Chapters+13+and+14

Chapters 13 & 14 Gas Laws  Period C    Introduction This page has been constructed on the thirteenth and fourteenth chapters of Prentice Hall's //Chemistry// textbook. The thirteenth focuses on the states of matter. It explains the nature of gases, liquids, and solids, as well as how elements undergo change between the three. The fourteenth chapter discusses the behavior of gases specifically. It goes more in depth into the properties of gases such as compressibility, pressure, volume, and pressure. It explains the gas laws that teach how to find volume, pressure, or temperature such as Boyle's Law, Charles' Law, Gay-Lussac's Law, the Combined Gas Law, and the Ideal Gas Law. Finally, it explains gas mixtures with Dalton's Law of Partial Pressures and Graham's Law. This wiki-page outlines the chapter in its entirety, and has included all [|information] crucial to the understanding of these chapters. Groups: Editor: Elizabeth Sieber Section 13.1-The Nature of Gases pg. 385-389: Marion Burdick(coeditor), Nina DeMeo, Eileen Corkery Section 13.2-The Nature of Liquids pg. 390-393: Tom DeMarco(coeditor), Steven Denison Section 13.3-4 -The Nature of Solids, Changes of State pg. 396-403: Courtney Gareau (coeditor), Kendyl Barron, Mike McShane Section 14.1-Properties of Gases pg. 413-417: Adam Shanahan (coeditor), Heather Bowditch Section 14.2-The Gas Laws pg. 418-425: Andrew Sciotti (coeditor), Christos Anastos Section 14.3-Ideal Gases pg. 426-429: Maggie Bie (coeditor), Evan Grandfield Section 14.4-Density pg. 89-83: Katherine Perry (coeditor), Nick Brault

Section 13.1 The Nature of Gases Marion Burdick (coeditor): pg. 385-386 Nina DeMeo: pg. 386-387 Eileen Corkery: pg. 388-389

**Marion Burdick (Coeditor)- pg. 385-386** __// ** Kinetic Theory and a Model for Gases ** //__ Picture: Marion Burdick media type="youtube" key="fIMdIMACyN4?rel=0" height="349" width="425" Video: Marion Burdick
 * Kinetic comes from a word meaning “to move”
 * **kinetic energy:** the energy of an object has because of its motion
 * **kinetic theory:** a theory explaining the states of matter, based on the concept that the particles in all forms of matter are in constant motion
 * ** Key Concept: A gas is composed of particles, usually molecules or atoms. **
 * These “particles” are small and spherical. They have no measurable volume and are far apart (because they are neither attracted nor repulsed from one another unless in high pressure or low temperature)
 * ** Key Concept: The particles in a gas move rapidly in constant random motion. They travel in straight paths and move independently of each other. **
 * Gases fill their container no matter the size or shape of the container
 * If a gas is uncontained (or in a container with an opening) it diffuses into space with no limits!
 * Gas particles change direction if a.) they collide with each other or b.) they collide with another object
 * ** Key Concept: All collisions between gas molecules are perfectly elastic. **
 * At each point of elastic collision energy is transferred from one particle to the next BUT the total kinetic energy remains constant
 * Video attribution: Marion Burdick

**Nina Demeo- pg. 386-387** //__Gas Pressure__// particles in a gas with an object - a helium filled balloon maintains its shape because of the gas pressure within it o gas pressure results from the force exerted by a gas per unit surface area of an object - if there are no particles, there can be no collisions and consequently there is no pressure - an empty space with no particles and no pressure is called a **vacuum** - a familiar gas pressure is air! o Air exerts pressure on earth because gravity holds the particles in air in earth’s atmosphere o **Atmospheric pressure** results from the collisions of atoms and molecules in air with objects; it depends on weather and altitude - A barometer measures atmospheric pressure - The SI unit of pressure is the **pascal** - Other pressure units include the atmosphere, and mm Hg (torr) - 1 atm = 760 mm Hg (torr) = 101.3 kPa Example of Converting: Pressure = 450 kPa Atm = ? atm = 450 **×** __1atm__ = 4.4 atm 101.3 kPa
 * Key Concept :** Gas pressure is the result of simultaneous collisions of billions of rapidly moving

media type="youtube" key="qv81QCGNnVo" height="390" width="480" Video: Nina DeMeo

