1. Energy of Chemical Reactions -OR-
Thermochemistry
Energy of Chemical Reactions -OR-
The study of heat changes that occur during chemical reactions and physical changes of state

2. Heat Transfer and Changes of State
Changes of state involve energy

3. Heating Curve Gas - KE  Boiling - PE  Liquid - KE  Melting - PE 
Solid - KE 

4. Heating/Cooling Curve for Water

5. Temp Changes?
How do you go from CelsiusFahrenheitKelvins

6. Energy and Chemistry
ENERGY is the capacity to do work or transfer heat.
can be: light, electrical, kinetic, potential, chemical
HEAT (represented by q) is the form of energy that flows between 2 samples because of a difference in temperature – will always go from a warmer object to a cooler one.
LAW OF CONSERVATION OF ENERGY energy is neither created nor destroyed

7. (1kcal = 4186J, 1 cal = 4.186 J and 1J = 0.239 cal)
Energy and Chemistry
CALORIE quantity of heat that raises the temperature of 1g of pure water 1°C
1 Cal or 1 kcal = 1000 cal
JOULE SI unit of heat and energy
(1kcal = 4186J, 1 cal = J and 1J = cal)
HEAT CAPACITY the amount of heat it takes to change an object’s temperature by exactly 1°C
Depends on an object’s mass
Ex. A cup of water has a greater heat capacity than a drop of water.

8. Chemical Potential Energy: Chemical bonds are a source of energy
CHEMICAL ENERGY
Chemical Potential Energy:
Energy stored in chemicals because of their compositions
Chemical bonds are a source of energy
BOND BREAKING - requires energy
BOND MAKING releases energy
In a chemical reaction:
if more energy is released in forming bonds than is used in
breaking bonds then
reaction is EXOTHERMIC
Energy is released as
HEAT, LIGHT, SOUND, WORK
if more energy is used in breaking bonds than is released in
forming bonds then
DEMONSTRATIONS:
Thermite reaction
Al + Fe2O3 -> Al2O3 + Fe
SPONTANEOUS HEAT, LIGHT reaction
Luminol
SPONTANEOUS LIGHT EMITTING REACTION
reaction is ENDOTHERMIC
- LIGHT - photochemistry
- WORK - electrochemistry
- COOLING of surroundings
Energy can be provided by

9. Specific Heat Capacity
A.K.A. Specific Heat: (represented by C) the amount of heat it takes to raise temperature of 1g of a substance 1°C (measured in J/(g x °C)
A difference in temperature leads to energy transfer.
Specific heat capacity =
heat lost or gained by substance (J)
(mass, g)
(T change, K)

10. Specific Heat q = mCT q = heat (j) m = mass (g)
C = specific heat (j/gC)
T = change in temperature = Tf-Ti (C)

11. Specific Heat Capacity - an example
If 25.0 g of Al cool from 310 K to 37 K, how many joules of heat energy are lost by the Al?
0.902 J/g.K =
heat gain/lost = q = (mass)(specific heat)(DT)
where DT = Tfinal - Tinitial = = -273 K
q = (0.902 J/g•K)(25.0 g)(-273 K)
q = J
negative sign of q  heat is “lost by” or transferred from Al

12. Specific Heat
A 10.0 g sample of iron changes temperature from 25.0C to 50.4 C while releasing 114 joules of heat. Calculate the specific heat of iron.

13. Example
q = mcT
114 = (10.0)c( )
c = j/g C

14. Another example
If the temperature of 34.4 g of ethanol increases from 25.0 C to 78.8 C how much heat will be absorbed if the specific heat of the ethanol is 2.44 j/g C

15. Another example
q = mcT
q = (34.4)(2.44)( )
q = 4.53 x 10 3 j

16. Yet another example
4.50 g of a gold nugget absorbs 276 J of heat. What is the final temperature of the gold if the initial temperature was 25.0 C & the specific heat of the gold is 0.129j/g C

17. Yet another example q = mcT 276 = (4.50)(0.129)T T = 475C
T = Tf-Ti
475 = Tf-25
Tf = 500 C

18. A. Kinetic Molecular Theory
KMT
Particles of matter are always in motion.
The kinetic energy (speed) of these particles increases as temperature increases.
Kelvin Temperature scale represents the relationship between temperature and average kinetic energy.
This relationship between temp. and energy is directly proportional
Example: particles of He at 200K have twice the average KE as particles of He at 100K

19. B. Four States of Matter Gases
high KE - particles can separate and move throughout container
Particles move in constant random motion
All collisions are perfectly elastic
variable shape
variable volume

20. B. Four States of Matter Liquids
low KE - particles can move around but are still close together
Particles flow past one another
Denser than gasses
variable shape
fixed volume

21. Evaporation
Definition: the conversion of a liquid to a gas or vapor below its boiling point (a.k.a. vaporization)
What happens…
Molecules at the surface of the liquid break away and go into the gas or vapor state
The liquid will evaporate faster as it is heated because KE increases
Particles with the highest KE escape first
Evaporation in a closed container is different

22. Particles collide with the walls of the sealed container and produce a vapor pressure above the liquid.
As the container stands, the number of particles turning into vapor increases
Some of the particles will condense and turn to liquid
After some time vapor particles condensing will equal particles vaporizing to create an equilibrium

23. B. Four States of Matter Solids
very low KE - particles vibrate but can’t move around
fixed shape
fixed volume
When heated, particles vibrate more rapidly and begin to spin as KE increases

24. B. Four States of Matter Plasma
very high KE - particles collide with enough energy to break into charged particles (+/-)
gas-like, variable shape & volume
stars

25. ENTHALPY DH = Hfinal - Hinitial Process is ENDOTHERMIC
For systems at constant pressure, the heat content is the same as the Enthalpy, or H, of the system.
Heat changes for reactions carried out at constant pressure are the same as changes in enthalpy (∆H)
DH = Hfinal - Hinitial
If Hfinal > Hinitial then DH is positive
Process is ENDOTHERMIC
If Hfinal < Hinitial then DH is negative
Process is EXOTHERMIC

26. Standard Enthalpy Values
Most DH values are labeled DHo
o means measured under standard conditions
P = 1 atmosphere ( = 760 torr = kPa)
Concentration = 1 mol/L
T = usually 25 oC
with all species in standard states
e.g., C = graphite and O2 = gas

27. Endo- and Exothermic ENDOTHERMIC EXOTHERMIC Heat Heat Surroundings
System
Surroundings
System
Heat
DEMONSTRATION
example of an ENDOTHERMIC reaction
- hydration of ammonium nitrate
(NH4)+(NO3)- (s) + H2O -> NH4 (aq) + NO3 (aq)
(cools surroundings enough to freeze water)
WHY SPONTANEOUS ?
ENDOTHERMIC
EXOTHERMIC

28. Heat and Changes of State
Heat of combustion (∆H)= the heat of reaction for the complete burning of one mole of a substance
Molar heat of fusion (∆Hfus)= the heat absorbed by one mole of a substance in melting from a solid to a liquid at a constant temperature
Molar heat of solidification (∆Hsolid)= heat lost when one mole of a liquid freezes to a solid at a constant temperature (equal to the negative heat of fusion)
Molar heat of vaporization (∆Hvap)= the heat absorbed by one mole of a substance in vaporizing from liquid to a gas
Molar heat of condensation (∆Hcond)= heat released by one mole of a vapor as it condenses