🌡️ Kinetic Molecular Theory

Part 1 of 7 — Particle Motion in Solids, Liquids, and Gases

Matter exists in three common phases — solid, liquid, and gas — and the Kinetic Molecular Theory (KMT) explains their properties by focusing on the behavior of individual particles (atoms, molecules, or ions).

The central idea: all particles are in constant motion, and the type and extent of that motion determines the phase of matter.

Postulates of Kinetic Molecular Theory

The KMT was originally developed for ideal gases, but its principles extend to all phases:

1. All matter is composed of tiny particles (atoms, molecules, or ions) that are in constant, random motion.

2. Temperature is a measure of the average kinetic energy of the particles:

$KE_{\text{avg}} = \frac{3}{2} k_B T$

where $k_B = 1.38 \times 10^{-23}$ J/K is Boltzmann's constant and $T$ is the absolute temperature in kelvin.

1. Collisions between gas particles and with container walls are perfectly elastic — no kinetic energy is lost.

2. The volume of individual gas particles is negligible compared to the volume of the container (for ideal gases).

3. There are no attractive or repulsive forces between ideal gas particles (real gases deviate from this).

Key Takeaway

At a given temperature, all gases have the same average kinetic energy, regardless of molar mass. Heavier molecules move more slowly; lighter molecules move faster.

Test your understanding of the basic postulates of Kinetic Molecular Theory.

How Particles Move in Each Phase

Solids 🧊

• Particles are tightly packed in fixed positions (usually a regular lattice).
• Particles vibrate about their fixed positions but do not translate or rotate freely.
• Strong intermolecular forces hold particles in place.
• Have a definite shape and definite volume.

Liquids 💧

• Particles are close together but can slide past one another.
• Particles have translational, rotational, and vibrational motion.
• Moderate intermolecular forces — strong enough to keep particles close, but not strong enough to fix them in place.
• Have a definite volume but take the shape of their container.

Gases 💨

• Particles are far apart with large distances between them.
• Particles move rapidly in random, straight-line paths until they collide.
• Very weak or negligible intermolecular forces (ideal gas assumption).
• Have no definite shape and no definite volume — expand to fill their container.

Complete each statement about the phases of matter.

The Relationship Between KE and Temperature

The average kinetic energy of particles depends only on temperature:

$KE_{\text{avg}} = \frac{3}{2} k_B T$

For a mole of particles, we can write:

$KE_{\text{avg per mole}} = \frac{3}{2} RT$

where $R = 8.314$ J/(mol·K) is the ideal gas constant.

Root-Mean-Square Speed

The rms speed ($v_{\text{rms}}$) relates KE to the molar mass $M$:

$v_{\text{rms}} = \sqrt{\frac{3RT}{M}}$

where $M$ is the molar mass in kg/mol (not g/mol!).

Important Consequences

• Doubling the temperature (in kelvin) doubles the average KE.
• Doubling the temperature increases $v_{\text{rms}}$ by a factor of $\sqrt{2}$ (not 2!).
• At the same temperature, a gas with 4× the molar mass has half the rms speed.

Use the equation $v_{\text{rms}} = \sqrt{3RT/M}$ to answer these questions. Use $R = 8.314$ J/(mol·K).

Maxwell-Boltzmann Distribution

Not all particles in a gas move at the same speed. The Maxwell-Boltzmann distribution shows the spread of molecular speeds at a given temperature:

Key features of the distribution curve:

• The curve is not symmetric — it is skewed to the right.
• Most probable speed ($v_p$): the peak of the curve (most common speed).
• Average speed ($v_{\text{avg}}$): slightly higher than $v_p$.
• Root-mean-square speed ($v_{\text{rms}}$): highest of the three, $v_{\text{rms}} > v_{\text{avg}} > v_p$.

Effect of Temperature

• Higher temperature → the curve shifts right (faster speeds) and becomes broader and flatter.
• Lower temperature → the curve shifts left (slower speeds) and becomes taller and narrower.
• The area under the curve is always the same (= total number of particles).

Effect of Molar Mass (at constant T)

• Lighter molecules → broader curve shifted to the right (faster).
• Heavier molecules → narrower curve shifted to the left (slower).

Test your understanding of the Maxwell-Boltzmann distribution.