Eileen Corkery pgs.388-389
//__ Kinetic Energy and Temperature __// à Kinetic Energy= the energy an object has because of its motion. //__ Average Kinetic Energy __// à Average Kinetic Energy is used when we talk about a collection of molecules or atoms in a substance. At any given temperature, the average will always be the same, even if in different physical states. · One example of this is the ions in table salt, the molecules in water, and the atoms in helium. Even though they are all in different physical states, they all have the same average kinetic energies (at room temperature). à An increase in kinetic energy of the particles in a substance causes the temperature of the substance to rise. A decrease in kinetic energy (when the particles move more slowly) makes the temperature to go down. · From hearing about the relationship between kinetic energy and temp, you would think that at a certain temperature, kinetic energy would not exist. This is what scientists call //absolute zero// (0 K, or 273.15 C). Right now, absolute zero has never been produced in a laboratory, but it has nearly been achieved in vacuum chambers. //__ Average Kinetic Energy and Kelvin Temperature __// à The Kelvin Temperature Scale reflects the relationship between temperature and average kinetic energy. à The Kelvin temperature of a substance is directly proportional to the average kinetic energy of the particles of the substance. Photo credits: Eileen Corkery Section 13.2 The Nature of Liquids Tom DeMarco (coeditor): pg. 390, 393-395 Steven Denison: pg. 391-393
 * When a substance is heated, the particles absorb energy and store it either inside themselves or outside. The energy stored within the particles doesn't raise the temperature, but the remaining absorbed energy does. When the particles speed up, the kinetic energy of the substance is increased, and the temperature rises.

**Thomas DeMarco (Coeditor) - Pg. 390, 393-395** //__A Model for Liquids__// Particles of both gases and liquids can flow past each other because they both have kinetic energy. This ability to flow means they can conform to the shapes of their containers. The main difference between gases and liquids is that liquid particles are attracted to each other, but gas particles are not. This means that liquids will have a definite volume. **The interplay between the disruptive motions of particles in a liquid and the attractions among the particles determines the physical properties of liquids.** The intermolecular attraction between liquid particles means there is less space in between them. That means liquids are much more dense than gases. Pressure on liquids and also solids has little impact. These states are, therefore, known as condensed states of matter.

//__Boiling Point__//
 * When a liquid is heated to a temperature at which particles throughout the liquid have enough kinetic energy to vaporize, the liquid begins to boil.** Bubbles will appear throughout the liquid and rise to the top and escape into the air. The temperature at which the vapor pressure of the liquid is just equal to the external pressure on the liquid is the **boiling point (bp)**.

//__Boiling Point and Pressure Changes__// Because liquids boil when the vapor pressure and external pressure are equal, liquids won’t boil at the same temperature. Boiling points decrease at higher altitudes because the atmospheric pressure decreases. In a pressure cooker, the vapor can’t escape so the vapor pressure increases. This means the water will boil at a temperature above 100oC and food can cook quicker. Normal Boiling Point Boiling is a similar cooling process to evaporation. The normal boiling point is defined as the boiling point of a liquid at a pressure of 101.3kPa. Picture: Tom DeMarco

// __Evaporation__ // The conversion of a liquid to a gas or vapor is called vaporization. When this process occurs on the surface of a liquid which is not boiling, it is called evaporation. Only molecules with enough kinetic energy can escape from the surface of the liquid. Even if the molecule has enough kinetic energy there is no guarantee that it will escape the liquid. Occasionally an escaping molecule will collide with a particle in the air and bounce back into the liquid. Heating a liquid will cause it to evaporate quicker due to the added amount of kinetic energy. Because those particles with the most kinetic energy escape first, the total amount of kinetic energy in the liquid decreases. This is why evaporation is considered a cooling process.
 * Steven Denison- Pg. 391-393 **

//__Vapor Pressure__// Evaporation of a liquid within a closed container produces a pressure known as vapor pressure. Vapor pressure is the measure of force exerted by a gas above a liquid. In a system at constant vapor pressure, a dynamic equilibrium exists between the vapor and the liquid. This happens because the rate of evaporation and condensation are equal. Vapor pressure increases when the liquid inside a container is heated. This is because, as stated above, with more kinetic energy added, more particles can escape and therefore exert force against the walls of a container. Vapor pressure can be measured using a manometer. A manometer is used to observe a change in vapor pressure as pressures on both sides of a U-shaped tube result in changing levels of mercury within the tube.

