Energy and Thermodynamics
Download
Report
Transcript Energy and Thermodynamics
ENERGY
Energy
Energy (E) is the ability to do work.
Many types, but we can say 3 main
types:
Radiant
Potential
Kinetic
US Energy Consumption by Source
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
Potential Energy (PE)
Stored Energy
Due
to position
Gravitational
Elastic
PE
PE
Chemical
bonds
Chemical
PE
Nuclear
energy
Fuels
Attractions
molecules
between
Kinetic Energy (KE)
Energy of motion
Atomic
vibrations
Molecular
movement
Vibration
Rotation
Translation
Movement
subatomic
particles
of
Kinetic Energy
Can be calculated:
How are each type shown here?
How are each type shown here?
Radiant
Rainbow
Kinetic
Windmill
= visible light
moving
Potential
All
molecules store
energy
Water
in clouds
Air
Materials
the windmill is
made from, the plants
at the bottom
Where is the PE? KE?
Temperature Scales: Measuring that
Thermal Energy
Boiling
Freezing
Fahrenheit (oF)
212
32
Celsius (oC)
100
0
Kelvin
373
273
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 32°F for freezing make sense, or 212°F for
boiling? 180 degrees separates them.
100 degrees, as in the Celcius scale (sometimes called the
Centigrade scale) makes much more sense.
Fahrenheit based 0°F on the freezing point of water mixed with
NH4Cl, and 32°F for freezing water, and 96°F 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 Rømer, whom he had met earlier. In Rømer’s 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.
Kelvin Temperatures
Based on absolute zero (0 K, -273 oC)
The
temperature at which ALL KE stops
NO molecular motion.
Lowest temperature theoretically possible
Can’t
really get there in real life
3rd Law of Thermodynamics in a few
slides)
(See
K = oC + 273
Technically
273.14, but we can stop at 3
significant digits
Why do we need the Kelvin scale?
Two reasons
We need a scale that is relative
to molecular motion
for certain topics
You
can’t use negative numbers to indicate motion when it
IS present
-20°C makes NO sense in light of indicating motion
And 40°C ISN’T twice as much motion as 20°C,
(40K IS twice the motion of 20K)
Because when working with equations, can’t 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
The 4 Es:
Energy, Exergy, Entropy, & Enthalpy
Energy (E): The
ability to do work
Entropy (S): The
measure of the
disorder of a
system
There will be more on these!
Exergy: The energy
available to do
work
No symbol
Enthalpy(H): The
thermal energy
(heat) content of a
system
Thermodynamics
The study of energy flow
inter-relation
between heat, work, and energy of a
system
Summary of the three laws:
1.
2.
3.
The energy in the universe is constant
Things get more disorganized over time in a system
until everything is equal
You can’t reach absolute zero
st
1
Law of Thermodynamics
The energy in the universe is constant
E=mc2
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 can’t get something for nothing because
energy and matter are conserved.
Time
Energy
Before the
nd
2
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
nd
2
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)
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.)
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 temeperature energy
back to the electricity company and ask for money back. They won't
accept it.
Energy and Matter Gain Entropy Over Time
Exergy: The Energy available to do
work
rd
3
Law of Thermodynamics
You can’t 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.
Law of Conservation of Energy
Energy cannot be created or
destroyed…but 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
have
forms of energy are less useful
less exergy
Radiant Energy:
EM Waves
Potential
Energy:
Stored
Kinetic
Energy:
Motion
The CPE in these items could:
Rio Summer Olympics
Proposed Solar
Waterfall
http://www.snopes.com/ph
otos/architecture/solartowe
r.asp
Combinations of PE and KE are very
common on a large scale
KE and PE animation
PE and KE
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 (2nd Law of thermodynamics)
It
is spread out (more entropy)
Often
in the forms of heat and light
Which are less useful (less exergy)
Energy Transformations
Thermal Energy: KE + PE on the small scale
• What’s up with Temperature vs Heat?
• Temperature is related to the average kinetic energy
of the particles in a substance.
Thermal energy relationships
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).
Heat
Cup gets cooler while hand
gets warmer
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
Heat and Temperature
Heat: the measure of the flow of
RANDOM kinetic energy
Temperature: the measure of heat
So…temperature 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)
Thermal Energy
•Thermal Energy is the
total of all the (kinetic and
potential) heat energy of
all the particles in a
substance.
•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
Exothermic and Endothermic Processes
Endothermic
Energy is being
gained/ absorbed by
the object or substance
(called the system)
from the surroundings
Have positive change
in enthalpy values
(+ΔH)
Exothermic
Energy is lost/
released from the
object or substance
(called the system) to
the surroundings
Have negative change
in enthalpy values
(-ΔH)
“Exothermic
Reactions? I
studied them
before they
were cool.”
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?
