MATTER
ENERGY
MATTER:
the&stuff
of the
universe
Matter
• Matter: anything that takes up space and
has mass
• Mass: the amount of “stuff” in an object
• Weight: gravity’s pull on mass
– on Earth, this is the same as mass
• Law of Conservation of Matter: “matter
cannot be created or destroyed”…but it can
be rearranged (via chemical or physical
reactions, the same as energy)
Atoms
Matter is composed of atoms
• Atoms: The smallest piece of an element that
still retains the properties of that element
• can not be further divided and still have those
properties
– Are composed of protons, neutrons, and electrons
Pure Substance
• A substance with uniform
and constant composition
(set formula)
– Ex: H2, Cl2, C6H12O6, NaCl,
H2O, Fe
• Elements and compounds
are both considered to be
pure substances
States of Matter
Solid
Liquid
Gas
Condensed States: Solids
• Solids
–Have a definite shape and volume
–Molecules are
• Tightly packed
• Moving VERY slowly in place or only
vibrating
• Very strongly attracted to each other
• Not compressible
Condensed States: Liquids
• Liquids

Have NO set shape but definite volume
–Molecules are
• moving faster than in a solid at the
same temperature
• Loosely packed (slip and slide)
• Strongly attracted to each other
• Not compressible
Expanded States: Gas
• Gas (vapor)
– Have no definite shape
• Depends on the container
– Have no definite volume
• Varies depending on temperature and pressure
– Molecules are
• Not very strongly attracted to each other
• Moving faster than those in a liquid at the same
temperature
• As a result are
– spread out as far as possible
– moving freely past each other
– compressible
Expanded States: Plasma
• Plasmas
– Occur at extremely high temperatures
– Are ionized matter
• May be some atoms or molecules but mostly
– Nuclei and electrons floating around not attached
– Usually are good conductors of electricity
– Most common form of matter in the universe
• Stars
– Here on Earth: flames, lightening, aurora borealis,
plasma TVs, neon signs
You WILL clean up
any mess you make
today and always!
What is a non-Newtonian fluid?
– You have to pull the trigger on a water pistol to get the
water to squirt out. To make the water to come out
faster, you have to pull the trigger harder. Fluids resist
flow. This phenomenon is known as viscosity.
– Newton devised a simple model for fluid flow that
could be used to relate how hard you have to pull the
trigger to how fast the liquid will squirt out of the
pistol. Picture a flowing liquid as a series of layers of
liquid sliding past each other.
– The resistance to flow arises because of the friction
between these layers. If you want one layer to slide over
another twice as fast as before, you'll have to overcome a
resisting force that's twice as great, Newton said.
– The slower one layer slides over another, the less
resistance there is, so that if there was no difference
between the speeds the layers were moving, there would
be no resistance.
– Fluids like water and gasoline behave according to
Newton's model, and are called Newtonian fluids.
– But ketchup, blood, yogurt, gravy, pie fillings, mud, and
cornstarch paste DON'T follow the model. They're nonNewtonian fluids because doubling the speed that the
layers slide past each other does not double the
resisting force. It may less than double (like ketchup),
or it may more than double (as in the case of quicksand
and gravy). That's why stirring gravy thickens it, and
why struggling in quicksand will make it even harder to
escape.
– For some fluids (like mud, or snow) you can push and
get no flow at all- until you push hard enough, and the
substance begins to flow like a normal liquid. This is
what causes mudslides and avalanches.
Walking on Non-Newtonian Fluids


Walking on cornstarch- en Español
Mythbusters- Ninja Walking
What is an element?
• Element: A substance that can not be
changed into a simpler substance(s)
under normal laboratory conditions
– A substance made up of only one type of
atom is an elemental substance
How do we represent elements?
