Editor: Julia McNamara

Group 1

The Nature of Gases

Coeditor: Marybeth Nametz

Group: Grayce Rose, Erika Paiva, and Marybeth Nametz

Kinetic Theory and a Model for Gases p.385
Grayce Rose

Kinetic Theory and a Model for Gases
  • The word kinetic refers to motion
  • The energy an object has because of its motion is called kinetic energy
  • According to the kinetic theory:
-all matter consists of tiny particles that are in constant motion (the particles in a gas are usually molecules or atoms)
  • The particles in a gas are considered to be small, hard spheres with an insignificant volume
Within a gas, particles are farther apart as opposed to particles in a liquid or solid. Between the particles is empty space.

external image states.gif Grayce Rose
  • the motion of the particles in a gas is rapid, constant, and random
Because of this, gases fill their containers despite the shape and volume of their containers. Uncontained gas can spread out into space withought limit Grayce Rose
  • All collisions between particles in a gas are perfectly elastic
During an elastic collision, kinetic enrgy is transferred withought loss from one particle to another, and the total kinetic nergy remains constant

external image boing.gif

A totally elastic collision is one in which the Kinetic Energy is conserved.

by Grayce Rose

Gas Pressure p. 386-387
Erikia Paiva

Summary of the Section with Key terms defined

- Gas Pressure:
is the pressure exerted by a gas. "Gas molecules inside a volume (e.g. a balloon) are constantly moving around freely. During this molecular motion they frequently collide with each other and with the surface of any enclosure there may be."
- gas pressure results from the collision of particles in a gas with an object.

-is an empty space that does not contain pressure and particles. Since there is no particles collisions can not take place and there is no pressure.

Atmospheric pressure:
- Pressure caused by the weight of the atmosphere. At sea level it has a mean value of one atmosphere but reduces with increasing altitude.
- Atmospheric pressure is caused by molecules colliding in air with objects. It decreases as you climb becuase of the density of Earth's atmosphere.
- depends on the weather and the altitude.

- is a tool that is used to measure atmospheric pressure.
-"It can measure the pressure exerted by the atmosphere by using water, air, or mercury. Pressure tendency can forecast short term changes in the weather. Numerous measurements of air pressure are used within surface weather analysis to help find surface troughs, high pressure systems, and frontal boundaries."

- is the SI unit for pressure.
-"It is a measure of force per unit area, defined as one newton per square metre. In everyday life, the pascal is perhaps best known from meteorological barometric pressure reports."

Standard atmosphere:
-(atm) is the pressure required to support 760mm of mercury in a mercury barometer at 25 degrees C.

1 atm = 760 mm Hg = 101.3 k Pa

Converting between Units of Pressure

A pressure gauge records a pressure of 450 kPa. What is the measurement expressed in atmospheres and millimeters of mercury?

1. Analyze: List the knowns and the unknowns

Pressure = 450 kPa
1 atm = 101.3 kPa
1 atm= 760 mm Hg

pressure= ? atm
Pressure= ? mm Hg

For converting kPa ------atm, the conversion factor is 1 atm/101.3

For converting kPA------ mm Hg, the conversion factor is 760 mm Hg/101.3 kPa

2. Calculate
450 kPa X 1 atm/101.3 kPa = 4.4 atm

450 kPa X 760 mm Hg/101.3 kPa = 3400 mm Hg = 3.4X10 ^3 mm Hg

3. Evaluate: Does this make sense?

external image images?q=tbn:ANd9GcSBe4c14EN4ZcUdGcFGGNT6yJFhGebbPjIi2vX7N2fdZrINRZDahYxk4Q
By Erika Paiva

external image pressure_altitude.jpg
Erika Paiva

Kinetic Energy and Temperature p.388-389
Marybeth Nametz

An increase in kinetic energy results in an increase in temperature.