•If the cup is the system, it is
undergoing an exothermic
process because it is losing
heat to the surroundings
(hand)
•If the ice is the system, it
is undergoing an
endothermic process
because it is absorbing
heat from the surroundings
(hand)
Cup gets cooler while
hand gets warmer
Ice gets warmer
while hand gets
cooler
Which process is endothermic? Which is
exothermic?
Trophic Levels and Energy
Consumers
are all
heterotrophs
3°Consumers:
Carnivores and
Omnivores
2°Consumers:
Carnivores and
Omnimores
1°Consumers: Herbivores
Producers: Autotrophs
Energy Out;
90% per level
Trophic Levels and Energy
Producers:
Trophic Levels and Energy
3°Consumers
2°Consumers
1°Consumers
While this shows only a 1° consumer, all animals “lose” E the same way; high
levels lose more to motion than do lower levels
Class Question:
How does a car exhibit all types
of energy?
Explain!
Energy Loss in a Car
http://www.consumerenergycenter.org/transportation/consumer_tips/vehicle_energy_losses.html
Can the world really run out of
Energy?
World-Wide Energy Sources, (2007)
PHASE
CHANGES &
ENERGY
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
Phase Diagrams
Phase Diagram for Water
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
Measuring Vapor Pressure
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/vaporv3.swf
Energy and Phase Changes
Phase Changes
Enthalpy(H):
The heat
(thermal
energy) content
of a system
States of Matter and Entropy
The states are NOT plateaus because entropy is NOT constant.
This isn’t a phase change diagram.
Energy and Matter and Connected
Any change in
matter
ALWAYS is
accompanied
by a change in
energy
Phase Changes and Energy
Temperature, ̊C
Heating Curve
Time,
min
Matter on a Heating
States States
on aof heating
curveCurve
GAS
Temperature, ̊ C
LIQUID
SOLID
Time, minutes
Changes
of State
on
aheating
Heating Curve
Changes
of
state
on
a
curve
Two states of matter exist on the plateaus
LIQUID
becoming
VAPOR
Temperature, ̊ C
SOLID
becoming
LIQUID
Time, minutes
Remember: Phase Changes and Energy
Why does temperature remains
constant when melting or boiling?
During melting or boiling, energy is
from the surroundings
absorbed
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
The E’s and Heating
•Endothermic process
•Energy is absorbed from surroundings
•Entropy increases
•Enthalpy is positive (+ΔH) since heat added
•Exergy decreases
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
The E’s and Cooling
•Exothermic process
•Energy is lost to surroundings
•Entropy decreases
•Enthalpy is negative (-ΔH) since heat is lost
•Exergy increases
Processes of a Heating Curve
GAS
VAPORIZATION
Temperature, ̊ C
LIQUID
FUSION (MELTING)
SOLID
Time, minutes
What happens during each segment
Cooling Curve: The Reverse of a
Heating Curve
Temperature, ̊ C
Time, minutes
Temp, ̊C
Cooling Curve:
The Reverse of a Heating Curve
Cooling
curve
erature, ̊ C
gas
Gas and liquid present
condensation
liquid
Liquid and solid present
Freezing
solid
Time, min
Energy and Phase Change in Nature:
Energy and Phase Change in Nature:
Measuring the Energy of Phase
Changes
The math of thermal energy flow
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
Specific Heat : c
• Things heat up or cool down at different
rates.
Land heats up and cools down faster than water,
and aren’t we lucky for that!?
•Specific heat is the amount of heat required to
raise the temperature of 1 kg of a material by one
degree °C
•cwater = 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 doesn’t 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.
Why does water have such a high
specific heat?
water
metal
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.
Specific Heat Capacities of Selected
Substances
cwater
= 4.184 J / g °C
cice = 2.09 J / g °C
csteam = 1.99 J / g °C
csand = 0.664 J / g °C
cAl = 0.90 J / g °C
cFe = 0.449 J / g °C
Heat can be Transferred even if there is No
Change in State
q = mc∆T
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.
Calculating Changes in Energy: The
Calorimetry Equation
q = mcT
•q = change in thermal energy
•(+) value means heat is absorbed
•(-) value means heat is released
•m = mass of substance
•T = change in temperature (Tfinal – Tinitial)
•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
Specific Heat Capacity Problems
If 25.0 g of Al cool from 310 oC to 37 oC, how many joules of heat energy are lost
by the Al?
•heat gained or lost = q = mc∆T
•where ∆T = Tfinal - Tinitial
•q = (25.0g) (0.897 J/g•oC)(37 - 310)oC
•q = - 6120 J
Notice that the negative sign on q signals heat “lost by” or
transferred OUT of Al.
Was this an endothermic or exothermic process?
Specific Heat Capacity Problems
If 25.0 g of Al cool from 310 oC to 37 oC, how many joules of heat energy are lost
by the Al?
•heat gained or lost = q = mc∆T
•where ∆T = Tfinal - Tinitial
•q = (25.0g) (0.897 J/g•K)(37 - 310)K
•q = - 6120 J
Notice that the negative sign on q signals heat “lost
by” or transferred OUT of Al.
