Download ch 12- states of matter

CHAPTER 12
STATES OF MATTER
4 STATES OF MATTER

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Gases
Liquids
Solids
Plasma
(Ice Ice Baby Rap)
Dance
KINETIC MOLECULAR THEORY
(KMT)- describes the behavior of particles
in terms of their motion
 Explains the effect of temp. and pressure
on matter

KMT video
 Molecules in Motion Song

Assumptions for KMT- gases

1. Particle size- all matter is composed of
small particles
 Between
the particles is empty space
 There are no attractive or repulsive forces
between particles

2. Particle motion- particles are in
constant random motion
 Travel
in straight line paths
 Change direction upon collision
 Collisions are elastic (no net loss of kinetic
energy [KE])

3. Particle energy- kinetic energy (KE) is a
factor of mass and velocity of a particle
 KE=
energy an object possesses because of
its motion
Formula: KE = 1/2mv2
 Depends on mass and velocity

PRESSURE
Dependent on force of collision and the #
of collisions
 Pressure = force
area
 Atmospheric pressure- (Barometric)
results from the collisions of air molecules
w/objects on earth

 Varies
depending on location
Tools for measuring pressure

Manometer = device use to measure
pressure of enclosed gas

Barometer = special manometer used to
measure atmospheric pressure
 Developed
Aneroid
barometer
by Torricelli
Standard Atmospheric Pressure
Average pressure at sea level
 = 760 mmHg (millimeters of mercury)
 Units for pressure:

 Pascal
(Pa) = 1 Newton
m2
 1000 Pa = 1kPa
 Standard atmospheric pressure = 101.3 kPa

Also known as 1 atm (atmosphere)
Pressure Conversion Factor
101.3 kPa = 1 atm= 760 mmHg = 760 torr = 14.7 psi
Example:
216.9 kPa = _____________ mmHg
216.9 kPa
760 mmHg
101.3 kPa
1627 mmHg
763 mmHg = ______________kPa
763 mmHG
101.3 kPa
760 mmHg
102 kPa
TEMPERATURE
DEFINITION: measure of the average KE
of the particles in a substance (> temp. = >
KE)
 If you have 2 substances at same temp.,
they must have the same average KE, so
the molecule w/the less mass will move
faster
 Ex: O2
and
H2
at 20°C
32 g/mole
2 g/mole
-so H2 has the > velocity

ABSOLUTE ZERO

Temperature at which the motion of
particles ceases (stops moving)

Absolute zero = -273°C
Absolute zero
KE
KE = 0
-273
0
Temp. °C
100
Supernova remnent
(star death)

Spitzer Space Telescope




Launched August 24, 2003
With this new infrared telescope, only the science instrument
chamber and a compact cryostat will be cold at launch, chilled to
about 1.5 Kelvin (-272 Celsius, or -457 Fahrenheit). Following
launch from Cape Canaveral Air Force Station in Florida, the
spacecraft cooled in the deep recesses of space for about five
weeks. The observatory uses the vapor from the boil-off of its
cryogen fluid to cool the telescope assembly down to its optimal
operating temperature of 5.5 Kelvin (-268 Celsius, or -450
Fahrenheit).
cryogen depletion date 5/15/09 (now in warm mission- two of its
arrays still working)
James Webb
Vega (dust
cloud
around)
A Joke Break…

Q: What did the thermometer say to the
graduated cylinder?

A: "You may have graduated but I've got
many degrees"
 Ha, Ha, Ha, again I crack myself up.
TEMPERATURE SCALES
Farenheit- 32°F= freezing water; 212°F=
boiling water
 Celcius- based scale on freezing point of
water = 0°C; boiling = 100°C
 Kelvin- based on absolute zero;
0 K = -273°C
(a change of 1 degree on K scale is same
on °C)

Conversion factor K = °C + 273
Example:
298
25 ° C = ___________K
127
400 K = ____________
°C
Use for temperature:
Determine direction of energy flow
 When a cool object (less KE) is in contact
w/a warmer object (higher KE) the energy
of the warmer object will transfer to the
cooler until the KE is equal or same temp.

