Warm-up #6 2. The balloon in the previous problem will burst if its volume reaches 400.0 L. Given the initial conditions specified in that problem, determine at what temperature, in degrees Celsius, the balloon will burst if its pressure at that bursting point is 0.475 atm. Identify what you know V1=250.0L T1= 22°C =273 + 22= 295K P1= 740.0 mm Hg V2=400.0L T2= ? P2= .475 atm = .475 atm x 760 mmHg P2 =361 mmHg 1 atm Solve for what you don’t know T2= 230 K or -43 C T2= T1 x P2V2 P1V1 T2=230.3 K

Review Homework See overhead

Ideal Gases & Ideal Gas Law

Ideal Gas Law- not in notes When dealing with gas, an equation was used to relate all of the factors needed in order to solve a gas problem. This equation is known as the Ideal Gas Equation. However, anything ideal does not exist and 2 assumption were made: 1) the gas particles have no forces acting among them 2) the particles do not take up any space (volume is ignored) An ideal gas is a hypothetical gas dreamed by chemists because it is be easier if things like intermolecular forces do not exist to complicate the simple Ideal Gas Law. This definition of an ideal gas contrasts with the Non-Ideal Gas definition, which describes how gases actually behaves in reality. For now, let us focus on the Ideal Gas laws.

No kinetic energy is lost in elastic collisions Ideal Gases Ideal gases are imaginary gases that perfectly fit all of the assumptions of the kinetic molecular theory. Gases consist of tiny particles that are far apart relative to their size. Collisions between gas particles and between particles and the walls of the container are elastic collisions No kinetic energy is lost in elastic collisions

Ideal Gases (continued) Gas particles are in constant, rapid motion. They therefore possess kinetic energy, the energy of motion There are no forces of attraction between gas particles The average kinetic energy of gas particles depends on temperature, not on the identity of the particle.

Real Gases Do Not Behave Ideally Real gases DO experience inter-molecular attractions Real gases DO have volume Real gases DO NOT have elastic collisions

Deviations from Ideal Behavior Likely to behave nearly ideally Gases at high temperature and low pressure Small non-polar gas molecules Likely not to behave ideally Gases at low temperature and high pressure Large, polar gas molecules

Variables that Describe a Gas Pressure (atm, mmHg, kPa) Volume (L) Temperature (˚C or K) Number of moles (mol)

The Ideal Gas Law Allows us to solve for a property of an ideal gas when properties are constant! PV=nRT P=pressure V=volume (Liters!!) T=temperature (Kelvin!!) n=number of moles R= 8.31 L x kPa or 0.08206 L x atm K x mol K x mol

Lets Practice Determine the volume occupied by 0.582 mol of a gas at 15˚C if the pressure is 81.1 kPa. Identify what you know n=0.582 mol T=15˚C + 273 = 288 K P=81.1 kPa. R= 8.31 L x kPa / K x mol V=?

PV=nRT 81.1 Kpa x V=(0.582 mol) (8.31 L x kPa / K x mol) (288K) 81.1 KPa x V=1392.89 L x kPa V=17.2 L

Two rules! Temp= Kelvin, Volume =Liters, and pressure= atm or kPa If you are given grams convert into moles. 36 grams of H2O x 1 mole H2O =2 mol H2O 18 g H2O

Dalton’s Law of Partial Pressures the total pressure of a mixture of gases is equal to the sum of the partial pressures of all the gases present. (at constant Temp & Volume) Formula: Ptotal= P1 + P2 + P3 + P4…

Let’s Practice 1) What is the total pressure of a gas mixture if it contains CO2 at 40.8 torr and O2 at 1009.9 torr & H2 at 791.4 torr? 2) What is the pressure of Ne gas if the total pressure of the gas is 100.6 atm, and the mixture contains 40.4 atm of He and 22.6 atm of HCl? PT=P1 +P2+P3 PT= 40.8 + 1009.9 + 791.4 = 1841.2 torr PT=P1 +P2+P3 100.6 atm = PNe + 40.4 atm + 22.6 atm = 37.6 atm

Water Displacement Gases produced in the laboratory are often collected over water. Thus, the gas is not pure but is always mixed with water vapor. The collection process makes it so the total pressure inside the bottle would be the same as the atmospheric pressure. Patm=Pgas + PH2O

Lets practice – (Hw check: pg 367 #1) 3) Some hydrogen gas is collected over water at 20.0 °C. The levels of water inside and outside the gas-collection bottle are the same. The partial pressure of hydrogen is 742.5 torr (mmHg). What is the barometric pressure at the time the gas is collected? Formula:Patm=Pgas + PH2O (to find the P use table A-8 pg. 859) Patm= 742.5 torr + 17.5 torr= 760.0 torr

Lets Practice Continued…. 4. Some CO2 gas is collected over water at 15.0 °C. The levels of water inside and outside the gas-collection bottle are the same. The barometric pressure at the time the gas is collected is 780 mmHg. What is the partial pressure of CO2? Formula: Patm=Pgas + PH2O 780 mmHg = PCO2 + 12.8 mmHg= 767.2mmHg

Dalton’s Law- not on notesheet Example #1: What is the partial pressure of oxygen at 101.3kPa of total pressure if the partial pressure of nitrogen is 79.10 kPa and carbon dioxide is 0.040kPa?

Solution: Ptotal= Po+ PN2 + PCO2 Po = Ptotal – (PN2 + PCO2 ) Po = 101.30kPa – (79.10kPa + 0.040 kPa) =22.16 kPa

Dalton Example #2 What is the partial pressure of each gas if you have a mixture of 2.0 moles of O2 and 2.0 moles of CO2 that exert at total pressure of 700 torr? O2= 2.0 mol O2 x 700 torr = 350 torr 4.0 total mol CO2= 2.0 mol CO2 x 700 torr = 350 torr 4.0 total mol

Now you try! Show all your work on worksheet-pg 16 Ideal Gas Law: PV=nRT Make sure: 1) Temp is in Kelvin (C+273= K) 2) Volume is in L (1000mL=1L) 350 mL x 1 L . 3) Pressure is in atm or kPa 1000 mL 4) 101.3Kpa=1 atm=760 mmHg=760 torr 5) grams-> mol (use molar mass) 36 grams of H2O x 1 mole H2O =2 mol H2O 18 g H2O

When to use the equations If the conditions of the gas change Then use the combined gas law If the conditions of the gas are fixed Then use the ideal gas law