Pictures: Steven Dension Section 13.3 The Nature of Solids Courtney Gareau (coeditor): pg. 402-403 Kendyl Barron: pg. 396-398 Mike McShane: pg. 399-401

Kendyl Barron pgs.396-398
__//A Model for Solids//__  à The atoms, ions, or molecules of a solid substance are orderly arranged in a tight structure.  à Solids are dense and not easy to compress and due its structure, solids do not flow. àHeating a solid causes its particles to vibrate rapidly, increasing the kinetic energy of the particles and breaking them down causing the solid to melt.  à The **melting point (mp)** of a solid is the temperature at which it melts into a liquid phase.  à At a solid’s melting point, the vibrations of a structure’s particles can overcome their strict crystalline structure.

__//Crystal Structure and Unit cells// __  à Solid substances are usually arranged in a crystalline structure, or an orderly, repeating, three-dimensional pattern known as a **crystal lattice.**  à The arrangement of a solid’s particles determines its crystal system and shape. <span style="font-family: Arial,Helvetica,sans-serif; margin: 0in 0in 0pt;"> à The bonds between particles of a crystal determine the solid’s melting points. <span style="font-family: Arial,Helvetica,sans-serif; margin: 0in 0in 0pt;"> à Ionic solids, with strong bonds, tend to have higher melting points than molecular solids. <span style="font-family: Arial,Helvetica,sans-serif; margin: 0in 0in 0pt;"> à However, some solids such as wood or some sugars do not melt but decompose when heated.

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<span style="font-family: Arial,Helvetica,sans-serif; margin: 0in 0in 0pt;">//Crystal systems// <span style="font-family: Arial,Helvetica,sans-serif; margin: 0in 0in 0pt;"> à The angles of intersection of a crystals faces are constant in a given substance. <span style="font-family: Arial,Helvetica,sans-serif; margin: 0in 0in 0pt;"> à The crystal classification of a solid is a defining characteristic. <span style="font-family: Arial,Helvetica,sans-serif; margin: 0in 0in 0pt;"> à Crystal systems include cubic, tetragonal, orthorhombic, monoclinic, triclinic, hexagonal, and rhombohedral.

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<span style="font-family: Arial,Helvetica,sans-serif; margin: 0in 0in 0pt;"> à Crystal shape depends on the organization of the particles within a substance. <span style="font-family: Arial,Helvetica,sans-serif; margin: 0in 0in 0pt;"> à A **unit cell** is the smallest group of particles of a crystal that still retains the crystal’s geometric shape and properties. <span style="font-family: Arial,Helvetica,sans-serif; margin: 0in 0in 0pt;"> à A crystal lattice is an array of one of several structures of a unit cell.

<span style="font-family: Arial,Helvetica,sans-serif; font-size: 120%; margin: 0in 0in 0pt;">**Mike McShane- pg. 399-401**

<span style="font-family: Arial,Helvetica,sans-serif; font-size: 120%; margin: 0in 0in 0pt;">__//<span style="font-family: 'Arial','sans-serif'; font-size: 14.5pt;">Non-Crystalline Solids //__ <span style="font-family: 'Arial','sans-serif'; font-size: 90%;">Non-crystalline solids are also known as __amorphous solids__ which means that they lack an ordered internal structure. They are this was because their atoms are arranged randomly Ex: Rubber or plastic.

One type of amorphous solids is __glass__. A glass is a see-through mixture of inorganic substances that are cooled to a solid aka a supercooled liquid. They do not melt at a definite temperature but rather soften when heated.

When an amorphous solid is shattered, the fragments have irregular angles and jagged edges as opposed to crystalline solids that, when shattered, have fragments with the same surface angles as the original solid. -Shattered Glass: Mike McShane

<span style="font-family: 'Arial','sans-serif'; font-size: 14.5pt;">//__Sublimation__// Sublimation is the change of a substance from a solid to a vapor without going through a liquid state. It is when a solid changes directly to a gas. It occurs in solids with vapor pressures that exceed atmospheric pressure at or near room temperature. One common example of sublimation is when dry ice disappears it is sublimating. Iodine is an example of a substance that undergoes sublimation.

