Chapter #5: Gases5.2: Pressure- The Result of Molecular Collisions- Force from the constant collision of molecules or atoms of a given area in gas creates pressure.- Pressure depends on concentration. Higher the concentration, higher the pressure.- More molecules equal more collisions and more force.- Pressure decrease with increasing altitude.Pressure Units- Pressure can be measured in millimeter of mercury (mmHg).- Barometer measures pressure.- mmHg can also be called torr after physicist Torricelli who invented the barometer.- Atmosphere is another unit for pressure.- 1 atm = 760 mmHg- 1 atm = 101325 Pa (pascal)- 1 atm - 29.92 inHg (inches of mercury)- 1 atm = 14.7 psi (pounds per square inch)Manometer- U-shaped tube with one side open to atmospheric pressure and other side with gas sample.- Measures pressure relative to atmospheric pressure.5.3: Simple Gas Laws- Boyle’s Law, Charles’s Law, and Avogadro’s LawBoyle’s Law- Volume and Pressure- There is an inverse relationship between volume and pressure.- Increase in one cause a decrease in other.- Volume decrease, same number of gas particles in smaller space has more collisions and thus increase pressure.- Temperature and amount of gas must remain constant.Charles’s Law- Volume and Temperature- Pressure kept constant.- Volume increase as temperature increase.- If linear relationship graph of volume and temperature is extrapolated, gas have zero volume at 0 K. This is coldest possible temperature and there can be no negative volume.- Increase in temperature result in increase in energy of particles and more collision, but for the pressure to remain the same, volume increases to accommodate the collisions.Avogadro’s Law- Volume and Moles- At constant temperature and pressure.- As concentration increase, volume increases.- More gas particles needs more space to occupy.5.4: Ideal Gas Law- Combine Boyle’s, Charles’s. and Avogadro’s Law to create ideal gas law.V ∝nTP- Replace proportionality sign with ideal gas constant to create equation: PV=nRT.- Ideal gas law for an ideal gas.- R= 0.08206 L*atm/mol*K (ideal gas constant)- Gay-Lussac’s law is as temperature increase, pressure also increases.- For ideal gas law, pressure (atm), volume (L), moles (mol), and temperature (K).5.5: Applications of Ideal Gas LawMolar Volume at STP- Molar volume is volume occupied by one mole.- Standard temperature and pressure is 273K and 1 atm.- Under STP a gas is 22.4L/1 mol.Density of Gas- Density (g/L) = Molar Mass / Molar Volume- Rearrange ideal gas law to n/V = P/RT to calculate density.- Density = PM/RTMolar Mass of Gas- Find moles using ideal gas law and divide mass of sample by moles to fine molar mass.5.6: Mixtures of Gases and Partial Pressures- Many gas samples are mixtures of different gases but ideal gas law assumes the different gas molecules do not interact.- Partial pressure is pressure of individual component in mixture.- Assume each gas component acts independently and calculate partial pressure of each component using ideal gas law and number of moles of component in mixture.- Dalton’s law of partial pressures states the sum of partial pressures equal total pressure.- All particles have the same average kinetic energy at a given temperature and exert the same force.- Mole fraction = mole of component/ total moles- Mole fraction of a component times total pressure equals partial pressure of component.5.7: Gases in Chemical Reactions- Use stoichiometry to calculate pressure, volume, or temperature of a product or reactant in a chemical equation.- Given mass or moles, convert to moles of another reactant or product.- Then use the ideal gas law to solve for pressure, volume, or temperature.Molar Volume and Stoichiometry- 1 mol of gas at STP equals 22.4L so use this conversion to convert between volume and grams or moles.5.8: Kinetic Molecular Theory- Kinetic molecular theory models gases as particles in constant motion and moves in a straight line until it a collision occurs.- Kinetic molecular theory assumptions include:- Size of a particle is negligibly small. - Kinetic energy is proportional to temperature (K).- Collision of particles is completely elastic.- Kinetic molecular theory assumption valid for ideal gas law.- Pressure = Force / Area ; Force = mass * acceleration- Ideal gas law is derived from kinetic molecular theory.Temperature and Molecular Velocities- Kinetic Energy = 0.5(mass)(velocity)^2- Of a gas mixture at given temperature, on average lighter particles travel fast than heavier particles.-Root mean square velocity = 3RTM. (R = 8.314 J/molK; M (kg/mol); T (K))- Root mean square velocity is an average velocity and some particles move faster or slower.- As velocity increases, velocity distribution becomes broader.- Increase in temperature or light particles create broader distribution.5.9: Mean Free Path, Diffusion, and Effusion of Gases- Gaseous particles travel at high speed but haphazard paths. Only travels short distance before it collides and changes direction.- Mean free path is the average distance a molecule travel between collision.- Mean free path increase with decreasing pressure.- Diffusion is when gas molecules spread out to concentration gradient.- Lighter molecules diffuse faster than heavier molecules because it has higher velocity.- Effusion is when gas escapes container through small hole.- Heavier molecules effuse more slowly than lighter ones.rate ∝1M- Graham’s law of effusion is the ratio of effusion rates of two different gases.rateArateB=MBMA5.10: Real Gases- Size and Intermolecular Forces- Most gases act very nearly to ideal gases if:- Volume of gas particles is smaller than space between them.- Forces between has particles are not significant.Effect of Finite Volume of Gas Particles- At low pressure, molar volume of a gas is identical to that of an ideal gas but as pressure increases, molar volume of gas becomes greater than ideal gas.- At high pressure, volume of atom occupy significant portion of gas volume making actual volume greater that ideal gas law prediction.- Ideal gas law predicts a volume that is too small.- Van der Waals suggest adding nb (n= moles; b= constant depending on gas) to account for volume of gas particle.Effect of Intermolecular Forces- Intermolecular forces do not matter much at high temperature and low pressure because attraction is not felt strongly.- Attraction can potentially affect collisions.- Temperature decrease, pressure