Unit+7+Gas+Laws

=Unit 7: Gas Laws =

(images and summary below: T. Lynch) EDITO﻿R: TERESA LYNCH

The Gas Laws unit in our books is made up of Chapters 13 and 14. In this unit we will learn about the properties of gas and the laws pertaining to gases. Some of the laws that will be included in this unit are: Charles's Law, Graham's Law of Effusion, Ideal Gas Law, Boyle's Law, Combined Gas Law, Dalton's Law of Partial Pressure, and Gay-Lussac's Law. From these laws we will learn the characteristics of gases and how to find volume, pressure, and temperature (and other properties) of a gas in certain situations.

Group 1 – p.385-389 Nature of Gases Coeditor: Kayla Anghinetti Lindsey Chou

Kayla Anghinetti p. 385-386

I. Kinetic Theory and a Model for Gases A. Kinetic energy- the energy an object has because of its motion B. Kinetic Theory – says that all matter consist of tiny particles that are in constant motion C. Particles in a gas are considered to be small, hard spheres with an insignificant volume. 1. Within a gas particles are spread far apart D. The motion of the particles in a gas is rapid, constant, and random 1. Gas fills a container regardless of the shape and volume 2. Uncontained gas can spread out into space with no limit 3. Particles change direction only when they rebound from collisions 4. The aimless path a molecule takes is called a //random walk// E. All collisions between particles in a gas are perfectly elastic 1. During an elastic collision, kinetic energy is transferred without loss from one particle to another- total kinetic energy remains constant



II. Gas Pressure A. Gas pressure – result of simultaneous collisions of billions of rapidly moving particles in a gas with an object 1. Vacuum - empty space with no particles a. No particles = no collisions = no pressure B. Air exerts pressure on Earth because gravity holds the particles in air in Earth’s atmosphere C. Atmospheric pressure – results from the collisions of atoms and molecules in air with objects D. Barometer – device used to measure atmospheric pressure 1. Height of mercury column in the tube depends on the pressure exerted by particles in the air colliding with the surface of the mercury in the dish. 2. Atmospheric pressure depends on weather and altitude 3. SI unit of pressure is the pascal (Pa) 4. One standard atmosphere (atm) is the pressure required to support 760 mm of Hg in a barometer a. 1atm = 760mm Hg = 101.3k Pa



**Pages 387-389** **Lindsey Chou**

**Converting Between Units of Pressure** There are 3 different units of pressure: kPa, atm, and mmHg (torr) Conversion factors: 1 atm = 101.3 kPa = 760 torr

Steps to Convert: 1) List the knowns and unknowns 2) Solve for the unknowns 3) Check if the results make sense

Problems: 1. Convert 200 kPa to atmospheres. 2. Convert 31 torr to kPa.

(Answers: 1) 200 kPa = 1.97 atm, 2) 31 torr = 0.041 kPa)

**Average Kinetic Energy** · When a substance is heated its particles absorb energy. Some is stored as potential energy, and some speeds up the particles, which is kinetic energy. · Increasing kinetic energy = increasing temperature / decreasing kinetic energy = decreasing temperature · At the same temperature, all particles of a substance have the **same average kinetic energy** whether they are a solid, liquid or gas · For example, a t the same temperature, particles of liquid water, water vapor, and ice all have the same average kinetic energy (Lindsey Chou)

· Particles of different elements in different states also have the same average kinetic energy at the same temperature · Absolute zero or 0 K is the temperature at which particles would have no kinetic energy because they would have no motion. It has never been achieved in a lab.

**Average Kinetic Energy and Kelvin Temperature** · Kelvin temperature is directly proportional to the average kinetic energy of a substance’s particles. · For example, particles of helium gas at 200 K has twice the average kinetic energy and twice the Kelvin temperature as particles in helium gas at 100 K  · The effects of temperature on particle motion are more complicated in solids/liquids than in gases

Group Two: Nature of Liquids Co-editor: Jessen Foster Chris Kelly Jessen Foster Pages 390-392 Model for Liquids - the kinetic energy of liquids allows the particles to flow past each other - particles in liquids are attracted to one another - the interplay between the disruptive motions of particles in a liquid and the attractions among the particles determines the physical properties of liquids Evaporation - conversion of liquids to a gas is called vaporation - when it occurs at the surface of a non-boiling liquid, it is called evaporation - during evaporation, only those molecules with a certain minimum kinetic energy can escape from the surface of the liquid - heating a liquid adds kinetic energy, which increases the rate it evaporates Vapor Pressure -vapor pressure is the measure of the force exerted by a gas above a liquid - in a system at constant vapor pressure, a dynamic equilibrium exists between vapor and the liquid - the system is in equilibrium because the rate of evaporation of liquid equals the rate of condensation of vapor Vapor Pressure and Temperature Change - the increase of the temperature of a contained liquid increases the vapor pressure - the particles of the warmed liquid have increased kinetic energy, so more of the particles have the kinetic energy necessary to escape the surface of the liquid

