Download Gases Honors

Energy of Phases
Do Now
• What happens to the chemicals in ice as it
melts?
• What is your body’s source of energy?
• What is energy?
• What are some different forms of energy?
Energy
• Without energy, your body will not function,
technology will not exist, and there probably
would not be life on earth.
• We eat food to obtain energy, part of this
source is through the combustion of glucose.
• C6H12O6 + 6O2  6CO2 + 6H2O + energy
Energy
• Energy = the capacity to do some kind of work
– For example: moving an object, forming a new
compound, generating light
• SI unit = joule (J)
Remember changes in matter
• Physical change = only the physical properties of
matter are effected
– Ex: ice melting, water boiling
• Chemical change = a new substance is made
– Ex: chemical rxns
• Every change in matter involves a change in
energy!
– Sometimes energy is supplied (melting/evaporating)
– Sometimes energy is released (condensation/freezing)
Endothermic vs. Exothermic
– Endothermic = energy is absorbed
• Ex: melting/evaporating
• Ex: cold pack
• Ex: baking
– Exothermic = energy is released
• Ex: condensation/freezing
• Ex: hot pack
• Ex: 2 H2 + O2  2 H2O (and an explosion)
Conservation of Energy
• During any physical or chemical change, the
total quantity of energy remains constant.
• Energy can not be destroyed or created!
• Energy is being transferred between the
reaction system and its surroundings
The Law of the Conservation of Energy
• Energy is neither created nor destroyed during
a chemical reaction or a physical change. It is
just changed from one form to another.
Forms of Energy
• Energy can be chemical (KE/PE), mechanical,
light, heat, electrical, sound, etc.
– Ex: photosynthesis – the chlorophyll in the plant’s
cells (system) absorb light energy from the sun
(surroundings) – endothermic
– Ex: light stick – chemicals react inside the stick
(system) to produce light energy (surroundings) –
exothermic
Chemical energy
• Chemical energy = total energy stored in
matter as kinetic energy or potential energy
• Kinetic energy = energy of motion
• Potential energy = stored energy
Kinetic Energy (KE)
• Kinetic Energy = the energy an object has due
to its motion
– For Ex: A hockey puck gains KE through a tap, or a
slam. The puck has energy that is why goalies
wear a mask!
• The amount of kinetic energy depends on the
velocity of the object and it’s mass.
Potential Energy (PE)
• Potential Energy = energy possessed by objects due to their
position (No apparent evidence of energy is observed)
• One form of energy can be transformed into other forms of
energy.
– For example: a spring :
• Pushed down, all coiled up = _________
• Let go = _________
– For example: water dam
• Water behind a dam has PE because of its elevated position.
• As the H2O level increases, so does its PE.
• When the water is released from the dam, its PE is converted to KE as
it falls to a lower level.
• This KE increases with speed
Energy as Heat
• Most chemical reactions, physical changes of
state, and dissolving processes involve energy
changes that can be measured.
Heat
• Heat = the energy transferred between objects
that are at different temperatures
• Heat of Reaction = heat energy absorbed or
released during chemical reaction (∆H)
• Heat energy can be measured as a result of
temperature change.
• Heat is different from temperature!
Temperature
• Temperature = measurement of the average kinetic energy of the
random motion of particles in a substance
• What is the SI Unit for temperature?
– 0 K = absolute zero = temperature at which all particles have the
minimum average kinetic energies
• What do we use to measure temperature?
• In lab, our thermometers use a Celsius scale, therefore we need to
convert values to Kelvin when performing calculations.
C = K - 273
K = C + 273
Try these conversions:
• Boiling Point of water is 100C, what is the
temperature in Kelvin degrees?
• Room Temperature is 25C, what is it in Kelvin
degrees?
Endothermic vs. Exothermic
• Endothermic Reactions - usually results in a
temperature decrease (absorb energy)
• Energy is written before the arrow
2NaHCO3 + Heat  Na2CO3 + H2O + CO2
• Exothermic Reactions - result in a temperature
increase (release energy)
• Energy is written after the arrow
2H2 + O2  2H2O + Heat
Heat vs. Temperature
• Transfer of heat does not always change the
temperature
– Ex: As ice melts in a closed container, the
temperature of the ice-water mixture remains
0C, even though heat energy is being transferred.
– Ex: As water boils in a closed container and water
vaporizes, the gas-liquid mixture remains 100C,
even though heat energy is being transferred.
• The energy is being used to move molecules!
Water Phase Changes
Boiling
point
Melting
point
Temperature remains __________
during a phase change.
http://www.kentchemistry.com/links/Matter/HeatingCurve.htm
Phase Changes
S
L
G
G
L
S
Endothermic or
Exothermic
Review: Phase Changes
Type of Change
Solid
Liquid
Liquid
Gas
Solid
Gas
Gas
Liquid
Liquid
Solid
Gas
Solid
Name
Heating Curve
Label the phases/changes:
25
Temperature (C)
e
15
d
c
5
a
b
-5
Energy
Heating Curve Energy Changes
• During a phase change
•
the kinetic energy remains the same but the
potential energy increases.
• During a phase
• The kinetic energy increases and the potential
energy remains the same
http://zonalandeducation.com/mstm/physics/mechanics/energy/heatAndTem
perature/changesOfPhase/changeOfState.html
Heating Curve
Label the KE/PE:
Kinetic energy
f
d
b
c
a
Time
e
Cooling Curves Energy Changes
• During a phase change
•
the kinetic energy remains the same but the
potential energy decreases.