Average Kinetic Energy

  • Average kinetic energy is used when discussing the kinetic energy of a collection of particles in a substance
  • At any given temperature the particles of all substances (despite any physical state they may be in) will have the same average kinetic energy
  • There is a relationship between the average kinetic enrgy of the particles in a substance and the substance's temperature:
~ An increase in the average kinetic energy of the particles causes the temperature of a substance to rise
~Therefore as a substance cools, the particles will move slower and their average kinetic enrgy will decline
  • At absolute zero (0 K or -273.15 oC) particles in motion will theoretically cease.
~There is no temperature below absolute zero.
~Absolute zero has never been produced in a laboratory

external image boltz1.gif(This shows the distribution of molecular kinetic energy)
Marybeth Nametz-

Average Kinetic Energy and Kelvin Temperature
  • The Kelvin temperature scale reflects the relationship between temperature and average kinetic energy
  • The Kelvin temperature of a substance is directly proportaional to the average kinetic energy of the particles of the substance
  • The effects of temperature on particle motion in liquids and solids are mor complex than in gases

Group 2

The Nature of Liquids

Coeditor: Kim Kogut

Group: Kim Kogut and David O'Brien

A Model for Liquid, Evaporation, Vapor Pressure p.390-392
David O'Brien

Model for Liquids
- Substances that can flow are fluids
- There is no attraction between particles in a gas
- There is an attraction between particles in a liquid
o Attraction causes less space between particles
o Therefore liquids are more dense than gas
- Liquids and solids are condensed states of matter
- Liquid to gas/vapor = vaporation
- When not boiling the process is called evaporation
- During evaporation molecules have to have a certain minimum amount of kinetic energy in order to escape from the liquids surface
- Heat adds more kinetic energy which causes a faster evaporation
- This heat and kinetic energy breaks the attraction between particles
- The more kinetic energy the faster the evaporation
Vapor Pressure
- Vapor pressure is a measure of force exerted by gas above a liquid
- In a closed container
- Liquid evaporates into vapor(gas) which is then condensed back into a liquid
- In a system at constant vapor pressure, a dynamic equilibrium exists between the vapor and the liquid
o Rate of evaporation of liquid = rate of vapor condensation
Temperature Change:
- An increase in temperature = an increase in vapor pressure
- This creates more kinetic energy
- Manometers measure vapor pressure

Boiling Points p.393-395
Kim Kogut

Boiling Points
-Heating liquids make their particles move faster and overcome forces that keep them in a normal liquid state.
-As they absorb more heat, the kinetic energy increases.
-When heated high enough, the particles will have enough energy to vaporize. This is when boiling happens.
-The boiling point is the temperature when the vapor pressure of the liquid is the same as the external pressure.

Boiling Point and Pressure Change
-Because boiling points involve external pressure, boiling points can be different depending on what the external pressure is.
-Sea level can affect atmospheric pressure.
-Boiling points decrease at lower pressures, and increases in higher ones. This is the particles may need more or less kinetic energy, depending on the atmospheric pressure.
-A liquid's temperature never surpasses its boiling point; therefore, adding more heat just makes it boil faster.

Normal Boiling Point
-We know that a liquid can have various boiling which one is official?
-The normal boiling point is decided as being the boiling point of a liquid at a pressure of 101.3 kPa (kilopascals)

Group 3

Changes of State

Coeditor: ?

Group: Lindsey Bedrosian, Shannon Degnan, and Abby John

p. 396-397
Lindsey Bedrosian

Model for Solids:
- The general properties of solids reflect the orderly arrangement of their particles.
- In most solids atoms, ions, or molecules packed tightly together.
- Solids are dense and not easy to compress.
- Particles in solids vibrate but don't move.
- Melting point: the temperature at which a solid changed into a liquid (The same a freezing

