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1/12/2011
Energy
If it’s not matter, it must be energy…..
• Energy is the capacity to do work or transfer heat
– Work is the energy needed to move a mass
– Heat is the energy needed to make temperature
rise
• Energy does not have mass and volume.
• Energy can cause physical and/or chemical
changes in matter.
Law of Conservation of Energy
• Total amount of energy in the universe is constant.
• All matter possesses some energy.
• Energy can be transferred from one place to another
and its form can be changed
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Forms of Energy
• Kinetic: Mass and speed of an object
– Electricity
– Heat or Thermal Energy
– Light Energy
• Potential: Energy stored within a physical system as a result of the
position or configuration of the different parts of that system.
– Electrostatic: Energy from the interactions of charged particles
– Nuclear: Potential energy in the nucleus of atoms.
– Chemical: Potential energy in the attachment of atoms
Conversion of Energy
Can you give examples of energy conversions?
• Chemical energy to heat and kinetic energy
• Chemical energy to kinetic energy
• Kinetic energy to potential energy
• Thermal energy to potential energy
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More on Kinetic Energy
Kinetic energy: the energy of motion
Kinetic energy depends on mass and speed of object:
Ekinetic = ½ mass x velocity2
2 objects at same speed; heavier one has more Ekinetic
2 objects w/same mass; faster one has more Ekinetic
Atoms and molecules have mass and are always in
motion ⇒ kinetic energy
Gravity
Potential Energy
Energy stored within a physical system as a result of the
positions of the components of the system
Potential energy comes from a force acting on an object:
• For gravity: Epotental = mass x gravity constant x height
• Electrostatic potential energy comes from the
interaction of charged particles
• Chemical energy comes from the energy in the bonds
that hold atoms together.
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Energy Transfers: Work and Heat
Work
• Energy used to move an object over a distance is work.
work = Force  distance
• Work can be done on a system by the surroundings or
by the surroundings on the system
Heat
• Energy transferred from hotter object to cooler object
• Heat can be transferred into or out of a system
First Law of Thermodynamics:
Energy is neither created nor destroyed.
Tracking energy in systems:
Exothermic/Exergonic:
Energy is given off to surroundings
Endothermic/Endergonic:
Energy is absorbed from surroundings
 Change in energy of the system:
∆E = Efinal – E initial
E is positive when energy is gained by the system
E is negative when energy is lost to the surroundings
Δ = delta = “change in”
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Energy Diagrams
E
E
E = Efinal − Einitial
Exothermic reaction:
Products have less
energy than reactants;
excess energy went to
surroundings
E is negative
Endothermic reaction:
Products have more
energy than reactants;
system took energy
from surroundings
E is positive
Measuring Energy Changes: Joules and Calories
• SI derived unit of energy is the joule with units of
1 J = 2 kg mass moving at 1 m/s
kg ∙ m2
s2
• 1 calorie = energy to raise temp of 1 g of water by 1°C.
• 1 calorie (cal) = 4.184 joules (J) exactly
• kcal = energy to raise 1 kg of water 1°C.
• Food calories = kcal or Cal.
How many joules in each pack?
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Another Unit for Energy: The Watt
1 Joule is the work required to produce 1 watt of power
for one second
1J=1W∙s
or
1 W = 1 J/s
Electricity is measured in kWh or kilowatt hours
1kWh = 1000 W x 60 min x 60 sec = 3.60 x 106 J
1h
1 min
Convert 225 Cal to Joules
Given:
225 Cal
Find:
?J
Conversion
Factors:
1 Cal = 1000 cal
1 cal = 4.184 J
Solution
Map:
Cal
cal
1000 cal
1 Cal
Solution:
225 Cal 
Check SFs &
Round:
Check:
3 sig figs
J
4.184 J
1 cal
1000 cal 4.184 J
 9.41400  105 J

1 Cal
1 cal
225 Cal = 9.41 x 105 J
3 significant figures
Units and overall size of number are correct.
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Exothermic Processes
• When a change results in the release of energy , it is
called an exothermic process.
• An exothermic chemical reaction: reactants have more
chemical potential energy than the products.
Na2O2 + 2Zn → Na2O + 2 ZnO
The reaction is initiated by the
addition of water
• The excess energy is released into the surrounding
materials often as heat or light.
Endothermic Processes
• Endothermic process: A physical or chemical change that
requires the absorption of energy.
• Endothermic chemical reaction: products have more
chemical potential energy than the reactants.
• The required energy is absorbed from the surroundings
Ba(OH)2∙8H2O(s) + 2NH4SCN(s) → Ba(SCN)2(s) + 2NH3 (g) + 10 H2O(s)
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Fahrenheit and Celsius Scales
• Important temperatures on the Fahrenheit scale are the
freezing point (32 °F) and boiling point (212 °F) of water.
– Room temperature is about 72 °F
• Reference points for the Celsius scale are the freezing (0 °C)
point and boiling point (100 °C) of distilled water.
– Room temperature is about 22 °C.
• 0° F is a lower temperature than 0° C
• A Celsius degree is 1.8 times a Fahrenheit degree.
C 
F - 32
1.8
What is 55° F on the Celsius scale?
