Transcript Energy

Energy
Chapter 16
Energy: Ability to do Work

Potential Energy = Energy of position

AKA STORED ENERGY

Kinetic Energy = Energy of motion

Radiant Energy = Electromagnetic

Ex: Sunlight
Types of Energy
Energy
Mechanical
Kinetic
Potential
(Not a complete list!)
Non-mechanical
Chemical Electrical
Magnetic
Radiant
Units of Energy
 SI
system - unit of energy is the
JOULE (J)
1
Joule = amount of energy required to
lift a golf ball 1 meter

Other Energy Units:
calorie, Calorie, BTU’s

1 calorie = 4.18 Joules

1 Calorie = 1000 calories = 1 kilocalorie
Kinetic Energy
 KE
 So
= ½ x Mass x Velocity2 = ½ mV2
KE depends on how heavy and how fast
Potential Energy
 stapler
 Rubberband
 Popper
 Anything
can have PE =
energy of position
= stored energy
 Potential
Energy can
be converted to
Kinetic Energy
Magnets
The
potential energy in the
system of 2 magnets depends
on their relative position
Electromagnetic Radiation
 Sunlight
– Visible radiation
 Ultraviolet radiation
 Infrared radiation
 Gamma rays
 X-rays
 Microwaves
 Radiowaves
Applet spectrum
Energy in Chemistry
Chemical
bonds
energy = energy stored in
Heat
– form of energy that flows
from warmer object to cooler
object
(Macroscopic)
Heat Energy
Heat:
energy associated with the
motion of atoms & molecules in
matter
(Microscopic)
Symbol
for heat energy = Q or q
Heat Energy
Heat
depends on amount of
substance present
We
measure heat changes
Temperature

= measure of average kinetic energy of
particles of substance
 Swimming
Pool vs. Mug
 Temperature

is NOT energy
Temperature does NOT depend on amount
of substance; energy does
Law of Conservation of Energy
 Energy
is neither created nor destroyed
in ordinary chemical or physical change
Energy before = Energy after
Energy can be converted from one form to
another
- potential to kinetic - radiant to electric
- electric to heat
- chemical to kinetic
- chemical to electrical
All physical & chemical
changes are accompanied by
change in energy
The chemistry of energy changes
is known as Thermochemistry!
Energy Transfer
 Measure
changes in heat
 amount
of energy transferred from one
substance to another
 You
can measure energy lost somewhere or
the energy gained somewhere else
 Cannot measure absolute heat content of
system
Energy of Universe is conserved
Universe
EnvironmentEnvironment
System
Energy
Energy can
move between
the system
and the
environment
Exothermic Change
 System

releases heat to environment
What happens to the temperature of the
environment?
 EXO
- energy leaves system (exits)
 What
happens to the energy level of the
system?

What happens to temperature of system?
EXO - energy leaves system
(exits)
Temperature of
environment 
Environment
Temperature of
system 
System
Energy
Exothermic Change
System
has net energy loss!
Environment has net energy
gain!
Energy
lost = Energy gained
Endothermic Change
 System

absorbs heat from environment
What happens to temperature of
environment?
 Endo
- Energy enters system
 What
happens to the energy level of the
system?

What happens to temperature of system?
Endo - Energy enters system
(entrance)
Temperature of
environment 
Environment
System
Energy
Temperature of
system 
Endothermic Change
System
has net energy gain!
Environment has net energy
loss!
Energy
lost = Energy gained
Heat Flow
Heat
flows from hotter object
to cooler object
Cold
pack on leg: Heat flows
from the leg to the cold pack!

Leg cools down; cold pack warms up
Quantity of heat transferred
Quantity
(amount) of heat
transferred depends on
Temperature change
 Mass of substance
 Specific Heat of substance

Calculating Heat Transferred
Simple system:
•pure substance in a 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
 Different
for each pure substance
Calorimeter
Another
example
source
Calorimetry
 Changes
in heat energy are measured by
calorimetry
 “universe” is contained in styrofoam cup
 “enviroment”
 “system”
water
is water****
is whatever we put in the
Calorimetry
 Energy
lost = Energy gained
 Difficult
to monitor “system”
 Easy to monitor “environment” (water)
 Energy
lost/gained by environment =
Energy gained/lost by system
Calorimetry
10 grams of NaOH is dissolved in 100 g of
water & the temperature of the water
increases from 22C to 30C
 was dissolving process endothermic or
exothermic
 how do you know?
Exothermic – temperature of environment ↑
Dissolving
 What’s
happening when NaOH dissolves?
Add H2O
molecules close together,
not interacting
molecules pulled apart &
interacting with H2O
Calorimetry
Calculate energy released by NaOH as it
dissolved 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
22C to 30C
 30C
 What
of water increased from
-22C = 8C = T
mass to use? Well, temp change
was for water, so want mass of water
m = 100 g
 Same goes for specific heat capacity;
calculate heat absorbed by water
cH 0 = 4.18J/g
2
Q = mCT
Q
= (100 g)(4.18 J/g)(8C)
Q
= 3344 Joules
Stability and Energy
 If
energy is high, stability is low
 If
energy is low, stability is high