• During a phase
• The kinetic energy decreases and the potential
energy remains the same
Cooling Curve
Label the phases/changes:
60
A
B
40
KE
C
D
20
E
0
TIME
Table I
• Heat of Reaction =
heat of the products –
heat of the reactants
• ∆H = Hproducts – Hreactants
• + ∆H value =
endothermic (energy
absorbed)
• - ∆H value = exothermic
(energy released)
Heat of Fusion
• The amount of heat needed to MELT (solid to
liquid) 1 gram of a given substance (J/g)
Temp.
Heat of fusion
Energy
Heat of Vaporization
• The amount of heat needed to VAPORIZE
(liquid to gas) 1 gram of a given substance
(J/g)
Heat of vaporization
Temp.
Energy
Specific Heat
• Different substances are effected by heat in
different ways, resulting in different temperature
changes
• Specific heat = the amount of heat energy
required to raise the temperature of 1 g of a
substance 1C or 1 K (J/gK)
• Specific Heat of water is high and found on your
reference tables (Table B)
– Which is larger heat of fusion or heat of vaporization?
FORMULAS
q = heat
m = mass
C = specific heat
ΔT= change in temperature (final - initial)
• q = m x C x ΔT
• q = m x Heat of Fusion (Hf)
• q = m x Heat of Vaporization (Hv)
Using: q = m x C x ΔT
• q (heat) is negative during EXOTHERMIC
REACTIONS (energy released)
• q (heat) is positive during ENDOTHERMIC
REACTIONS (energy absorbed)
Using: q = m x C x ΔT
Temp
Heat Energy Added
Using: q = m x C x ΔT
http://chemistry.bd.psu.edu/jircitano/heatcurv.html
Q = mc∆T
Temp
Q = mHv
Q = mc∆T
Q = mHf
Q = mc∆T
Heat Energy Added
Using: q = m x C x ΔT
Specific heat of water = 4.18 J/g∙K
• How many kilojoules of heat are needed to
raise the temperature of 500. grams of water
from 10.0˚C to 30.0˚C?
Using: q = m x C x ΔT
Specific heat of water = 4.18 J/g∙K
• If 4.0 grams of water at 1˚C absorbs 33 joules
of heat, what will be the change in
temperature of the water?
Using: q = m x C x ΔT
Specific heat of water = 4.18 J/g∙K
• When 84 joules of heat are added to 2.0
grams of water at 15˚C, what will be the final
temperature of the water?
Using: q = m x C x ΔT
Specific heat of water = 4.18 J/g∙K
• The temperature of 50. grams of water was
raised to 50.˚C by the addition of 4,180 joules
of heat energy. What was the initial
temperature of water?
Types of Change
Physical Change
A change in a
substance that does
not involve a change in
the identity of the
substance.
Solid Liquid Gas
Chemical Change
A change in which one
or more substances are
converted into
different substances.
Can be detected
through:
•Energy changes
•Change in color
•Emission of gases
States of Matter
Solid
Liquid
Heat of ReactionAmount of heat
released or
absorbed during a
chemical reaction
Exothermic
reactions that release
Energy to their
surroundings.
-result in a temperature
increase
-Energy is written after
the arrow
Endothermic Reactions
reactions that absorb
energy from their
surroundings
Usually results in a
temperature decrease
-Energy is written
before the arrow
Gas
Forms of Energy
Phase Changes
Heating Curve
Endothermic Reaction
Cooling Curve
Exothermic Reaction
Potential Energy
Kinetic Energy
Energy possessed
by objects through
their position
type of energy that only
moving objects have.
Energy in motion.
What did you learn today?
What did you learn today?
• Heat is a transfer of energy (usually thermal energy) from a body of higher
temperature to a body of lower temperature. Thermal energy is the
energy associated with the random motion of atoms and molecules.
• Temperature is a measurement of the average kinetic energy of the
particles in a sample of material. Temperature is not a form of energy.
• The concepts of kinetic and potential energy can be used to explain
physical processes that include: fusion (melting), solidification (freezing),
vaporization (boiling, evaporation), condensation, sublimation, and
deposition.
• A physical change results in the rearrangement of existing particles in a
substance. A chemical change results in the formation of different
substances with changed properties.
• Chemical and physical changes can be exothermic or endothermic.
• The structure and arrangement of particles and their interactions
determine the physical state of a substance at a given temperature and
pressure.
Lab
• Fire and Ice lab
Review Phases
Do Now:
What can you infer from this?
States of Matter
• Remember: There are 3 main states of matter:
– Solids: have fixed positions and fixed volumes
• Solids can exist in crystalline form (hard and brittle like salt
or soft like lead)
– Liquids: have unfixed positions, but fixed volumes
• Review: cohesive vs. adhesive properties
– Cohesion = liquid particles can have attraction for each other
– Adhesion = attraction to other particles (ex. Solid)
• Surface tension = decrease their surface area to the smallest
possible by pulling particles on the surface down into the
liquid
– Gases: have unfixed positions and unfixed volumes
Changes of Matter
• If you add energy as heat to ice, the ice will melt
and change from solid to liquid water.