Was this an endothermic or exothermic process?
Or… Heat Transfer can cause a
Change of State
Changes of state involve energy changes at constant T
Ice + 333 J/g (heat of fusion) -----> Liquid water
Is there an equation? Of course!
Or… Heat Transfer can cause a
Change of State
Changes of state involve energy at constant T
H20(s) +333J/g H20(l)
Ice + 333 J/g (heat of fusion) Liquid water
q = mΔHfusion
•m= mass
•ΔHfusion = the enthalpy of melting
• the change in thermal energy associated
with melting
•Units are J/g or KJ/Kg
q = mΔHfusion
WHY DO I NEED THIS WHEN I HAVE
q = mc∆T?
Well, when a phase changes THERE IS NO
change in temperature… but there is definitely
a change in energy!
Sample Problem:
How much heat energy is required to melt 25.0g of ice,
(assuming constant temperature of O°C)?
q = mΔHfusion
•m= 25.0g
•ΔHfusion =333J/g for water
q= (25.0g)333J/g =8385J
Value is positive, which means heat is absorbed, which makes
sense!
Latent heat* and the PE of particles
molecule
strong attraction
Regular
arrangement
breaks up
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.
Latent heat and the PE or particles
Energy has to be supplied to oppose the
attractive force of the particles.
PE as
molecules
separate
PE related to the forces of attraction between the particles
solid liquid or liquid gas
average potential energy
Latent heat and PE
The transfer of energy does not change the KE.
Temperature does not change.
latent heat = change in PE between
molecules during change of state
Video and song:
http://www.youtube.com/watch?v=jaaGqui9NVY
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
Latent Heats
You have a certain energy change associated with
changing state. These values are usually reported
for fusion and vaporization as:
ΔHfusion= (latent)
Heat of fusion (melting)
Δ Hvaporization = (latent) Heat of vaporization
Δ Hsublimation =(latent) Heat of sublimation
Different materials have different values for each
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= -ΔHfusion
(latent) Heat of condensation= -Δ Hvaporization
Enthalpy changes with phase changes
Enthalpy values for H2O
∆Hfusion= 334 J/g
∆Hvaporization= 2259 J/g
∆Hsublimation = 25940 J/g
From:
http:/
/hype
rphysi
cs.phy
astr.gs
u.edu
/hbas
e/tabl
es/ph
ase.ht
ml#c1
http://hy
perphysic
s.phyastr.gsu.e
du/hbase
/tables/p
hase.html
#c2
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=mcΔT.
Also, check to see if there is a phase change.
•
If yes, you need to use
•
•
•
•
q= Δ Hfusionmass
• or
q= Δ Hvaporizationmass
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 Δ Hfusion or Δ Hvaporization
How much energy is required to change 0.5 kg of
water at 0 °C to ice?
Things you know:
•m = 0.5 kg= 500.g
•There is no temperature change, and there is
change of state (freezing)
•The water is going to freeze
So…..
this all tells you to use –ΔHfusion (negative of
melting value) in
q= –ΔHfusionm
q= (-334J/g)(500.g)= -1.67E5J
(The negative value makes sense since you are
cooling the water, so energy leaves)
How much energy is required to melt 0.5 kg of ice at 0 °C
temperature raised to 80 °C?
Total energy required
= latent heat (ice at 0 °C → water at 0 °C)
+ energy (water: 0 °C → 80 °C)
= mΔHf + mc ΔT
= (500g)(334 J/ g) + (500.g x 4.184J/g°C x 80°C)
= (167000 J) + (167360J) =
=334360J= 3.34E5 J
(Value is positive, makes sense since you add E to heat water)
Heat & Changes of State
What quantity of heat is required to melt 500. g of ice and
heat the water to steam at 100. oC?
Heat of fusion of ice = 334 J/g
Specific heat of water = 4.184 J/g•°C
Heat of vaporization = 2259 J/g
+2257
J/g
+334 J/g
Putting it all together…
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
2.
• q = ∆Hfusionm
Then, there is moving the temperature from 0°C to 100°C.
• For this there is a change in temperature so we use
•
q= mc∆T
3.
That just takes us to 100°C, what about vaporizing the molecules?
•
We need q=∆Hvaporizationm
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!)
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 oC?
1.
To melt ice : q = m∆Hfusion
q = (500.g)(334 J/g) = 1.67 x 105 J
2.
To raise water from 0 oC to 100 oC : q = mc∆T
q = (500. g)(4.184 J/g• oC)(100 - 0oC) = 2.1 x 105 J
3.
To evaporate water at 100 oC: q = m∆Hvaporization
q =(500.g)(2259 J/g) = 1.13 x 106 J
4. Total heat energy = 1.51 x 106 J = 1510 kJ
Maybe a picture can help….
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?
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