Heat (em cee delta tee song)
The energy transferred due to a temp.
difference
 Physical and chemical changes are
accompanied by energy changes
 Heat (Q) – energy transferred from a
hotter object to a cooler object due to a
temp. difference

Exothermic- released energy; gives off
heat or light
 Endothermic- absorbs energy; cool to the
touch

LAW OF CONSERVATION OF
ENERGY

ENERGY CAN BE CONVERTED FROM
ONE FORM TO ANOTHER BUT IT IS
NOT CREATED NOR DESTROYED

Activation energy = the minimum amount
of energy needed to get a reaction started
UNITS FOR HEAT
Joule (J)- quantitative measurement of an
energy change or heat
 English system uses a calorie (cal)
 Conversion factors

1
cal = 4.18 J
 1000 cal = 1 Cal (food value) = 1 kcal
little”c”
big “C”
 Ex. 1 tic tac = 4180 J or 1 Cal
Orange Juice
340 kJ (80 Cal)
SPECIFIC HEAT Constant
The heat needed to raise the temp. of 1 g
of a substance 1°C or 1 K
 Represented with the letter “c”
 Unit= J/g°C
 Ex: c water = 4.18 J/g°C

MEASURING ENERGY CHANGES

CALORIMETER- device used to measure
energy changes
 Usually
contains water
 Measures heat absorbed or released
 Follows Law of conservation of Energy
1. Formula for heat Q = m c rT
 Heat lost or gained =

(mass in grams)(specific heat constant)(change in
temp)

How much heat is needed to heat up 145g
of water from 25.0 C to 95.0 C? Specific
heat of water= 4.18 J/g C.
2. Finding specific heat of unknown
substance
 Heat lost = heat gained by water
 Q lost = Q gained
 Q lost = mcrT
Q gained = mcrT
 so, mcrTlost = mcrTgained


A piece of metal with a mass of 35.0 g and
a temperature of 100.0 C is placed into
105.0 g of water at a temperature of
25.0C. After the metal cools the final
temperature of the system is 31.5 C. What
is the specific heat of the metal?
Lab- Specific Heat
Data (w/ units)
 Calculations: (do for each metal)

 1.
Mass of water (data #3- data #2)
 2. Change in temp. water (data #6- data #5)
 3. Change in temp metal (data #4 – data # 6)
mcrTlost = mcrTgained
 (data #1) c (calc #3)=(calc # 1) 4.18 (calc #2)
 (solve for c)
 % Error= O (above) – A x 100

A
 Also: Questions & Conclusion
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What is the pressure in atmospheres if the
pressure is 742 mm Hg?
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If a book with a force or weight of 25 N is
laying flat on a table covers a surface that
is .20 m by .35 m, what is the pressure it is
applying to the table?
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A piece of unknown metal with mass of
14.9 g is heated to 100.0 C and dropped
into 75.0 g of water that was at 20.0 C.
The final temperature of the system is 28.5
C. What is the specific heat of the metal?
Lab- Calories of Food
How to use the Lab Pro
Data (w/ units)
 Calculations (one for each food item)
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 1.
Change in temp (data 7- data 8)
 2. . mass of water (data #4- data #5)
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3. Heat lost by food= Heat gained water
Q
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lost= mc Δ t (water)
Q = (calc. #2) (4.18) (calc #1)
4. Convert heat lost (joules to cal)
 Calc
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#3/ 4.18
5. Mass of food burned
 (data
#1- data#2)
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6. Heat lost by 1 gram of food
 calories

(calc #4) / mass of food (calc #5)
7. % Error
 Calc
#6- accepted value x 100
Accepted value
Questions (1,2) and conclusion
STATES OF
MATTER
Describing the kinetic theory
as it applies to the states of
matter Dance
INTRAMOLECULAR FORCES

Attractive forces that hold particles
together in ionic or covalent bonds (intra=
within)
INTERMOLECULAR FORCES
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Forces of attraction between particles (holding
similar particles close together like in a solid or
liquid)
Intermolecular forces video
Types:
 van
der Waals
 Dispersion forces
 Dipole-dipole forces (polar molecules)
 Hydrogen bond (special type of dipole-dipole force)
GAS
Gas video
Independent particles moving in straight
lines
 Change direction with collision
 Travel randomly
 Assume shape and volume of container
 Large amount of empty space
 No attraction force
 compressible

Gas properties:
Fluidity- gas particles glide and flow past
each other
 Expansion- fill any container
 Compressibility- can decrease volume
 Diffusion- spontaneous mixing of 2 gases;
flow until evenly dispersed; flow from area
of higher concentration to lower

Graham’s Law of Diffusion- proportion
comparing diffusion rate
 Rate A =
molar mass B
Rate B
molar mass A
 Ex: HCl
and
NH3