Chris Kelly p. 393-395

Boiling Point -When a liquid is heated to a temperature at which particles throughout have enough kinetic energy to vaporize, the liquid begins to boil -bubbles form, rise, and escape into the air -boiling point (bp) is when the vapor pressure of the liquid is equal to the external pressure

Boiling Point and Pressure Changes -Liquids don’t always boil at the same temperature because of pressure changes -the temperature of a boiling liquid never rises above boiling point

Normal Boiling Point -Normal boiling point is the boiling point of a liquid at a pressure of 101.3 kPa

Here is a list of normal boiling points

Carbon disulfide (CS 2 ) - 46.0 Chloroform (CHCl 3 ) - 61.7 Methanol (CH 4 O) - 64.7 Tetrachloromethane (CCl 4 ) - 76.8 Ethanol (C 2 H 6 O) - 78.5 Water (H 2 O) - 100.0
 * Name and formula Boiling Point ( ** ° ** C) **

Group 3 Pg. 396-403 Allie Chabot- Coeditor Ellie Kawa

13.3 The Nature of Solids (Pg, 396-400) Allie Chabot

· The general properties of solids reflect the orderly arrangement of their particles and the fixed locations of their particles. · In most solids; the atoms, ions, and molecules are packed tightly together. · Melting point is the temperature at which a solid changes into a liquid. · In a crystal the particles are arranged in an orderly repeating three dimensional pattern called a crystal lattice. · The shape of a crystal reflects the arrangement of the particles within the solid. · The type of bonding that occurs between particles in crystals determines their melting points. · The shape of a crystal depends on the arrangement of the particles within it. (A ﻿llie Chabot) · The smallest group of particles within a crystal that retains the geometric shape of the crystal id known as a unit cell. · Some solid substances can exist in more than one form. · Allotropes are two or more different molecular forms of the same element in the same physical state. · Not all solids area crystalline in form; some solids are amorphous. · Amorphous solid lacks an ordered internal structure. · Examples of amorphous solids are glasses. A glass is a transparent fusion product of inorganic substances that have cooled to a rigid state without crystallizing.

[[image:http://www.hcc.mnscu.edu/chem/V.17/amorphous_solid.jpg width="223" height="158"]](AllieChabot)
**__Changes of State__** **__By: Ellie Kawa (401-403)__** __Sublimation__ · Sublimation - change of a substance from a solid to a vapor without passing through the liquid state o can occur because solids have a vapor pressure · sublimation occurs in solids with vapor pressures that exceed atmospheric pressure at or near room temperature o ex: iodine- violet-black solid changes into a purple vapor without passing through a liquid state · sublimation has many useful applications o freeze-dried coffee o solid carbon dioxide (dry ice) o solid air fresheners o organic chemists use sublimation to separate mixtures and to purify compounds

__Phase Diagrams__ · Phase Diagram - a graph that gives the conditions of temperature and pressure at which a substance exists as solid, liquid, and gas (vapor) · Triple Point - point on the diagram at which all three curves met o describes the only set of conditions at which all the three phases can exist in equilibrium with one another § ex: water- temperature of 0.016ºC and pressure of 0.61 kPa (0.0060 atm) o decrease in pressure lowers the boiling point and raises the melting point o an increase in pressure raises the boiling point and lowers the melting point § with an increase in pressure, water vapor begins to behave more like a solid and is no longer easily compressed § an increase in pressure affects the melting point of ice = 14. 1 Properties of Gases = = Co-editor: Elena Conroy = = Bryan Dextradeur =