– Melting = solid to liquid
– Freezing = liquid to solid
• If you add more energy to the liquid water, it will
change from liquid to gas
– Evaporation = liquid to gas
– Condensation = gas to liquid
– Sublimation = solid to gas (ex. Ice on windshield)
– Deposition = gas to solid (ex. Frost from water vapor)
Temperature and Energy
• All matter has energy related to the random
motion of particles.
• This energy increases as temperature
increases.
• Remember: Temperature= measure of the
average kinetic energy of particles
Temperature vs. Heat
• Heat = the energy transferred between
objects that are at different temperatures
• Heat or energy can be measured in Joules (J).
• Temperature = a measure of the average
kinetic energy of the particles in an object
Points
• Boiling point = temperature and pressure at
which bubble of vapor rise to the surface and
the temperature of the liquid remains
constant.
• Melting point = the temperature and pressure
at which a solid becomes a liquid
• Freezing point = the temperature at which a
liquid substances freezes
– Melting point and freezing point are the same for
pure substances!
Endothermic vs. Exothermic
• Remember:
– Endothermic = absorbs energy
• Which state changes are endothermic?
• Evaporation, melting, and sublimation
– Exothermic = release energy
• Which state changes are exothermic?
• Condensation, freezing, and deposition
Phase Diagrams
• Every phase change occurs at a specific
combination of temperature and pressure.
• The direction of the phase change depends on
whether heat energy is being added or
removed.
• If no heat energy is added or removed, the
two phases remain in a state of dynamic
equilibrium at that temperature and pressure.
General Phase Diagram
• Each line represents the
pressure-temperature
combination at which
two phases are in
dynamic equilibrium.
• At any combination of
pressure and
temperature that
doesn’t fall on one of
the lines, only one
phase can exist.
General Phase Diagram
• Point A = triple point,
the only point where all
three phases can
coexist in dynamic
equilibrium.
• Point B = critical point,
above this critical
pressure and critical
temperature only the
gas phase can exist.
The Phase Diagram for Water: Triple point
(A), 0.0098 oC, 4.58 torr
• Normal melting point (B), 0 oC, 1 atm
• Normal boiling point (C), 100 oC, 1 atm
• Critical point (D), 374.4 oC, 217.7 atm
The Phase Diagram for Carbon Dioxide:
Triple point (X), −56.4 oC, 5.11 atm
• Normal sublimation point (Y), −78.5 oC,
1 atm
• Critical point (Z), 31.1 oC, 73.0 atm
• Notice that CO2 does not exist in the
liquid phase at 1 atm of pressure.
Gases
Do Now
• What are some unique properties among
gases?
Nitrogen dioxide gas
Hot air balloon
Helium balloon
Chlorine gas
Characteristics of Gases
• Gases are considered fluids, because the distance
between the particles is great enough for the
substance to flow.
• Gases have low density, because of the relatively
large distances between gas particles and long
distance before particles collide with each other.
• Gases are highly compressible, because gas particles
can be pushed closer together.
Characteristics of Gases
• Gases completely fill a container, because gas
particles do not have fixed positions or
volumes since the particles are constantly
moving at high speeds and collide with each
other expanding to fill the entire volume
available.
Properties of Gases
• No definite shape and no definite volume
• Even though most gases are “invisible,” they are
made up of matter and atoms that take up space
(NO2 gas is brown and Cl2 gas is green).
• Gases will completely and evenly fill their container.
• Gases have mass, although the density of a gas is
much less than the density of a liquid or solid.
• Gases high compressibility (think of a syringe)
• Gases will diffuse (spread out) evenly in all directions
http://www.biosci.ohiou.edu/introbioslab/Bios170/d
iffusion/Diffusion.html
Gas Pressure
• As gas molecules are pulled toward the
surface of Earth, they collide with each other
more often.
• Collisions of gas molecules cause pressure.
Gas Pressure
• What pressures do you think are involved in
allowing a balloon to keep its shape?
Atmospheric Pressure
• Gases in the atmosphere exert a pressure with
everything it is in contact with.
• It counter balances the inside pressure.
• What happens to pressure as you go further
out into the atmosphere?
It decreases!
Measuring Pressure
• Pressure = the amount of force exerted per unit area
of surface
• Pressure is commonly defined as “force divided by
area”
– Force is measured in newtons.
• Pressure is measured in pascal, Pa = the force of one
newton applied over an area of one meter squared
1 Pa = 1 N / 1 m2
• Pressure can be measured with a barometer
Measuring Pressure
• At sea level, the atmosphere keeps the
mercury in a barometer at an average height
of 760mm, which is 1 atmosphere of pressure.
• One millimeter of mercury is also called a torr.
STP
• Scientists have specified a set of standard
conditions called STANDARD TEMPERATURE
and PRESSURE or STP, which is equal to 0°C
and 1 atm
• At STP, the atmospheric pressure is balanced
by exactly 1 atmosphere (atm) or 760
millimeters of Mercury (mmHg)
Barometers
An aneroid barometer is
an alternative to a
mercurial barometer; it
is easier to read and
transport.
Mercury Barometer
Atmospheric pressure is
typically measured in
inches of mercury (in. Hg.)
by a mercurial barometer.
Aneroid barometer
Since weather stations are located around the globe, all local barometric
pressure readings are converted to a sea level pressure to provide a standard
for records and reports. To achieve this, each station converts its barometric
pressure by adding approximately 1 inch of mercury for every 1,000 feet of
elevation gain. For example, a station at 5,000 feet above sea level, with a
reading of 24.92 inches of mercury, reports a sea level pressure reading of
29.92 inches.
