Chapter 8: The Gas Phase Flashcards Preview

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Flashcards in Chapter 8: The Gas Phase Deck (25):

4 variables which define the state of a gaseous sample?

Pressure, volume, temperature and moles


Relationship between different units of gas pressure?

1 atm = 760mm Hg = 760 torr = 101.325 kPa


Characteristics which make the gas phase unique?

Compressible fluids, rapid molecular motion, large intermolecular distances, weak intermolecular forces


What is an ideal gas? When do real gases deviate from this?

Hypothetical gas with molecules which have no intermolecular forces and occupy no volume. At high pressures and low temperatures


Ideal gas law (formula)?

PV = nRT, -> pressure, volume, moles of gas, ideal gas constant (0.0821 L*atm/mol*K or 8.314J/K*mol), temp in K


Ideal gas law rearranged to find density of a gas?

density = m/v and PV = m/MRT, so density = PM/RT


How much volume does a mole of ideal gas at STP occupy? What can this be used for?

22.4L. Calculating the density by dividing the mass by the predicted volume


Combined gas law?

P1V1/T1 = P2V2/T2, where state one is at one set of conditions (eg. STP) and state two is at another


How can molar mass be calculated from combined gas law?

determine V2 (at STP) by making the formula V2= V1(P1/P2(1atm))(T2(273K)/T1), then find density by dividing mass of sample by volume at STP. Finally, multiply density of sample by V2 at STP


Avagadro's principle (formula)?

n/V = k, where n is moles of gas, V is volume and k is a constant or n1/V1 = n2/V2


Boyle's Law? When is this true?

P1V1 = P2V2. Only in isothermal conditions


What is Charles's Law? When is this true?

V1/T1 = V2/T2. Only in isobaric conditions


What is Gay-Lussac's Law? When is it true?

P1/T1 = P2/T2. Only in isovolumetric conditions


Dalton's Law of partial pressures?

When 2 or more gases that do not chemically interact are found in one vessel, each will behave independently of the others. Total pressure (Pt) = Pa + Pb + Pc etc. where each are a partial pressure of different gases


What is partial pressure related to?

Mole fraction, as Pa = XaPt, where Xa = (moles of gas A/total moles of gas)


What is vapor pressure? Henry's Law to calculate it?

Pressure exerted by evaporated particles above the surface of a liquid.
[A] = Kh * Pa or [A]1/P1 = [A]2/P2 = Kh

where [A] is the concentration of A in solution, Kh is Henry's constant, and Pa is the partial pressure of A


5 assumptions made by the kinetic molecular theory?

1. Gases made of particles w volumes that are negligible compared to container volume
2. Gas atoms/molecules have no intermolecular forces
3.Gas particles re in continuous random motion and undergo collisions with walls and other particles
4.Collisions of particles and w walls are elastic - conservation of momentum and kinetic energy
5. Average kinetic energy of gas particles is proportional to temp of the gas in kelvin and is the same for any gas at a given temp, irrespective of atomic mass


Equation for the kinetic energy of a gas particle?

KE = 1/2mv^2 = 3/2KbT, where Kb is the Boltzmann constant (1.38x10^-23) and T is temp


How can the average speed of a gas particle be determined?

Find the root-mean-square speed = sqrt(3RT/M),
where R is the ideal gas constant(8.314 in this case), T is temp in K and M is molar mass in kg


What does a Maxwell-Boltzmann distribution curve show as temperature increases?

Number of molecules on Y, molecular speed on X. Bell curve flattens out and shifts right as T increases


Graham's Law of diffusion?

Under isothermal and isobaric conditions, rates at which 2 gases diffuse are inversely proportional to the sqrt of their molar masses. r1/r2 = sqrt(m1/M2), where r1 and 2 are the diffusion rates (can also use average speed)


What is effusion? Graham's law of effusion?

The flow of gas particles under pressure from one compartment to another through a small opening. 2 gases at the same temp have rates of effusion proportional to the average speeds


Why do real gases deviate from ideal gases at high pressure (low volume) or low temperature?

Intermolecular forces and particle volumes become significant. Gas volume is larger than expected at moderately high pressure (due to attraction). At extremely high pressure, particle volume causes the gas volume to be larger than expected


Deviations in gas volume when temp approaches boiling point/condensation point?

Intermolecular attraction causes gas to have a smaller volume than what is predicted from the ideal gas law


Van der waals equation of state?

Used to correct for deviations from ideal gas law.
P = (nRT/V-nb) - (n^2a/V^2), where a and b are physical constants determined experimentally, where a corrects for attractive forces (larger for larger and more polar gases) and b corrects for volume of molecules (larger for larger molecules)