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Table I
0 Exothermic reactions release heat and have negative
values.
0 Example: When Carbon and Oxygen react they release
393.5kJ of heat per mole reacted.
0 Endothermic reactions absorb heat and have positive
values.
0 Example: When Nitrogen and Oxygen react they absorb
182.6kJ of heat per mole.
Table I examples
0 When 2 moles of CH4 burn in oxygen, how much heat
is released?
0 When C2H4 is formed, is heat released or absorbed?
0 Reactions that release the most energy are the most
stable. Which reaction becomes the most stable?
0 Where did these values come from?
Calculating Heat of reactions
0 q is the symbol for heat.
0 If q is positive, the heat is endo.
0 If q is negative, the heat is exo.
0 q is measure in Joules, (J) or kilojoules (kJ).
0 The heat of a reaction is based on the mass of the
substance, the temperature change it undergoes and
specific heat.
Specific Heat
0 Specific heat is the heat needed to raise the temperature of
one gram of a substance one degree Celsius.
To calculate heat:
q = mcΔT
1. The temperature of 95.4g of
copper increases from 25 to 48C
and absorbed 849J. Calculate
copper’s specific heat.
Q= (95.4) (4.18) (48-25)
Q = 0.387 J/gC
q = mcΔT
2. How much heat is needed to raise the
temperature of 100g of water 50C?
Q = (100) (4.18) (50)
Q = 20900J or 20.90 KJ
q = mcΔT
3. If 600J are needed to heat 50g of water to
100C, what is the initial temperature?
600 = (50) (4.18) (x)
X = 2.87
97.13C
Kinetic Molecular Theory of Gases
KMT describes perfect gases:
0 Gases move in constant, random, straight-line paths.
0 Gases are separated by large distances, much larger than
their particle size. Therefore, gases volume is negligible. And
gases are easily compressed.
0 Gases do not have attractive or repulsive forces between
molecules.
0 Collisions between molecules can transfer energy but the
total energy of the system is constant. This is called an
elastic system.
Kinetic Molecular Theory of Gases
In summary, Perfect gases:
0 Have no mass
0 Have no volume
0 Have no intermolecular forces
Kinetic Molecular Theory of
Gases
0 But we don’t have perfect gases. How do real gases
deviate from ideal gases?
0 They have a volume, mass and small IMF under
high pressure and low temperature.
0 So, a real gas must be hot and under low pressure
to behave like an ideal gas.
Pressure
0 Gases exert a pressure on
surrounding substances
because they are constantly
moving and colliding with other
surfaces.
0 Only in a vacuum, where there
are no molecules, there is no
pressure.
0 Gas pressure can be measured
in atmospheres or kilopascals,
according to reference table A.
Liquids
0 No definite shape
0 Definite volume
0 Constant motion
0 No arrangement
0 Molecules are closer
together than a gas
Solids
0 Definite shape
0 Definite volume
0 Constant vibration
0 Molecules are packed
tightly in a geometric
(crystalline) pattern
Phase Changes
Identify the phase change and if it’s
endothermic or exothermic:
Liquid to gas
endothermic
0 Evaporation
0 Condensation
Gas to liquid
exothermic
0 Melting
Solid to liquid endothermic
0 Freezing
Liquid to solid exothermic
0 Sublimation
Solid to gas
endothermic
0 Deposition
gas to solid
exothermic
Thermochemistry
0 The study of energy changes that occur in chemical
reactions.
0 Kinetic energy refers to energy of motion.
(Temperature)
0 Potential Energy refers to stored energy.
Phase Change Diagrams
Cooling Curve
A
C
B
E
D
F
When can you use q=mcT?
0 Only on the solid,
liquid and gas only
lines. (Where the
temperature
changes)
0 So, what equations
do we use if the
temperature is not
changing?
Two more equations from Table T
0 Heat of vaporization: heat needed to change a
substance from gas to liquid or liquid to gas.
q=mHv
0 Heat of fusion: heat needed to change a substance
from solid to liquid or liquid to solid.
q=mHf
0 If the IMF is strong, the heats of vaporization and
fusion is high.
Q=mHv
1.
Calculate the number of joules needed to vaporize
423g of H2O.
Q = (423) (2260)
955, 980J or 955.98KJ
Q=mHf
0 How much heat is needed to melt ice at 0C if the
sample weighs 255g?
Q = (255) (334)
85,170J or 85.17 KJ
0 When you finish # 1-4, try these:
5. How much heat is absorbed by 550g block of ice to raise the
temperature from -15 to 0C?
6. How much heat energy must be absorbed to raise the temperature
of a 200 gram block of ice from -10 to 0C and then completely melt
it to a liquid at the same temperature?
7. How much energy would be required to heat the same 200g of
liquid water in #6 (at 0C) to the normal boiling point of water and
then vaporize it?
8. If the temperature of the 200 grams of steam generated in #7
were heated to a new temperature of 120C, how much energy
would be absorbed?
BONUS: What is the total amount of energy needed to heat 150g of
ice at -10C to gas at 140C? (Use a heating curve to help you).
Measuring heat in the lab
You can measure the heat of
physical and chemical
changes in a calorimeter.
The calorimeter acts like a
styrofoam cup, it insulates
the reaction (doesn’t let
the overall heat change).
Measuring heat in the lab
0 The heat released by the
reaction equals the heat
absorbed by the water.
0 You will measure the
change in heat of the
water using q=mcT.
Measuring heat in the lab
0 You will use a calorimeter
more like this.
0 You must make sure you
always stir the solution to
make the heat equal
throughout the cup.
A student places a 68.4g piece of metal at 99.5C in a calorimeter
filled with 103g of water at 25.2C. The temperature changes to
27.6C.
1.
2.
3.
4.
5.
In terms of the metal, is the reaction endothermic or
exothermic?
Calculate the heat absorbed by the water.
Calculate the heat released by the metal.
Calculate the specific heat of the metal.
Using the following specific heats, determine the identity of
the metal and calculate the % error.
Aluminum: 0.21 J/gC
Copper: 0.090 J/gC
Gold: 0.030 J/gC
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