Manometers
• Closed tube manometers are used when the pressure of the
enclosed gas is less than atmospheric pressure.
• Open tube manometers are used when the pressure of the
enclosed gas is at or near atmospheric pressure
Standard Temperature and
Pressure (STP)
• Due to the difference in pressure and
temperature in different areas standards are
used.
• Standard Temperature = 0 °C or 273 K
• Standard Pressure =1 atm or 101.3 Kpa
or 760 torr or 760 mmHg
Summary
• Units of pressure = atm, KPa, torr, mmHg
• STP = Standard Temperature and Pressure
• 760 torr = 1 atm
• 101.3 KPa = 1 atm
Convert the following pressures
1. 1.3 atm = ___________ KPa
1.3 atm x 101.3 KPa =
1 atm
2. 124.6 KPa = ____________ atm
124.6 KPa x 1 atm =
101.3 KPa
3. 0.98 atm = ___________ KPa
0.98 atm x 101.3 KPa =
1 atm
Convert the following Pressures:
1. 1.3 atm = _____ kPa
2. 124.6 kPa = _____ atm
3. 0.98 atm = ______ kPa
22
Do Now
• Why are gases considered fluids?
• Describe the particles in gas?
• What do you think the Kinetic Molecular
Theory of Matter (KMT) explains?
Kinetic Molecular Theory (KMT)
• This theory is used to explain the properties of
matter (solids, liquids, gases).
• It is based on the idea that particles are in
constant, random, straight line motion.
• The theory can be used to explain the
behaviors and the properties of gas molecules.
Ideal vs. Real gas
• The theory describes the behavior of an ideal
gas.
• What does ideal mean?
• IDEAL GAS = an imaginary gas that obeys all
the assumptions of the kinetic molecular
theory (KMT)
• REAL GAS = an actual gas that does not
behave completely according to the kinetic
molecular theory (KMT)
Ideal vs. Real gas
• There is no such thing as an ideal gas because
real gases have mass and have attractions
between them (intermolecular forces).
• Real gases have:
– Mass and volume
– Forces of attraction between them
(intermolecular forces)
• IMF increase as the distance between them decreases
• The higher the molecular mass, the stronger the IMF
http://antoine.frostburg.edu/chem/senese/101
/liquids/faq/h-bonding-vs-london-forces.shtml
Are gases attracted to each other?
Kinetic Molecular Theory (KMT)
• 1. Gases consist of large numbers of tiny particles
usually molecules (ex. CO2) or atoms (Ex. Ne), that
are far apart relative to their size.
– Gases have an insignificant volume (almost zero)
and the particles are separated from each other
by large distances.
• 2. Collisions between gas particles and between
particles and the container walls are elastic collisions.
– Elastic collisions = no net loss of energy to the
surrounds.
Elastic vs. Inelastic collisions?
http://www.chm.davidson.edu/ChemistryApplets/KineticMolecularTheory/BasicConc
epts.html
Kinetic Molecular Theory (KMT)
• 3. Gas particles are in constant, rapid, random
motion. They travel in straight line paths and
fill their containers. They possess kinetic
energy, which is energy of motion.
Analogy: pool balls when hit travel
In straight line paths.
Kinetic Molecular Theory (KMT)
• 4. Energy may be transferred between
colliding particles. There is no net loss of
energy as the result of these collisions. These
collisions are perfectly elastic.
• 5. There are no forces of attraction or
repulsion between gas particles.
Kinetic Molecular Theory (KMT)
• 6. The average kinetic energy of the gas
particles depends on the temperature of the
gas.
– Lower temperature = gas molecules have lower
kinetic energy
– Higher temperature = gas molecules have higher
kinetic energy
http://www.epa.gov/apti/bces/module1/kinetics/animation/kani1/kani104.htm
Review
• What does the KMT explain?
• Describe parts of the KMT.
Review KMT
1.
2.
3.
4.
5.
6.
A gas is composed of particles, usually molecules (ex. CO2) or atoms (Ex.
Ne), that are far apart relative to their size. (They have insignificant
volume and are separated by large distances).
Collisions between gas particles and between particles and the container
walls are elastic collisions.
Elastic collisions- no net loss of energy
Energy may be transferred between colliding particles. There is no net
loss of energy as the result of these collisions. These collisions are
perfectly elastic.
Gas particles are in constant, rapid, random motion. They travel in
straight line paths and fill their containers. They possess kinetic energy,
which is energy of motion.
There are no forces of attraction or repulsion between gas particles.
The average kinetic energy of the gas particles depends on the
temperature of the gas.
What properties of gases can be
explained in terms of the KMT?
1.
2.
3.
4.
5.
EXPANSION- Gases have no definite shape and no definite volume. They
expand to fill their container. This is because gases have insignificant
attractive forces and are in constant motion.
FLUIDITY- Gas particles glide past each other, this is because gases have
insignificant attractive forces.
LOW DENSITY- Gases have low mass per unit volume. This is because gas
particles are so far apart from each other compared to their size.
COMPRESSIBILITY- The volume of a gas can be compressed into a much
smaller volume. This is because gas particles are so far apart from each
other.
DIFFUSION and EFFUSION
– DIFFUSION= the spontaneous mixing of gas particles even when the gas is not stirred.
