Download principles of reactivity: energy and chemical reactions

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HONORS UNIT 5 NOTES: THERMODYNAMICS
Part I: Principles of Reactivity: Energy & Chemical Reactions
OBJECTIVES



Describe various forms of energy and energy transfer.
Understand the terms reactant-favored, product-favored, and thermodynamics.
Differentiate between kinetic and potential energy and know the SI unit used to
measure thermal energy.
 Understand the term specific heat capacity and know how to calculate amount of
thermal energy transferred from one object to another.
 Use heat of fusion and heat of vaporization to solve simple thermal problems.
 Recognize and correctly use vocabulary related to thermodynamics: system,
surroundings, endothermic, exothermic, enthalpy, first law of thermodynamics, and
calorimetry.
 Calculate the enthalpy change of a system in several ways: graphically,
experimentally, and using Hess’s Law and the summation equation.
ENERGY: SOME BASIC PRINCIPLES (pg. 197 – 200 in textbook)
1. Definitions
a. Thermodynamics
b. Energy
***Note: Work is force acting over a distance.
c. Kinetic Energy
d. Potential Energy
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2. Law of Conservation of Energy (a.k.a. First Law of Thermodynamics)
Law of Conservation of Energy states that energy can be converted from one
form to another but CANNOT be created or destroyed.
This mean, the total energy of the universe is _____________________.
3. Heat vs. Temperature
a. Temperature – measure of the average kinetic energy in a substance
b. Heat – (
) flow of energy due to a temperature difference
c. Important aspects of thermal energy and temperature:
 ___________ is NOT the same as _________________________.
 The more __________________________a substance has, the greater
the temperature of its atoms and molecules.
 The total thermal energy in an object is the sum of its individual energies
of all of the molecules.
 For any given substance, its thermal energy depends not only on its
composition, but also on the amount of substance.
4. System vs. Surroundings
A system is:
The surroundings include:
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5. Direction of Heat Flow: Thermal Equilibrium
o Heat transfer occurs when two objects are at _________________________

Eventually, the two objects reach the same temperature

We say that the system has reached __________________
o Heat transfer always occurs from a ________ object to a __________ object.

Transfer of heat continues until both objects are at the same temp.
o The quantity of heat lost by a hotter object and the quantity of heat gained by
a cooler object are ___________________.
6. Exothermic vs. Endothermic Processes

EXOTHERMIC PROCESS: heat transferred from system to the surroundings
o Heat is ___________ from the system
o Temperature in the system ___________________

