Energy Energy (E) is the ability to do work. Many types, but we
can say 3 main types: Radiant Potential Kinetic
Slide 3
US Energy Consumption by Source
Slide 4
Radiant Energy Light Energy Visible and Invisible Travels in
waves over distances Electromagnetic waves Waves that spread out in
all directions from the source Visible light, UV light, Infra Red
Radiation, X-rays, microwaves, radio waves
Slide 5
Potential Energy (PE) Stored Energy Due to position
Gravitational PE Elastic PE Chemical bonds Chemical PE Nuclear
energy Fuels Attractions between molecules
Slide 6
Kinetic Energy (KE) Energy of motion Atomic vibrations
Molecular movement Vibration Rotation Translation Movement of
subatomic particles
Slide 7
Kinetic Energy Can be calculated :
Slide 8
How are each type shown here? Radiant Rainbow = visible light
Kinetic Windmill moving Potential All molecules store energy Water
in clouds Air Materials the windmill is made from, the plants at
the bottom
Slide 9
Temperature Scales: Measuring that Thermal Energy
BoilingFreezing Fahrenheit ( o F)21232 Celsius ( o C)1000
Kelvin373273
Slide 10
A Note on the Fahrenheit Scale NEVER use it in this class.
Ever. Only Belize and the US use this scale. Gabriel Fahrenheit
made great thermometers. His scale was replicated the world over
because of this. But if you stop and think about it, does 32F for
freezing make sense, or 212F for boiling? 180 degrees separates
them. 100 degrees, as in the Celcius scale (sometimes called the
Centigrade scale) makes much more sense. Fahrenheit based 0F on the
freezing point of water mixed with NH 4 Cl, and 32F for freezing
water, and 96F for human body temperature (he was off by 2.6). Why?
Because he felt like it and it was easy to draw lines at those
intervals. ( According to a letter Fahrenheit wrote to his friend
Herman Boerhaave, [8] his scale was built on the work of Ole Rmer,
whom he had met earlier. In Rmers scale, brine freezes at 0
degrees, ice melts at 7.5 degrees, body temperature is 22.5, and
water boils at 60 degrees. Fahrenheit multiplied each value by four
in order to eliminate fractions and increase the granularity of the
scale. He then re-calibrated his scale using the melting point of
ice and normal human body temperature (which were at 30 and 90
degrees); he adjusted the scale so that the melting point of ice
would be 32 degrees and body temperature 96 degrees, so that 64
intervals would separate the two, allowing him to mark degree lines
on his instruments by simply bisecting the interval six times
(since 64 is 2 to the sixth power). I took this from Wikipedia. (
According to a letter Fahrenheit wrote to his friend Herman
Boerhaave, [8] his scale was built on the work of Ole Rmer, whom he
had met earlier. In Rmers scale, brine freezes at 0 degrees, ice
melts at 7.5 degrees, body temperature is 22.5, and water boils at
60 degrees. Fahrenheit multiplied each value by four in order to
eliminate fractions and increase the granularity of the scale. He
then re-calibrated his scale using the melting point of ice and
normal human body temperature (which were at 30 and 90 degrees); he
adjusted the scale so that the melting point of ice would be 32
degrees and body temperature 96 degrees, so that 64 intervals would
separate the two, allowing him to mark degree lines on his
instruments by simply bisecting the interval six times (since 64 is
2 to the sixth power). I took this from Wikipedia.
Slide 11
Kelvin Temperatures Based on absolute zero (0 K, -273 o C) The
temperature at which ALL KE stops NO molecular motion. Lowest
temperature theoretically possible Cant really get there in real
life (See 3 rd Law of Thermodynamics in a few slides) K = o C + 273
Technically 273.14, but we can stop at 3 significant digits
Slide 12
Why do we need the Kelvin scale? Two reasons We need a scale
that is relative to molecular motion for certain topics You cant
use negative numbers to indicate motion when it IS present -20C
makes NO sense in light of indicating motion And 40C ISNT twice as
much motion as 20C, (40K IS twice the motion of 20K) Because when
working with equations, cant use zero We get undefined answers if
we divide We get answers of 0 if we multiple And those answers
would NOT make sense if compared to answers calculated with a
positive or negative number
Slide 13
Energy (E): The ability to do work Exergy: The energy available
to do work No symbol Entropy (S): The measure of the disorder of a
system Enthalpy(H): The thermal energy (heat) content of a system
The 4 Es: Energy, Exergy, Entropy, & Enthalpy There will be
more on these!
