Energy and Chemical Change

1
Energy and Chemical Change

• Chapter 16

2
Energy

• Section 16.1

3
Energy

• The ability to do work or produce heat.
• It exists in two basic forms, potential energy
  and kinetic energy
• Potential energy due to the composition or
  position of an object
• Kinetic energy of motion
• Chemical potential energy is the energy stored in
  a substance because of its composition.

4
Energy flow

• Exothermic when heat flows from the object to
  its surroundings heat is released
• Endothermic when heat flows from the
  surroundings to the object heat is absorbed
• Was making ice cream exothermic or endothermic?

5
Law of conservation of energy

• In any chemical reaction or physical process,
  energy can be converted from one form to another,
  but is neither created nor destroyed.

6
Heat (q)

• Energy that is in the process of flowing from a
  warmer object to a cooler object.
• The amount of heat required to raise one gram of
  water by one degree Celsius is defined as a
  calorie (cal).
• 1 Calorie 1 kcal 1000 cal
• The SI unit of heat and energy is the joule (J).

7
Converting energy units

• 1 J 0.2390 cal 1 cal 4.184 J
• A granola bar contains 142 Calories. Convert
  this to joules.

8
Specific heat

• The amount of heat required to raise the
  temperature of one gram of that substance by one
  degree Celsius.
• q m c ?T
• q the heat absorbed or released
• m the mass of the sample in g
• c the specific heat of the substance
• ?T the change in temperature

9
Calculations with specific heat

• The temperature of a sample of iron with a mass
  of 10.0g changed from 50.4C to 25.0C with the
  release of 114J heat. What is the specific heat
  of iron?
• If the temperature of 34.4g of ethanol increases
  from 25.0C to 78.8C, how much heat has been
  absorbed by the ethanol? The specific heat of
  ethanol is 2.44J/(gC)

10
Homework

• After learning about section 16.1, you should be
  able to do Heat And Its Measurement
• On the back
• Pick a favorite food that has a nutrition label
• Write down what the food is and the Calorie
  content for one serving
• Convert this to cal, J, kJ

11
Problem Solving Lab (p. 503)

• Create the graph and answer questions 1-3

12
Heat in Chemical Reactions and Processes

• Section 16.2

13
Thermochemistry

• The study of heat changes that accompanies
  chemical reactions and phase changes.
• System the specific part of the universe that is
  being studied (chemical reaction)
• Surroundings everything else
• Universe system surroundings

14
Measuring heat

• A calorimeter is an insulated device used for
  measuring the amount of heat absorbed or released
  during a chemical or physical process.

15
Bomb Calorimeter
16
Enthalpy (H)

• The heat content of a system at constant pressure
• Enthalpy of reaction(?Hrxn) is the change in
  enthalpy for a reaction
• ?Hrxn Hfinal Hinitial
• ?Hrxn Hproducts - Hreactants
• If ?H is positive, the reaction is endothermic
• If ?H is negative, the reaction is exothermic

17
Think-Pair-Share

• Compare energy changes in chemical reactions to
  profits and losses in a business.
• Each month the business has receipts (positive
  dollar amounts) and expenses (negative dollar
  amounts).
• If the total receipts exceed expenditures, a
  positive dollar amount or profit occurs.
• If expenditures are greater than receipts, money
  is lost and a negative dollar amount results.

18
Practice Problem

• A 75.0g sample of a metal is placed in boiling
  water until its temperature is 100.0C. A
  calorimeter contains 100.00g of water at a
  temperature of 24.4C. The metal sample is
  removed from the boiling water and immediately
  placed in the water in the calorimeter. The
  final temperature of the metal and water in the
  calorimeter is 34.9C. Assuming that the
  calorimeter provides perfect insulation, what is
  the specific heat of the metal?

