=
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PROPERTIES OF GASES
"The main distinguishing property of gases is their uncanny ability to be compressed into smaller and smaller spaces. Gases are also the least complex state of matter. Don't get it wrong, just because they are the simplest doesn't mean that they are not one of the most interesting and useful states of matter."
"Gases are easily expandable and compressible unlike solids and liquids. Gases have a measurement of pressure. Pressure is defined as force exerted per unit area of surface. It can be measured in several units such as kilopascals (kPa), atmospheres (atm), and millimeters of Mercury (mmHg). Gas has a low density because its molecules are spread apart over a large volume. A gas will fill whatever container that it is in. An example of this is a bottle of ammonia being opened in a room and the smell traveling throughout the room."
(http://library.thinkquest.org/10429/low/gaslaws/gaslaws.htm)

Chapters 13-14: The Gas Laws

PAGES: 385-437

Editor: Dan McCormack

GROUP 1:

PAGES: 385-389

Co-EDITOR: CAROLINE RUBINO(388-389)

Kinetic Energy and Temperature

• As a substance is heated, its particles absorbe energy

- some of this energy is stored within the particles

• Potential energy

- the stored portion of the energy within the particles
- does not raise the temperature of the substance

• Kinetic energy

- the remaining absorbed energy speeds up the particles (speeds up the kinetic energy)
- This increase in the kinetic energy results in an increase in temperature

• Average Kinetic

- 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, regardless of physical state, have the same average kinetic energy.
For Example: The ions in table salt, the molecules in water, and the atoms in helium all have the same average kinetic energy at room temperature even though the three substances are in difent physical states.
This diagram shows the distribution of kinetic energies of water molecules at two different temperatures.

• An increase in the average kinetic energy of the particles causees the temperature of a substance to rise.
• As a substance cools, the particles tend to move more slowly, and theire average kinetic energy declines.
• Absolute Zero = temperature at which the motion of particles stops (0k, or -273.15 C)
• 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 proportional to the average kinetic energy of the particles of the substance.

ZOEY KILLION (385-387)

Kinetic Theory & a Model for Gases (pages 385-386):

• Kinetic Energy - the energy has because of its motion
• Kinetric Theory - states that the tiny particles in all forms of matter are in constant motion
• A gas is composed of particles, usually molecules or atoms.
• The particles in a gas move rapidly in constant random motion. They travel in straight paths &move independently of each other.
• All collisions between gas molecules are perfectly elastic.

Gas Pressure (pages 386-387):

• Gas Pressure - the force ecerted by a gas per unit surface area of an object
• Gas pressure is the result of simultaneous collisions of billions of rapidly moving gas particles with an object.
• Vacuum - an empty space, with no particles and no pressure
• Atmospheric Pressure - results from the collisions of air molecules with objects
• Barometers - devices commonly used to measure atmospheric pressure

• Pascal (Pa) - SI Unit of Pressure
• Standard Atmosphere (atm) - the messure required to support 760 mm of mercury in a mercury barometer at 25 degrees Celcius
  • 1 ATM = 760 mm Hg = 101.3 kPa

Video on Kinetic Theory:
http://www.youtube.com/watch?v=nrSt_fl1q6w

GROUP 2:

PAGES: 390-395

CO-EDITOR: SHANNON LAMY(390-391)

A Model for Liquids (Shannon Lamy)
· Kinetic theory states particles in gases and liquids have kinetic energy.
o This is why gas and liquid particles can flow past each other.
REMINDER –

• Substances that can flow = liquids
• Both gases and liquids take the shape of their container

• Particles in gas = NO ATTRACTION

VS

• Particles in liquid = ATTRACTED TO EACHOTHER

o Keeps molecules close together
§ Definite volume
*The interactions between the disruptive motions of particles in a liquid and the attraction between molecules determine its physical properties.**

• Liquids = greater density then gases

• Increasing pressure on liquid doesn’t effect it

o Compared to a gas like the demo’s Mr. D has shown us in class.
Liquids and solids = CONDENSED STATES OF MATTER
· Because pressure does not effect volume.
Evaporation (Shannon Lamy)

• VAPORIZATION = the conversion of a liquid to a gas or vapor.
• EVAPORATION = when a conversion from liquid to gas or vapor occurs at the SURFACE of a liquid that is NOT boiling.

Evaporation in water cycle

*During evaporation, only those molecules with a certain minimum kinetic energy can escape from the surface of a liquid.**

• Liquids evaporate faster when heated

o Increases average kinetic energy of particles
§ More particles can overcome attraction that keeps them a liquid and can evaporate.

