Solids, liquids, and gases — AP Chemistry Study Guide

For: AP Chemistry candidates sitting AP Chemistry.

Covers: Kinetic molecular theory comparisons of solids, liquids, and gases, intermolecular force strength correlation to phase, comparative bulk properties, ideal vs real gas behavior, phase density comparisons, and how particle arrangement drives bulk properties.

You should already know: Definition of intermolecular forces, basic kinetic molecular theory postulates for gases, units of absolute temperature and pressure.

A note on the practice questions: All worked questions in the "Practice Questions" section below are original problems written by us in the AP Chemistry style for educational use. They are not reproductions of past College Board / Cambridge / IB papers and may differ in wording, numerical values, or context. Use them to practise the technique; cross-check with official mark schemes for grading conventions.

1. What Is Solids, liquids, and gases?

This topic describes the three common macroscopic phases of matter, differentiated by the arrangement and motion of their constituent particles (atoms, molecules, or ions), tied directly to intermolecular force (IMF) strength and thermal kinetic energy. Per the AP Chemistry Course and Exam Description (CED), Unit 3 (Intermolecular Forces and Properties) makes up 18–22% of the total exam score, and this topic is a core foundational strand that appears in both multiple-choice (MCQ) and free-response (FRQ) sections. It is almost always combined with other Unit 3 topics, including IMF identification, gas law calculations, and phase diagram interpretation. Unlike introductory treatments that only list surface properties, AP Chemistry requires connecting microscopic particle behavior to macroscopic bulk properties: for example, why liquids are incompressible but flow, why solids have fixed shape, and why gases expand to fill their container. This topic also sets up the distinction between ideal and real gases, a frequent exam question, and links IMF strength to the phase a substance adopts at given temperature and pressure conditions.

2. Kinetic Molecular Theory (KMT) Comparison of Phases

Kinetic molecular theory models all matter as particles in constant random motion, with average kinetic energy directly proportional to absolute (Kelvin) temperature. The phase of a sample at given conditions is determined by the balance of two competing factors: intermolecular attractive forces that pull particles together, and thermal kinetic energy that pushes particles apart. For each phase, this balance produces consistent microscopic behavior:

• Gases: Kinetic energy >> IMF strength. Particles are far apart, move freely at high speed, fill the entire container, and are highly compressible due to large amounts of empty inter-particle space.
• Liquids: IMF strength ≈ kinetic energy. Particles are nearly touching (almost no empty space, so incompressible) but have enough energy to slide past one another, so they flow and take the shape of their container while retaining fixed volume.
• Solids: IMF strength >> kinetic energy. Particles are locked in a fixed lattice arrangement, only vibrating around their fixed positions, so they retain fixed shape and volume and are nearly incompressible.

Worked Example

Problem: For equal molar amounts of F₂, Br₂, and I₂ at 25°C and 1 atm, one is solid, one liquid, one gas. Match each substance to its phase and justify your answer using KMT and intermolecular force strength.

1. All three are nonpolar diatomic halogens, so the only intermolecular force present is London dispersion force (LDF), whose strength increases with increasing molar mass (due to greater polarizability of larger electron clouds).
2. Relative molar masses give the order of LDF strength: $M_{{\text{F}}_2}=38\text{g/mol}<M_{{\text{Br}}_2}=160\text{g/mol}<M_{{\text{I}}_2}=254\text{g/mol}$, so LDF strength increases as F₂ < Br₂ < I₂.
3. At a fixed temperature of 25°C, all three samples have the same average kinetic energy per particle, so the only difference is IMF strength.
4. Weakest IMFs (F₂): kinetic energy exceeds IMF attraction → F₂ is gas. Intermediate IMFs (Br₂): IMF strength is comparable to kinetic energy → Br₂ is liquid. Strongest IMFs (I₂): IMF strength exceeds kinetic energy → I₂ is solid.

Exam tip: AP exam graders always require you to explicitly connect IMF strength to the kinetic energy balance, not just state the matching. Always mention that temperature is fixed so average kinetic energy is equal for both substances to earn full justification points.

3. Comparative Bulk Properties of Phases

Bulk properties (measurable macroscopic properties) of each phase are direct consequences of their microscopic particle arrangement. Key properties compared across phases include compressibility, density, diffusion rate, and ability to flow:

• Compressibility: A measure of how much volume decreases under increased pressure. Gases have very high compressibility because ~99% of a gas sample is empty space between particles. Liquids and solids have particles touching, so there is almost no empty space to squeeze out, making them nearly incompressible.
• Density: Mass per unit volume, defined as $\rho =\frac{m}{V}$. For most pure substances, density follows the order ${\rho }_{\text{solid}}>{\rho }_{\text{liquid}}>{\rho }_{\text{gas}}$, because particles are most tightly packed in solids and most spread out in gases. Gas density is typically ~1000x lower than solid/liquid density for the same substance. The key exception to the solid > liquid density rule is water: hydrogen bonding creates an open crystal lattice in ice, so ${\rho }_{\text{ice}}=0.92{\text{g/cm}}^3<{\rho }_{\text{liquid water}}=1.0{\text{g/cm}}^3$, which is why ice floats.
• Diffusion: Spontaneous mixing of particles due to random motion. Diffusion rate is fastest in gases, slower in liquids, and extremely slow in solids, due to differences in free particle motion and inter-particle spacing.