– Represented by symbols on the periodic table
• Usually comes from the name, a person, or a
place
• One or two letters for those with official names
–First letter is ALWAYS capitalized
–Second letter is never capitalized
• Unofficially named elements have three letters,
starting with a capital U
Meet the Elements
 118 Elements
 94 are naturally occurring
 90 of those were discovered in nature
 27 were discovered by being made by
man; of those
 Tc, Pm, and Pu were found to exist
after being synthesized
 Es found in residue from atomic
testing
 everything is made from these elements
Elements in Nature
• Everything in nature is made up of elements
• Some elements are never found in their pure
form in nature, but commonly are found in
compounds
– ex: Na (sodium) is virtually non-existent in pure
form, but found in common NaCl, a compound
we call table salt.
Elements
When elements were discovered
Allotropes
•Allotropes are different forms of an
elemental material
• Shown here are allotropes of carbon
•Diamond, graphite, and
buckminsterfullerene
•Difference is how the atoms are
connected to each other
MOLECULES
Composition of molecules is given
by a MOLECULAR FORMULA
H2O
C8H10N4O2 - caffeine
Molecule
• Is defined as two (2) or more atoms
sharing electrons in definite
proportions
– O2
– H2
– H2O
– C6H12O6
• is the smallest unit of a compound that
retains the chemical characteristics of the
compound.
ELEMENTS THAT EXIST AS DIATOMIC
MOLECULES
Remember:
BrINClHOF
These elements
only exist as
PAIRS. Note that
when they combine
to make
compounds, they
are no longer
elements so they
are no longer in
pairs!
CHEMICAL COMPOUNDS are
composed of atoms and so can be
decomposed to those atoms.
The red compound is composed
of
• nickel (Ni) (silver)
• carbon (C) (black)
• hydrogen (H) (white)
• oxygen (O) (red)
• nitrogen (N) (blue)
Compounds
• Two (2) or more DIFFERENT elements
combined in definite proportions
– Has own, unique:
• Formula (law of definite proportions)
• properties
– Boiling point, freezing point, density, etc
– Need a chemical reaction to separate them into
the elements it is made from
• 2 H2O -> 2H2 and O2
Formulas of Compounds (and
Molecules)
•
•
•
Tells how many atoms of each element are
present.
Uses symbols and subscripts
Rules:
– each element is represented by its symbol
– The number of each type of atom is written as a
subscript next to the element
•
1 is not written
Properties of Compounds
 properties
differ from
those of individual
elements under the
same conditions

EX: table salt (NaCl)


Na is a metal
Cl2 is a gas
 Ex: water (H2O)
 H2 is a gas
 O2 also a gas
Law of Definite Proportions
• In samples of any chemical compound, the
masses of the elements are always in the
same proportions
– H2O is always H2O, and never H2O2
– Once the formula is different, the materials and
the properties are different
What is the difference between
compounds and molecules?
Molecules: two
or more atoms
bonded
Compounds: two
or more atoms of
different
elements bonded
•The term molecule is
more general,
compound is more
specific.
•Molecules can be
elemental – like O2,
or Cl2
•Compounds by
definition can not be
elemental.
Separation of compounds into their
elements
• Compounds can
be broken into
their elements by
chemical
reactions, such as
• Electrolysis. This is
the decomposition
of a substance
(here, water) into
it’s elements by an
electric current.
Physical and
Chemical Properties
and Changes
Physical Property
• Characteristic of a
substance that can
be observed
without changing
the identity
(formula) of the
substance
• Examples:
– state of matter
– density
– color
– melting pt.
– odor
– boiling pt.
Chemical Property
• Ability of a substance to change to a
different substance
• You must change the identity
(formula) of the substance to observe
a chemical property
• Examples:
– flammability
– reactivity with acids and bases
Intrinsic v. Extrinsic Properties
• Intrinsic (intensive) Property: A
property that does not change when
the amount of substance changes
–Color
–Odor
–Boiling point
–Melting point
Intrinsic v. Extrinsic Properties
• Extrinsic (extensive) property: A
property that changes depending
upon the amount of substance
present
–Mass
–Volume
–Shape
–Size
Density…Intrinsic or Extrinsic?