Sublimation can be useful. It can be used to make freeze-dried coffee, dry ice is used as a coolant during shipping, and air fresheners contain substances that undergo sublimation. It is also useful for separating substances such as mixtures and to purify compounds. -Dry Ice: Mike McShane

<span style="font-family: Arial,Helvetica,sans-serif; font-size: 120%; margin: 0in 0in 0pt;">**Courtney Gareau (Coeditor) - pg. 402-403** <span style="font-family: Arial,Helvetica,sans-serif; font-size: 130%; margin: 0in 0in 0pt;">//__Phase Diagrams__// <span style="font-family: Arial,Helvetica,sans-serif; font-size: small; margin: 0in 0in 0pt;"> **Phase Diagram** – a graph showing the conditions at which a substance exists as a solid, liquid, or vapor <span style="color: #000000; font-family: Arial,Helvetica,sans-serif; font-size: small; margin: 0in 0in 0pt;">Key Concept – The conditions of pressure and temperature at which two phases exist in equilibrium are indicated on a phase diagram by a line separating the phases.

media type="youtube" key="Qp87Z4m8R-w" height="349" width="560" <span style="font-family: Arial,Helvetica,sans-serif; margin: 0in 0in 0pt;">Video on Phase Diagrams - Courtney Gareau

<span style="font-family: Arial,Helvetica,sans-serif; margin: 0in 0in 0pt;">In the diagram below, each of the curving lines illustrates at what place in the diagram the abutting phases are in equilibrium. For example, the line between the pink “Solid” section and the orange “Liquid” shows the conditions when the phases are in equilibrium. The lines also show how vapor pressure of water varies with temperature.

<span style="font-family: Arial,Helvetica,sans-serif; font-size: small; margin: 0in 0in 0pt;"> **Triple Point** – the point on a phase diagram that represents the only set of conditions at which all three phases exist in equilibrium with one another <span style="font-family: Arial,Helvetica,sans-serif; margin: 0in 0in 0pt;">For water, the triple is a temperature of 0.016°C and a pressure of 0.61 kPa (0.0060 atm).

<span style="font-family: Arial,Helvetica,sans-serif; font-size: 110%; margin: 0in 0in 0pt;">The phase diagram can also show that a decrease in pressure lowers the boiling point and raises the melting point. An increase in in pressure will raise the boiling point and lower the melting point.

<span style="font-family: Arial,Helvetica,sans-serif;"> <span style="font-family: Arial,Helvetica,sans-serif;">Phase Diagram - Courtney Gareau

<span style="color: #d73cd7; font-family: Arial,Helvetica,sans-serif; font-size: 200%; margin: 0in 0in 0pt;">Section 14.1 Properties of Gases <span style="font-family: Arial,Helvetica,sans-serif; margin: 0in 0in 0pt;">Heather Bowditch: pg. 413-414 <span style="font-family: Arial,Helvetica,sans-serif; margin: 0in 0in 0pt;">Adam Shanahan (coeditor): pg. 414-417

**Heather Bowditch pg 413-414** // __Compressibility__ // Compressibility--measure of how much volume of matter decreases under pressure Kinetic Theory explains that **gases are easily compressed because of the space between the particles in a gas.** Gases are compressed more easily than liquids and solids. Particles in a gas at STP have spaces between them of about 10 times the size of their diameter. media type="youtube" key="apOSDqZd6Fg" height="349" width="425" Video- Heather Bowditch
 * [[image:gas.gif width="147" height="158"]] ||
 * particles in a gas- Heather Bowditch ||

**Adam Shanahan (Coeditor) pg 414-417**
//__Factors Affecting Gas Pressure__// There are four variables used to describe a gas. They are pressure (P) in kilopascals, volume (V) in liters, temperature (T) in kelvins, and the 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. This is how aerosol cans work. Inside the areosol 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 whatever is inside the can (whipped cream, spray paint, ect.) through the opening. Picture- Adam Shanahan

Decreasing the amount of volume increases the amount of pressure. Pistons are used in vehicles to reduce the volume of a container and increase the pressure.

An increase in temperature increases the pressure. This is because it increases the everage 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 an aerosol can explodes in a fire.

