1. 13.1 – NOTES Energy and Entropy

2. Energy ability to do work or produce heat; kinetic vs. potential
A. The Nature of Energy
Law of conservation of energy similar to conservation of matter; energy can not be created or destroyed in a chemical reaction, but may be converted from one form to another;
Chemical potential energy - energy stored in a substance b/c of its composition
potential energy can often be thought of as heat, q
heat is energy that is in the process of flowing from a warmer object to a cooler object;

3. Measuring heat calorie – heat required to raise the temp of one gram of a pure water by 1°C
food Calories = 1,000 calories; 1 calorie = joules

4. B. Specific Heat
Specific heat, C, is the amount of heat necessary to raise the temp of 1 gram of the substance by 1°C
unique to each substance
water is considered to have a HIGH specific
heat; why large bodies of water remain cool
long into summer and warm long into fall
swimming pools are cooler in the
morning;

5. The specific heat of water = 4. 184 J/gºC, while that of Al = 0
The specific heat of water = J/gºC, while that of Al = J/gºC and Au is J/gºC
Specific heats: ice: J/gºC steam: J/gºC
Calculating heat evolved and absorbed
q = mC∆T

6. Examples: a. 10. 0 g of water is heated from 21. 5ºC to 35. 6ºC
Examples: a g of water is heated from 21.5ºC to 35.6ºC. How much heat is required?
q = mCΔT = (10.0)(4.184)(14.1) = 590. J

7. b. A 145 g sample of an unknown substance was heated from 25.0ºC to 45.0ºC with the absorption of 5750 J of energy. What is the specific heat of the substance?
C = q/mΔT = 5750/145/20.0 = 1.98J/g°C

8. C. Calculations for water
Heat of fusion for water: 6.0 kJ/mole ΔHfus = 6.0 kJ/mol x 1000J/kJ / 18g/mol = 333 J/g
Definition: energy required to melt a substance;
solid  liquid = 6.0 kJ/mol
liquid  solid = -6.0 kJ/mol (heat of solid);

9. Heat of vaporization for water: 40. 6 kJ/mole ΔHvap = 40
Heat of vaporization for water: kJ/mole ΔHvap = 40.6 kJ/mol x 1000J/kJ / 18g/mol = 2260 J/g
Definition: energy required to boil a substance
liquid  gas = 40.6 kJ/mol
gas  liquid = kJ/mol (heat of condensation)

10. Examples: a. What energy is involved when 20
Examples: a. What energy is involved when 20.0 grams of ice is heated from −7.0ºC to 0ºC, then melted and heated to 25ºC. -7.0°C = 0°C in sig figs or do sig figs count???
Step #1 = q = mCΔT = (20.0)(2.077)(7.0) = J
(qsolid)
Step #2 = q = mΔHfus = (20.0)(333) = 6,660 J
(ΔHfus)
Step #3 = q = mCΔT = (20.0)(4.184)(25.0) = 2092 J
(qliquid) = 9040J

11. How much energy is involved when 50
. How much energy is involved when 50.0 g of steam at 109ºC is cooled, condensed, and cooled down to 10ºC?
Step #1 = q = mCΔT = (50.0)(2.042)(9.0) = J
Step #2 = q = mΔHvap = (50.0)(2260) = ,000 J
Step #3 = q = mCΔT = (50.0)(4.184)(90.0) = 18,828 J
133, 000J

12. Using the sun’s energy could supply all energy needs, but practicality has delayed development (amount of daylight & cloud cover)
storage is critical
– solar ponds (disadv: large quantities of land and not 100% efficient)
photovoltaic cells – convert solar energy to electricity (adv – clean and efficient; disadv – costly)
hydrates – when heated dissolves in own water of hydration, releases energy when cooled to re-crystallize;

13. II. Heat in Chemical Reactions and Processes
A. Measuring heat
A calorimeter insulated device used for measuring amount of heat absorbed or released during a chemical or physical process
place a known mass of water in an insulated chamber in the calorimeter to absorb the energy released by the process – increase/decrease in temperature
can use a Styrofoam cup – this type of device is open to the atmosphere and occurs at a constant pressure;