Worked Example

Problem: A 1.0 g sample of liquid ethanol has a volume of 1.27 mL at 25°C. The same mass of ethanol vapor at 25°C and 1 atm has a volume of 530 mL. Calculate the ratio of the density of liquid ethanol to gaseous ethanol, and explain what this ratio reveals about inter-particle distance.

1. For equal mass samples, density ratio simplifies to $\frac{{\rho }_{\text{liquid}}}{{\rho }_{\text{gas}}}=\frac{m/V_l}{m/V_g}=\frac{V_g}{V_l}$, since mass cancels out.
2. Substitute the given volumes: $\frac{V_g}{V_l}=\frac{530\text{mL}}{1.27\text{mL}}\approx 417$, so liquid ethanol is ~420 times denser than gaseous ethanol at these conditions.
3. A 420x density ratio means the gas phase occupies 420 times more volume for the same number of particles. The average distance between particles in a gas is proportional to the cube root of the volume ratio, so $\sqrt[3]{420}\approx 7.5$, meaning average inter-particle distance in gaseous ethanol is ~7.5 times larger than in liquid ethanol.
4. This confirms that most of the volume of a gas is empty space between particles, while particles are nearly touching in the liquid phase.

Exam tip: When asked to explain density differences across phases, always link the difference to particle spacing, not just particle mass. Even heavy molecules have much lower density in the gas phase than the same substance in liquid form.

4. Ideal vs Real Gases: Deviations from KMT Postulates

The KMT model for ideal gases relies on two key postulates that are only approximately true for real gases: (1) ideal gas particles have negligible intrinsic volume compared to the total container volume, and (2) there are no attractive or repulsive intermolecular forces between ideal gas particles. For real gases, both postulates are false, leading to deviations from the ideal gas law $PV=nRT$. Deviations become significant under two conditions:

1. High pressure: When pressure is high, gas molecules are squeezed close together, so the intrinsic volume of the particles themselves becomes a significant fraction of the total container volume. The postulate of negligible particle volume breaks down here, leading to a measured volume larger than the ideal prediction.
2. Low temperature: When temperature is low, average kinetic energy is low, so intermolecular attractive forces are significant compared to kinetic energy. The postulate of no IMFs breaks down here, leading to a measured pressure lower than the ideal prediction.

Worked Example

Problem: 1.0 mol samples of ammonia (NH₃) are tested at 1 atm 25°C and 5 atm -40°C. Which sample shows a larger deviation from ideal gas behavior, and what is the main source of the deviation?

1. Ammonia is a polar molecule with strong hydrogen bonding between molecules, so IMFs are much stronger than in nonpolar gases of similar molar mass.
2. The first sample is at moderate pressure and high (room) temperature: particles are far apart, kinetic energy is high enough that IMFs are negligible, and particle volume is negligible compared to total volume, so deviation is small.
3. The second sample is at lower temperature (-40°C = 233 K), so average kinetic energy is much lower. Even at moderate pressure of 5 atm, the strong hydrogen bonding IMFs between NH₃ molecules are significant compared to kinetic energy.
4. The main source of deviation here is the presence of significant intermolecular attractive forces, violating the second KMT postulate for ideal gases, so the -40°C sample has much larger deviation.

Exam tip: Always link the source of deviation to the conditions: low temperature causes deviations from non-negligible IMFs, while very high pressure causes deviations from non-negligible particle volume. Do not mix these two up on FRQ justifications.

5. Common Pitfalls (and how to avoid them)

• Wrong move: Claims that gases are less dense than liquids because gas molecules have less mass than liquid molecules of the same substance. Why: Students confuse total mass of the sample with mass per unit volume, misremembering the definition of density. Correct move: Always start from the definition $\rho =m/V$, and compare mass per unit volume, or for equal mass compare inverse volume.
• Wrong move: States that ice is less dense than liquid water because ice molecules are larger than liquid water molecules. Why: Students confuse the open lattice structure from hydrogen bonding with a change in molecular size. Correct move: Always attribute lower ice density to the open hydrogen-bonded crystal lattice that leaves more empty space between water molecules than in liquid water.
• Wrong move: Claims all deviations of real gases from ideal behavior are caused by intermolecular forces, regardless of conditions. Why: Students memorize that IMFs cause deviation but forget the particle volume postulate violation that dominates at high pressure. Correct move: For any deviation question, first check conditions: low temperature = dominant deviation from IMFs; high pressure = dominant deviation from non-negligible particle volume.
• Wrong move: Justifies a phase difference between two substances only by saying "one has stronger IMFs", without linking to kinetic energy at the given conditions. Why: Students skip the core KMT balance that AP requires for full justification points. Correct move: Always explicitly state that at a given temperature, average kinetic energy is the same for both substances, so stronger IMFs shift the balance toward a more condensed phase.
• Wrong move: Assumes solids have no particle motion at all, only liquids and gases have motion. Why: Introductory courses often oversimplify solid particle behavior. Correct move: Recall that solid particles vibrate around their fixed lattice positions, so they do have kinetic energy proportional to temperature, just no large-scale translational motion.