• Particles with highest kinetic energy escape first

Highest energy particles are releases first.

• Liquid temp decreases with evaporation

o EVAPORATION =COOLING PROCESS

Alex Nunan, pg. 392-393

Vapor Pressure:

vapor pressure- measure of the force exerted by a gas above a liquid in a sealed container; a dynamic equilibrium exists between the vapor and the liquid:

• Where there is constant vapor pressure, there is a dynamic equilibrium between the vapor and the liquid. This is because the rate of evaporation of liquid is equal to the rate of condensation of vapor.

Vapor Pressure and Temperature Change:

• increased temperature of a contained liquid = increased vapor pressure
  • warmed liquid has increased kinetic energy
  • This means more of the particles will have the minimum kinetic energy necessary to escape the surface of the liquid. These particles will escape the liquid and collide with the walls of the container at a greater frequency.
• vapor pressure data indicates how volatile (or how easily it evaporates) the liquid is

Vapor Pressure Measurements:

• a manometer is a device that measures the vapor pressure of a liquid
• check out this site to see how they [[The Gas Laws#|work]]: http://www.chm.davidson.edu/vce/gaslaws/pressure.html
  • (This [[The Gas Laws#|website]] has a lot of other information about all the different gas laws, too; just go to the home page!)

http://www.elmhurst.edu/~chm/vchembook/163boilingpt.html

Boiling Point:

boiling point- the temperature at which the vapor pressure of a liquid is just equal to the external pressure on the liquid:

• evaporation rate increases with heat
• Liquid begins to boil when it is heated to a temperature at which particles throughout the liquid have enough kinetic energy to vaporize.

OLIVIA RICHARDSON (394-395)

Boiling Point and Pressure Changes

• A liquid boils when its vapor pressure is equal to the external pressure so liquids boil at different temperatures.

For example changing the altitude of a liquid will change the boiling point of water. At sea level the boiling point of water is around 100 C and atop Mount Everest, a higher altitude, the atmospheric pressure is lower than it is at sea level so the water boils at a lower temperature.

• This graph shows how the boiling point of a liquid is related to the vapor pressure; at a lower external pressure, the boiling point decreases and vise versa.

• Boiling is a cooling process similar to evaporation
• When a liquid is boiling, the particles with the highest kinetic energy escape first.
• The more heat that is supplied, the more particles are escaped from the liquid but the temperature never rises above the boiling point.
• The vapor produced during boiling is the same temperature as the liquid.
• The potential energy is much higher in the vapor then in the liquid but the kinetic energy is about the same
• normal boiling point- the boiling point of a liquid at a pressure of 101.3kPa. For example the normal boiling point for water is 100 C

_

GROUP 3: PAGES:396-403

- PJ Hamill
- Mike Hanley
- Brandon Boisclair

PJ HAMILL( 396-398)

13.3 the nature of solids

Melting point – is the temperature in which a solid turns to a liquid
Crystal – the particles are arranged in an orderly repeating three-dimensional pattern called crystal lattice
A Crystal has sides or faces and the angles at which the faces of crystal intersect are always the same for a given substance
Allotropes - one of many forms in which a chemical element occurs, each differing in physical properties, e.g. diamonds and coal as forms of carbon

Non- Crystalline solids

MIKE HANLEY(399-401)

• An Amourphous solid lacks an ordered internal structure.

• Rubber, plastic, and asphalt are all examples of amorphous solids
• A glass is a transparent fusion of inorganic substances that have cooled to a rigid state without crystallizing
• The irregular internal structure of glasses are intermediate between those of a crystalline solids and those of a free flowing liquid
• Glasses don't melt at a definite temperature, they gradually soften when being heated.

Picture of glass being heated for glass blowing,
it gradually becomes softer rather than immediately
melting.

Sublimation

• The change of a substance from a solid to a vapor without passing through the liquid state is called sublimation.
• Sublimation can occur because like liquids, solids have a vapor pressure.
• Key Concept: Sublimation occurs in solids with vapor pressures that exceed atmospheric pressureat or near room temperature
• Iodine is an example of a substance that undergoes sublimation, it turns into a purple vapor without passing through the liquid state.