Complete these key statements from Part 1.

🧊 Properties of Solids

Part 2 of 7 — Types of Solids and Their Properties

Solids have a definite shape and definite volume because their particles are held in fixed positions by strong interparticle forces. But not all solids are the same — the type of particles and bonding within a solid determine its physical properties.

We classify solids into two broad categories:

• Crystalline solids — particles arranged in a regular, repeating 3D pattern (a crystal lattice)
• Amorphous solids — particles arranged in a random, disordered pattern (no long-range order)

Crystalline vs. Amorphous Solids

Crystalline Solids

• Have a well-defined melting point (sharp transition from solid to liquid).
• Particles arranged in an orderly, repeating lattice.
• Examples: NaCl, diamond, quartz, iron, ice.

Amorphous Solids

• Have no definite melting point — they soften gradually over a range of temperatures.
• Particles arranged randomly, without long-range order.
• Often called "supercooled liquids" because their structure resembles a frozen liquid.
• Examples: glass, rubber, plastics, chocolate.

Distinguish between crystalline and amorphous solids.

Types of Crystalline Solids

1. Ionic Solids

Lattice particles: Cations (+) and anions (−) held together by electrostatic (ionic) bonds.

Properties:

• High melting points (strong ionic bonds; typically > 500°C)
• Hard but brittle — displacing a layer brings like charges together, causing repulsion and fracture
• Do not conduct electricity as solids (ions locked in place)
• Conduct electricity when melted or dissolved in water (ions free to move)
• Many are soluble in polar solvents like water

Examples: NaCl (801°C mp), MgO (2852°C mp), CaF₂

Why is MgO's melting point so much higher than NaCl's?

The lattice energy (strength of ionic bonding) depends on charge and ionic radius:

$\text{Lattice Energy} \propto \frac{q^+ \times q^-}{r^+ + r^-}$

MgO has 2+/2− charges vs. NaCl's 1+/1−, and smaller ions, giving MgO much stronger ionic bonds.

2. Molecular Solids

Lattice particles: Discrete molecules held together by intermolecular forces (LDF, dipole-dipole, or hydrogen bonds).

Properties:

• Low melting points (weak IMFs; typically < 300°C)
• Soft — easy to deform
• Do not conduct electricity in any phase (no free ions or delocalized electrons)
• Solubility follows "like dissolves like"

Examples: Ice (H₂O, 0°C mp), dry ice (CO₂, sublimes at −78°C), sugar (C₁₂H₂₂O₁₁), I₂

Key Insight

The intramolecular bonds (covalent bonds within each molecule) are strong, but the intermolecular forces (between molecules) are relatively weak. It's the IMFs that determine the melting point — you're separating molecules from each other, not breaking covalent bonds.

3. Metallic Solids

Lattice particles: Metal cations surrounded by a "sea" of delocalized electrons (metallic bonding).

Properties:

• Variable melting points (from −39°C for Hg to 3422°C for W/tungsten)
• Malleable (can be hammered into sheets) and ductile (drawn into wires)
• Excellent conductors of heat and electricity (delocalized electrons carry charge and energy)
• Lustrous (shiny) — delocalized electrons absorb and re-emit light
• Insoluble in most solvents

Examples: Fe, Cu, Al, Au, Na

Why are metals malleable instead of brittle?

When layers of a metal are displaced, the delocalized electrons shift to maintain bonding in the new arrangement. In ionic solids, displacement brings like charges together → shattering.

4. Network Covalent (Atomic) Solids

Lattice particles: Atoms connected by a continuous network of covalent bonds extending throughout the entire solid.

Properties:

• Extremely high melting points (strong covalent bonds; often > 1000°C)
• Very hard (diamond is the hardest natural substance)
• Do not conduct electricity (no free electrons or ions) — exception: graphite
• Insoluble in virtually all solvents

Examples:

• Diamond (C, mp 3550°C) — each carbon bonded to 4 others in a tetrahedral network
• Silicon dioxide / Quartz (SiO₂, mp 1713°C) — Si and O atoms in a 3D network
• Silicon carbide (SiC, mp 2730°C)
• Graphite (C) — carbon atoms in 2D sheets with delocalized electrons between layers; conducts electricity!