Density data for a sample of gold.
 Density= mass/ volume
 Units are g/mL or g/cm3

Mass (g)Mass (g)
Volume (mL)
Volume Density
(mL)
46.71 46.71
2.42
2.42
297.80 297.80
15.43
15.43
662.76 662.76
34.34
34.34
147.07 147.07
7.62
7.62
431.162431.162
22.34
22.34
85.69 85.69
4.44
4.44
Density…Intrinsic or Extrinsic?
Density data for a sample of gold.
 Density= mass/ volume
 Units are g/mL or g/cm3

Mass (g)Mass (g)
Volume (mL)
Volume Density
(mL)
46.71 46.71
2.42
2.42 19.3
297.80 297.80
15.43
15.4319.30
662.76 662.76
34.34
34.3419.30
147.07 147.07
7.62
7.62 19.3
431.162431.162
22.34
22.3419.30
85.69 85.69
4.44
4.44 19.3
Physical Change
• Change in the form of the substance, not in its
chemical nature
– No chemical bonds broken or made
• The chemical formula is the same before and after the
change
– The attractions between the molecules are made or broken, not
the molecules themselves
– Examples:
• cutting/tearing
• change in state: melting, boiling,
evaporating, condensing, freezing
• bending
Chemical Change (Reaction)
• Change one substance into another
• Atoms are reorganized
– Bonds are broken and reformed in new
ways
• Examples:
– Burning
– Mixing baking soda and vinegar
– Digestion
– Rusting
Signs of Chemical Change
• Change in:
– state
• Forming a precipitate (solid made from liquids)
• Bubbling/ fizzing
– color
– temperature
• Hot (exothermic)
• Cold (endothermic)
– odor
• giving off light
Mixture
• A blend of 2 or more pure substances
• Have variable compositions
–No chemical formula
• Examples:
soda
coffee
salad dressing
alloys
air
salt water
Air: A Mixture
Types of Mixtures
• Homogeneous
– uniform
– evenly mixed
– cannot see different parts
– Examples
•
•
•
•
salt water
coffee
air
milk
Types of mixture, con’t
• Heterogeneous
–not uniform
–not evenly mixed
–can see different parts
–Examples:
• salad dressing
• sand
• concrete
• salsa
Subtypes of mixtures...
• Solutions: one thing dissolved in another;
they do not separate upon standing or when
filtered (particles are molecules)
– Two parts:
• Solvent- the thing that does the dissolving
• Solute- the thing that gets dissolved
Ex: Salt water – water is the solvent, salt the solute
– Can be any state of matter
• Alloys are solutions of metals (a solid solution)
Subtypes of mixtures...
• Suspensions:
– Heterogeneous mixture of two or more
substances
– may look like a solution but can
separate when standing/ filtering since
– particles are not dissolved
• particles are larger than molecules;
sometimes they can be seen by the naked
eye
Subtypes of mixtures...
• Colloid- microscopically mixed homogeneous solutions
that will disperse light (Tyndall effect)
– They will not separate out upon standing
– Emulsion- a type of colloid made from of two liquids that
normally will not mix (are immiscible; like oil and water)
Tyndall Effect
Tyndall Effect in Nature
• Twenty-four-karat gold is an element
• Eighteen-karat gold is an alloy.
• Fourteen-karat gold is an alloy.
