Unit 12 Thermodynamics Ppt

THERMODYNAMICS: REACTION ENERGY

Unit 12

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

SYSTEMSurroundings

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 FLOWMATCH 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 youcan visualize the potentialenergy all the way acrossthe graph

Reactants

Products

Activation Energy

ASSESSMENT – POTENTIAL ENERGY DIAGRAM

1. Is the above reaction endothermic or exothermic?2. What letter represents the potential energy of the

reactants?3. What letter represents the potential energy of the

products?4. What letter represents

the change in energy forthe reaction?

b

f

d

ASSESSMENT – POTENTIAL ENERGY DIAGRAM

5. What letter represents the activation energy of the forward reaction?

6. What letter represents the activation energy of the reverse reaction (read the chart backwards)?

7. What letter represents thepotential energy of theactivated complex?

8. Is the reverse reactionendo or exothermic?

9. If a catalyst were added,what letter(s) would change?

a

e

c

a and c

THERMODYNAMICS: REACTION ENERGY

Day 2

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 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

Heat Change

EXAMPLE PROBLEMS - ENTHALPYDEFINE 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 or exothermic

2. 2H2(g) + O2(g) 2H2O(l); ΔH = -571.6 kJ

endothermic or exothermic

3. 4NO(g) + 6H2O(l) 4NH3(g) + 5O2(g);    ΔH = +1170 kJ

 endothermic or exothermic

4. SO2 (g) +296 kJ S(s)+ O2 (g)

endothermic or exothermic

Heat is shown as a product

Heat is negative

Heat is positive

Heat is shown as a reactant

ENERGY WS: PROBLEM #6

Potential energy of reactants = 350 kJ/mol

Potential energy of products = 250 kJ/mol

Activation energy = 100 kJ/mol = 450 kJ/mol

100

200

300

400

EXOTHERMIC REACTION

ΔH = - 100 kJ

Reaction Pathway (timeline)

Pote

nti

al Energ

y in k

J

POTENTIAL ENERGY DIAGRAM WS

1. Which of the letters a–f in the diagram represents the potential energy of the products? _______2. Which letter indicates the potential energy of the activated complex? ____3. Which letter indicates the potential energy of the reactants? ________4. Which letter indicates the activation energy? _____5. Which letter indicates the heat of reaction? ______6. Is the reaction exothermic or endothermic? ______7. Which letter indicates the activation energy of the reverse reaction? ________8. Which letter indicates the heat of reaction of the reverse reaction? ________9. Is the reverse reaction exothermic or endothermic? __________

e

ca

b

fendo

f

d

exothermic

READING A CHART WITH NUMBERS!1. The heat content of the reactants of the forward reaction is about ______ kJ.2. The heat content of the products of the forward reaction is about _______kJ.3. The heat content of the activated complex of the forward reaction is about ______ kJ.4. The activation energy of the forward reaction is about ______ kJ.5. The heat of reaction (ΔH) of the forward reaction is about ______ kJ.6. The forward reaction is _______________ (endothermic or exothermic).7. The heat content of the reactants of the reverse reaction is about ________ kJ.8. The heat content of the products of the reverse reaction is about _______ kJ.9. The heat content of the activated complex of the reverse reaction is about _______kJ.10. The activation energy of the reverse reaction is about _______ kJ.11. The heat of reaction (ΔH) of the reverse reaction is about _______ kJ.12. The reverse reaction is __________________ (endothermic or exothermic).

80160

240160

+80endothermic

exothermic

160

80

80

240

- 80

THERMODYNAMICS: REACTION ENERGY

Day 3

REACTION ENERGY CONTINUED, HEAT OF FORMATION AND REACTION

Calculating Heat Changes Standard Heat of Formation (ΔHf

0) – 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) - ΔHf

0 (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 ΔHf

0 = 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. Find the balanced chemical equation for the reaction;

must have states of matter2. Find the sum of the Heats of Formation for the

reactantsa. Multiply the ΔHf

0 for each reactant by its corresponding number of moles (coefficient) from the balance equation

b. Sum the reactants

3. Find the sum of the Heats of Formation for the productsa. Multiply the ΔHf

0 for each product by its corresponding number of moles (coefficient) from the balance equation

b. Sum the products

4. Subtract the ΔHf0 (reactants) from the ΔHf

0 (products)

EXAMPLE: 2CO(G) + O2(G) 2CO2(G)

1. Find the balanced chemical equation for the reaction; must have states of matter

2. Find the sum of the Heats of Formation for the reactantsa. Multiply the ΔHf

0 for each reactant by its corresponding number of moles (coefficient) from the balance equation

b. Sum the reactants

3. Find the sum of the Heats of Formation for the productsa. Multiply the ΔHf

0 for each product by its corresponding number of moles (coefficient) from the balance equation

b. Sum the products

REMEMBER! Elements in their natural state are understood to have a ΔHf

0 = 0

OTHERWISE, use the table provided to lookup individual ΔHf0

EXAMPLE: 2CO(G) + O2(G) 2CO2(G)

4. Subtract the ΔHf0 (reactants) from the

ΔHf0 (products)

ΔH0 = ΔHf0 (products) - ΔHf

0 (reactants)

(show work here)