Diamond vs. Graphite

Both are pure carbon, but:

• Diamond: 3D network of $sp^3$ bonds → extremely hard, does not conduct
• Graphite: 2D sheets of $sp^2$ bonds with delocalized $\pi$ electrons → slippery layers, conducts electricity

Identify the type of solid based on its properties.

For each substance, select the correct type of crystalline solid.

Predicting Relative Melting Points

The melting point of a solid depends on the strength of the forces holding particles in the lattice:

Ranking by typical melting points (from lowest to highest):

$\text{Molecular} < \text{Metallic (varies)} < \text{Ionic} < \text{Network Covalent}$

Within each category:

Molecular solids: Stronger IMFs → higher mp

• H-bonding > dipole-dipole > LDF (for similar size)
• Larger molar mass → stronger LDF → higher mp

Ionic solids: Greater charge, smaller ions → higher lattice energy → higher mp

• MgO (2+/2−) ≫ NaCl (1+/1−)

Metallic solids: More valence electrons and smaller atomic radius → stronger metallic bonding → higher mp

• Tungsten (W, 3422°C) ≫ Sodium (Na, 98°C)

Network covalent solids: Shorter, stronger covalent bonds → higher mp

• Diamond (C–C bonds, 3550°C) > SiC (2730°C) > SiO₂ (1713°C)

Apply your knowledge of solid types to rank melting points.

Complete these key facts about types of solids.

💧 Properties of Liquids

Part 3 of 7 — Surface Tension, Viscosity, Capillary Action, and Vapor Pressure

Liquids occupy a middle ground between solids and gases. Their particles are close together (like solids) but can move past one another (like gases). This unique combination gives liquids several distinctive properties that are directly tied to the strength of their intermolecular forces (IMFs).

The four key liquid properties we'll explore:

1. Surface tension — the "skin" on a liquid surface
2. Viscosity — resistance to flow
3. Capillary action — liquid climbing up narrow tubes
4. Vapor pressure — tendency of molecules to escape to the gas phase

Surface Tension

What Is It?

Surface tension is the energy required to increase the surface area of a liquid. It arises because molecules at the surface experience an unbalanced pull — they are attracted to neighboring molecules on the sides and below, but not above (where there is air).

This net inward pull causes the surface to contract to the smallest possible area, behaving like an elastic "skin."

Why Does It Happen?

• Interior molecules are pulled equally in all directions → net force = 0.
• Surface molecules are pulled inward and sideways but not upward → net inward force.
• The liquid minimizes its surface area to minimize the number of molecules in this unfavorable surface position.

Factors Affecting Surface Tension

• Stronger IMFs → higher surface tension

  • Water (H-bonding) has much higher surface tension than ethanol
  • Mercury (metallic bonding) has extremely high surface tension

• Higher temperature → lower surface tension

  • Increased KE allows molecules to overcome surface forces more easily

Examples

Test your understanding of surface tension.

Viscosity

What Is It?

Viscosity is a liquid's resistance to flow. A "thick" liquid like honey has high viscosity; a "thin" liquid like water has low viscosity.

Molecular Explanation

For a liquid to flow, molecules must slide past one another. Anything that makes this more difficult increases viscosity:

1. Stronger IMFs → higher viscosity — molecules cling to each other more tightly
2. Larger, more complex molecular shapes → higher viscosity — molecules get tangled and entangled
3. Higher temperature → lower viscosity — more KE helps molecules overcome IMFs and slide past each other

Examples

Temperature Dependence

This is why warming honey makes it flow more easily — increased thermal energy helps molecules overcome intermolecular attractions.

Complete each statement about viscosity.

Capillary Action

What Is It?

Capillary action is the ability of a liquid to flow in narrow spaces without the assistance of gravity (and even against it). You see it when water climbs up a thin glass tube, or when a paper towel soaks up a spill.