– (Alloys are homogeneous mixtures- specifically
solutions- of metals)
– What is the solvent in 18K and 14K gold? Au
– What is/ are the solute(s) in 18K and 14K gold? Cu
and Ag
Classification of Matter
Separating Mixtures
• Mixtures can be separated by
physical means
– No chemical reactions or changes needed
– 5 main methods are based on physical properties
•
•
•
•
•
Evaporation
Crystallization
Chromatography
Filtration
Distillation
Separating Mixtures
• Evaporation: boiling off a liquid
• Crystallization: solid left behind from
a solution crystallized when the
liquid leaves
–Usually paired
• sweat when it dries
• seawater
• rock candy formation
• geode formation
Chromatography: “color writing”
• Using how quickly things move to separate them
– Use specific ratios of movement through a medium
(paper, water, column) based on how some molecules
move through the medium
– Phases:
• Stationary phase
– stays in place (paper, the column)
– Hangs onto parts of the mixture at different rates
• Mobile phase
– moves over the stationary phase
– carries the parts of the mixture
– Rf values constant for component of a mixture
• Ratio of how far one component moves as compared to the
whole
• Ex: running colors in ink when wet
chromatograms
Column
Chromatography
Gas Chromotography
Filtration
• Use a funnel
and filter paper
• Separates a
solid from a
liquid
• Based on
particle size
• Separates 2 or more
liquids
Distillation
• One liquid has to
have a lower boiling
point
– Heat both; the one
with the lower BP is
boiled out first
– Can use condensation
to collect materials as
they boil out
• Used with crude oil
(cracking) to isolate
separate products
Distillation of Crude Oil:
Cracking
Other methods…
• Use polarity (charge)
– Magnetism
• Use density
– For solids: mixture of pieces is placed in a
solution
• Some float and can be skimmed off
• Some sink
– For liquids: use a separatory funnel
Separatory Funnel
•Used for
liquids that are
immiscible
(will not mix,
like oil and
water)
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, Xrays, microwaves, radio
Potential Energy (PE)
• Stored Energy
– Due to position
• Gravitational PE
• Elastic PE
– Chemical bonds
• Chemical PE
–Nuclear energy
–Fuels
–Attractions
between molecules
Kinetic Energy (KE)
• Energy of motion
– Atomic
vibrations
– Molecular
movement
• Vibration
• Rotation
• Translation
– Movement of
subatomic
particles
Kinetic Energy
Can be calculated:
How are each type shown here?
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
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
– (See 3rd Law of Thermodynamics in a few
slides)
• K = oC + 273
– Technically 273.14, but we can stop at 3
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
• 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
There will be more on these!
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 can’t reach absolute zero
1st 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
Time
Energy
Before the 2nd 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
2nd 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 toothpaste 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
3rd 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 forms of energy are less useful
• have 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
Matter)
– 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 lost/ released
from the object or
substance (called the
system) to the
surroundings
• Have positive change in
enthalpy values (+ΔH)
Exothermic
• Energy is being gained/
absorbed by the object or
substance (called the
system) from the
surroundings
• Have negative change in
enthalpy values (-ΔH)
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 is 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
States ofon
Matter
on a Heating curve
Curve
States
a heating
GAS
Temperature, ̊ C
LIQUID
SOLID
Time, minutes
Changes of State on a Heating Curve
ChangesTwo
ofstates
state
on exist
a heating
curve
of matter
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
Cooling Curve: The Reverse of a Heating Curve
Temp, ̊C
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 = 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 (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•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?
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 = 2594 kJ/g
From:
http:/
/hype
rphysi
cs.phy
astr.gs
u.edu
/hbas
e/tabl
es/ph
ase.ht
ml#c1
•
http://h
yperphy
sics.phyastr.gsu.
edu/hba
se/table
s/phase.
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 T
= (500g)(334 J/ g) + (500.g  4.184J/g°C  80°C)
= (167 J) + (167360J) = 167,527J
=168, 000J= 1.68E5 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)(333 J/g) = 1.67 x 105 J
2.
To raise water from 0 oC to 100 oC : q = mc∆T
q = (500. g)(4.2 J/g• oC)(100 - 0oC) = 2.1 x 105 J
3.
To evaporate water at 100 oC: q = m∆Hvaporization
q =(500.g)(2260 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