ENDOTHERMIC PROCESS: heat transferred from surroundings to system
o Heat is _____________ to the system
o Temperature in the system _____________________
7. Units of Energy
a. Joule (J) – the SI unit of energy & heat
One kilojoule (kJ) =
b. calorie (cal) = heat required to raise the temp of ____________ by ________
1 calorie = _____________Joules (J)
c. dietary Calorie (Cal) – a.k.a. food calorie or kilocalorie
1 Cal =________ kcal = ____________calories (cal)
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Example #1: 3800 cal = _____________ Calories = ________________ Joules
Example #2: The label on a cereal box indicates that 1 serving provides 250 Cal. What
is this energy in kJ?
SPECIFIC HEAT CAPACITY AND HEAT TRANSFER
1. Direction and sign of heat flow, (q) — MUST MEMORIZE!!
 Case 1: q is __________________ when heat is added to the system and
the temperature increases (reaction is __________________________)
 Case 2: q is ___________________ when heat is lost from the system
and the temperature decreases (reaction is _______________________)
2. Heat capacity – extensive property
The quantity of heat required to raise an object’s temperature by 1°C (or by 1 K)
3. Specific heat (also called specific heat capacity) – intensive property
The quantity of heat required to raise the temperature of one gram of a
substance by 1°C.
J
Units: ⁄g°C
4. Examples of specific heat: sand vs. water at the beach
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5. Equation and units:
q=
m=
C=
ΔT =
 Example #1: How much heat in J is given off by a 75.0 gram sample of pure
aluminum when it cools from 84.0oC to 46.7oC? The specific heat of aluminum is
0.899 J/goC.
 Example #2: What is the specific heat of benzene if 3450J of heat is added to a
150.0g sample of benzene and its temperature increases from 22.5 oC to 35.8oC?
 Example #3: A 50.0 g sample of water gives off 1.025 kJ as it is cooled. If the initial
temperature of the water was 85.0oC, what was the final temperature of the water?
The specific heat of water is 4.18 J/goC.
-----------------------------END OF MATERIAL ON MIDTERM----------------------Dr. Mihelcic
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CALORIMETRY (pg. 200 – 203 in textbook)
1. Calorimetry – measurement of quantities of heat
 A ________________________ is the device in which heat is
measured.
 Assumptions:
i. Heat lost = - heat gained by the system
ii. In a simple calorimeter, assume no heat is lost to the surroundings.
2. Heat flow in a coffee cup calorimeter - diagram and equation
Steps:
1. Add hot solid metal to cool water.
2. Water will heat up as metal cools.
3. Eventually, water & metal are at the same
temperature.
qmetal +qwater = 0
qmetal = - qwater
***Heat lost by the metal = -heat gained by the water
Example A 358.11 g piece of lead was heated in water to 94.1oC. It was removed from
the water and placed into 100.0 mL of water in a Styrofoam cup. The initial temperature
of the water was 18.7oC and the final temperature of the lead and water was 26.1 oC.
What is the specific heat of lead according to this data?
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3. Bomb Calorimeter
o Constant volume calorimeter
o Burns sample in oxygen (O2)
o Some heat from reactions warms the water,
some warms the bomb.
o qrxn + qwater + qbomb = 0
ENERGY AND CHANGES OF STATE
1. Changes of state: All changes of state involve energy changes!!! (More in Unit 9)
2. State Functions – a property where the change from initial to final state ______
_______________________________________________________________
 Example: Change in elevation from the top to bottom of a ski slope is
independent of the path taken to go down the slope.
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ENTHALPY CHANGES FOR CHEMICAL REACTIONS
(pg. 203 – 206 in textbook)
1. Heat of reaction – the heat absorbed or given off when a chemical reaction
occurs at constant temperature (T) and pressure (P)
2. Enthalpy (
) – the heat content of a reaction (chemical energy)
 ΔH =
o The amount of energy absorbed by or lost from a system as heat
during a chemical process as constant P
o ΔH =
3. Properties of enthalpy
 Enthalpy is an ______________________ property (depends on quantity)
 Enthalpy is a ____________________________________
i. Depends only on the final & initial values
 Every reaction has a unique enthalpy value since ΔH = Hproducts - Hreactants
4. Representation of enthalpy using a graph
EXOTHERMIC
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ENDOTHERMIC
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5. Two Ways to Designate Thermochemical Equations
Endothermic:
a) H2 (g) + I2 (s)  2 HI (g)
ΔH = 53.0 kJ
b) H2 (g) + I2 (s) + 53.0 kJ  2 HI (g)
Exothermic:
ΔH = -445.2 kJ
a) ½ CH4 (g) + O2 (g)  ½ CO2 (g) + H2O (l)
b) ½ CH4(g) + O2(g)  ½ CO2(g) + H2O(l) + 445.2 kJ
****Note the meaning of the sign of H in the equations above!
Endothermic:
Exothermic:
 Note the importance of designating the physical state or phase of matter. Why?
 What do the coefficients stand for?
 How can they differ from the ones we have used before?
 What is standard state? How do we designate conditions of temperature and
pressure that are not at standard state?
o Standard state = 1 atm of pressure and 25°C
o ΔH° = ΔH at standard state