Two Competing Forces

Capillary action involves two types of forces:

1. Cohesion — attraction between like molecules (liquid–liquid)

• Example: water molecules attracting each other via H-bonds

2. Adhesion — attraction between unlike molecules (liquid–surface)

• Example: water molecules attracted to the glass surface (SiO₂ has polar O–H groups)

Meniscus Shape

The shape of the liquid surface (meniscus) in a tube reveals the relative strength of adhesion vs. cohesion:

• Adhesion > Cohesion → liquid "climbs" the walls → concave meniscus (curves up)
  • Example: Water in glass
• Cohesion > Adhesion → liquid "pulls away" from the walls → convex meniscus (curves down)
  • Example: Mercury in glass

Capillary Rise

The height a liquid rises in a capillary tube depends on:

• Stronger adhesion → greater rise
• Smaller tube diameter → greater rise (more surface area relative to volume)
• Lower liquid density → greater rise
• Lower surface tension can reduce the effect

Test your understanding of capillary action and meniscus shape.

Vapor Pressure

What Is It?

Vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid in a closed container. At any temperature, some liquid molecules have enough KE to escape the surface and enter the gas phase (evaporation). In a closed system, gas molecules also return to the liquid (condensation).

Equilibrium is reached when the rate of evaporation = rate of condensation. The pressure of the vapor at this point is the equilibrium vapor pressure.

Factors Affecting Vapor Pressure

1. Strength of IMFs (most important)

• Weaker IMFs → higher vapor pressure (molecules escape more easily)
• Diethyl ether (weak LDF/dipole) has much higher vapor pressure than water (H-bonding)

2. Temperature

• Higher temperature → higher vapor pressure (more molecules have enough KE to escape)
• The relationship is exponential, described by the Clausius-Clapeyron equation:

$\ln P = -\frac{\Delta H_{\text{vap}}}{RT} + C$

or in its two-point form:

$\ln\frac{P_2}{P_1} = -\frac{\Delta H_{\text{vap}}}{R}\left(\frac{1}{T_2} - \frac{1}{T_1}\right)$

Boiling Point Connection

A liquid boils when its vapor pressure equals the external (atmospheric) pressure.

• Normal boiling point: the temperature at which vapor pressure = 1 atm (101.3 kPa)
• At higher altitudes (lower atmospheric pressure), liquids boil at lower temperatures
• In a pressure cooker (higher pressure), liquids boil at higher temperatures

Apply your understanding of vapor pressure.

Use the Clausius-Clapeyron equation: $\ln\frac{P_2}{P_1} = -\frac{\Delta H_{\text{vap}}}{R}\left(\frac{1}{T_2} - \frac{1}{T_1}\right)$ with $R = 8.314$ J/(mol·K).

Complete these key statements about liquid properties.

🔥 Phase Changes

Part 4 of 7 — Melting, Boiling, Sublimation, and Heating Curves

A phase change (or phase transition) occurs when matter converts from one phase to another. Energy is either absorbed or released during the process, but the temperature remains constant during the phase change itself — all the energy goes into overcoming intermolecular forces, not increasing kinetic energy.

The Six Phase Changes

Key Relationships

Since enthalpy is a state function, sublimation can be thought of as melting + vaporization:

$\Delta H_{\text{sub}} = \Delta H_{\text{fus}} + \Delta H_{\text{vap}}$

Why Does $\Delta H_{\text{vap}} > \Delta H_{\text{fus}}$?

• Melting only partially disrupts IMFs — particles go from fixed positions to sliding past each other, but remain close.
• Vaporization completely overcomes IMFs — particles go from close together to far apart.
• Therefore, vaporization always requires more energy than melting for the same substance.

For water: $\Delta H_{\text{fus}} = 6.01$ kJ/mol vs. $\Delta H_{\text{vap}} = 40.7$ kJ/mol

Identify the correct phase change.

Heating Curves

A heating curve plots temperature vs. heat added as a substance is heated from solid to gas at constant pressure. It has five distinct regions:

The Five Regions

Region 1: Heating the Solid (temperature rises)

• $q = m \times c_{\text{solid}} \times \Delta T$
• Particles vibrate faster; temperature increases

Region 2: Melting (Plateau) (temperature constant at $T_{\text{m}}$)

• $q = n \times \Delta H_{\text{fus}}$ (or $q = m \times \Delta H_{\text{fus, per gram}}$)
• Energy breaks IMFs to convert solid → liquid
• Temperature stays constant — all energy disrupts the lattice

Region 3: Heating the Liquid (temperature rises)

• $q = m \times c_{\text{liquid}} \times \Delta T$
• Particles move faster; temperature increases

