Ch. 10.4 and 11
Thermodynamics
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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
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)

Specific Heat Practice Problem #1
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)
Practice Problem #2
How much heat is lost when 50.0
grams of Al is cooled from 130.0 °C to
62.0 °C? The specific heat of Al is
0.897 J/g°C
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


Heating/Cooling Curves
Change of
Temperature
Change of Phase
Know which formula to use when!
Change of state Practice Problems
1.
2.
3.
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.
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?

Exothermic reactions
 chemicals
react and give off heat (feel
hot); H is negative; products are
more stable
 4Fe + 3O2  2Fe2O3 + 1625 kJ
reactants
products
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
Practice problems #1
 CO
(g) + NO (g)  CO2 (g) + N2 (g)
Practice #2
 CH4 (g)+
O2 (g)  CO2 (g)+ H2O (g)
Practice #3
 N2
(g) + O2 (g)  NO2 (g)
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

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.
Entropy Examples

Water (liquid) - water (solid)

KCl(s)  KCl (l)

C(s) + O2 (g)  CO2 (g)
Gibbs’ Free Energy (G)
Energy available to do work
 Relates enthalpy (H) and entropy (S),
using the equation:



ΔG = ΔH – TΔS
ΔG = GP - GR
Still confused?
Try these online notes
 http://www.sciencegeek.net/Chemistry/Po
werpoint/Unit7/Unit7_files/frame.htm

Spontaneous Reactions
If ΔG is negative, the reaction will occur
spontaneously.
 If ΔG is positive, the reaction will NOT
occur spontaneously.

Quick Spontaneous Chart
ΔH
-
+
+
-
ΔS
+
-
+
-
ΔG
-
+
Depends
on temp
Depends
on temp
spontaneous
Nonspontaneous