– EFFUSION= the passing of gas particles through a tiny opening. This is because gas
particles are in constant motion and there are insignificant attractions between them.
http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/animations/Effusion2.html
http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/gaslaw/micro_effusio
n.html
Diffusion
DO NOW:
• How do real gases differ from ideal gases?
Real vs. Ideal gas
• Real gases behave most like ideal gases under
the following conditions:
– Low pressure
– High temperature
Real vs. Ideal gas
• Why do real gases behave like ideal gases
under high temperatures and low pressure?
• Under these conditions, gas particles are
furthest apart from each other and have very
weak forces of attraction (IMF) between them.
Real vs. Ideal gas
• What makes real gases deviate from ideal
gases?
• If a molecule is polar = assymetrical
• Weighs a lot
• Strong IMF
• Hydrogen bonding between molecules
• Dipole-dipole attractions between molecules
Real vs. Ideal gas
• What properties do some real gases have that
allow them to behave like ideal gases?
• If a molecule is nonpolar = have symetry
• Does not weigh a lot
• Weak IMF
• Noble gases
REAL GASES
IDEAL GASES
•_____follow the
Kinetic Molecular
Theory
•_____follow the
Kinetic Molecular
Theory
•_____have volume
•_____have volume
•_____ have forces of
attraction between
them
•_____ have forces of
attraction between
them
•Exist under ___ pressure
and ____ temperature
•Exist under ___ pressure
and ____ temperature
REAL GASES
IDEAL GASES
•_____follow the
Kinetic Molecular
Theory
•_____follow the
Kinetic Molecular
Theory
•_____have volume
•_____have volume
•_____ have forces of
attraction between
them
•_____ have forces of
attraction between
them
•Exist under ___ pressure
and ____ temperature
•Exist under ___ pressure
and ____ temperature
14
Which of the following behave most
like ideal gases and which deviate the
most?
1.
2.
3.
4.
5.
6.
NH3 (g) _____
O2 (g) ______
He ______
H2O ______
N2 ______
Ne ______
Which of the following behave most like
ideal gases and which deviate the most?
1.
2.
3.
4.
5.
6.
NH3 (g) _____
O2 (g) ______
He ______
H2O ______
N2 ______
Ne ______
15
Do Now
• What happens to the pressure of a gas if
volume increase?
• What happens to the volume of a gas if
pressure increases?
Gas Laws
• Gas laws are simple mathematical relationships between the
volume, temperature and pressure of a gas.
• 1. VOLUME (V) – amount of space occupied by a sample of
matter. Units = L, mL, cm3
– 1 mole = 22.4 L
– 1 mL = 1 cm3
• 2. TEMPERATURE (T) - average kinetic energy of the particles
in a sample of matter. Measured with a thermometer. ALWAYS
USE KELVIN!
– K = C + 273
– 0 °C =
-25 ° C=
100 ° C =
• 3. PRESSURE (P) – Force per unit area on a surface. Force
produced by the collision of gas particles with the container
walls. Units = atm
Boyle’s Law
• Relates pressure and volume at a constant
temperature.
• At a constant temperature, pressure varies
inversely with volume
• Formula: P1V1 =P2V2
P x V = k (constant)
this is a closed
• 1 = initial Since
container, the number of
does not change
• 2 = final particles
from the initial and final states
Boyle’s Law Example
• A sample of oxygen gas has a volume of 150
mL when it’s pressure is 0.947 atm. What will
the volume be at a pressure of 0.987 atm if
the temperature remains constant?
Boyle’s Law Example
• A gas has a pressure of 1.26 atm and occupies
a volume of 7.40 L. If the gas is compressed to
a volume of 2.93 L, what will its pressure be,
assuming constant temperature?
Pressure
Volume k (constant)
1 atm
X
10L
=
2 atm
X
5L
=
4 atm
X
2.5 L
=
8atm
X
1.25L
=
Pressure
Volume k (constant)
1 atm
X
10L
=
2 atm
X
5L
=
4 atm
X
2.5 L
=
8atm
X
1.25L
=
25
Volume
(L)
Pressure (atm)
26
Review Boyle’s Law
• At constant temperature, if there’s an increase
in pressure, there’s a decreases in volume.
http://www.epa.gov/apti/bces/module1/kinetics/animation/kani3/kani3a.htm
http://academic.pgcc.edu/psc/chm101/ideal_gas/animatation_1.htm
http://www.edumedia-sciences.com/a257_l2-kinetic-pressure.html
Do Now
• What happens to the volume of a gas if
temperature increase?
• What happens to the volume of a gas if
temperature decrease?
Charles’ Law
• Relates volume and temperature at a constant
pressure.
• At a constant pressure, volume varies directly
with temperature.
• Formula: V1 = V2
T1 T2
V = k (constant)
T
Since this is a closed
Temperature must be in
• 1 = initial container, the number of
Kelvin!!!
particles does not change
K = C + 273
• 2 = final from the initial and final states
Charles’ Law Example
• A sample of neon gas occupies a volume of
752 mL at 25°C. What volume will the gas
occupy at 50°C if the pressure remains
constant?
Charles’ Law Example
• A helium balloon has a volume of 2.75L at
20°C. The volume of the balloon decreases to
2.46L after it is placed outside on a cold day.
What is the outside temperature in K? In
degrees Celsius?