Pages 413-414
=Compressibility =
 * Unlike a Solid or a Liquid, a Gas can expand to fill its container.
 * The reverse of this is also true: Gases are easily compressed, or squeezed into a smaller volume.
 * ** Compressibility ** is a measure of how much the volume of matter decreases under pressure.
 * An example of this is the airbags in a car. When a car experiences a sudden and large decrease in speed, a chemical reaction is triggered in airbags, causing nitrogen gas to be releases and the airbags to expand. A collision with an airbag is much less harmful than a collision with the steering wheel or dashboard because the impact causes the gas molecules in the inflated airbag closer together, and the compression of the gas absorbs the energy of the impact.
 * Kinetic Theory can explain why Gases are compressed more easily than liquids or solids.
 * **Gases are easily compressed because of the space between the particles in a gas.**
 * The volume of the particles in a Gas is small compared to the overall volume of the Gas.
 * As a result, the distance between the particles in a Gas is much greater than the distance between particles in a Liquid or Solid.
 * Under pressure, the particles in a Gas are forced closer together, or compressed.
 * At room temperature, the distance between particles in an enclosed Gas is about 10 times the diameter of the particle.

= Affecting Gas Pressure =
 * Kinetic Theory can help explain other properties of Gases, such as their ability to expand and take the shape and volume of their containers.
 * Gas particles move along straight-line paths until they collide with other particles or the walls of their container.
 * The motion of the particles is constant and random.
 * Because there are no significant forces of attraction or repulsion among particles in a Gas, particles in a Gas can move freely.
 * Four variables are generally used to to describe a Gas.
 * Pressure (//P//) in Kilopascals.
 * Volume (//V//) in Liters.
 * Temperature (//T//) in Kelvins.
 * Number of Moles (//n//).
 * **The amount of Gas, Volume, and Temperature are factors that affect Gas Pressure.**



Above Images: Bryan Dextradeur

Pages 415-417
**__ Amount of Gas __**
 * You can use the kinetic theory to predict and explain how gases will respond to a change of conditions
 * Adding gas to a container increases the number of particles which increases the number of collisions. The higher number of collisions among particles, the higher gas pressure
 * Doubling the particles of gas doubles the gas pressure
 * If the gas pressure exceeds the strength of the container, the container will burst
 * If the amount of gas is reduced, the gas pressure is reduced
 * If the pressure of the gas in a sealed container is lower than outside air pressure air will rush into a container eg. Vacuum packs
 * When the pressure of a gas in a sealed container is higher than the outside air pressure, the gas will flow out of the container when it is unsealed eg. Aerosol cans

Images: Elena Conroy **__ Volume of Gas __**

When a gas is heated, the temperature increases and the average kinetic energy of the particles in the gas increases. They will strike the edges of the container with force.

 * When the temperature of an enclosed gas decreases, the gas pressure decreases

**David Monti (Co-Editor)**


 * Neal McGovern**
 * pg. 418- 421**

Boyle’s Law – Pressure and Volume · According to Boyle’s Law, as the pressure of a gas increases at a constant temperature, the volume of the gas decreases · Boyle proposed this law in 1662 to describe the relationships between the two. · P1 x V1 = P2 x V2 Charles Law – Temperature and Volume · According to Charles’s Law, as the temperature of an enclosed gas increases at constant pressure, then the volume increases · In 1787, Jacques Charles discovered that the temperature of a fixed mass of gas has a direct relationship to the volume of the gas (Neal McGovern)

David Monti (422-425)

__Gay-Lussac’s Law: Pressure and Temperature__

· States: As the temperature of an enclosed gas increases, the pressure increases, if the volume is constant · In other words, this French Chemist’s law states that the pressure of a gas is directly proportional to the Kelvin temperature of that gas if the volume remains constant. This can be represented by the following equation: P/T = P/T So the initial pressure divided by the initial temperature is equal to the resulting pressure and the resulting temperature.



__The Combined Gas Law__ · Describes the relationship among 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 This can be represented by the following equation:

__P X V__ __P X V__ T = T



So the initial pressure multiplied by the initial volume and divided by the initial temperature is equal to the resulting pressure multiplied by the resulting volume divided by the resulting temperature. · Basically, the combined gas law incorporates the different components from all of the gas laws into one universal equation.