Molar mass = 36.5
Rate NH3
=
molar mass HCl
Rate HCl
molar mass NH3
= 36.5
17.0
= 1.5
So NH3 diffuses 1.5 x faster
molar mass = 17.0
LIQUID
Properties video
 Form of matter that flows, has constant
volume and takes the shape of its
container
 Particles are in motion but slower than
gases (slip/slide motion)

Particles are held together by weak
intermolecular forces (don’t have enough
KE to break away from the attraction)
 Reduced amount of empty space

Properties of Liquids

Density and compression
 Liquids
are denser than their gas phase
 Liquids can be compressed, but an enormous
amount of pressure must be applied to reduce
the volume by just a small proportion

Fluidity
 Liquids
flow and can be diffused (yet not as well as
gases)

Viscosity
 Measure
of a liquids resistance to flow (stronger the
attractive force, the higher the viscosity)
 Viscosity increases w/a decrease in temp. (cold oil
doesn’t flow well as warm)
 Molecules w/longer chains have higher viscosity
Surface Tension
A property of liquid surfaces that causes
the surface layer to behave like a thin
elastic 'skin'.
 Molecules in a liquid have attractive forces
that hold them together. Molecules on the
surface are attracted to molecules from all
sides and below, but not from above .

 Surfactants-
(soap/detergent); compounds
that lower the surface tension of water
Capillary action
Force of adhesion between a liquid and a
solid
 Ex:

 meniscus
in graduated cylinder
 Cellulose fibers in paper towel wicking up
water
 (water suspension)
SOLID
Properties video-solid
 Particles are packed against one another
in a highly organized fashion
 Move much slower, don’t slide from place
to place but vibrate & rotate about fixed
points (straight line paths w/neighbors)

Definite pattern in
arrangement of particles
Definite shape and definite volume
 Dense and incompressible

Crystals
All true solids are crystals
 A substance in which the particles are
arranged in an orderly, geometric,
repeating pattern
 Have flat faces that meet at definite angles

Crystal Structure
3-D pattern of small units repeating over
and over
 Determined by the type of bond between
particles

Crystal Systems
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Cubic- salt
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Tetragonal

Orthorhombic
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Rhombohedral
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monoclinic
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Triclinic
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Hexagonal- water, quartz
Water Example (hexagonal)
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Water molecules are further apart in solid state than
liquid (solid-> less dense)
Water is different because of hydrogen bonding. A water
molecule is made from one oxygen atom and two
hydrogen atoms, strongly joined to each other with
covalent bonds.
Water molecules are also attracted to each other by
weaker chemical bonds (hydrogen bonds) between the
positively-charged hydrogen atoms and the negativelycharged oxygen atoms of neighboring water molecules.
As water cools below 4°C, the hydrogen bonds adjust to
hold the negatively charged oxygen atoms apart.
Amorphous Materials
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
Appears to be a solid but its particles have a
disorderly arrangement (no crystal form)
Ex:
 glass
no defined melting point, super cooled liquid
When shattered breaks at irregular angles, where a crystal
when shattered will break along the unit cell
 Butter-
 Plastic
also super cooled

METASTABLE- substance that can occur
in long-lasting amorphous form
(crystallization will eventually occur but not
in your lifetime- millions of years)
PLASMA
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Occurs when matter is heated to a very high
temperature > 5000ºC
The collisions have so much KE and are so
violent that electrons are knocked away from
the atoms
Plasma is those electrons and the left over
positive ions from the collisions
Behaves generally like a gas

PARTIAL PLASMA- Only a few of the
atoms are ionized
 Neon

signs, lightning, fluor. lights
Highly ionized plasma= (50000 K100000K)
 Stars,
sun
PHASE CHANGE
Occurs whenever a physical state of a
substance changes (at the melting and
boiling pts.)
Ex: phase change of water
(0°C- melt; 100°C boil)

On the graph following, get a leveling off at
phase change pts because all of the
energy is going into breaking the attractive
forces (intermolecular)not raising
temperature
 Physical state depending on bonding
structure (ionic, covalent)

* Requires energy
w Releases energy
GAS
*MELTING
SOLID
wFREEZING
LIQUID
Sublimation = goes directly from solid to
gas (dry ice, moth balls, Glade plug in)
 Deposition= goes from gas to solid

 frost
PHASE DIAGRAMS
p.429-430 picture
 Diff for each substance because each
substance has different boiling and
freezing point
 Variables that control the phase of a
substance are temp. and pressure

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Phase diagrams show the relationship between
temperature and pressure
Triple point = temp. and pressure at which all 3
phases can coexist
Critical Point- liquids can no longer exist