Volume
Temperature
constant
100L
50K
=
200L
100K
=
400L
200K
=
800L
400K
=
Volume
Temperature
constant
100L
50K
=
200L
100K
=
400L
200K
=
800L
400K
=
32
Volume
(L)
Temperature (K)
33
Review Charles’ Law
• At constant pressure, if there’s an increase in
volume, there’s an increase in temperature.
http://www.epa.gov/apti/bces/module1/kinetics/animation/kani2/kani2c.htm
http://academic.pgcc.edu/psc/chm101/ideal_gas/animatation_2.htm
Do Now
• Now that you know the relationship between
pressure and volume, and volume and
temperature, do you think there is a
relationship between pressure and
temperature?
Gay-Lussac's law
•
P=k
T
P1 = P2
T1 T2
– P is the pressure of the gas
– T is the temperature of the gas (Kelvin).
– k is a constant
• Remember: temperature is a measure of the
average kinetic energy of a substance; as the
kinetic energy of a gas increases, its particles
collide with the container walls more rapidly,
thereby exerting increased pressure.
Gay-Lussac’s Examples
• A gas has a pressure at 2.0 atm at 18°C. What
is the new pressure when the temperature is
62°C? (V and n constant)
Calculation with Gay-Lussac’s
Law
1. A gas has a pressure at 2.0 atm at 18°C. What
is the new pressure when the temperature is
62°C? (V and n constant)
2. A gas has a pressure of 645 torr at 128°C. What is
the temperature in Celsius if the pressure increases
to 1.50 atm (n and V remain constant)?
3.The gas in an aerosol can is at a pressure of 3.00atm
at 25oC. Directions on the can warn the user not to
keep the can in a place where the temperature
exceeds 52oC. What would the gas pressure in the
can be at 52oC?
Gay-Lussac’s Review
• Think about tires bursting when they get too
hot!
Do Now:
• Explain the relationship between pressure and
temperature.
• Explain the relationship between volume and
temperature.
• Explain the relationship between pressure and
volume.
Do Now:
• Explain the relationship between pressure and
temperature.
• Explain the relationship between volume and
temperature.
• Explain the relationship between pressure and
volume.
Combined Gas Law
• Relates pressure, volume and temperature of
a fixed amount of gas.
Since this is a closed
• Formula: P1 x V1 = P2 x V2 container, the number of
particles does not change
T1
T2 from the initial and final states
• STP = 273 K
= 1 atm or 101.3 KPa
Combined Gas Law Example
• A helium filled balloon has a volume of 50.0L
at 25 °C and 1.08atm. What volume will it
have at 0.855 atm and 10 °C?
Combined Gas Law Example
• A helium filled balloon has a volume of 50.0L
at 25 °C and 1.08atm. What volume will it
have at 0.855 atm and 10 °C?
37
Combined Gas Law Example
• A 700 mL gas sample at STP is compressed to
a volume of 200mL, and the temperature is
increased to 30.0 °C. What is the new pressure
of the gas in kiloPascals?
• A 700 mL gas sample at STP is compressed to
a volume of 200mL, and the temperature is
increased to 30.0 °C. What is the new pressure
of the gas in kiloPascals?
38
Review Combined Gas Law
• http://www.epa.gov/apti/bces/module1/kinet
ics/animation/kani2/kani2a.htm
The Ideal Gas Law
• All of the gas laws can be combined into a
single, universal relationship.
Pressure x Volume
= Constant
Number of particles x Temperature
• or:
PV=nRT
*R = the molar gas constant that depends on the units of
pressure and volume (0.0821 L atm/mol K or 8.31 J/mol K)
Practice Problem
• How many moles of an ideally behaving gas
occupy 400 L at 0.821 atm and 200 K?
Practice Problem
• How many moles of an ideally behaving gas
occupy 400 L at 0.821 atm and 200 K?
PV=nRT
P = 0.821 atm
V = 400 L
n=?
R = 0.0821 L atm/mol K
T = 200 K
Practice Problem
• How many moles of an ideally behaving gas
occupy 400 L at 0.821 atm and 200 K?
PV=nRT
(0.821 atm)(400L) = n(0.0821)(200K)
(0.0821)(200K)
(0.0821)(200K)
n = 20 mol
Practice Problem
• What volume will an ideally behaving gas
containing 1 mole of particles occupy at STP?
PV=nRT
(1atm) V = (1mol)(0.0821L atm/mol K)(273K)
(1atm)
(1 atm)
V = 22.4 L
Do Now
• What does partial mean?
• Define pressure.
• What do you think partial pressure is?
Dalton’s Law of Partial Pressures
• John Dalton showed that in a mixture of
gases, each gas exerts a certain pressure as if
it were alone with no other gases mixed with
it. The pressure of each gas in a mixture is
called partial pressure.
• The total pressure of a mixture of gases is the
sum of the partial pressures of the gases.
• Formula: Ptotal = Pa + Pb + Pc + ….
Pressure
Pressure of individual
gases
Dalton’s Law of Partial Pressures
Scuba Diving
• When a scuba diver dives, the
increased pressure causes N2(g)
to dissolve in the blood.
• If a diver rises too fast, the
dissolved N2 will form bubbles
in the blood, a dangerous and
painful condition called "the
bends".
• Helium, which does not
dissolve in the blood, is mixed
with O2 to prepare breathing
mixtures for deep descents.
Learning Check
• A scuba tank contains O2
with a pressure of 0.450atm
and He at 855mmHg. What
is the total pressure in
mmHg in the tank?
Gases We Breathe
• The air we breathe is a gas
mixture.
