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Energy
Chapter 16
Energy: Ability to do Work
Potential Energy (PE) = Energy of position

aka STORED energy
Kinetic Energy (KE) = Energy of motion
Radiant Energy = Electromagnetic radiation

ex: sunlight
Some Types of Energy
Energy
Mechanical
Kinetic
Potential
Non-mechanical
Chemical Electrical
(not a complete list!)
Magnetic
Radiant
Units of Energy
 SI
1
system - unit of energy is JOULE (J)
Joule ≅ amount of energy required to
lift 1 golf ball about 1 meter

other energy units:
 calorie (cal)
 Calorie (Cal)
 british thermal units (BTU)
1 calorie = 4.18 Joules
1 Calorie = 1000 cal = 1 kilocalorie = 4180J
Kinetic Energy
 KE
 so
= ½ x mass x velocity2 = ½ mv2
KE of matter depends on:
how heavy and how fast
Potential Energy
 stapler
 rubberband
 popper
 anything
can have PE
= energy of position
= stored energy
 PE
can be
converted to KE
Magnets
PE
in the system of 2
magnets depends on their
relative position


when magnets get close together they
will pull together due to attraction
when magnets are far apart they can’t
attract each other
Electromagnetic Radiation

sunlight – visible
radiation

ultraviolet radiation

infrared radiation

gamma rays

x-rays

microwaves

radiowaves
Energy in Chemistry
chemical energy:
energy stored within
chemical bonds between atoms
heat:
form of energy
 flows from warmer object to cooler
object

Heat Energy
heat: energy associated with motion
of atoms/molecules in matter
symbol
for heat energy = Q or q
Heat Energy
heat
depends on amount of
substance present
can
only measure changes in
heat energy

not absolute value of heat energy
Temperature
measure of average KE of particles in sub
 temperature is NOT energy
 TEMP does NOT depend on amount
 ENERGY does !
Law of Conservation of Energy
 energy
is neither created nor destroyed
in ordinary chemical or physical change
energy before = energy after
 reminder:
nuclear rxn a small amount of
mass is converted to energy
Energy can be converted from
one form to another
 potential
to kinetic
golf ball hit off tee
 radiant
to electric
solar heat to electricity
 electric
to heat
electric stove cooking food
 chemical
to kinetic
burning charcoal on grill
 chemical
to electrical
batteries creating electricity
ALL physical & chemical
changes are accompanied by
change in energy
Thermochemistry:
chemistry of energy changes
Energy Transfer
 measure
changes in heat (amount energy
transferred from one substance to another)
 measure
energy lost somewhere
or
energy gained somewhere else
energy of universe is conserved
universe
(room)
environment Environment
(container)
system
(substances)
energy
energy can
move between
system and
environment
(goes in or out)
EXothermic Change

system releases heat to environment
 what happens to temperature of environment?
temperature of system 
 EXo

- heat is EXiting
what happens to temperature of system?
environment
temperature of environment 
system
what happens to energy level of system?
↓
Exothermic Change
energy lost = energy gained
(system)
(environment)
Endothermic Change
 system

absorbs heat from environment
what happens to temperature of environment?
temperature of environment 
 ENdo

- heat ENters system
what happens to temperature of system?
temperature of system 
environment
system
what happens to energy level of system?
↑
Endothermic Change
energy lost = energy gained
(environment)
(system)
Heat Flow
heat
ALWAYS flows from
hotter object to cooler object
cold

pack on leg:
heat flows from leg to cold pack!
• leg cools down; cold pack warms up
Calorimetry
 changes
in heat energy are measured using
calorimeter (energy lost = energy gained)
 difficult
to monitor “system”
 easy to monitor “environment” (water)
 energy
lost/gained by environment =
energy gained/lost by system
calorimeter:
used to measure
heat changes
“universe” = styrofoam cup
“enviroment” = water****
“system” is whatever put in water (reactants)
quantity
(amount) of heat
transferred depends on:

amount temperature change

mass of substance

specific heat of substance
Calculating Heat Transferred
simple system:
•pure substance in single phase
•calculate heat gained or lost using:
Q = mcT
Q = amount of heat transferred
m = mass of substance
c = specific heat capacity of the substance
T = temperature change = Tfinal – Tinitial
Specific Heat
 amount
heat energy required to
raise temp of 1 gram of substance by 1oC
 symbol
=c
 specific
heat = a physical constant
 unique
 see
for each pure substance
Table B for water (4.18J/g˚C)
typical word problem
10 grams of NaOH(s) is dissolved in 100 g of
water & the temperature of the water
increases from 22C to 30C.
 dissolving


process:
exothermic
was it endothermic or exothermic?
how do you know?
temperature ↑
Solid Dissolving in Water
 What’s
happening when NaOH dissolves?
Add H2O
NaOH molecules close
together, not interacting
NaOH molecules pulled apart
& interact with H2O molecules
Calorimetry
calculate energy released by NaOH as it
dissolves in water
energy lost by NaOH = energy gained by water

• easier to calculate from H2O perspective
Q = mcT
Q = energy (Joules)
m = mass (grams)
c = specific heat capacity (Table B)
T = temperature change = Tf - Ti
Calorimetry & Q = mCT
 temperature
to 30C
 what
of water increased from 22C
30C -22C = 8C = T
mass to use?
temp change was for water, so use mass H2O
m= 100 g
 same
goes for specific heat capacity;
calculate heat absorbed by water
C H 0 = 4.18J/gC
2
Q = mcT
Q
= (100 g)(4.18 _J )(8C)
g C
Q
= 3344 J
Stability and Energy
 if
energy is high, stability is low
 if
energy is low, stability is high