2. Energy Part II Conservation of Energy First Law of Thermodynamics
Enthalpy of Reaction

3. Law Of Conservation Of Energy
Energy cannot be created or destroyed but can be transformed from one form of energy to another
Also known as the first law of thermodynamics

4. Internal Energy is Conserved
1st Law of Thermodynamics: For an isolated system the internal energy (E) is constant:
Δ E = Ef - Ei = 0
Δ E = Eproduct - Ereactant = 0
We can’t measure the internal energy of anything, so we measure the changes in energy
E is a state function
E = work + heat

5. Calorimeters A calorimeter is used to measure heat transfer.
can be made with a coffee cup and a thermometer.
indicates the heat lost by a sample
indicates the heat gained by water.

6. State Function
A property whose value depends only on the present state of the system, not on the method or mechanism used to arrive at that state
Position is a state function: both train and car travel to the same locations although their paths vary
The actual distance traveled does vary with path
New York
Los Angeles

7. ΔE= q + w
internal energy changes are state functions

8. Work and Pistons Pressure = force/area. force=P x area
Work = P x (area x d)
Work = -P×ΔV
In expansion, ΔV>0,
and is exothermic
Work is done by
the system in
expansion
W= force x d

9. How does work relate to reactions?
Work = Force · Distance
is most often due to the expansion or contraction of a system due to changing moles of gas.
Gases push against the atmospheric pressure , so Psystem = -Patm
w = -Patm×ΔV W=work done by system
Work = Force X distance

10. H = E + W; W= work done by system on surrounding
Enthalpy is the internal energy of the system plus the work done by the system on surroundings (H)
E= is the internal energy and it is the amount of heat absorbed or release by a system plus the work done by the surroundings
H = E + W; W= work done by system on surrounding
W = -P×ΔV (Isobaric system)
∆H = ∆E -P∆V (1)
E= Q + W ; W=work down by surroundings ∆E = ∆Q + ∆PV ∆Q = q
∆E = q + P∆V (2) Plugging (2) in equation (1)
∆H= q + P ∆V -P∆V AH = q
Isobaric: Constant pressure
ΔP =0

11. ∆H- Change in Enthalpy
The difference in enthalpy is the maximum amount of thermal energy derivable from a thermodynamic process in which the pressure is held constant.
Most processes involve transfers of thermal energy at a constant atmospheric pressure
∆H= q
Such changes can be easily measured and the heat evolved or absorbed is called the change in enthalpy and given the symbol ∆H

12. The Sign Convention
Endothermic systems require energy to be added to the system, thus the q is (+)
Exothermic reactions release energy to the surroundings. Their q is (-)
Energy changes are measured from the point of view of the system

13. Entropy Change Entropy (S)-disorder or randomness in a system or
universe
∆S =Change in disorder or randomness
For a change, as the system proceeds from…
order disorder
Entropy increases, system looses energy
disorder order
Entropy decreases, system gains
energy

14. Energy changes States Changes of Matter The heat of fusion
is the amount of heat released when 1 gram of liquid freezes (at its freezing point).
is the amount of heat needed to melt 1 gram of a solid (at its melting point).
for water (at 0 °C) is
80. cal
1 g water

15. Sublimation Sublimation
occurs when particles change directly from solid to a gas.
is typical of dry ice, which sublimes at -78 C.
takes place in frost-free refrigerators.
is used to prepare freeze-dried foods for long-term storage.

16. Evaporation and Condensation
Water
evaporates when molecules on the surface gain sufficient energy to form a gas.
condenses when gas molecules lose energy and form a liquid.

17. Boiling
At boiling,
all the water molecules acquire enough energy to form a gas.
bubbles appear throughout the liquid.

18. Heat of Vaporization The heat of vaporization is the amount of heat
absorbed to vaporize 1 g of a liquid to gas at the boiling point.
released when 1 g of a gas condenses to liquid at the boiling point.
Boiling Point of Water = 100 °C
Heat of Vaporization (water)
= cal
1 g water

19. Heating Curve

20. - Heat + Heat GAS LIQUID SOLID - Heat - Heat + Heat + Heat Deposition
Sublimation
+ Heat
GAS
LIQUID
Condensation
SOLID
Crystallization
- Heat
- Heat
+ Heat
+ Heat
Liquefaction
Vaporization

21. B Concept Question A solid with a high heat of fusion has:
a) high melting point and weak attractive forces between particles
b) high melting point and large attractive forces between particles
c) low melting point and weak attractive forces between particles
d) low melting point and large attractive forces between particles B

22. Concept Question
Why is the enthalpy from liquid to gas usually larger than the enthalpy from solid to liquid in a heating curve?

23. Chemical Potential Energy
Chemical bond: net attractive forces that bind atomic nuclei and electrons together
Exothermic reactions form stronger bonds in the product than in the reactant and release energy (-E) Reactions release energy to the surroundings. Their q is (-)
Endothermic reactions break stronger bonds than they make and require energy (+E) Requires energy to be added to the system, thus the q is (+)
Energy changes are measured from the point of view of the system

24. Energy changes Endothermic HEAT + Reactants Products Exothermic
Heat is the most common energy form observed in chemical reactions.
Most processes involve transfers of thermal energy at a constant atmospheric pressure
∆H= Q
Endothermic
HEAT + Reactants Products
the q is (+)
Exothermic
Reactants Products + HEAT
q is (-)

25. Physical / Chemical Changes:
A B + Heat Exothermic
X + Heat Y Endothermic
AH =q (sys)
Exothermic
q (sys) (-)
AH (-)
Endothermic
q (sys) (+)
AH (+)