<span style="color: #d73cd7; font-family: Arial,Helvetica,sans-serif; font-size: 200%; margin: 0in 0in 0pt;">Section 14.2 The Gas Laws <span style="font-family: Arial,Helvetica,sans-serif; margin: 0in 0in 0pt;">Christos Anastos: pg. 418-421 <span style="font-family: Arial,Helvetica,sans-serif; margin: 0in 0in 0pt;">Andrew Sciotti (coeditor): pg. 422-425

Ø <span style="font-family: Arial,Helvetica,sans-serif; font-size: 120%;">//__ Boyle's Law __// <span style="font-family: Arial,Helvetica,sans-serif; font-size: 120%; line-height: 115%;">· <span style="font-family: Arial,Helvetica,sans-serif; font-size: 110%;">** Boyles Law ** : if the temperature is constant, as the pressure of a gas increases, the volume decreases. o Pressure is inversely proportional to volume in a gas o the equation that shows Boyle's Law is P1 x V1 = P2 x V2 o in the equation pressure is represented by P and volume is represented by V   Ø <span style="font-family: Arial,Helvetica,sans-serif; font-size: 120%;">//__ Charles' Law __// //<span style="font-family: Arial,Helvetica,sans-serif; font-size: 120%;">· // <span style="font-family: Arial,Helvetica,sans-serif; font-size: 110%;">** Charles' Law ** **:** As the temperature of an enclosed gas increases, the volume increases if the pressure is constant o the volume of a gas directly proportional to its temperature o the equation that shows Charles' Law is (V1 / T1) = (V2 / T2) o in the equation temperature is represented by T and volume is represented by V o Temperature should be done in Kelvin Scale media type="youtube" key="ikr83oIJiOU" height="349" width="425" Video: Christos Anastos
 * <span style="font-family: Arial,Helvetica,sans-serif; font-size: 120%; margin: 0in 0in 0pt;">Christos Anastos- pg. 418-421 **

**<span style="font-family: 'Times New Roman',serif; line-height: 115%;">Andrew Sciotti (Coeditor)- Pg 422-425 ** __//<span style="font-family: 'Times New Roman',serif; line-height: 115%;">Gay-Lussac's law: Pressure and Temperature //__ v <span style="font-family: 'Times New Roman',serif; font-size: 12pt; line-height: 115%;">Gay-Lussac’s Law (Used to find pressure or temperature when a change in either occurs) Ø <span style="font-family: 'Times New Roman',serif; font-size: 12pt; line-height: 115%;">As the temperature increases the pressure increases Ø <span style="font-family: 'Times New Roman',serif; font-size: 12pt; line-height: 115%;">The volume must be constant for this to be correct as if the volume changes the pressure changes also, thereby changing the conditions Ø <span style="font-family: 'Times New Roman',serif; font-size: 12pt; line-height: 115%;">The formula for Gay-Lussac’s Law is P1/T1 = P2/T2 Ø <span style="font-family: 'Times New Roman',serif; font-size: 12pt; line-height: 115%;">The P is pressure (liters, milliliters, ect) and the T is temperature (in Kelvins ALWAYS)

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__//The Combined Gas Law//__ v <span style="font-family: 'Times New Roman',serif; font-size: 12pt; line-height: 115%;">Combined Gas Law (combines Boyle’s, Charles’s, and Gay- Lussac’s Laws.   Ø <span style="font-family: 'Times New Roman',serif; font-size: 12pt; line-height: 115%;">Only the amount of gas has to be constant    Ø <span style="font-family: 'Times New Roman',serif; font-size: 12pt; line-height: 115%;">[(P1 * V1) / T1] = [(P2 * V2) / T2]    Ø <span style="font-family: 'Times New Roman',serif; font-size: 12pt; line-height: 115%;">This formula describes the relation between pressure, volume, and temperature.    Ø <span style="font-family: 'Times New Roman',serif; font-size: 12pt; line-height: 115%;">A great way to manipulate the formula to help you find your answers is to isolate the variable you are looking for. So if you were looking for the ending pressure of a gas then you would do P2 = [(P1 * V1 * T2) / (T1* V2)



media type="youtube" key="bftkRnTcFj8" height="349" width="560" Video: Andrew Sciotti

<span style="color: #d73cd7; font-family: Arial,Helvetica,sans-serif; font-size: 200%; margin: 0in 0in 0pt;">Section 14.3 Ideal Gases <span style="font-family: Arial,Helvetica,sans-serif; margin: 0in 0in 0pt;">Evan Grandfield: pg. 426-427 <span style="font-family: Arial,Helvetica,sans-serif; margin: 0in 0in 0pt;">Maggie Bie (coeditor): pg. 428-429