o Conditions would be shown over the arrow if not at standard conditions
 How can we find the enthalpy of reaction when we reverse it?
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Example:
Given Reaction #1 below, find the ΔH for Reactions #2 and #3.
Reaction #1: 2 SO2 (g) +
Reaction #2:
O2 (g  2 SO3 (g)
ΔH = ???
SO2 (g) + ½ O2 (g)  SO3 (g)
Reaction #3: 4 SO3 (g)  4 SO2 (g)
+
ΔH = +197.8 kJ
2 O2 (g)
ΔH = ???
6. H as a Stoichiometric Quantity
Example #1: Given the reaction below, how much heat is produced when 15.0 g
of NO2 are produced?
2 NO (g)
Example #2:
+ O2 (g)
 2NO2 (g)
ΔH = -114.1 kJ
Given: CO (g) + 1/2 O2(g)  CO2(g)
H = -283 kJ
a. Calculate the enthalpy of the above reaction (formation of CO 2) when 3.00 grams
of product are formed.
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Given: CO (g) + 1/2 O2(g)  CO2(g)
H = -283 kJ
b. If only 10.0 grams of oxygen and an unlimited supply of CO are available to run
this reaction, how much heat will be given off?
c. How many grams of carbon monoxide are necessary (assuming oxygen is
unlimited) to produce 500 kJ of energy in this reaction?
d. Calculate the heat of decomposition of two moles of carbon dioxide.
HESS’S LAW (pg. 207 – 208 in textbook)
1. Statement and usage:
a. The heat of a reaction (H) is constant, whether the reaction is carried out
directly in one step or indirectly through a number of steps.
b. The heat of a reaction (H) can be determined as the sum of heats of
reaction of several steps.
Σ ΔH along one path = Σ ΔH along another
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Σ ΔH along one path = Σ ΔH along another
Example #1:
Given:
C(s) + O2(g)  CO2(g)
H = -393.5 kJ
2CO(g) + O2(g)  2CO2(g)
H = -577.0 kJ
Determine the heat of reaction for: C(s) + ½ O2(g)  CO(g)
H = ???
Example #2:
Given:
C(s) + O2(g)  CO2(g)
H = -393.5 kJ
C2H4 (g) + 3 O2(g) 
2 CO2(g) + 2H2O (l)
H2(g) + ½ O2(g)  H2O (l)
Determine the heat of reaction for: 2C(s) + 2H2(g)  C2H4 (g)
Dr. Mihelcic
H = -1410.9 kJ
H = -285.8 kJ
H = ???
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STANDARD ENTHALPIES OF FORMATION (pg. 208 – 212 in textbook)
1. Standard molar heat of formation (ΔHf°) – the heat content or enthalpy change
when one mole of a compound is formed at 1.0 atm pressure and 25°C from its
elements under the same conditions.
a. Examples:
H2(g) + ½ O2(g)  H2O(g)
ΔHf°(H2O, g) = -241.8 kJ/mol
C(s) + ½ O2(g)  CO(g)
ΔHf°(CO, g) = -111 kJ/mol
***Formation equations must show elements forming 1 mole of a compound
(Units are per mole)
b. Can look up standard formation values in your reference book!
i. By definition, ΔHf° = 0 for all elements in their standard states!
1. Example:
Cl2 (g)
H2 (g)
Ca (s)
2. Summation equation for calculation of H from enthalpies of formation:
ΔHrxno =
Σ Hof (products) - Σ Hof (reactants)
Example:
Use the summation equation & values from the reference book to determine the
enthalpy of the following reaction:
4NH3(g) + 5O2(g)  4NO(g) + 6H2O(g)
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Part II. Thermodynamics:
Spontaneity, Entropy, and Free Energy
Objectives:
1. Students will be able to define spontaneous changes, entropy, and free energy.
2. Students will be able to predict the sign of ΔS by observing a chemical equation.
3. Students will be able to calculate ΔSo and ΔGo for a reaction using the tables of
standard values of So and ΔGof.
4. Students will be able to recognize a spontaneous or nonspontaneous reaction by the
sign of free energy, using the Gibbs equation.
5. Students will be able to calculate ΔGo for a reaction by using the Gibbs equation.
Spontaneous Change and Heat (pg. 451 – 452 in textbook)
A. What is a spontaneous (thermodynamically favored) process?
A process that occurs by itself without an outside force helping it.
B. Which of the following are spontaneous processes?
1. Snowman melting in the sun
2. Assembling a jigsaw puzzle
3. Rusting of an iron object in humid air
4. Recharging of a camera battery
C. Spontaneous Reactions and Energy
a. Many spontaneous reactions are exothermic, but not all!
Example: H2O (s)  H2O (l) is spontaneous and ENDOTHERMIC
b. What other factor influences spontaneity?
The “Randomness factor” – Nature tends to move spontaneously
towards a more random state.
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ENTROPY & THE SECOND LAW (pg. 453 – 458 in textbook)
1. Entropy – a measure of the randomness (disorder) of a system
o Entropy is a state function
2. The Second Law of Thermodynamics states:
In a spontaneous process, there is a net increase of entropy (taking into account
system and surroundings).
3. Spontaneous processes result in more random states (more disorder).
 Examples:
Example: Predict which of the following processes have a positive change in entropy
(an increase in the randomness/disorder = +ΔS°)
a. Taking dry ice from a freezer and allowing it to warm from -80oC to room
temperature
b. Dissolving blue food coloring in water
c. Freezing water into ice cubes
ENNTROPY & THE THIRD LAW
1. Entropy is used to quantify _____________________________________________
2. Like enthalpy, entropy is also a _______________function.
3. Standard Molar Entropies
a. The Third Law of Thermodynamics states: A completely ordered pure
crystalline solid has an entropy of zero at 0 Kelvin (K).
b. Standard Molar Entropies So (1 mole, standard conditions):