Slide 14
Thermodynamics The study of energy flow inter-relation between
heat, work, and energy of a system Summary of the three laws: 1.
The energy in the universe is constant 2. Things get more
disorganized over time in a system until everything is equal 3. You
cant reach absolute zero
Slide 15
1 st Law of Thermodynamics The energy in the universe is
constant E=mc 2 Law of Conservation Matter Matter can not be
created or destroyed Law of Conservation of Energy Energy can not
be created or destroyed However, matter and energy can both change
forms in chemical reactions Can also interconvert between matter
and energy in NUCLEAR reactions (more on this later this year.)
Summed up: You can not win. You cant get something for nothing
because energy and matter are conserved.
Slide 16
Energy Time
Slide 17
Before the 2 nd Law Entropy (S) is a measure of DISORGANIZATION
in a system (this simply put; there is a much more complicated
description about the unavailable energy to do work) Anything
disorganized has higher entropy than something organized Exergy is
the Energy available to do work
Slide 18
2 nd Law of Thermodynamics Things get more disorganized over
time in a system until everything equilibrium is reached
(everything is equal) Heat flows from hot to cold, not the reverse
Law of Entropy By nature, things get more disorganized to spread
out energy and matter The quality of the energy (which is exergy)
decreases over time Summed up: You can not break even. You can not
return to the same energy state because things get more
disorganized (gain entropy)
Slide 19
Exergy and Energy The energy of the universe is constant, but
exergy is constantly consumed. This can be compared with a
tooth-paste tube: When you squeeze the tube (= conduct any process)
the paste (= exergy) comes out. You can never put the paste back in
the tube again (try!), and in the end you have only the tube itself
(= low-exergy) left. When you squeeze the tube, the depressions (=
entropy) will increase. (The entropy of a system increases when
exergy is lost) But you can never take the depressions in the tube
and 'un-brush' your teeth. (I.e. entropy is not negative
exergy.)entropy When you buy energy from the electricity network,
you actually buy exergy. You can find the energy as room
temperature heat after some time, but you can not take that room
temperature energy back to the electricity company and ask for
money back. They won't accept it.
Slide 20
Energy and Matter Gain Entropy Over Time
Slide 21
Exergy: The Energy available to do work
Slide 22
3 rd Law of Thermodynamics You cant reach absolute zero and
expect things to happen At absolute zero, all kinetic motion
ceases. And that energy needs to go somewhere. It goes to something
else. And gets transferred back until everything is at an equal
temperature. Summed up: You can not get out of the game, because
absolute zero is unobtainable.
Slide 23
Law of Conservation of Energy Energy cannot be created or
destroyedbut it CAN change forms. Example: Burning wood in a fire
The energy in chemical bonds is released as heat (KE and PE), light
(RE), sound (KE) These forms of energy are less useful have less
exergy
Slide 24
Radiant Energy: EM Waves Potential Energy: Stored Kinetic
Energy: Motion
Slide 25
The CPE in these items could:
Slide 26
Rio Summer Olympics Proposed Solar Waterfall
http://www.snopes.com/ph otos/architecture/solartowe r.asp
Slide 27
Combinations of PE and KE are very common on a large scale KE
and PE animation
Slide 28
PE and KE
Slide 29
When E changes forms The amount of energy one thing loses is
gained somewhere else. E lost = E gained (Law of Conservation of
Energy) But the E gained is usually not all in one place (2 nd Law
of thermodynamics) It is spread out (more entropy) Often in the
forms of heat and light Which are less useful (less exergy)
Slide 30
Energy Transformations
Slide 31
Whats up with Temperature vs Heat? Temperature is related to
the average kinetic energy of the particles in a substance. Thermal
Energy: KE + PE on the small scale
Slide 32
As temperature increases, so does thermal energy (because the
energy of the particles increased). If the temperature stays the
same, the thermal energy in a more massive substance is higher
(because it is a total measure of energy). Thermal energy
relationships
Slide 33
Heat The flow of thermal energy from one object to another.