Region 4: Boiling (Plateau) (temperature constant at $T_{\text{b}}$)

• $q = n \times \Delta H_{\text{vap}}$
• Energy completely separates particles from the liquid
• Temperature stays constant — all energy overcomes remaining IMFs
• This plateau is longer than the melting plateau because $\Delta H_{\text{vap}} > \Delta H_{\text{fus}}$

Region 5: Heating the Gas (temperature rises)

• $q = m \times c_{\text{gas}} \times \Delta T$
• Gas particles move faster; temperature increases

Key Features

• Flat regions = phase changes (temperature constant, energy disrupts/forms IMFs)
• Sloped regions = single phase (temperature changes, energy increases KE)
• The slope of each region depends on the specific heat capacity $c$ of that phase

Complete each statement about heating curves.

Calculating Total Energy for Heating Curves

To calculate the total energy required to heat a substance through a phase change, you must add the energy for each region separately.

Example: Heating ice at −20°C to steam at 120°C

For water ($m$ = 18.015 g/mol):

• $c_{\text{ice}} = 2.09$ J/(g·°C)
• $\Delta H_{\text{fus}} = 6.01$ kJ/mol = 334 J/g
• $c_{\text{water}} = 4.18$ J/(g·°C)
• $\Delta H_{\text{vap}} = 40.7$ kJ/mol = 2,260 J/g
• $c_{\text{steam}} = 2.01$ J/(g·°C)

For 1 mole (18.015 g) of water:

Step 1: Heat ice from −20°C to 0°C: $q_1 = 18.015 \times 2.09 \times 20 = 753$ J

Step 2: Melt ice at 0°C: $q_2 = 1 \times 6010 = 6{,}010$ J

Step 3: Heat water from 0°C to 100°C: $q_3 = 18.015 \times 4.18 \times 100 = 7{,}530$ J

Step 4: Boil water at 100°C: $q_4 = 1 \times 40{,}700 = 40{,}700$ J

Step 5: Heat steam from 100°C to 120°C: $q_5 = 18.015 \times 2.01 \times 20 = 724$ J

Total: $q_{\text{total}} = 753 + 6{,}010 + 7{,}530 + 40{,}700 + 724 = 55{,}717$ J $\approx 55.7$ kJ

Calculate the energy for individual steps. Use: $c_{\text{ice}} = 2.09$ J/(g·°C), $\Delta H_{\text{fus}} = 334$ J/g, $c_{\text{water}} = 4.18$ J/(g·°C), $\Delta H_{\text{vap}} = 2260$ J/g.

Cooling Curves

A cooling curve is the reverse of a heating curve — it plots temperature vs. time (or energy removed) as a substance cools.

• The same plateaus appear at the same temperatures (freezing point and condensation point).
• Energy is released (exothermic) during phase changes: condensation and freezing.
• The magnitudes of $\Delta H$ are the same, but the sign is negative (energy released).

Supercooling

Sometimes a liquid can be cooled below its freezing point without solidifying — this is called supercooling. When crystallization finally begins, the temperature may briefly rise back to the freezing point as the exothermic freezing process releases heat.

Test your overall understanding of phase changes and heating curves.

Complete these statements about phase changes.

📊 Phase Diagrams

Part 5 of 7 — Triple Points, Critical Points, and Reading Phase Diagrams

A phase diagram is a graph that shows which phase of a substance is most stable at any given combination of temperature (x-axis) and pressure (y-axis). Phase diagrams encode an enormous amount of information about a substance's behavior in a single image.

Anatomy of a Phase Diagram

A typical phase diagram has three regions (areas) and three lines (boundaries):

The Three Regions

• Solid region — upper left (high pressure, low temperature)
• Liquid region — middle area
• Gas region — lower right (low pressure, high temperature)

The Three Boundary Lines

Each line represents conditions where two phases coexist in equilibrium:

1. Solid-Liquid line (fusion curve) — separates solid and liquid regions

   • Follows the equation related to the Clausius-Clapeyron relation
   • Slope is usually positive (slants right) — increased pressure favors the denser solid phase

2. Liquid-Gas line (vaporization curve) — separates liquid and gas regions

   • Ends at the critical point
   • Corresponds to the vapor pressure vs. temperature relationship