Tells you the entropy at 25°C and 1 atm

Units =
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4. For a substance, entropy generally increases as:
a. Phase changes occur from solid  liquid  gas
b. The total number of moles of gas increase from reactants to products
c.
Temperature increases (K.E. increases)
5. For a reaction, entropy generally increases as:
a. Reactants ( __________ ) Products (______________)
b. Total # of moles of products > Total # of moles of reactants
c. Total # of moles of ________________ products > Total # of moles of
_________________ reactants
d. T is increasing
+ ΔS = entropy increases
- ΔS = entropy decreases
Example: Predict the sign of ΔS in each reaction. Explain your prediction!!!
a. NH3 (g) + HCl
(g)
 NH4Cl (s)
b. 2 KClO3 (s)  2 KCl
(s)
+ 3 O2 (g)
c. CO (g) + H2O (g)  CO2 (g) + H2 (g)
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F. Calculation of Standard entropy changes in a reaction using standard entropies of
products and reactants (Equation):
ΔSo =
Σ So (products) - Σ So (reactants)
Example: Calculate ΔSo for the reaction below using the values in the reference book
and the equation above. Does the sign of your calculated value match the prediction?
2 H2 (g)
+
O2 (g)

2H2O (l)
GIBBS FREE ENERGY CHANGE (pg. 458 – 462 in textbook)
1. The Gibbs (also known as Gibbs-Helmholtz) Equation shows the relationship
between Energy, Entropy and Spontaneity:
ΔG
= ΔH - T ΔS
2. What is “free” energy?
3. The relationship between ΔG and Spontaneity
 If ΔG is positive, the reaction is _________________________________
 If ΔG = 0, the reaction is _______________________________________
 If ΔG is negative, the reaction is ________________________________
***Spontaneous processes MUST have a negative free energy!
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Case #
ΔH
ΔS
ΔG
1
-
+
-
-
-
-
-
-
+
+
+
+
+
+
-
+
-
+
Result
spontaneous at all T
spontaneous toward low T
HOWEVER
2
nonspontaneous toward high T
nonspontaneous toward low T
HOWEVER
3
4
spontaneous toward high T
nonspontaneous at all T
***Exothermic reactions with increasing entropy are always spontaneous!
Example: Predict which of the four cases in the table above will apply to the following
reactions, and if the reaction will be spontaneous. Use ΔH as given and your estimate
of the sign of ΔS.
a. C6H12O6 (s) + 6 O2 (g)  6 CO2 (g) + 6 H2O (g)
ΔH = -2540 kJ
b. Cl2 (g)  2 Cl (g)
ΔH is positive
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4. Equation at standard conditions:
ΔGo = ΔHo - T ΔSo
Note: the units for ΔHo are generally in ____________
and the units for ΔSo generally are in ___________
***You must convert FIRST before beginning the problem!
(Temperature must in _______________; K = °C + 273)
Example: Calculate ΔGo for the reaction below from the given information.
C (s) +
ΔSo = -80.8 J/mol K
2H2 (g) 
CH4 (g)
ΔHo = -74.8 kJ/mol
T= 298 K
5. Standard Free Energy of Formation
ΔGoreaction =
Σ Gof (products) - Σ Gof (reactants)
***Same summation equation as ΔH° equation – use reference book for values!!
Dr. Mihelcic