The Gas Phase

General Chemistry

Atomic Structure
The Periodic Table
Bonding and Chemical Interactions
Compounds and Stoichiometry
Chemical Kinetics
Equilibrium
Thermochemistry
The Gas Phase
Solutions
Acids and Bases
Oxidation-Reduction Reactions
Electrochemistry
High-Yield Cheat Sheet

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MCAT Books › General Chemistry › The Gas Phase

Chapter 8: The Gas Phase

5 min read Updated Mar 26, 2026

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🎯 Diagnostic: Test Your Starting Level 24 questions (2 per section). No prior reading required, see what you already know.

Aim to answer every question before checking. Missed questions point you to the sections you need most.

1. (8.1) Gases differ from liquids and solids in that they:

A) Have a defined shape
B) Have strong intermolecular forces
C) Fill any container completely and are highly compressible
D) Have high density

C. Gas molecules move with significant free space between them. Density is tiny and the space is easy to compress.

2. (8.1) Gases behave most ideally at:

A) Low pressure and high temperature
B) High pressure and low temperature
C) High pressure and high temperature
D) Near the boiling point

A. Low P keeps molecules far apart (molecular volume negligible); high T keeps them moving fast (attractive forces negligible).

3. (8.2) Standard temperature and pressure (historical IUPAC) is:

A) 298 K, 1 atm
B) 0 K, 0 atm
C) 100 °C, 0 atm
D) 0 °C (273.15 K) and 1 atm

D. On the MCAT, treat STP as 273 K and 1 atm unless told otherwise.

4. (8.2) The molar volume of an ideal gas at STP is approximately:

A) 1 L
B) 22.4 L
C) 100 L
D) 0 L

B. Worth memorizing: 1 mol of any ideal gas at STP occupies 22.4 L.

5. (8.3) Boyle's law (constant T, n) says:

A) V rises linearly with T
B) V × T is constant
C) P and V are inversely proportional; PV = constant
D) P is proportional to V

C. The earliest gas law (1662). Squeeze a gas and it resists by increasing pressure.

6. (8.3) If a gas's pressure is doubled at constant T:

A) Volume halves
B) Volume doubles
C) Volume is unchanged
D) Volume quadruples

A. P₁V₁ = P₂V₂; doubling P means V must halve.

7. (8.4) Charles's law (constant P, n) says:

A) V is inversely proportional to T
B) P is proportional to V
C) V stays constant
D) V is proportional to absolute T; V/T = constant

D. Always use absolute temperature (K). Heating a balloon expands it; cooling shrinks it.

8. (8.4) Heating a gas at constant pressure causes the volume to:

A) Decrease
B) Increase in proportion to absolute temperature
C) Stay constant
D) Halve

B. Doubling T doubles V at constant P.

9. (8.5) Gay-Lussac's law (constant V, n) says:

A) V ∝ T only
B) P ∝ 1/T
C) P is proportional to absolute T; P/T = constant
D) PV = nRT only

C. The reason a pressurized can explodes if tossed in a fire: at constant V, doubling T doubles P.

10. (8.5) In a rigid container, doubling the absolute temperature of a gas:

A) Doubles the pressure
B) Halves the pressure
C) Leaves pressure unchanged
D) Quadruples the pressure

A. P/T = constant at fixed V; 2T ⇒ 2P.

11. (8.6) Avogadro's law states:

A) P = nRT alone
B) V is inversely related to n
C) P and V are inversely related
D) At the same T and P, equal volumes of gas contain equal moles (V ∝ n)

D. This links macroscopic volume to particle count, independent of species.

12. (8.6) Two gas samples at the same T and P but in different-sized containers:

A) Have equal moles but different pressures
B) Have different moles because volumes differ; bigger V → more moles
C) Have identical moles regardless of volume
D) Cannot be compared

B. With T and P fixed, V scales with n.

13. (8.7) The ideal gas law is:

A) PV = k
B) PV ∝ n only
C) PV = nRT
D) PV = T / n

C. Combines the previous gas laws with the molar quantity n. R is the universal gas constant.

14. (8.7) The gas constant R has value:

A) 0.0821 L·atm/(mol·K) or 8.314 J/(mol·K)
B) A value that depends on the gas species
C) 1 for all gases
D) 273 K

A. Same gas constant across all ideal gases. Use 0.0821 if working in atm and L; use 8.314 if working in SI units.

15. (8.8) Dalton's law of partial pressures states:

A) Each gas contributes only a fraction
B) Partial pressures cancel
C) $P_{\text{total}}$ is the inverse sum
D) $P_{\text{total}}$ equals the sum of the partial pressures of each component gas

D. Each gas acts as if it alone occupied the container. Handy when collecting a gas over water (subtract the water vapor pressure).

