Unit 12 THERMODYNAMICS PPT

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Unit 12
THERMODYNAMICS: REACTION ENERGY
IMPORTANT DATES
Tuesday, May 8 – EOC Field Testing for
Chemistry
 Wednesday, May 9 – Unit 12 Test Review
 Thursday, May 10 – Unit 12 Test

DAY 1 NOTES
 The
Flow of Energy
 Energy
– the capacity to do work or supply
heat
 Chemical Potential Energy – energy stored
within the bonds of chemical compounds
 Activation Energy – the minimum energy
colliding particles must have in order to react
THERMODYNAMICS
Thermodynamics – the study of energy in chemical
reactions; literally means “changes in heat.”
 All chemical reactions either release or absorb
energy when they occur. Another way to say this
is that reactions either “give off” or “take in”
energy.
 The Law of Conservation of Energy – 1st Law of
Thermodynamics; states energy can neither be
created nor destroyed, only transformed
 ALL energy is either work performed, stored
potential, or heat lost.

REACTION SYSTEM, SURROUNDINGS, AND
UNIVERSE
 System
– the chemical reaction under
study
 Surroundings – every place in the
universe except the system
 Universe – the system and the
surroundings
REACTION SYSTEM, SURROUNDINGS, AND
UNIVERSE
TYPES OF HEAT FLOW

Exothermic Reaction – a reaction in which heat
is released by the system to the surroundings;
from the perspective of the system, “heat is
given off”
TYPES OF HEAT FLOW

Endothermic Reaction - is a reaction in which
heat is absorbed by the system from the
surroundings. From the perspective of the
system, “heat is taken in.”
QUALITATIVE PRACTICE – HEAT FLOW
MATCH THE REACTION DIAGRAM WITH THE APPROPRIATE DESCRIPTION OF THE FLOW OF
ENERGY.
Energy level diagram for an exothermic chemical reaction without showing the activation energy;
it could also be seen as quite exothermic with a highly unlikely zero activation energy, but reactions between
two ions of opposite charge usually has a very low activation energy.
Very endothermic reaction with a large activation energy.
Moderately exothermic reaction with a moderately high activation energy.
A small activation energy reaction with no net energy change; this is theoretically possible if the total energy
absorbed by the reactants in bond breaking equals the energy released by bonds forming in the products.
Very exothermic reaction with a small activation energy.
Energy level diagram for an endothermic chemical reaction without showing the activation energy; it could also
be seen as quite endothermic with zero activation energy.
Moderately endothermic reaction with a moderately high activation energy.
HOW TO READ THE GRAPH!
Potential Energy – Read from the x-axis to the reaction line
EX: b is the potential energy of ….. (answer to question 2)
A+B and C+D are NOT answers on the questions. They represent the reactants (A+B)
and the products (C+D).
It might help to draw
dashed lines so that you
can visualize the potential
energy all the way across
the graph
Reactants
Activation Energy
Products
ASSESSMENT – POTENTIAL ENERGY DIAGRAM
1.
2.
3.
4.
Is the above reaction endothermic or exothermic?
What letter represents the potential energy of the
reactants? b
What letter represents the potential energy of the
products? f
What letter represents
the change in energy for
the reaction?
d
ASSESSMENT – POTENTIAL ENERGY DIAGRAM
5.
6.
7.
8.
9.
What letter represents the activation energy of the forward
reaction? a
What letter represents the activation energy of the reverse
reaction (read the chart backwards)? e
What letter represents the
potential energy of the
activated complex? c
Is the reverse reaction
endo or exothermic?
If a catalyst were added,
what letter(s) would
change? a and c
Day 2
THERMODYNAMICS: REACTION ENERGY
REACTION ENERGY

Measuring and Expressing Heat Changes



Calorimetry – the accurate and precise measurement of
heat change for chemical and physical processes
Calorimeter – the insulated device used to measure the
absorption or release of heat in chemical or physical
processes
Enthalpy (H) – heat energy content of a system at constant
pressure
Enthalpy is the heat absorbed or released by a system when
pressure is constant.
 It is impossible to record enthalpy directly, but change in enthalpy
(ΔH ) can be measured.
 Units of heat energy: calorie (cal), joules (J), or kJ (kJ).

REACTION ENERGY – EXOTHERMIC RXN

For an exothermic reaction, the sign of ΔH is
negative.
When a reaction is exothermic (ΔH is negative), that is a
favorable condition. Enthalpy is just one of the variables
involved when predicting whether or not a reaction will
occur, but, in general, reactions which release heat are
more likely to occur than ones in which heat is required.
 Heat EXITS the system so the energy (in kJ or J) is
shown as a PRODUCT

 AB
+ CD  AD + BC + ΔH
REACTION ENERGY – ENDOTHERMIC RXN

For an endothermic reaction, the sign of ΔH is
positive.
When a reaction is endothermic (ΔH is positive), that is
an unfavorable condition. Enthalpy is just one of the
variables involved when predicting whether or not a
reaction will occur, but reactions which absorb heat are
less likely to occur than ones in which heat is released,
all things being equal.
 Heat is PUT INTO the system so the energy (in kJ or J) is
shown as a REACTANT

 AB
+ CD + ΔH  AD + BC
REACTION ENERGY, CON’T

Thermochemical Equation – an equation that
includes the heat change

Ex. CaO(s) + H2O(l) → Ca(OH)2(s) + 65.2kJ
Heat Change
 Heat
Change as a reactant means endothermic
 Heat Change as a product means exothermic

