaneous Chart- + ΔS ΔG
Depends on temp
spontaneous
Non-