Heat always flows from warmer to cooler objects. Ice gets warmer
while hand gets cooler Cup gets cooler while hand gets warmer
Slide 34
Heat and Temperature Heat: the measure of the flow of RANDOM
kinetic energy Temperature: the measure of heat Sotemperature is a
measure of kinetic energy of the particles of a substance *
Sometimes heat is radiated as IR (infra-red radiation, a form of
radiant energy)
Slide 35
PE from how the molecules are placed relative to each other
(attractions) Farther = more PE, just like how something farther
off the ground has higher gravitational PE Thermal Energy Thermal
Energy is the total of all the (kinetic and potential) heat energy
of all the particles in a substance.
Slide 36
Energy is being gained/ absorbed by the object or substance
(called the system) from the surroundings Have positive change in
enthalpy values (+ H) Energy is lost/ released from the object or
substance (called the system) to the surroundings Have negative
change in enthalpy values (- H) EndothermicExothermic Exothermic
and Endothermic Processes
Slide 37
The big picture How do we see this energy cycling in the real
world, and not just as a part of Chemistry class? Around the house?
In the environment? While thinking about a car?
Slide 38
If the cup is the system, it is undergoing an exothermic
process because it is losing heat to the surroundings (hand) Ice
gets warmer while hand gets cooler Cup gets cooler while hand gets
warmer If the ice is the system, it is undergoing an endothermic
process because it is absorbing heat from the surroundings
(hand)
Slide 39
Which is process is endothermic? Which is exothermic?
Slide 40
3Consumers: Carnivores and Omnivores 2Consumers: Carnivores and
Omnimores 1Consumers: Herbivores Producers: Autotrophs Trophic
Levels and Energy Energy Out; 90% per level Consumers are all
heterotrophs
Slide 41
Can the world really run out of Energy? World-Wide Energy
Sources, (2007)
Slide 42
PHASE CHANGES & ENERGY
Slide 43
Phase Diagrams Tell what state of matter a material is in at a
given temperature and pressure The triple point is the pressure and
temperature when a solid, liquid, and a gas of the same substance
exist at equilibrium Equilibrium: When there is no net change Here
referring to changes in state Can also refer to temperature and
chemicals The critical point is the temperature above which a
substance will always be a gas, regardless of pressure Fullerton
Phase Diagram Explorer Link Fullerton Phase Diagram Explorer
Link
Slide 44
Phase Diagrams
Slide 45
Phase Diagram for Water
Slide 46
A few terms Freezing Point - The temperature at which the solid
and liquid phases of a substance are in equilibrium at atmospheric
pressure. The same temperature as the melting point Boiling Point -
The temperature at which the vapor pressure of a liquid is equal to
the pressure on the liquid. Vapor Pressure- The pressure at which
the vaporization rates are equal to condensation rates
Slide 47
Phase Changes Enthalpy(H): The heat (thermal energy) content of
a system
Slide 48
States of Matter and Entropy The states are NOT plateaus
because entropy is NOT constant. This isnt a phase change
diagram.
Slide 49
Energy and Matter and Connected Any change in matter ALWAYS is
accompanied by a change in energy
Slide 50
Phase Changes and Energy
Slide 51
Heating Curve Temperature, C Time, min
Slide 52
Why does temperature remains constant when melting or boiling?