• It contains mostly N2 and O2
and small amounts of other
gases.
• Each gas exerts a pressure
proportional to the
percentage or number of
molecules.
Dalton’s Law of Partial Pressures
Example
• What is the total pressure exerted by a
mixture of gases where the partial pressure of
nitrogen gas is 100 atm, oxygen gas is 300
atm, and carbon dioxide is 150 atm?
Dalton’s Law of Partial Pressures
Example
• What is the total pressure exerted by a
mixture of gases where the partial pressure of
nitrogen gas is 100 atm, oxygen gas is 300
atm, and carbon dioxide is 150 atm?
Dalton’s Law of Partial Pressures
Example
• What is the partial pressure of nitrogen in a
mixture of nitrogen and oxygen where the
total pressure is 760 torr and the mixture is
70% nitrogen?
Dalton’s Law of Partial Pressures
Example
• What is the partial pressure of nitrogen in a
mixture of nitrogen and oxygen where the
total pressure is 760 torr and the mixture is
70% nitrogen?
41
Gas Law
What kind of
relationship?
Boyle’s
Law
Charles’s
Law
Gay-Lussac’s
Combined Gas
Law
N/A
Dalton’s Law of
Partial
Pressures
N/A
Ideal Gas Law
N/A
Formula
Gas Law
What kind of
relationship?
Boyle’s
Law
Charles’s
Law
Combined
Gas Law
N/A
Dalton’s Law of
Partial
Pressures
N/A
Ideal Gas
Law
N/A
Formula
Do Now
• Recall the meaning of a mole.
• How many particles are in a mole?
• What is the volume occupied by a mole of
gas?
Gases and Moles
• Avogadro’s Law: equal volumes of gases at the
same temperatures and pressure contain
equal numbers of molecules
– One mole of any gas contains 6.02 x 1023 particles
(1 mol = 6.02 x 1023 particles)
– One mole of any gas occupies 22.4 L
(1 mol = 22.4 L)
– 22.4 L of any gas contains 6.02 x 1023 particles
(22.4 L = 6.02 x 1023 particles)
So, if you have two gases of the same volume, they will have the same number of particles!
Gases and Moles
• 22.4 L = 6.02 x 1023 particles
• So, all of the following have the same number
of particles, because they represent the same
volume of different gases at the same
temperature and pressure.
Moles of Gas Example
• A chemical reaction produces 0.0680 mol of
oxygen gas. What volume in liters is occupied
by this gas sample at STP?
Moles of Gas Example
• A chemical reaction produces 0.0680 mol of
oxygen gas. What volume in liters is occupied
by this gas sample at STP?
Moles of Gas Example
• How many moles of methane gas are there in
135 L of methane gas?
Moles of Gas Example
• How many moles of methane gas are there in
135 L of methane gas?
44
Moles of Gas Example
• A chemical reaction produces 98.0 mL of
sulfur dioxide gas, SO2, at STP. What was the
mass (in grams) of the gas produced?
Moles of Gas Example
• A chemical reaction produces 98.0 mL of
sulfur dioxide gas, SO2, at STP. What was the
mass (in grams) of the gas produced?
45
Density at STP
• Since the mole is associated with the molar
mass and molar volume at STP, we can
calculate gas density for ideal gases at STP.
dSTP = m
Vm
Vm = 22.4 L/mol at STP
Density Practice Problem
• Calculate the density of CO gas at STP.
dSTP = m molar mass CO = 28.01g/mol
Vm
dSTP = 28.01 g/mol
22.4 L/mol
dSTP = 1.25 g/L at STP
Do Now
• Balance the following equation:
CO (g) +
O2 (g)  CO2 (g)
• What do the coefficients represent?
Stoichiometry of Gases
• For gaseous reactants and products in
chemical equations, the coefficients not only
represent molar ratios, but also volume ratios.
2 liters,
1 liter,
2 liters of each gas
2CO (g) + O2 (g) →
2CO2 (g)
Volume Ratio
• Write volume ratios for the equation:
N2 + 3H2 → 2NH3
1 L N2
2 L NH3
1 L N2
3 L H2
Volume-Volume Calculations
• Propane gas, C3H8, combusts according to the
following equation:
C3H8 (g) + 5O2 (g) → 3CO2 (g) + 4H2O (g)
• What will be the volume, in liters, of oxygen
required for the complete combustion of
0.35L of propane?
Volume-Volume Calculations
• Propane gas, C3H8, combusts according to the
following equation:
C3H8 (g) + 5O2 (g) → 3CO2 (g) + 4H2O (g)
• What will be the volume, in liters, of oxygen
required for the complete combustion of
0.35L of propane?
0.35 L propane x 5 L oxygen = 1.75 L O2
1 L propane
Volume-Volume Calculations
• What will be the volume of carbon dioxide
produced in the reaction again combusting
0.35L of propane?
C3H8 (g) + 5O2 (g) → 3CO2 (g) + 4H2O (g)
Volume-Volume Calculations
• What will be the volume of carbon dioxide
produced in the reaction again combusting
0.35L of propane?
Volume-Volume Calculations
• What volume of hydrogen gas is needed to
react completely with 4.55L of oxygen gas to
produce water vapor?
O2 (g) + 2H2 (g) → 2H2O (g)
Volume-Volume Calculations
• What volume of hydrogen gas is needed to
react completely with 4.55L of oxygen gas to
produce water vapor?