3. Solid-Gas line (sublimation curve) — separates solid and gas regions

   • Below the triple point

Two Special Points

Triple Point — where all three boundary lines meet

• The unique temperature and pressure where solid, liquid, and gas coexist simultaneously
• For water: $T = 0.01°C$, $P = 611.7$ Pa (0.00604 atm)

Critical Point — the end of the liquid-gas boundary line

• Above this temperature and pressure, liquid and gas become indistinguishable → a supercritical fluid
• For water: $T_c = 374°C$, $P_c = 218$ atm
• For CO₂: $T_c = 31.1°C$, $P_c = 73$ atm

Test your understanding of phase diagram features.

How to Read a Phase Diagram

Determining the Phase

To find the phase at a specific temperature and pressure:

1. Locate the point $(T, P)$ on the diagram.
2. Determine which region the point falls in → that's the stable phase.

Tracing a Path

Heating at constant pressure (horizontal line from left to right):

• Cross the solid-liquid line → melting
• Cross the liquid-gas line → boiling

Increasing pressure at constant temperature (vertical line going up):

• Cross the gas-liquid line → condensation
• Cross the liquid-solid line → freezing

Example: What happens when you heat CO₂ at 1 atm?

At 1 atm pressure, trace a horizontal line across the CO₂ phase diagram:

• You start in the solid region.
• You cross the solid-gas line (sublimation curve) — but not the solid-liquid or liquid-gas lines!
• CO₂ goes directly from solid → gas (sublimation).

Why? Because 1 atm is below the triple point pressure of CO₂ (5.11 atm). At pressures below the triple point, the liquid phase does not exist — the substance sublimes.

This is why dry ice sublimes at atmospheric pressure instead of melting!

Complete each statement about interpreting phase diagrams.

Water vs. CO₂: The Anomaly

Typical Substances (e.g., CO₂)

For most substances, the solid-liquid line slopes to the right (positive slope). This means:

• Increasing pressure at constant temperature favors the solid phase.
• The solid is denser than the liquid (which is the norm).

Water — The Exception

Water's solid-liquid line slopes to the left (negative slope)! This means:

• Increasing pressure at constant temperature can melt ice → favors the liquid.
• Ice is less dense than liquid water — an anomaly among substances.

Why Is Ice Less Dense?

In ice, water molecules form a hexagonal crystal lattice stabilized by hydrogen bonds. This open structure has empty space, making ice less dense than liquid water (where H-bond network is partially disrupted and molecules are more randomly packed).

Consequences of Water's Anomaly

1. Ice floats — lakes freeze from the top down, insulating aquatic life below.
2. Ice skating — high pressure under the blade can melt ice (though friction is the main factor).
3. Glacial movement — high pressure at the base of glaciers can cause localized melting.

Comparison Table

Test your understanding of water's unusual phase diagram.

Supercritical Fluids

Above the critical temperature ($T_c$) and critical pressure ($P_c$), a substance enters the supercritical fluid region. In this state:

• The boundary between liquid and gas disappears — there is no distinct phase transition.
• The substance has properties of both a liquid (dissolving power, density) and a gas (fills container, low viscosity).

Supercritical CO₂ — An Important Application

Because CO₂'s critical point is relatively accessible ($T_c = 31.1°C$, $P_c = 73$ atm):

• Supercritical CO₂ is widely used as a "green" solvent in industrial processes.
• It is used to decaffeinate coffee — it dissolves caffeine selectively, then the CO₂ is depressurized and the caffeine is recovered.
• It leaves no toxic residue (CO₂ simply evaporates when pressure is released).

Complete these key statements about phase diagrams.

Final check on phase diagram concepts.

🔧 Problem-Solving Workshop

Part 6 of 7 — Predicting States and Comparing Properties Based on IMFs

In this workshop, we'll practice the most important skill for AP Chemistry: connecting the type and strength of intermolecular forces to observable physical properties of substances.

The chain of reasoning is: $\text{Structure} \rightarrow \text{IMFs} \rightarrow \text{Physical Properties}$

Stronger IMFs → higher melting/boiling points, higher surface tension, higher viscosity, lower vapor pressure.