During melting or boiling, energy is absorbed from the surroundings
Due to the increase in the thermal energy of the particles from the
increase in PE of the particles Molecules are moving apart breaking
attractions which Absorbs latent (hidden) heat can not be measured
on a thermometer Substance (system) gets warmer
Slide 53
The Es and Heating Endothermic process Energy is absorbed from
surroundings Entropy increases Enthalpy is positive (+H) since heat
added Exergy decreases
Slide 54
Why does temperature remains constant when freezing or
condensing? During freezing or condensing, energy is released to
the surroundings Due to the decrease in the thermal energy from the
decrease in PE of the particles Molecules are moving closer forming
new attractions that are Releasing latent (hidden) heat can not be
measured on a thermometer Substance (system) gets colder
Slide 55
The Es and Cooling Exothermic process Energy is lost to
surroundings Entropy decreases Enthalpy is negative (-H) since heat
is lost Exergy increases
Slide 56
What happens during each segment
Slide 57
Cooling Curve: The Reverse of a Heating Curve
Slide 58
Measuring the Energy of Phase Changes The math of thermal
energy flow
Slide 59
REMEMBER: Energy and Matter and Connected Any change in matter
ALWAYS is accompanied by a change in energy This includes changes
in temperature and/ or phase
Slide 60
Things heat up or cool down at different rates. Land heats up
and cools down faster than water, and arent we lucky for that!?
Specific Heat : c
Slide 61
Specific heat is the amount of heat required to raise the
temperature of 1 kg of a material by one degree C c water = 4.184 J
/ g C the number is high; water holds its heat c sand = 0.664 J / g
C less E than water to change it; it doesnt hold heat as well as
water does This is why land heats up quickly during the day and
cools quickly at night and why water takes longer.
Slide 62
Why does water have such a high specific heat? Water molecules
form strong attractions with other water molecules; it takes more
heat energy to break those attractions than other materials with
weaker forces of attraction between them. water metal
Slide 63
Specific Heat Capacities of Selected Substances c water = 4.184
J / g C c ice = 2.09 J / g C c steam = 1.99 J / g C c sand = 0.664
J / g C c Al = 0.90 J / g C c Fe = 0.449 J / g C
Slide 64
Heat can be Transferred even if there is No Change in State q =
mcT
Slide 65
Remember this? Which is process is endothermic? Which is
exothermic? Now we care about how much energy is being transferred,
and are ready to calculate that change.
Slide 66
Calculating Changes in Energy: The Calorimetry Equation q = mc
T q = change in thermal energy (+) value means heat is absorbed (-)
value means heat is released m = mass of substance T = change in
temperature (T final T initial ) c = specific heat of substance
Each substance has a different c (see CRH, p__) Different states of
matter for the same substance may have a different c
Slide 67
Specific Heat Capacity Problems If 25.0 g of Al cool from 310 o
C to 37 o C, how many joules of heat energy are lost by the Al?
Notice that the negative sign on q signals heat lost by or
transferred OUT of Al. Was this an endothermic or exothermic
process?
Slide 68
Or Heat Transfer can cause a Change of State Changes of state
involve energy changes Changes of state involve energy changes at
constant T Ice + 334 J/g (heat of fusion) -----> Liquid water Is
there an equation? Of course!
Slide 69
Or Heat Transfer can cause a Change of State Changes of state
involve energy Changes of state involve energy at constant T H 2 0
(s) +334 J/g H 2 0 (l) Ice + 334 J/g (heat of fusion) Liquid water
q = mH fusion m = mass H fusion = the enthalpy of melting the
change in thermal energy associated with melting Units are J/g or
KJ/Kg
Slide 70
q = m H fusion WHY DO I NEED THIS WHEN I HAVE q = mcT? Well,
when a phase changes THERE IS NO change in temperature but there is
definitely a change in energy!
Slide 71
Sample Problem: How much heat energy is required to melt 25.0g
of ice, (assuming constant temperature of OC)? Value is positive,
which means heat is absorbed, which makes sense!
Slide 72
molecule Latent heat* and the PE of particles Regular
arrangement breaks up strong attraction weak attraction *Latent
means hidden. Latent heat is the thermal energy (potential energy)
associated with the attractions between molecules, and can not be
measured with a thermometer.