Volume-Volume Calculations
• What volume of oxygen gas is needed to react
completely with 0.626L of carbon monoxide
gas, CO, to form gaseous carbon dioxide?
Volume-Volume Calculations
• What volume of oxygen gas is needed to react
completely with 0.626L of carbon monoxide
gas, CO, to form gaseous carbon dioxide?
2CO (g) + O2 (g) → 2CO2 (g)
0.626 L CO x 1 L O2 = 0.313 L O2
2 L CO
Mass-Volume Calculation
• How many liters of CO2(g) are produced at STP
with 64 grams of O2(g)?
2CO(g) + O2(g)  2CO2(g)
mass O2  mol O2  mol CO2  L CO2
64 g O2 x 1mol O2 x 2 mol CO2 x 22.4L CO2
32g O2
1mol O2 1mol CO2
= 89.6 L CO2
Mass-Volume Calculation
• How many liters of O2(g) are produced at STP
with 28 grams of CO(g)?
2CO(g) + O2(g)  2CO2(g)
mass CO  mol CO  mol O2  L O2
28g CO x 1mol CO x 1mol O2 x 22.4L O2
28g CO 2mol CO 1mol O2
= 11.2 L O2
Not “Ideal” Situations
• Not at STP
Gas Stoichiometry & Ideal Gas Law
• Moles  Liters of a Gas:
– STP - use 22.4 L/mol
– Non-STP - use ideal gas law
• Non-STP
– Given liters of gas
• start with ideal gas law
– Looking for liters of gas?
• start with stoichiometry conversion
Gas Stoichiometry Problem
• What volume of CO2 forms from 5.25 g of
CaCO3 at 103 kPa & 25ºC?
CaCO3
5.25 g

CaO
+
Looking for liters: Start with
stoich and calculate moles of CO2.
g
mol
CaCO3 CaCO3
mol
CO2
g
mol
CaCO3 CaCO3
=
CO2
?L
nonSTP
mol CO2
Plug this into the Ideal
Gas Law to find liters.
Gas Stoichiometry Problem
• What volume of CO2 forms from 5.25 g of
CaCO3 at 103 kPa & 25ºC?
GIVEN:
WORK:
P=
V=
n=
T=
R=
Gas Stoichiometry Problem
• What volume of CO2 forms from 5.25 g of
CaCO3 at 103 kPa & 25ºC?
GIVEN:
WORK:
P=
V=
n=
T=
R=
Gas Stoichiometry Problem
• How many grams of Al2O3 are formed from 15.0
L of O2 at 97.3 kPa & 21°C?
4 Al
GIVEN:
P=
V=
n=
T=
R=
+
3 O2
15.0 L
non-STP

WORK:
2 Al2O3
?g
Given liters: Start
with Ideal Gas Law and
calculate moles of O2.
NEXT 
Gas Stoichiometry Problem
• How many grams of Al2O3 are formed
from 15.0 L of O2 at 97.3 kPa & 21°C?
4 Al
+
Use stoich to convert
moles of O2 to grams
Al2O3.
mol O2
3 O2
15.0L
non-STP

mol
Al2O3
g
Al2O3
mol O2
mol
Al2O3
2 Al2O3
?g
= g Al2O3
• Tungsten, W, a metal used in light-bulb filaments, is produced
industrially by the reaction of tungsten oxide with hydrogen.
WO3(s) + 3H2(g) → W(s) + 3H2O(l)
How many liters of hydrogen gas at 35oC and 0.980atm are needed
to react completely with 875g of tungsten oxide?
• Tungsten, W, a metal used in light-bulb filaments, is produced
industrially by the reaction of tungsten oxide with hydrogen.
WO3(s) + 3H2(g) → W(s) + 3H2O(l)
How many liters of hydrogen gas at 35oC and 0.980atm are needed
to react completely with 875g of tungsten oxide?
THE END
• http://www.chm.davidson.edu/ChemistryAppl
ets/KineticMolecularTheory/Pressure.html
Avogadro’s Law
• The volume of a gas at constant temperature and
pressure is directly proportional to the number of
moles of the gas.
• Mathematically, this means
V = kn
Gases are one of the most pervasive aspects
of our environment on the Earth.
We continually exist with constant exposure
to gases of all forms.
The steam formed in the air during a hot
shower is a gas.
The Helium used to fill a birthday balloon is
a gas.
The oxygen in the air is an essential gas for
life.
A windy day or a still day is a result of the
difference in pressure of gases in two
different locations.
A fresh breeze on a mountain peak is a study
in basic gas laws.
Deviations from Ideal Behavior
The assumptions made in the kinetic-molecular
model break down at high pressure and/or low
temperature.
Real Gases
Real gases animation at different temperatures
• http://www.mpcfaculty.net/mark_bishop/KM
T.htm
Why do real gases behave like ideal gases
under high temperatures and low
pressure real gases:
• Under these conditions gas particles will
be furthest apart from each other and
therefore will have the weakest
intermolecular forces.
13
What makes real gases deviate from
ideal gases?
• The greater the intermolecular forces
between the gas molecules, the more its
properties will deviate from those of an
ideal gas.
• The heavier a gas is the more it
deviates from an ideal gas.
Properties of real gases that allow them
to behave like ideal gases:
Gases most like ideal gases are nonpolar,
small have weak intermolecular forces.
Noble gases behave like ideal gases because
they are nonpolar.
Nonpolar molecules have symmetry.
16