16. (8.8) The partial pressure of gas A in a mixture equals:

A) $x_{A}$ × R × T
B) $x_{A}$ × $P_{\text{total}}$ (mole fraction × total pressure)
C) $n_{A}$ × T
D) 1 / $x_{A}$

B. Derive from PV = nRT applied to each species. Mole fraction

x_{A}

n_{A}

n_{\text{total}}

17. (8.9) Kinetic molecular theory assumes that gas particles:

A) Attract each other strongly
B) Have finite volume and continuous interactions
C) Have negligible volume and negligible intermolecular forces, collide elastically, and move randomly
D) All have the same kinetic energy at a given instant

C. These assumptions give PV = nRT. Real gases deviate when these assumptions fail at high P or low T.

18. (8.9) The average kinetic energy of gas molecules depends on:

A) Absolute temperature (KE_avg = $\frac{3}{2}$ $k_{B}$ T per molecule)
B) Pressure
C) Volume
D) Number of molecules

A. Temperature IS molecular kinetic energy. Two different gases at the same T have the same average KE per molecule.

19. (8.10) The Maxwell-Boltzmann distribution describes:

A) Ions in solution
B) Electronic transitions
C) Bond vibrations
D) The distribution of molecular speeds in a gas at a given temperature

D. The curve has a right-skewed shape. Most-probable, mean, and rms speeds differ from each other.

20. (8.10) As temperature rises, the Maxwell-Boltzmann curve:

A) Narrows and shifts to lower speed
B) Broadens and shifts to higher speed, with a lower peak
C) Stays the same
D) Moves only downward

B. Area under the curve is conserved. A broader, flatter distribution means more molecules with speeds greater than any given threshold (including reaction threshold).

21. (8.11) Graham's law of effusion states:

A) Rates are equal for all gases
B) Rate increases with molar mass
C) Rate of effusion is inversely proportional to √(molar mass)
D) Effusion happens only in vacuum

C. Rate₁ / Rate₂ = √(M₂ / M₁). Lighter gases effuse faster.

22. (8.11) A gas with one-quarter the molar mass of another will effuse:

A) Twice as fast (√4 = 2)
B) Four times as fast
C) Half as fast
D) At the same rate

A. Take the square root because the scaling is with √M.

23. (8.12) Real gases deviate from ideal behavior:

A) Mostly at low P and high T
B) Only far from the critical point
C) At low density
D) At high P and/or low T, where molecular volumes and intermolecular attractions matter

D. Under these conditions the KMT assumptions break down. Liquefaction is the extreme case.

24. (8.12) The van der Waals equation adjusts the ideal gas law to account for:

A) Gravitational attraction
B) Molecular volume and intermolecular attractions, via the b and a parameters
C) Radioactive decay
D) Only pressure changes

B. (P + an²/V²)(V - nb) = nRT. The "a" term corrects pressure for attractions; "b" corrects volume for particle size.

Previous Attempts

Take a deep breath. You just inhaled roughly 10 sextillion gas molecules - and every single one of them is bouncing around at hundreds of meters per second, slamming into the walls of your lungs to create the pressure that keeps you alive. Gases are invisible, but their behavior follows simple, predictable rules that show up constantly on the MCAT.

This chapter is built around one master idea: gas behavior is governed by the relationships between pressure, volume, temperature, and the number of particles. Every gas law is just a snapshot of what happens when you hold some of these variables constant and change the others. If you understand the ideal gas law (PV = nRT), every other gas law falls out as a special case.

We will start with the basic characteristics of gases, then build up through the individual gas laws to the ideal gas law that unifies them all. From there, we will cover Dalton’s law for gas mixtures, the kinetic molecular theory that explains WHY gas laws work, the Maxwell-Boltzmann distribution that describes molecular speeds, Graham’s law for diffusion and effusion, and finally real gases - where the ideal model breaks down and the van der Waals equation picks up the slack.

💡 Real World

Gas laws are like the rules of a crowded dance floor. The dancers (gas molecules) move randomly in all directions. Pack more dancers into the room (increase pressure or decrease volume) and they bump into the walls more often. Turn up the music (increase temperature) and they move faster, pushing harder against the walls. The ideal gas law assumes every dancer is a point that never interacts with others - which works great for a spacious club but breaks down when the room gets too crowded or too cold.

In This Chapter

Characteristics of Gases
Standard Temperature and Pressure (STP)
Boyle’s Law
Charles’s Law
Gay-Lussac’s Law
Avogadro’s Law
The Ideal Gas Law
Dalton’s Law of Partial Pressures
Kinetic Molecular Theory
Maxwell-Boltzmann Distribution
Graham’s Law of Effusion and Diffusion
Real Gases
Section Test

← PreviousSection Test Next →Characteristics of Gases - Compressibility, Expansion, and Low Density

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