14. B. Chemical energy and the universe
Thermochemistry: study of heat changes that accompany chemical reactions and phase changes
Thermochemical equations: balanced chemical equation that includes states and energy (q) expressed as ΔH;

15. System and surroundings:
system = specific part of the universe that contains the reaction or process being studied
surroundings = everything else
universe = system + surroundings;

16. Exothermic reaction heat is released and energy is considered a product
heat flows from the system to the surroundings
heat pack – heat produced by the reaction (sodium acetate supersaturated solution that precipitates) flows from the pack into cold hands
G  L  S

17. Endothermic reaction heat is absorbed and energy is considered a reactant
heat flows from the surroundings to the system
cold pack – heat produced by the reaction (ammonium nitrate dissolving in water) flows from the pack into cold hands
S  L  G

18. ∆Hrxn = Hproducts - Hreactants
Enthalpy and enthalpy changes - heat content of a system at constant pressure, H
can’t measure the absolute value, but can measure the ΔH;
∆Hrxn = heat of reaction = q
change in enthalpy for a rxn is enthalpy (heat) of reaction; change in enthalpy that exists at the end compared to the initial enthalpy;
∆Hrxn = Hproducts - Hreactants

19. Signs of ∆H:
Exothermic – 4Fe(s) + 3O2(g)  2Fe2O3(s) kJ
reactants are losing heat, therefore, Hprod < Hreact; when Hreactants is subtracted from smaller Hproducts a negative value is obtained;

20. Endothermic – 27 kJ + NH4NO3(s)  NH4+(aq) + NO3-(aq)
reactants are gaining heat, therefore, Hprod > Hreact; when Hreactants is subtracted from larger Hproducts a positive value is obtained;

21. III. Energy as a driving force
In nature, processes tend to move toward lower energy states. This means that MOST reactions are exothermic (in an attempt to release heat = low energy)
Spontaneous reactions _physical or chemical changes that occur with no outside intervention. Ice melts at room temperature – endothermic spontaneous process

22. A. Reaction pathways
Activation energy: the minimum energy needed to initiate a reaction (the height of the energy barrier to formation of products)
exothermic reactions – molecules collide with enough energy to overcome the activation energy barrier, form an activated complex, then release energy and form products which are at a higher energy level
endothermic reactions – the reactant molecules must absorb energy to overcome the activation energy barrier and form high-energy products

23. Exothermic: Endothermic:

24. IV. Entropy as a driving force
In nature, processes tend toward higher entropy. This means molecules are more likely to exist in a high state of disorder (mixed) than in a low state (unmixed) _
Why do alcohol and water mix easily? There is no energy change, so the driving force must be entropy! Disorder!;
ΔS = Sprod – Sreact; ΔS > 0

25. Entropies of the phases: compare solid, liquid, gas
ΔSgas > ΔSliquid > ΔSsolid;
H2O(l)  H2O(g) ΔSsys > 0
H2O(s)  H2O(l) ΔSsys > 0

26. The dissolving of a gas in a solvent ALWAYS results in a decrease in entropy
CO2(g)  CO2(aq) ΔSsys < 0
Assuming no change in physical state, entropy of system usually increases when the # of gaseous product particles is greater than the # of gaseous reactant particles
2SO3(g)  2SO2(g) + O2(g) ΔSsys > 0
With some exceptions, you can predict the change in entropy when a solid or liquid dissolves to form a solution
NaCl(s)  Na+(aq) + Cl-(aq) ΔSsys > 0
An increase in temp of a substance is ALWAYS accompanied by an increase in random motion of the particles

27. Complete the following.
For each of the following reactions does the entropy increase or decrease? Explain each:
KClO3(s) → KCl(s) /2 O2(g)
Increase, gas produced
H2(g) O2(g) → H2O(l)
Decrease, liquid produced from gas
C8H18(l) O2(g) → CO2(g) H2O(g)
Increase, MORE gas produced
H2(g) Cl2(g) → HCl(g)
Stays the same, no change

28. For each of the following processes, does the entropy increase or decrease?
1. Organizing silverware in a drawer. Decrease
2. Dissolving sugar in iced tea. Increase
3. Melting an ice cube. Increase
4. Crystallizing iodine vapor. Decrease