<span style="font-family: Arial,Helvetica,sans-serif; margin: 0in 0in 0pt;">** Maggie Bie (Coeditor) - pg 428-429 ** //__ Ideal Gases and Real Gases __// <span style="font-family: Arial,Helvetica,sans-serif; font-size: 110%; margin: 0in 0in 0pt;">Ideal gas – a gas that follows the gas laws at all the conditions of pressure and temperature *Needs to conform precisely to the assumptions of kinetic theory *No volume *No attraction between particles in the gas <span style="font-family: Arial,Helvetica,sans-serif; font-size: 110%; margin: 0in 0in 0pt;">*No gas that fulfills this <span style="font-family: Arial,Helvetica,sans-serif; font-size: 110%; margin: 0in 0in 0pt;">-but at some conditions and pressures some real gases are very much like an ideal gas

<span style="font-family: Arial,Helvetica,sans-serif; font-size: 110%; margin: 0in 0in 0pt;">Real gas- gases that have volume with attractions between the particles <span style="font-family: Arial,Helvetica,sans-serif; font-size: 110%; margin: 0in 0in 0pt;">-Can condense or solidify if compressed or cooled <span style="font-family: Arial,Helvetica,sans-serif; font-size: 110%; margin: 0in 0in 0pt;">Temperature and pressures for gases to condense varies with each individual gas

<span style="font-family: Arial,Helvetica,sans-serif; font-size: 110%; margin: 0in 0in 0pt;">Key Concept: Real gases differ most from an ideal gas at low temperatures and high pressures

<span style="font-family: Arial,Helvetica,sans-serif; font-size: 110%; margin: 0in 0in 0pt;">The ratio PV/ nRT changes as pressure increases *In ideal gas, when graphed, the line is flat *If ratio is greater than 1, the line curves upward *If ratio is less than 1, the line curves downward <span style="font-family: Arial,Helvetica,sans-serif; font-size: 110%; margin: 0in 0in 0pt;">Why? When attractive forces force distance between particles, gas occupies less volume than expected, causing to be below one. Actual volume of molecules causes it to be greater than one

<span style="font-family: Arial,Helvetica,sans-serif; font-size: 110%; margin: 0in 0in 0pt;">In portions of the curves below the ideal gas line, intermolecular attractions dominate <span style="font-family: Arial,Helvetica,sans-serif; font-size: 110%; margin: 0in 0in 0pt;">In portions of the curves above the ideal gas line, molecular attractions dominate

<span style="font-family: 'Times New Roman','serif'; font-size: 110%; line-height: 0px; margin: 0in 0in 0pt; overflow: hidden;">media type="youtube" key="wVVjaJmMWoQ" height="349" width="560"

<span style="font-family: 'Times New Roman','serif'; font-size: 110%; line-height: 0px; margin: 0in 0in 0pt; overflow: hidden;">Video- Maggie Bie

__//Section 3 Ideal Gases//__ Vocab 1. **ideal gas constant**-the constant in the ideal gas law with the symbol R and the value 8.31(L x kPa)/(K x mol) 2. **ideal gas law**-the [|relationship] PV=nRT, which describes the behavior of an ideal gas
 * Evan Grandfield- Pg. 426-429 **

Notes I. Ideal Gas Law A. The combined gas law solves problems with 3 variables: 1 pressure, 2 volume, and 3 temperature. 1. However, the combined gas law makes the assumption that the amount of gas does not vary. 2. For this reason, you CANNOT use the combined gas law to find number of moles of a gas in a fixed volume at a known temperature and pressure. 3. To calculate the number of moles of a contained gas requires an expression that contains the variable “n.” 4. The # of moles of gas is directly proportional to the number of particles. a. Volume occupied by a gas at a specified temperature and pressure also depends on the # of particles. 5. Hence, moles are DIRECTLY proportional to volume. 6. Here is how the combined gas law looks when we divide each side by “n.” (P1 X V1)(T1 x n1)=(P2 x V2)(T2 x n2)

7. PV/Tn is constant. 8. Given P, V,T, and n for 1 set of conditions, you can find a value for the constant. 9. The ideal gas constant is called R. 10. This is identical to the lab we did.