19
Homework

• P. 500 (14-18)
• P. 525 (79-81)

20
Thermochemical Equations

• Section 16.3

21
Thermochemical Equation

• A balanced equation that includes the physical
  states and energy change.
• Example
• 4Fe(s) 3O2(g) ? 2Fe2O3(s) ?H -1625 kJ

22
Enthalpy (heat) of combustion (?Hcomb)

• The enthalpy change for the complete burning of
  one mole of a substance
• Example
• C6H12O6(s) 6O2(g) ? 6CO2(g) 6H2O(l)
  ?Hcomb-2808kJ

23
Changes of State

• Molar enthalpy (heat) of vaporization (?Hvap)
• H2O(l) ? H2O(g) ?Hvap 40.7kJ
• H2O(g) ? H2O(l) ?Hcond -40.7kJ
• Molar enthalpy (heat) of fusion (?Hfus)
• H2O(s) ? H2O(l) ?Hfus 6.01kJ
• H2O(l) ? H2O(s) ?Hsolid 6.01kJ
• What would be the molar enthalpy of sublimation?

24
Practice Problems

• Calculate the heat required to melt 25.7 g of
  solid methanol at its melting point. ?Hfus3.22
  kJ/mol
• How much heat is evolved when 275 g of ammonia
  gas condenses to a liquid at its boiling point?
  ?Hvap23.3 kJ/mol
• What mass of methane must be burned in order to
  liberate 12,880 kJ of heat? ?Hcomb-891 kJ/mol

25
Homework

• P. 505 (26-27)
• P. 525 (82-84)

26
Calculating Enthalpy Change

• Section 16.4

27
Hesss Law

• Read section on p. 506
• Add the following equations
• A B ? C
• C D ? E B

28
Example

• Ex) 2S(s) 3O2(g) ? 2SO3(g) ?H ?
• S(s) O2(g) ? SO2(g) ?H -297kJ
• 2SO3(g) ? 2SO2(g) O2(g) ?H 198kJ

29
Reaction Spontaneity

• 16.5

30
Spontaneous Process

• A spontaneous process is a physical or chemical
  change that occurs with no outside intervention.
• Ex. Iron forming rust, combustion of methane,
  melting of ice cream
• A non-spontaneous change is a change that occurs
  only when driven
• e.g. forcing electric current through a metal
  block to heat it
• What makes a reaction spontaneous?

31
What determines if a reaction is spontaneous?

• Entropy Measure of the disorder or randomness of
  the particles that make up a system.
• Spontaneous processes always proceed in such a
  way that the entropy of the universe increases.
  (?S)
• ?Suniverse ?Ssystem
  ?Ssurroundings
• Larger entropy value, larger degree of
  randomness.
• Law of Disorder (second law of thermodynamics)
  Spontaneous processes always proceed in such a
  way that the entropy of the universe increases.

32
Law of Disorder (Second Law of Thermodynamics)

• Law of disorder Spontaneous processes always
  proceed in such a way that the entropy of the
  universe increases.
• Entropy gas gt entropy liquid gt entropy solid
• Dissolving of a gas in a solvent always results
  in a decrease in entropy
• An increase in temperature results in an increase
  in entropy.
• Entropy increases when the number of product
  particles is greater than the number of reactant
  particles.
• 2SO3 (g) ?2SO2 (g) O2 (g) ?Ssystem gt 0

33
Entropy, the Universe and Free Energy

• For any spontaneous reaction
• ?Suniverse gt 0
• ?Suniverse is positive when,
• 1. the reaction is exothermic
• 2. The entropy of the system increases, so
  ?Ssystem is positive
• Gibbs Free Energy (G) or Free energy
• -the energy that is available to do work.
• ?Gsystem ?Hsystem
  -T?Ssystem

34
Type of reaction or process ?Gsystem ?Suniverse
Spontaneous Negative Positive
Nonspontaneous Positive negative
How ?Hsystem and ?Ssystem Affect Reaction Spontaneity How ?Hsystem and ?Ssystem Affect Reaction Spontaneity How ?Hsystem and ?Ssystem Affect Reaction Spontaneity
-?Hsystem ?Hsystem
?Ssystem Always spontaneous Spontaneity depends upon temperature
-?Ssystem Spontaneity depends upon temperature Never spontaneous
35
Practice Problem

• For a process, ?Hsystem 145 kJ and ?Ssystem322
  J/K. Is the process spontaneous at 382 K?
• Answer 22,000 J