.
Image of Iodine sublimation

• Sublimation often has many useful applications.
• Organic Chemists can use sublimation to separate mixtures and to purify compouds
• The video below gives an explanation and demonstration of sublimation:
• http://www.youtube.com/watch?v=dBNELFi5XiY

Brandon Boisclair ( 402 - 403 )

PHASE DIAGAMS

• The relationship among the solid, liquid, and vapor states (or phases) of a substance in a sealed container can be represented in a single graph.

- The graph is called a phase diagram

• A phase diagram gives the conditions of temperature and pressure at which a substance exists as solid, liquid, and gas (vapor)

___

GROUP 4: Properties Of Gases

PAGES: 413-417

CO-EDITOR: EVAN SOMMERICH

JAMES PAYNE 413-414 Compressibility

Compressibility-a measure of how much the volume of matter deceases under pressure

Compression of gas absorbs energy of impact with an airbag
Gases are easily compressed because of the space between particles in a gas
Under pressure, gas particles are forced closer together

• The amount of gas, volume, and temperature are factors that affect gas pressure.

DAKOTA PIMENTEL and EVAN SOMMERICH

Factors Affecting Gas Pressure There are four variables used to describe a gas.

• pressure (P) (atm,torr,pascal...)
• volume (V) (liters)
• temperature (T) in kelvin
• number of moles (n).

The amount of gas, volume, and temperature affect gas pressure.

• Increasing the amount of gas increases the pressure.
  • This is because it increases the number of gas particles which increases the number of collisions.
  • Example: This is how aerosol cans work. Inside the aerosol can is a gas with a higher pressure then the gas outside of the can. When the button on the top of the can is pressed, and an opening is created, the gas inside propels the product inside the can through the opening.

http://oujnews.in/aerosol-can-diagram&page=2

SECTION TWO: VOLUME AND PRESSURE

Key Point: Decreasing the amount of volume increases the amount of pressure.

• An increase in temperature increases the pressure.
• This is because it increases the kinetic energy of the particles in the gas. These faster moving particles impact the walls of the container with more energy.
  • This is the reason why a bag of potato chips explodes when placed in a sunny location, or why an aerosol can explodes in a fire.

GROUP 5

PAGES: 418-425

CO-EDITOR: CHRISTIAN COOKE (418-420)

MITCH"MOOSE" MARTIN (421-422)

NATE LYNCH (423-425)

Using Gay- Lussac's Law

- P1/T1 = P2/T2
make a chart
convert celcius to Kelvins
Substitute values and calculate

The Combined Gas Law


make a chart and substitute values
The combined gas law allows you to do calculations for situations in which only the amount of gas is present.
You can do problems when only two variables are changing.

GROUP 6: IDEAL GASES

PAGES: 426-429

CO-EDITOR: BECKY HYATT(426-427)

Ideal Gas Law:

• to calculate the number of moles of a contained gas requires an expression that contains the variable n
• the number of moles of gas is directly proportional to the number of particles
  • P1 x V1 / T1 x n1 = P2 x V2 / T2 x n2
• 1 mol of every gas occupies 22.4 L at STP
• the ideal gas constant (R) has the value 8.31 (L x KPa) / (K x mol)
• the gas law that includes all four variables--P, V, T, and n--is called the ideal gas law
  • P x V = n x R X T or PV = nRT

Sample Problem 14.5
Using the Ideal Gas Law to Find the Amount of a Gas:
A deep underground cavern contains 2.24 x 10^6 L of methane gas (CH4) at a pressure of 1.50 x 10^3 kPa and a temperature of 315 K. How many kilograms of CH4 does the cavern contain?

• Analyze- list the knowns and the unknown
  • calculate the number of moles (n) using the ideal gas law
• Calculate- solve for the unknown
  • rearrange the equation for the ideal gas law to isolate n
• Evaluate- Does the result make sense?
  • although the methane is compressed, its volume is still very large

Practice Problem:
When the temperature of a rigid hollow sphere containing 685 L of helium gas is held at 621 K, the pressure of the gas is 1.89 x 10^3 kPa. How many moles of helium does the sphere contain?
For more help with the Ideal Gas Law, see this video demonstration:
http://www.youtube.com/watchv=erjMiErRgSQ&NR=1&feature=fvwp

HALEY CONATSER(428-429)

Ideal Gases and Real Gases:

Ideal Gas: one that follows the gas laws at all conditions of pressure and temperature
- an ideal gas would have to conform precisely to the kinetic theory.
- particles have no volume
- no attraction between particles
- an idal gas does NOT exist.