R=PV/Tn= (101.3kPa x 22.4L)/(273.15K x 1mol) = 8.31(L x kPa)/(K x mol)

11. The ideal gas law includes variables P,V,T, and n. II. Ideal Gases and Real Gases A. Ideal Gas-a gas that follows the gas laws at all conditions of pressure and temperature. Ideal gases would follow kinetic theory exactly. 1. In a hypothetical ideal gas, the particles would have ABSOLUTELY NO volume, and ABSOLUTELY NO attraction between particles in the gas. 2. As with all things of the world, perfection does not exist in gases; consequently, ideal gases do not exist. a. [|Lucky] for us however, at a number of contingencies of temperature and pressure, real gases can act like ideal gases! 3. Contrasting the ideal gas, particles in real gases do have volume and there are attractions between particles. a. These attributes allow gases to condense into liquid and solidify when cooled to an even lower temperature. B. Real gases differ most from an ideal gas at low temperatures and high pressures. 1. The value of the ratio PV/nRT changes as pressure increases for real gases. a. An ideal gas would always keep the ratio equal to 1.

[|http://www.chem.ufl.edu/~itl/2045/lectures/lec_e.html]

b. When the ratio of PV/nRT is greater than 1, the curve rises above the ideal gas line. When the ratio PV/nRT is less than 1, the curve drops below the line. C. 2 [|Factors] that explain these deviations! 1. Ratio less than 1: As attractive forces lessen the distance between a real gas’s particles, a real gas occupies less volume than expected. 2. Ratio greater than 1: Actual volume of the molecules. D. When the line curves below the line, intermolecular attractions dominate. 1. More kinetic energy allows molecules to overcome intermolecular attractions. E. When the lines curves above the line, molecular volume dominates. media type="youtube" key="_dpnZ-Icd-k" height="349" width="425" Video: Made by Evan Grandfield

<span style="color: #d73cd7; font-family: Arial,Helvetica,sans-serif; font-size: 200%; margin: 0in 0in 0pt;">Section 14.4 Gases: Mixtures and Movements <span style="font-family: Arial,Helvetica,sans-serif; margin: 0in 0in 0pt;">Nick Brault: pg. 432-434 <span style="font-family: Arial,Helvetica,sans-serif; margin: 0in 0in 0pt;">Katherine Perry (coeditor): pg. 435-437

**<span style="font-family: Arial,Helvetica,sans-serif; margin: 0in 0in 0pt;">Nick Brault- pg. 432-434 ** //__<span style="font-family: 'Arial Bold','serif';">Dalton's Law Of Partial Pressure __// <span style="font-family: 'Times','serif';">• <span style="font-family: 'Arial','sans-serif';">This law relates to the pressures of each gas in a mixture. <span style="font-family: Times,serif; margin: 0in 0in 0pt;">P(total)= P(1) + P(2) + P(3) +.... etc. <span style="font-family: Times,serif; margin: 0in 0in 0pt 0.5in; text-indent: -25pt;">• Each P except for the total represents the "partial pressure" or contribution of each gas in the mixture in the container <span style="font-family: Times,serif; margin: 0in 0in 0pt 0.5in; text-indent: -25pt;">• It is useful for calculating the pressure of gases that were collected over water. <span style="font-family: Times,serif; margin: 0in 0in 0pt;">Sample Problem: <span style="font-family: Times,serif; margin: 0in 0in 0pt;">In a mixture of oxygen, nitrogen, and hydrogen, oxygen has a partial pressure of 2 atm, hydrogen 1 atrm, and nitrogen 3 atm. What is the total pressure of the system? <span style="font-family: Times,serif; margin: 0in 0in 0pt;">In this problem, P(total)= P(O2) + P(N2) + P(H2), so the equation would be P(total)= 2 atm + 1atm + 3atm. The answer would be 6 atm//.//

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**Katherine Perry (Coeditor) - pg. 435-437** //__Graham's Law__// Diffusion is the tendency of molecules to move toward areas of lower concentration until the concentration is uniform throughout. Effusion occurs when a gas escapes through a tiny hole in its container.

With effusion and diffusion, the type of particle is important: Key Concept: Gases of lower molar mass diffuse and effuse faster than gases of higher molar mass.

Thomas Graham formed Graham's law of effusion: Graham's law of effusion states that the rate of effusion of a gas is inversely proportional to the square root of the gas's molar mass. This law can also be applied to the diffusion of gas.

So: __Rate A__ = <span style="font-family: 'Times New Roman',serif; font-size: 110%; margin: 0in 0in 0pt;">__√____ molar mass B __ Rate B <span style="font-family: Times New Roman,serif;">√ <span style="display: inline !important; margin-bottom: 0in; text-decoration: overline;">molar mass A

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