Crystal Structure and Unit Cells:
- Crystal particles are arranged in an orderly, repeating pattern called crystal lattice.
- The shape of a crystal reflects the arrangement of the particles within the solid.
- The type of bonding determines the melting point.
- There are seven types of crystal:
- Galena, Zircon, Topaz, Gypsum, Amazonite, Tourmaline, Calcite

p. 398-399
Shannon Degnan

-a diamond is one of the most known allotropes
-the term allotrope refers to one or more forms of a substance
-substances that have allotropes include carbon, oxygen, sulfur and phosphorus

Amorphus solids and gases
-an amorphus solid is the name for a solid that does not have a crystalline structure
-there are no long-range patterns of crystal
-this term is usually associated with glasses

350px-Eight_Allotropes_of_Carbon.png images.jpeg

Shannon Degnan

p. 400-403
Abby John

Group 4

Properties of Gases

Coeditor: Hannah Valley

Group: Meghan Faber, Hannah Valley, and Lauren Altmeyer

Compressibility p. 413-414
Meghan Faber
  • Gas can expand to fill its container, unlike a solid or liquid.
  • Gases are easily compressed.
  • Compressibility: a measure of how much the volume of matter decreases under pressure.
  • Kinetic theory: gases are easily compressed because of the space between particles in a gas
  • The volume of the particles in a gas is smaller than the overall volume of the gas.

picture by Meghan Faber

Factors Affecting Gas Pressure p.414-416
Lauren Altmeyer

  • Increasing the number of gas particles increases the number of collisions, which makes the pressure increase
  • More molecules means more collisions
  • Gases move from areas of high pressure to low pressure
  • As volume decreases, pressure increases

Lauren Altmeyer

Volume and Temperature p. 416-417
Hannah Valley

  • you can raise the pressure exerted by a gas by reducing its volume
  • more compression, greater pressure
  • if volume is doubled, pressure is cut in half
  • increase in temperature of an enclosed gas causes an increase in its pressure
  • as gas is heated, temperature increases and average kinetic energy of the particles in the gas increases
  • as temperature of enclosed gas decreased, pressure decreases, because they are moving slower and have less kinetic energy

external image moz-screenshot-6.png
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Group 5

The Gas Laws

Coeditor: Kelsey Sullivan

Group: Kelsey Sullivan, Emily Taylor, and Anne O'Toole

Boyle's Law: Pressure and Volume p. 418-419
Anne O'Toole
  • if the temperature is constant, as the pressure of a gas increases, the volume decreases.
  • As the pressure decreases, the volume increases
  • This is Boyle's law
  • Boyle's law states that for a given mass of gas at constant temperature, the volume of the gas varies inversely with pressure.
  • In an inverse relationship, the product of the two variable quantities is constant.
  • P1 X V1 = P2 X V2
external image Boyles-Law.gif
Picture by : Anne O'Toole

Charles Law: Temperature and Volume p. 420
Emily Taylor
  • If the pressure is constant, the volume increases as the temperature of an enclosed gas increases.
  • Jacques Charles studied the effect of temperature on the volume of a gas at a constant pressure.
  • When he graphed his data, he saw that the glas volume vs, the temperature at 0C was a straight line for any gas, and when he extended his line to zero volume, the line intersected the temperature axis at -273.15C which is equal to 0C on the Kelvin scale.
  • Charles's law states that the volume of a fixed mass of gas is directly proportional to its Kelvin temperature if the pressure is kept constant.
  • V1 = V2
T1 T2
  • The ratio of the variables is always a constant in a direct relationship and the graph is always a straight line.
  • It is NOT a direct relationship if the temperatures are shown in Celsius; the temperatures must always be expressed in kelvins.