Quick Review: IMF Strength Ranking

From weakest to strongest:

1. London Dispersion Forces (LDF) — present in ALL molecules

   • Strength increases with molar mass and surface area
   • Only force in nonpolar molecules

2. Dipole-Dipole Forces — between polar molecules

   • Stronger than LDF of comparable size
   • Requires permanent dipoles

3. Hydrogen Bonding — special, strong dipole-dipole

   • Requires H bonded to F, O, or N and a lone pair on F, O, or N on another molecule
   • Much stronger than ordinary dipole-dipole

4. Ion-Dipole Forces — between ions and polar molecules

   • Important for dissolving ionic compounds in water

Critical Reminder

When comparing two substances:

• First identify the dominant IMF in each.
• If both have the same type of IMF, compare molar mass (affects LDF strength).
• The substance with stronger IMFs has: higher bp, higher surface tension, higher viscosity, lower vapor pressure.

Predict the state of matter at room temperature (25°C, 1 atm) based on IMFs.

Strategy for Comparing Boiling Points

Step-by-Step Method

Step 1: Determine the type of substance (ionic, molecular, metallic, network covalent).

• Ionic/metallic/network covalent → generally much higher bp than molecular.

Step 2: For molecular substances, identify the dominant IMF:

• Can the molecule form hydrogen bonds? (H bonded to F, O, or N?)
• Is the molecule polar? (dipole-dipole + LDF)
• Is the molecule nonpolar? (LDF only)

Step 3: If two substances have the same type of dominant IMF, compare:

• Molar mass (larger → stronger LDF → higher bp)
• Number of H-bonding sites (more → stronger H-bonding network → higher bp)
• Molecular shape (more elongated → more surface contact → higher bp)

Common Comparisons on the AP Exam

Rank these substances by boiling point.

For each pair, select which substance has the HIGHER vapor pressure at the same temperature.

Apply IMF reasoning to surface tension and viscosity.

Predict properties based on IMF reasoning.

Final comprehensive check on predicting properties from IMFs.

🎯 Synthesis & AP Review

Part 7 of 7 — Connecting IMFs to Physical Properties and AP-Style Problems

This final part brings everything together. On the AP Chemistry exam, you'll be expected to:

1. Identify the types of IMFs present in a substance from its structure.
2. Compare physical properties (bp, mp, vapor pressure, viscosity, surface tension) of different substances based on their IMFs.
3. Interpret heating curves and phase diagrams.
4. Explain macroscopic observations using particulate-level reasoning.

Let's practice with AP-style questions and synthesis problems.

The Central Chain of Reasoning

$\text{Molecular Structure} \xrightarrow{\text{determines}} \text{IMFs} \xrightarrow{\text{determines}} \text{Physical Properties}$

Complete Property Summary

The AP Exam Wants You To...

• Never say "stronger bonds" when you mean "stronger intermolecular forces"
• Distinguish IMFs (between molecules) from intramolecular bonds (within molecules)
• Be specific: say "hydrogen bonding" or "London Dispersion Forces," not just "IMFs"
• Explain trends using particle-level reasoning

Answer these AP-style questions about properties of states of matter.

More AP-style questions integrating concepts from all parts.

AP Free-Response Practice

Consider the following four substances:

Analysis

All four are hydrides of Period 2 or 3 elements. If only LDF mattered, boiling points would increase with molar mass: CH₄ < H₂O < HF < H₂S.

But the actual order is: CH₄ (−161) < H₂S (−60) < HF (19.5) < H₂O (100).

Why?

• CH₄: Nonpolar, tetrahedral → LDF only → very low bp
• H₂S: Polar, but S is not F/O/N → no H-bonding → dipole-dipole + LDF → moderate bp
• HF: H-bonding (H–F) → high bp, but only one H per molecule limits the H-bond network
• H₂O: H-bonding (O–H), and each molecule has 2 H atoms and 2 lone pairs → extensive 3D H-bonding network → highest bp

Key Lesson

Water's extraordinarily high boiling point (relative to its molar mass) is due to its ability to form an extensive hydrogen bonding network — each water molecule can form up to 4 hydrogen bonds.

Answer these AP-style short-answer questions.

Complete each statement connecting structure, IMFs, and properties.

Demonstrate mastery of all concepts from Parts 1–7.

Complete these final summary statements.