**Ideal Gas Law** - p426
===The combined gas law is used for solving problems that deal with the three variables pressure, volume, and temperature. This law is used when the amount of gas is assumed not to vary - not when you need to calculate how many moles of a gas in a fixed volume at a known temperature. To calculate the number of moles of a contained gas requires an expression that contains the variable n. Modify the combined gas law by dividing each side of the equation by n.===

The equation shows that (P x V)/(T x n) is a constant. This constant applies to all ideal gases (gases that conform to the gas laws).
===You can calculate the value for the constant if you know the values for P, V, T, and n for one set of conditions. Use STP - 1 mol occupies 22.4 L. Put the values of P, V, T, and n into (P x V)/(T x n). The value of R is: 8.31 (L x kPa)/(K x mol)===

To find moles or to start a mole-mass conversion, isolate n by rearranging the equation: n = (P x V)/(R x T)

 * Hannah Kumlin**
 * Pg. 428 - 429**

-An ideal gas is one that follows the gas laws at all conditions of pressure and temperature. -In order for a gas to be an ideal gas, it would have to conform precisely to the kinetic theory. To do so, the gas's particles would have no volume and there would be no attraction between particles in the gas. However, there is no gas that fits this criteria and therefore an ideal gas does not exist. -Real gases behave like an ideal gas at many conditions of temperature and pressure. -A gas can condense or solidify when it is compressed or cooled. For example: when water vapor is cooled below 100 degrees Celsius, it condenses to a liquid. Hannah Kumlin
 * Ideal Gases and Real Gases**

The above chart shows how the value of the ratio PV/nRT changes as pressure increases. -For an ideal gas, the result is a horizontal line because the ratio is always equal to 1. -For real gases at a high pressure, the ratio can deviate/depart from the ideal. This is because as attractive forces reduce the distance between particles, the gas occupies less volume than expected, causing the ratio to be less than 1. The actual volume of the molecules causes the ratio to be greater than 1.
 * When the ratio is greater than 1 the curve rises above the ideal gas line. When the ratio is less than 1 the curve drops below the line.
 * Portions of curves below the line - intermolecular attractions dominate.
 * Portions of curves above the line - molecular volume dominates.

Shannon Leavey(co-editor), Lauren Murphy, Nicole Sheehan

Comparing Effusion Rates Shannon Leavey, pg 436

Graham's law: rateA/rateB = the square root of molar massB/ molar mass A



images:Shannon Leavey

Lauren Murphy Pages 432-433

Dalton’s Law · Total pressure exerted by a gaseous mixture is equal to the sum of the partial pressures of each individual component in a gas mixture

Ptotal= P1 + P2 + P3 + …… · The contribution each gas makes is called partial pressure · If the percent composition of a mix of gases doesn’t change, the fraction of the pressure exerted by a gas doesn’t change as the total pressure changes

<span style="color: #000000; font-family: arial,helvetica,sans-serif; text-indent: -0.25in;">Graham’s Law <span style="color: #000000; font-family: arial,helvetica,sans-serif; text-indent: -0.25in;">Nikki Sheehan (435) <span style="color: #000000; font-family: arial,helvetica,sans-serif; text-indent: -0.25in;">Diffusion – the tendency of molecules to move toward areas of lower concentration until it is uniform throughout. <span style="color: #000000; font-family: arial,helvetica,sans-serif; text-indent: -0.25in;">Effusion – the process of gas escaping from a small hole in its container. <span style="color: #000000; font-family: arial,helvetica,sans-serif; text-indent: -0.25in;">-Both processes involve the movement of molecules in a gas. <span style="color: #000000; font-family: arial,helvetica,sans-serif; text-indent: -0.25in;">-Graham’s Law of effusion – the rate of effusion of a gas is inversely proportional to the square root of a gas’s molar mass. (This also applies to diffusion) <span style="color: #000000; font-family: arial,helvetica,sans-serif; text-indent: -0.25in;">-This law makes sense if you know how the mass, velocity, and kinetic energy are related (KE=1/2mv2) <span style="color: #000000; font-family: arial,helvetica,sans-serif; text-indent: -0.25in;">-If the kinetic energy of a moving gas is constant, an increase in mass must be balanced by a decrease in velocity. <span style="color: #000000; font-family: arial,helvetica,sans-serif; text-indent: -0.25in;">-Gases of lower molar mass diffuse and effuse faster than gases of higher molar mass. <span style="color: #000000; font-family: arial,helvetica,sans-serif; text-indent: -0.25in;">-For example, if a 2g ball and a 1g ball with the same kinetic energy are traveling, the 1g ball will move faster than the 2g ball. <span style="color: #000000; font-family: arial,helvetica,sans-serif; text-indent: -0.25in;">-In the same way, gases with different masses and the same kinetic energies move at different rates. <span style="color: #000000; font-family: arial,helvetica,sans-serif; text-indent: -0.25in;"> <span style="color: #000000; font-family: arial,helvetica,sans-serif; text-indent: -0.25in;">(images Nikki Sheehan)