The Kelvin Scale
• Both the Celsius and Fahrenheit scales have negative
numbers; Kelvin values are always positive.
• 0 K is called absolute zero. It is too cold for matter to exist
because all molecular motion would stop.
– 0 K = -273 °C = -459 °F.
– Absolute zero is an extrapolated theoretical value
• The size of a “degree” on the Kelvin scale is the same as on the
Celsius scale
K  C  273
What would –25 °C be on the Kelvin scale?
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Temperature vs. Heat
• Temperature is related to the random
motion of molecules
• More motion→ higher temperature
• Units are oF, oC or K
• Heat: exchange of thermal energy
caused by a difference in temperature
• Units of heat = energy units: J or cal
Heat energy in → increase in temp of water
Change in temp (ΔT) depends on:
How much heat energy comes in
What substance is in the pipe
How much of the substance there is
Energy & Temperature of Matter
• As an object loses or gains heat, its temp changes in
proportion to amount of thermal energy transferred.
– Double the heat energy, temp increases twice as much.
• The amount the temperature of an object increases
depends on its mass.
– Double the mass, need 2x as much heat energy to raise
the temperature the same amount.
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silicone: ~20 J
ceramic: ~17 J
Heat Capacity and Specific Heat
steel: ~10 J
wood: ~40 J
The amount of energy required to raise the temperature of an
object by 1 K (1C) is its heat capacity. Units are: cal/°C or J/°C
Specific heat capacity (or specific heat) is the energy required to
raise the temperature of 1 g of a substance by 1 K.
Other Substances
Substance
Specific Heat
(J/g∙K)
Sand
0.835
Soil
0.800
Wood
1.7 (1.2 to 2.3)
What substance would you rather have between you and a heat source?
Specific Heat Capacities
Substance
Specific Heat
J/g°C
Aluminum
0.903
Carbon (dia)
0.508
Carbon (gra)
0.708
Copper
0.385
Gold
0.128
Iron
0.449
Lead
0.128
Silver
0.235
Ethanol
2.42
Water (l)
4.184
Water (s)
2.03
Water (g)
2.02
Ethanol
Ice
Structure
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How Much Heat?
• The amount of heat it takes to heat an object depends on 3
factors:
– how much material there is (mass of the object)
– what the material is (specific heat capacity)
– how hot you want to heat it (T)
• Amount of Heat = Mass x Heat Capacity x Temperature Change
q = m x Cs x T
q is heat in joules
m is mass in grams
Cs is the specific heat capacity (J/g ∙ oC
T = Tfinal – Tinitial = change in temperature in oC
Do you believe:
• Freezing is an exothermic process?
• Can you protect an orange grove by spraying water on
the trees when freezing temperatures are expected?
• The amount of heat released when a substance goes
from liquid to solid state is called the heat of fusion (ΔHf)
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Tracking the Heat
NaC2H3O2 ∙ 3H2O(l) → NaC2H3O2 ∙ 3H2O(s) + heat
∆Hfusion = 289 kJ/kg
How much heat is given off by a hand warmer containing 16.74 g of
NaC2H3O2 ∙ 3H2O?
16.74 g x 289 kJ/kg x 1000j/kJ x 1 kg/1000 g = 4837.86 J or 4838 J
If you put the activated hand warmer in a foam cup containing 100.0 g
water at 21.0 oC, how much hotter would the water get? The specific
heat of water is 4.184 J/g∙ oC
q = Cs ∙ m ∙ T
T = q/ Cs ∙ m
T = 4837.86 J/(4.184 J/g∙oC x 116.74 g) = 9.905 oC hotter
Final T = 30.9 oC
Let’s Try It!
How much heat is released from a big hot pack?
Data needed:
•
•
•
•
•
q = Cs ∙ m ∙ T
Cs for water is 4.184 J/g ∙ oC
Hot pack weighs:
Water in cooler weighs:
Starting temp of water is:
Final temp of water:
m=
T =
q=
q = What is ΔHf in J/g?
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Calculate the Amount of Heat Released When 7.40 g of Water
Cools from 49° to 29°C
Given:
T1 = 49 °C, T2 = 29 °C, m = 7.40 g
Find:
q, J
Solution Map:
Cs m, DT
q
q  m  Cs  ΔT
q = m ∙ Cs ∙ T
Cs = 4.18 J/gC (Table 3.4)
Relationships:
Solution:
T  T2  T1
T  29 C - 49C 
 - 20 C
q  m  Cs  ΔT


 7.40 g   4.18 g C  - 20 C 
J
 618.64 J  6.2  10 2 J
Check:
The unit and sign are correct.
For Wednesday
(No Class on Monday)
• Finish Chapter 3 , First Lecture Quiz (Ch 1-3) is Wed
– Key Concepts: Classifying matter; energy and its units
– Density and heat capacity equations
– Come to Tuesday class (10 am to 11:50 am for review)
• Online Homework:
– Chapters 1 and 2
01/19/11 05:00 pm
– Chapters 3 and 4
01/26/10 11:00 pm*
*Do the Chapter 3 problems before the exam!
• Density Lab Pages will be due (Data and Calculations,
Conclusions, Post Lab questions)
• Next week’s lab: Prop 0603 Resolving a Mixture; Pre
Lab will be due.
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