How are they related?
CHEMISTRY AND ENERGY
What is Energy?
 Defined as the ability to do work or create
heat.
 Many types of energy
 Thermal
 Light
Light Energy
 The Dual Nature of Light (or electrons IN
light):
1.Sometimes it behaves like a particle
(called a photon), which explains how light
travels in straight lines
2. Sometimes it behaves like a wave, which
explains how light bends (or diffracts) around
an object
Electromagnetic Spectrum
 Light is a form of energy called electromagnetic energy
a. Other types of electromagnetic energy include radio,
microwave , heat (infrared), ultraviolet (UV) and many other
types
 Electromagnetic energy travels in waves
 The waves are different from each other in their lengths –
called wavelength
Electromagnetic Spectrum
Continued
 The shorter the wavelength and the greater the
frequency, the higher the energy
a. Radio waves are as long as soccer fields, low
frequency and low energy
b. Gamma rays are smaller than an atom, high
frequency and high energy
What Does This Energy Look
Like?
How Are They All Related
 Frequency, wavelength and the speed of light
are related by the following equation:
c = λν
 where c is the speed of light 3.00 x 108 m/s
 ν is the frequency measured in Hz (hertz) (s-1)
 λ is the wavelength measured in nm (nanometers)
Lets Look at an Example
 What is the frequency of a photon with λ =
550 nm
Light Energy and Bohr’s
Model
 the Bohr model depicts an atom as a small,
positively charged nucleus surrounded by
electrons. These electrons travel in circular
orbits around the nucleus—similar in
structure to the solar system, except
electrostatic forces rather than gravity
provide attraction.
Properties of Bohr’s Model
 Electrons in atoms orbit the nucleus.
 Electrons can only orbit stably in certain
orbits ("stationary orbits“- remember the
ladder rungs ).
 These orbits are associated with definite
energies and are also called energy shells or
energy levels. REMEMBER THAT ALL
ELECTRONS IN THE SAME ORBITAL ON THE
SAME ENERGY LEVEL ARE DEGENERATE.
So How Can Electrons Move?
 Electrons can only gain or
lose energy by jumping
from one orbit to another
 To go from a lower energy
to a higher energy,
electromagnetic radiation
is absorbed
 To go from a higher energy
to a lower energy
electromagnetic radiation
is emitted in the form of
photons (little packets of
light)
Electron Energy Calculations
 ΔE = hν
 1) E is the energy of the
particular quantum of energy
under study measured in
Joules (J)
 2) h stands for a fundamental
constant of nature known as
Planck's Constant.
 The value for Planck's
Constant is 6.626 x 10¯34
Joule second (Js).
 3) ν is the frequency of the
particular photon being
studied (still measured in s-1)
Lets Look at a Problem
 Remember the problem from a few slides
ago? Now lets calculate the energy in a
wavelength of that frequency
 What is the frequency of a photon with λ =
550 nm?
Thermochemistry
 The study of heat used or released in a
chemical reaction.
How Do Chemical Reactions
Create Heat energy?
 Consider the combustion of gasoline (octane)
2 C8H18 +25 O2  16 CO2 +18 H2O
 Potential Energy: Stored energy
 To break bonds of reactants, energy is required
 When bonds are broken, the energy is available
 When product bonds form, some energy is used
 Any excess energy is released as heat
Kinetic Energy
 Directly related to temperature
 Molecules move faster at higher temperatures
 Remember from states of matter - gases vibrate
faster than liquids and solids
Is Heat Used or Released?
 Endothermic reactions use heat from the
surroundings
 Sweating
 Refrigeration
 Exothermic reactions releases heat to the
surroundings
 Hot hands
 Combustion
 Exercise
Endothermic Reactions
 Decrease in kinetic energy  decrease in





temperature  heat will transfer from the
environment to the system resulting in a
cooler environment
Absorbs heat from its surrounding.
The system gains heat
Positive value for q
H = q = 0
Hproducts  Hreactants
Exothermic Reactions
 Increase in kinetic energy  increase in





temperature of system heat released to
the environment resulting in a hotter
environment
Releases heat to its surroundings
The system loses heat
Negative value for q
H = q = 0
Hproducts  Hreactants
Enthalpy
 Measures the change in heat/energy at
constant pressure
 Symbol is H
 Terms heat and enthalpy are used
interchangeably for this course
 H = q = m C T
 Heat moves from ____ to ____.
Enthalpy
 Measures the change in heat/energy at
constant pressure
 Symbol is H
 Terms heat and enthalpy are used
interchangeably for this course
 H = q = m C T
 Heat moves from hot to cold.