Real Gas: Have volume and there are attractions between particles
- because of this a gas can condense or even soldify when it is compressedd and cool

• Real gases differ from an ideal gas to low temps. and high pressures. *

Real Gases Deviate from the Ideal:

• For an ideal gas the result is a horizontal line because the ratio is 1
• When the ratio is greater then 1 the curve rises
• When the ratio is less then 1 the curve drops

• As attractive forces reduce distance between particles, a gas occupies less volume, causing the ratio to be less then 1
• The actual volume of the molecules cause the ratio to be greater then one
• In portions of the curve below the line, intermolecular attractions dominate
• In portions of the curve above the line, molecule volume dominates

GROUP 7:MIXTURES AND ELEMENTS

PAGES 432-437

CO-EDITOR: ANDREA LOUNGO

ELIZABETH HOWARD

Dalton's Law

Elizabeth Howard (pages 432-434)

• 
  
• Gas pressure depends only on the number of particles in a given volume and on their average kinetic energy.
• Partial Pressure is the contribution each gas in a mixture of gases makes to the total pressure.
• In a mixture of gases, the total pressure is the sum of the partial pressure of the gases.
• Dalton's Law of Partial Pressures states that at constant volume and temperature , the total pressure exerted by a mixture of gases is equal to the sum of the partial pressures of the component gases.

• If the percent composition does not change, the fraction of the pressure exerted by a gas does not change as the total pressure changes.
  • This is important for people operationg at high altitudes. For example, the atmospheric pressure at the top of Mount Everest is 33.83kPa, about one-third of its value at sea level. The partial pressure must then reduce to one-third as well, 7.06kPa.
    • Because this amount is not high enough to support the respiration of humans, oxygen masks are needed by humans to survive.

• Using Dalton's Law of Partial Pressure-Sample Problem:
  • Air contains oxygen, nitrogen, carbon dioxide, and trace amounts of other gases. What is the partial pressure of Oxygen (PO2) at 101.30kPa of total pressure if the partial pressure of nitrogen, carbon dioxide, and other gases are 79.10kPa, 0.040kPa, and 0.94kPa, respectively?
  • List the Known and Unknown:
    • Known:
    • PN2=79.10kPa
    • PCO2=0.040kPa
    • POthers=0.94kPa
    • PTotal=101.30kPa
    • Unknown:
    • PO2=?kPa
  • Solve for the Unknown:
    • Rearrange Dalton's Law to isolate PO2. Substitute the values for the partial pressures and solve the equation.
    • PO2=PTotal-(PN2+PCO2+POthers)
    • PO2=101.30kPa-(79.10kPa+0.040kPa+0.94kPa)
    • PO2=21.22kPa
  • Does the result make sense?
    • The partial pressure of oxygen must be smaller than that of nitrogen because PTotal is only 101.30kPa. The other partial pressures are small, so no answer of 21.22kPa seems reasonable.

Graham's Law

By: Andrea Luongo (Page 435-437)

Graham's Law

• Molecules in a room spread out to fill all the space within the room.

• Diffusion is the tendency of molecules to move toward areas of lower concentration until the concentration is uniform throughout the area it occupies.
• Effusion is the process that occurs when a gas escapes through a tiny hole in its container.

• Gases of lower molar mass diffuse and effuse faster than gases of higher molar mass.
• In effusion and diffusion, the type of particle is important.

Thomas Graham's Contributions
Thomas Graham

• Thomas Graham was a Scottish chemist who studied rates of effusion in the 1840s.
• During his studies, he proposed a law. This was called the law of effusion.

Law of Effusion

• Graham's Law of Effusion states the the rate of effusion of a gas is inversely proportional to the sqaure root of the gas's molar mass.
  • This is also applied to the Diffusion of gases.
• In order to understand Graham's Law, you must know how the mass, velocity, and kinetic energy of a moving object are related.
  • The formula to relate the mass (m), velocity (v), and kinetic energy (E) is
    • 

• To have a constant kinetic energy, mass increases to balance with velocity that decreases.
• If 2 objects with different masses has the same kinetic energy, then the lighter object will move faster.


Comparing Effusion Rates:

• The rate of effusion is related to a particle's speed.
• Graham's Law for two gases, 1 and 2, is written:
  • 
  • M = molar mass
• The rate of effusion of 2 gases are inversely proportional to the square roots of their molar masses.

• To learn more about the rate of effusion and to try some practice problems visit:
  • http://www.youtube.com/watch?v=TropqfFqUiI