Charles's Law shown in a demo:
- Emily Taylor

Gay-Lussac's Law: Pressure and Temperature p. 422
Kelsey Sullivan
  • If the volume of an enclosed gas is constant, as the temperature increases, the pressure increases
  • Joseph Gay-Lussac was a French Chemist
  • He discovered the relationship of temperature and pressure
  • Guy-Lussac's law says that the pressure of a gas is proportionalto the Kelvin temperature if the volume is constant
  • Guy-Lussacs formula
    • P1/T1 = P2/T2
  • Here is a video explaining how to use Guy-Lussac's law of pressure and temperature
-Kelsey Sullivan

The Combined Gas Law p.424-425
Anne O'Toole
  • The combined gas law describes the relationship among the pressure, temperature, and volume of an enclosed gas. The combined gas law allows you to do calculations for situations in which only the amount of gas is constant.
  • P1 x V1 = P2 x V2
------------- -----------
T1 T2
  • A drop in temperature causes the volume of an enclosed gas to decrease.
  • The drop in pressure must affect the volume more than the drop in temperature does

Picture by : Anne O'Toole

Group 6

Ideal Gases

Coeditor: Alex Fischbach

Group: Brendan Lynch, Lauren Bedard, and Alex Fischbach

Brendan Lynch
Ideal Gas Law
With the ideal gas law, you can solve problems with 3 different variables.
ex. P1 x V1 P2 x V2
------------ = -------------
T1 x n1 T2 x n2

the number of moles of gas is directly proportional to the number of particles.
therefore, moles must be directly proportional to volume as well.

The ideal gas law - P x V = n x R x T or PV= nRT

external image ideal_gas.png
Picture by Brendan Lynch
Lauren Bedard
Ideal Gases and Real Gases
  • ideal gas:
    • follows the gas laws at all conditions of pressure and temperature
    • conform to assumptions of kinetic theory
    • particles-- no volume/no attraction between them
    • DOES NOT EXIST, but at many conditions real gases behave a lot like ideal gases
  • real gas:
    • particles have volume/ attraction between them
    • can condense or solidify when compressed/cooled
  • real gases differ most from an idea gas at low temperatures and high pressures

external image ideal_gas.png

Alex Fischbach
Real Gases Deviate From the Ideal
  • The ratio of (PV/nRT) changes as pressure increases
  • For an ideal gas the ratio is always equal to 1
  • For real gases at high pressure the ratio may depart from the ideal
  • As attractive forces reduce the distance between particles, a gas occupies less volume then expected, causing the ratio to be less than 1. But the actual volume of the molecules causes the ratio to be greater than 1.


Picture by: Alexandra Fischbach

Group 7

Gases: Mixtures and Movements

Coeditor: Erin Garrity

Group: Erin Garrity, Mark Cuddy, Colleen Fitzgerald

Dalton's Law p.432-433
Erin Garrity

· When particles in a gas collide with an object, gas pressure results
· The number of collisions is proportional to the amount of particles in that when the number of particles increases the number of collisions increases
· The amount of pressure each gas in a gas mixture makes is the partial pressure
· In a gas mixture Ptotal=P1+P2+P3+P….
· Dalton’s Law of partial pressures tells us that when T (temperature) and V (volume) are constant the total exerted pressure of the gas is the sum of the partial pressures of each gas
o This value is represented either in torr or atm. (atmospheres)
o 1 atm. = 760 torr
Graham's Law p.434-435
Mark Cuddy

Graham's Law
- Law of effusion by Scottish chemist Thomas Graham

-The rate of effusion of a gas is inversely proportional to the square root of the gas's molar mass

Formula: r1 / r2 = Square Root (M2 / M1)
VIDEO on Graham's Law

-Diffusion is the process of gases moving to areas of low concentration
-Effusion is the process of gas escaping through tiny holes in its container
-Gases of lower molar mass diffuse and effuse faster than gases of higher molar mass

-Gas molecules bounce around their container until they make impact with something
-Gas molecules have the chance of escaping straight through holes in its container
- Thomas Graham observed the relation of the rate and mass of the particles
-Mass, velocity, and kinetic energy are key
-For the Kinetic energy of the gas to be constant, any increase in mass must be balanced by a decrease in velocity
-This observation helped Graham deduce his new law

Comparing Effusion Rates p.434-435
Colleen Fitzgerald