Heat of Reaction – the heat of change for the
equation exactly as it is written
 ΔH
= positive means endothermic
 ΔH = negative means exothermic
EXAMPLE PROBLEMS - ENTHALPY
DEFINE THE FOLLOWING EXAMPLES AS EITHER ENDOTHERMIC OR EXOTHERMIC BASED UPON THE CHANGE IN
HEAT.
1.
2NO(g) + O2(g)  2NO2(g) + 113.04 kJ
endothermic
2.
2H2(g) + O2(g)  2H2O(l);
endothermic
3.
exothermic
ΔH = -571.6 kJ
or
Heat is negative
exothermic
4NO(g) + 6H2O(l)  4NH3(g) + 5O2(g); ΔH = +1170 kJ
endothermic
4.
or
Heat is shown as a product
or
exothermic
Heat is positive
SO2 (g) +296 kJ  S(s)+ O2 (g)
endothermic
or
exothermic
Heat is shown as a reactant
ENERGY WS: PROBLEM #6
Potential Energy in kJ
Activation energy = 100 kJ/mol = 450 kJ/mol
400
Potential energy of reactants = 350 kJ/mol
300
Potential energy of products = 250 kJ/mol
200
EXOTHERMIC REACTION
100
ΔH = - 100 kJ
Reaction Pathway (timeline)
POTENTIAL ENERGY DIAGRAM WS
1. Which of the letters a–f in the diagram represents the potential energy of
e
the products? _______
c
2. Which letter indicates the potential energy of the activated complex? ____
3. Which letter indicates the potential energy of the reactants? ________
a
b
4. Which letter indicates the activation energy? _____
5. Which letter indicates the heat of reaction? ______
f
6. Is the reaction exothermic or endothermic? ______
endo
7. Which letter indicates the activation
d
energy of the reverse reaction? ________
8. Which letter indicates the heat of
f
reaction of the reverse reaction? ________
9. Is the reverse reaction exothermic or
exothermic
endothermic? __________
READING A CHART WITH NUMBERS!
80 kJ.
1. The heat content of the reactants of the forward reaction is about ______
160
2. The heat content of the products of the forward reaction is about _______kJ.
240 kJ.
3. The heat content of the activated complex of the forward reaction is about ______
160 kJ.
4. The activation energy of the forward reaction is about ______
+80 kJ.
5. The heat of reaction (ΔH) of the forward reaction is about ______
6. The forward reaction is _______________
endothermic
(endothermic or exothermic).
7. The heat content of the reactants of the reverse reaction
160 kJ.
is about ________
8. The heat content of the products of the reverse reaction
80 kJ.
is about _______
9. The heat content of the activated complex of the reverse
240
reaction is about _______kJ.
10. The activation energy of the reverse reaction is about
_______
80 kJ.
11. The heat of reaction (ΔH) of the reverse reaction is
- 80 kJ.
about _______
exothermic
12. The reverse reaction is __________________
(endothermic or exothermic).
Day 3
THERMODYNAMICS: REACTION ENERGY
REACTION ENERGY CONTINUED, HEAT OF
FORMATION AND REACTION

Calculating Heat Changes


Standard Heat of Formation (ΔHf0) – the change in enthalpy
that accompanies the formation of one mole of a compound
from its elements
Heat of Reaction (ΔH0)– the heat released or absorbed
during a chemical reaction, or Enthalpy
ΔH0 = ΔHf0 (products) - ΔHf0 (reactants)
 READ AS: Heat of Rxn EQUALS the SUM of Heat of Formation of the
Products minus the SUM of Heat of Formation of the Reactants


Standard Heats of Formation have been determined for
many common pure substances, both elements and
compounds. Elements in their natural state are understood
to have a ΔHf0 = 0
HEAT OF FORMATION AND REACTION

In order to calculate ΔH0, the standard heats of formation of
the reactants and products must be known; they can be found
on the following table:
CALCULATING HEAT OF RXN

Steps to Calculate Heat of Reaction:
1.
2.
Find the balanced chemical equation for the reaction;
must have states of matter
Find the sum of the Heats of Formation for the reactants
a.
b.
3.
Find the sum of the Heats of Formation for the products
a.
b.
4.
Multiply the ΔHf0 for each reactant by its corresponding number of
moles (coefficient) from the balance equation
Sum the reactants
Multiply the ΔHf0 for each product by its corresponding number of
moles (coefficient) from the balance equation
Sum the products
Subtract the ΔHf0 (reactants) from the ΔHf0 (products)
EXAMPLE: 2CO(G) + O2(G)  2CO2(G)
1.
2.
Find the balanced chemical equation for the reaction; must
have states of matter
Find the sum of the Heats of Formation for the reactants
a.
b.
3.
Multiply the ΔHf0 for each reactant by its corresponding number of moles
(coefficient) from the balance equation
Sum the reactants
Find the sum of the Heats of Formation for the products
a.
b.
Multiply the ΔHf0 for each product by its corresponding number of moles
(coefficient) from the balance equation
Sum the products
REMEMBER! Elements in their natural state are understood to have a ΔH
OTHERWISE, use the table provided to lookup individual ΔH
f
0
f
0=
0
EXAMPLE: 2CO(G) + O2(G)  2CO2(G)
4.
Subtract the ΔHf0 (reactants) from the ΔHf0
(products)
ΔH0 = ΔHf0 (products) - ΔHf0 (reactants)
(show work here)

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