Slide 73
PE related to the forces of attraction between the particles
Energy has to be supplied to oppose the attractive force of the
particles. PE as molecules separate solid liquid or liquid gas
average potential energy Latent heat and the PE or particles
Slide 74
The transfer of energy does not change the KE. Temperature does
not change. latent heat = change in PE between molecules during
change of state Latent heat and PE Video and song:
http://www.youtube.com/watch?v=jaaGqui9NVY
Slide 75
Remember Energy changes accompany changes in state; either:
Energy is added (endothermic) Gain thermal energy Molecules Move
more (gain KE) Separate (gain PE from broken attractions between
molecules) Have a higher entropy Are more disorganized Or Energy is
removed (exothermic) Molecules move less Lose thermal energy Move
less (lose KE) Move closer (lose PE from new attractions between
molecules) Have lower entropy Get more organized
Slide 76
Latent Heats You have a certain energy change associated with
changing state. These values are usually reported for fusion and
vaporization as: H fusion = (latent) Heat of fusion (melting) H
vaporization = (latent) Heat of vaporization H sublimation
=(latent) Heat of sublimation Different materials have different
values for each
Slide 77
What about freezing and condensation? Values for freezing and
condensation are not typically listed, but are the negative values
of those for fusion and vaporization because the energy transferred
is the same, but in the opposite direction (latent) Heat of
freezing= - H fusion (latent) Heat of condensation= - H
vaporization
Slide 78
Enthalpy changes with phase changes
Slide 79
Enthalpy values for H 2 O H fusion = 334 J/g H vaporization =
2259 J/g H sublimation = 2594 kJ/g
Summing it all up: How do you know what to do to calculate
energy changes? Check to see if there is a temperature change. If
yes, use q=mc T. Also, check to see if there is a phase change. If
yes, you need to use q= H fusion mass or q= H vaporization mass
depending on which one applies* or both if there are two phase
changes *If the material freezes or condenses. You can use the
negative value H fusion or H vaporization
Slide 83
How much energy is required to change 0.5 kg of water at 0 C to
ice? Things you know: m = There is____ temperature change, and
there is change of state (freezing) The water is going __________
So.. this all tells you to use _________(negative of melting value)
in q= (The negative value makes sense since you are cooling the
water, so energy leaves)
Slide 84
Total energy required How much energy is required to melt 0.5
kg of ice at 0 C temperature raised to 80 C?
Slide 85
Heat & Changes of State What quantity of heat is required
to melt 500. g of ice and heat the water to steam at 100. o C? Heat
of fusion of ice = 334 J/g Specific heat of water = 4.184 J/gC Heat
of vaporization = 2259 J/g +334 J/g +2257 J/g
Slide 86
So if I want the total heat to take ice and turn it to steam I
need to add values from 3 steps 1.To melt the ice I need to
multiply the heat of fusion with the mass q = H fusion m 2.Then,
there is moving the temperature from 0C to 100 C. For this there is
a change in temperature so we For this there is a change in
temperature so we use q= mcT 3.That just takes us to 100 C, what
about vaporizing the molecules? We need q= H vaporization m Add up
all the values, and you have it. (However, if you are taking it
from below the freezing point to above 100 C, you need to add in
the changes with q=mc T there, too!) Putting it all together
Slide 87
And now More! Heat & Changes of State How much heat is
required to melt 500. g of ice and heat the water to steam at 100 o
C? 1. To melt ice 2.To raise water from 0 o C to 100 o C : 3.To
evaporate water at 100 o C: 4. Total heat energy =
Slide 88
Maybe a picture can help.
Slide 89
Putting it all together: How are matter and energy related?
What influences does energy have on matter? What does this tell us
about the world as we know it?
Slide 90
Slide 91
Making Pizza: Changing Matter Describe the pizza making process
in terms of: Matter States (s, l, g) Elements, compounds, mixtures
Homogeneous and heterogeneous mixtures Properties and changes Both
chemical and physical Intrinsic (intensive) and extrinsic
(extensive) Energy