Intro to Gases Pick Up a New Unit Packet Write down the following Essential Question: How are Pressure, Temperature, and Volume related and calculated ?

First, remember the 3 states of matter… Particles in a SOLID: tightly packed, vibrate about a FIXED position Particles in a LIQUID: loosely packed, slide past one another (fluidity), no set shape Particles in a GAS: particles move about most empty space, no set shape or volume

Kinetic Molecular Theory (5 parts) 1.Gas particles are much smaller than the distance between particles, therefore, most of the volume of a gas is empty space. –Gas particles have LOW density (D = m/v) Solid > Liquids > Gas –Gasses can be compressed Compression – can be pushed into a smaller volume Expansion – can be pulled into a greater volume

2.The particles are in constant motion. The collisions of the particles with the walls of the container are the cause of the pressure exerted by the gas. –Particles move in a straight line until the collide with each other or the walls of the container Kinetic Molecular Theory Cont.

3.There are no attractive or repulsive forces between the molecules. –“Ideal” gases would NOT liquefy because the particles wouldn’t attract each other. –Usually the gas particles are to far apart, except when polar molecules are under high pressure & low temperature Ex. H 2 O (g) and H 2 O (l) Kinetic Molecular Theory Cont.

4.The average kinetic energy of a collection of gas particles is related to the temperature of the gas. –Kinetic energy: energy due to motion –KE = ½ mv 2 Mass & Velocity –Temperature: measure of the average KE 5.Gas have Elastic Collision: no Kinetic Energy is lost when particles collide –Energy is never lost b/c particles don’t encounter friction Kinetic Molecular Theory Cont.

Temp: Absolute Zero There is a theoretical temperature at which the motion of the atoms/molecules stops, absolute zero! –Occurs at -273 ˚C or 0 Kelvin –Kelvin Scale (K): direct measure of avg. KE K = ˚C

Gas Movement Remember: KE = ½ mv2 –Big molecules move slower than small molecules Gases move from areas of high to low concentration! –Effusion: gas escapes thru tiny holes Ex: hole in tire or He vs Air in Balloon –Diffusion: gas A moving thru gas B Ex: perfume being sprayed or smell of a rotten egg

Pressure Pressure is defined as the force the gas exerts on a given area of the container in which it is contained. To increase Pressure: 1.Increase # moles –Add more gas particles 2.Decrease Volume –Decrease the size of the container 3.Increase Temperature –Increase speed thus increasing collisions

Pressure Units SI unit of pressure is the Pascal (Pa). –Pressure is typically measured in kPa (kilo) Alternative units: –psi: pounds per square inch, ex: tires –atm: atmospheric pressure, measured with a barometer –mmHg: millimeters of mercury Conversions –101.3 kPa = 14.7 psi = 1 atm = 760 mmHg STP: s tandard temp and pressure –0˚C or 273 K and 1 atm or 760 mm

Pressure Conversion Problems 1.Convert 190 mmHg to atm 2.Pressure on Mt. Everest is 4.89 psi. How many mmHg is that? atm? 3.Air pressure in Death Valley is 1.08 atm. How many mmHg? kPa?

Volume Volume is the 3-D space inside the container holding the gas. SI unit for volume is the cubic meter, m 3. More common and convenient unit is the liter, L. –Remember: if units are given in mL and need to be in L…..convert by moving the decimal 3 places to the left! For a gas: 1 mole = 22.4 L

Relationships Directly proportional – both variables will either increase or decrease together Inversely proportional – one variable will increase and the other variable will decrease

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Temperature & Pressure Temperature and pressure are directly proportional As one goes up, the other goes up As one goes down, the other goes down

Temperature & Volume Temperature and volume are directly proportional As one goes up, the other goes up As one goes down, the other goes down

Pressure & Volume Pressure and volume are Inversely proportional As one goes up, the other goes down As one goes down, the other goes up

Gas Law Calculations

Boyle’s Law –Robert Boyle –Studied relationship between pressure and volume when temperature was constant. – Determined they were inversely proportional. –P 1 V 1 = P 2 V 2

Charles’ Law –Jacques Charles – Studied relationship between volume and temperature when pressure is held constant – Determined they were directly proportional –V 1 T 2 = V 2 T 1

Gay-Lussac’s Law –Joseph Gay-Lussac –Studied relationship between pressure and temperature when volume is held constant –Determined they were directly proportional –P 1 T 2 = P 2 T 1

Combined Gas Law P = pressure, V = volume T = Temperature P 1 V 1 = P 2 V 2 or P 1 V 1 T 2 = P 2 V 2 T 1 T 1 T 2 This is a combination of all 3 gas laws, can be used with all variables or when holding one constant.

Dalton’s Law: Partial Pressure The pressure of a mixture of gases is equal to the sum of the pressures of all of the constituent gases alone. P = pressure and n = moles Pressure total = Pressure 1 + Pressure 2... Pressure 100 P x = P Total ( n x / n Total ) –P x = partial pressure of gas x –P Total = total pressure of all gases –n x = number of moles of gas x –n Total = number of moles of all gases

Ideal Gas Law P = pressure, V = volume T = Temperature N = # of moles, R = PV = nRT

Avogadro’s Law At a given temperature and pressure, equal volumes of gas contain equal numbers of moles. N = # of Moles, V = volume N 1 V 1 = N 2 V 2