Law of Conservation of
Energy
 Energy is not lost or gained in a chemical
reaction
 In a chemical reaction potential energy is
transferred to kinetic energy
Thermochemical Equations
 An equation that includes the heat change
 Exothermic = product
 Endothermic = reactant
 Example: write the thermochemical equation
for this reaction
 CaO(s) + H2O(l) Ca(OH)2(s)
H = -65.2 kJ /mol
CaO(s) + H2O(l) Ca(OH)2(s) + 65.2 kJ/mol
Specific Heat Calculations
 q = mCΔT
q = heat (J)
m = mass (g)
C = specific heat (J/g oC)
ΔT = change in temperature (o C or K)
= T f - Ti
Specific Heat
 Specific heat of water = 1 J / goC
 Specific heat of most metals = < 1 J / goC
 Do metals heat slowly or quickly compared to
water?
 Do metals stay warm longer or shorter than
water?
Practice Problem
 How much energy is required to heat 120.0 g
of water from 2.0 oC to 24.0oC?
q = mCΔT
Practice Problem
 How much energy is required to heat 120.0 g
of water from 2.0 oC to 24.0oC?
q = mCΔT
m= 120.0 g
C = 4.184 J/goC
ΔT= (24.0 – 2.0)oC = 22.0oC
q = (120.0g)(4.184 J/goC)(22.0oC) = 524.08J
Practice Problem
 How much heat (in kJ) is given off when 85.0
g of lead cools from 200.0oC to 10.0 oC?
(Specific heat of lead = 0.129 J/g oC)
q = mCΔT
Practice Problem
 How much heat (in kJ) is given off when 85.0
g of lead cools from 200.0oC to 10.0 oC?
(Specific heat of lead = 0.129 J/g oC)
q = mCΔT
m = 85.0 g
C = 0.129 J/g oC
ΔT = (10.0 – 200.0)oC = - 190.0oC
q = (85.0 g)(0.129 J/g oC)(- 190.0oC) = -2083.35J
Stoichiometry and Thermochemistry
Tin metal can be extracted from its oxide
according to the following reaction:
SnO2(s) + 4NO2(g) + 2H2O(l) + 192 kJ 
Sn(s) + 4HNO3(aq)
How much energy will be required to
extract 59.5 grams of tin?
How to solve
1. Use your stoichiometry
2. Treat heat as a reactant or product
SnO2(s) + 4NO2(g) + 2H2O(l) + 192 kJ 
Sn(s) + 4HNO3(aq)
59.5 g Sn 1 mol Sn 192 kJ
1
g Sn 1 mol Sn
How does ice melt?
HEAT DURING A CHANGE OF
STATE
What Is a Phase Change?
 Is a change from one state of matter
(solid, liquid, gas) to another.
 Phase changes are physical changes
because:
- It only affects physical appearance, not
chemical make-up.
- Reversible
What Happens During a Phase
Change
 During a phase change, heat energy is
either absorbed or released.
 Heat energy is released as molecules
slow down and move closer together.
 Heat energy is absorbed as molecules
speed up and expand.
What Happens During a Phase
Change
Melting and Boiling Points
 Melting Point: The temperature at which a
solid changes into a liquid.
 Boiling Point: The temperature at which a
liquid changes into a gas.
 What is a Freezing point? Compare the
freezing and melting points of water.
Molar Heat of Fusion (or
Melting)
 Heat absorbed by one mole of a substance
during melting
 Constant temperature
 Hfus
 H2O(s)  H2O(l)
H = 6.01 kJ/mol
Molar Heat of Solidification
(or Freezing)
 Heat lost when 1 mole of a liquid solidifies
 Temperature is constant
 Hsolid
 Hfus = -Hsolid
 H2O(l)  H2O(s) H = -6.01 kJ/mol
Molar Heat of Vaporization
(or Boiling)
 Heat needed to vaporize 1 mole of a liquid
 Hvap
 H2O(l)  H2O(g)
Hvap = 40.7 kJ/mol
Molar Heat of Condensation
 Heat released when 1 mole of vapor
condenses
 Hcond
 H2O(g)  H2O(l)
 Hvap = -Hcond
Hcond = -40.7 kJ/mol
Heating Curve for Water
What Determines Phase
 Whether a substance is a solid, liquid or gas
depends on the temperature and pressure
 Keeping T constant while increasing P usually
produces a solid
 Water is an exception, increasing pressure on
ice produces water
 This causes ice skates to melt ice and
freezing water to expand and produce frost
heaves
Before We Look at a Phase
Diagram…
 Phase Diagram
 Plot of Pressure versus Temperature
 Triple Point
 A point on the phase diagram at which all
three phases exist (solid, liquid and gas)
 Critical Point
 A point on the phase diagram at which the
density of the liquid and vapor phases are
the same
Phase Diagram
Phase Diagram Features
 Beyond the critical point there is no distinction
between a liquid and gas
 Solid and liquid phases separated by a fusion
curve
 Solid and gas phases separated by a
sublimation curve
In Conclusion