6. Practice Questions (AP Chemistry Style)

Question 1 (Multiple Choice)

Xenon (Xe) is a monatomic gas at room temperature, but can be frozen into a solid at -118°C at 1 atm. Which of the following correctly ranks the compressibility of 1 mol Xe at 1 atm from highest to lowest at the following temperatures: -120°C (all solid), -100°C (all liquid), 25°C (all gas)? A) Solid > liquid > gas B) Gas > solid > liquid C) Gas > liquid > solid D) Liquid > gas > solid

Worked Solution: Compressibility depends on the amount of empty space between particles: more empty space = higher compressibility. Gases have large amounts of empty space between particles, so they have the highest compressibility. Liquids have particles almost touching, so they have very low compressibility, slightly higher than solids where particles are locked in a tightly packed lattice. The ranking from highest to lowest compressibility is gas > liquid > solid. The correct answer is C.

Question 2 (Free Response)

Sodium chloride (NaCl) has a melting point of 801°C at 1 atm, while oxygen (O₂) has a melting point of -218°C at 1 atm. (a) Identify the phase of each compound at 25°C and 1 atm, and justify each identification. (b) Explain the large difference in melting point between the two compounds in terms of attractive forces and the KMT balance between kinetic energy and attraction. (c) A 100 g sample of NaCl(s) and 100 g sample of O₂(g) at 1 atm 25°C have the same mass. Which sample has a larger volume? Justify your answer in terms of particle spacing.

Worked Solution: (a) 25°C is below the melting point of NaCl (25°C < 801°C), so NaCl is solid. 25°C is well above the melting point of O₂ and also above O₂'s boiling point of -183°C, so O₂ is gas at 25°C 1 atm. (b) NaCl is an ionic compound, so the attractive forces between oppositely charged ions are very strong ionic bonds. O₂ is a nonpolar molecular compound, so only weak London dispersion forces exist between O₂ molecules. At a given temperature, strong attractive forces in NaCl mean they exceed kinetic energy, holding the sample together as a solid, requiring a very high temperature (high kinetic energy) to melt. Weak IMFs in O₂ mean kinetic energy exceeds attractive forces at room temperature, so O₂ is gas. (c) The 100 g O₂(g) sample has a much larger volume. Gases have large average distances between particles, with most of the sample volume being empty space, while solid NaCl has ions nearly touching in a crystal lattice with almost no empty space. For the same mass, the large inter-particle spacing in gaseous O₂ leads to a much larger volume.

Question 3 (Application / Real-World Style)

Scuba divers breathe compressed air, which is mostly nitrogen gas. At very deep depths, high pressure causes nitrogen gas to dissolve in the diver's blood and body tissues. When a diver ascends too quickly, the pressure decreases, and nitrogen comes out of solution as small gas bubbles in the blood. Explain why these bubbles cause physical damage to tissues in terms of the properties of gases compared to liquid blood.

Worked Solution:

1. Nitrogen bubbles are gaseous nitrogen, while blood is a liquid. Gases have much lower density than liquids, and expand as external pressure decreases during ascent.
2. Because gas particles have large amounts of empty space between them, gases are highly compressible and expand when pressure decreases. The small nitrogen bubbles that form when pressure drops expand rapidly as the diver moves up, increasing in volume significantly.
3. The expanding gas bubbles push against and damage surrounding body tissue, causing the dangerous condition known as decompression sickness.

7. Quick Reference Cheatsheet

8. What's Next

This topic is the foundational prerequisite for all subsequent phase behavior and gas topics in AP Chemistry. Next, you will apply the KMT and phase balance concepts to phase diagrams, which describe the phase of a substance at different combinations of temperature and pressure, and to gas law calculations that relate pressure, volume, temperature, and moles of gas. Without a solid understanding of how particle arrangement and IMF strength drive phase properties, you will struggle to earn full points for justifications on phase change and gas deviation questions on the exam. This topic also feeds into the broader study of solution properties, where you will compare intermolecular forces between solute and solvent to predict solubility.

Phase Diagrams Ideal Gas Law Intermolecular Forces Solutions and Solubility