Unit 3: Intermolecular Forces and Properties

Intermolecular Forces and Molecular Polarity

Chemists separate the idea of a chemical bond (an intramolecular attraction that holds atoms together within a particle) from an intermolecular force (an attraction between separate particles, such as between two molecules or between an ion and a molecule). This distinction matters because many “everyday” properties you observe—whether a substance is a gas or liquid at room temperature, how easily it evaporates, whether two liquids mix, how viscous a liquid feels—are controlled much more by intermolecular forces (IMFs) than by the covalent bonds inside each molecule.

A useful way to think about IMFs is that they are the “stickiness” between particles. Stronger stickiness makes it harder to pull particles apart, so it usually raises boiling point, lowers vapor pressure, and increases viscosity and surface tension. To predict which substances have stronger IMFs, connect three ideas:

1. How charge is distributed within a particle (polarity)
2. How close particles can get (size and shape)
3. What kinds of particles are interacting (molecule–molecule, ion–molecule)

Electron density, bond polarity, and dipoles

When a covalent bond is formed and electrons are shared, they may be shared unequally. The more electronegative element in the bond has the stronger pull on electrons. This unequal sharing creates a polar covalent bond, which gives each atom a partial charge. Partial charges are indicated using the lowercase Greek letter delta: delta plus and delta minus.

A dipole is opposite charges separated by some distance, and a dipole moment measures the degree of that charge separation. When atoms in a covalent bond have equal (or very similar) electronegativities, electrons are shared more evenly, there is essentially no dipole, and the bond is nonpolar covalent.

Molecular polarity (geometry matters)

A key point that causes mistakes: a molecule can have polar bonds but still be nonpolar overall if the bond dipoles cancel because of molecular geometry. To decide whether a molecule is polar overall, you typically:

1. Determine the molecular shape (from Lewis structure and VSEPR)
2. Identify polar bonds (electronegativity differences)
3. Consider symmetry: do dipoles cancel?

For example, carbon dioxide has polar C=O bonds, but linear geometry makes the dipoles cancel, so the molecule is nonpolar. Similarly, in carbon tetrafluoride, the individual C–F bonds are polar, but because there are four identical bonds arranged symmetrically in a tetrahedral shape, the dipoles cancel and the molecule is nonpolar.

In NF3, only three bonding locations have fluorine, and the remaining region is a lone pair on nitrogen. Fluorine is more electronegative than nitrogen, so the fluorines tend to carry partial negative character. The presence of a lone pair also makes the shape asymmetric, so the bond dipoles do not cancel completely and the molecule is polar.

A good symmetry shortcut is: if the central atom has no lone pairs and all bonded atoms are the same, the molecule is often symmetrical and nonpolar. If the central atom has lone pairs and/or different surrounding atoms, the molecule is often asymmetrical and polar.

However, lone pairs do not always guarantee polarity. In some electron geometries, lone pairs can be arranged symmetrically so that the overall dipole cancels. Two important examples:

• Linear molecules derived from trigonal bipyramidal electron geometry with three equatorial lone pairs (for example, XeF2) can be nonpolar.
• Square planar molecules derived from octahedral electron geometry with two opposite lone pairs (for example, XeF4) can be nonpolar.

Also, because in Lewis structures the least electronegative atom is often placed in the center, many polar molecules have a central atom that is partially positive relative to surrounding atoms. A common exception is when hydrogen is a terminal atom: hydrogen’s low electronegativity can make the atom it is bonded to carry partial negative character.

The major types of intermolecular forces

In AP Chemistry, you focus on a small set of IMFs that explain most observable trends.

London dispersion forces (LDF)

London dispersion forces are attractions caused by temporary fluctuations in electron density. Even in a nonpolar molecule, electrons are moving; at any instant the electron cloud can be slightly uneven, creating an instantaneous dipole. That instantaneous dipole can induce a dipole in a neighboring particle, leading to attraction.

LDF are present in all substances (polar or nonpolar). In nonpolar substances, LDF are the only IMFs, so they control boiling point trends among nonpolar molecules.

What makes LDF stronger:

• More electrons / larger molar mass: bigger electron clouds are more easily distorted (more polarizable)
• Greater surface area of contact: long, less-branched molecules can pack with more contact area, strengthening dispersion attractions

Even though LDF are often weak for small molecules, they can become very large for big, highly polarizable molecules. In large-enough molecules, dispersion forces can rival or even outweigh the effect of dipole–dipole attractions and can be comparable in impact to hydrogen bonding when electron counts and polarizability are very high.

Dipole–dipole forces

Dipole–dipole forces occur between polar molecules: the positive end of one molecule is attracted to the negative end of another.

When two substances are similar in size, the polar one generally has a higher boiling point because dipole–dipole attractions add to dispersion forces. Dipole–dipole forces depend on molecules being able to orient so opposite partial charges align; increased temperature disrupts that alignment, which is one reason heating makes liquids evaporate.

Hydrogen bonding

Hydrogen bonding is a particularly strong form of dipole–dipole attraction that occurs when a hydrogen atom is covalently bonded to nitrogen, oxygen, or fluorine and is attracted to a lone pair on N, O, or F of a neighboring molecule.

Hydrogen bonding is unusually strong because N, O, and F are very electronegative, so the H bonded to them is highly electron-poor (strong partial positive). Also, hydrogen is extremely small and has no inner (core) electrons, so the partially positive region on H is very concentrated—often described as the “nucleus being more exposed” compared with larger atoms.

Hydrogen bonding is often tested through property predictions:

• Water’s unusually high boiling point for its molar mass
• Alcohols having higher boiling points than similar-mass hydrocarbons
• Water and ammonia having higher melting and boiling points than similar molecules that cannot hydrogen bond
• Biological structure (DNA base pairing) as a real-world context

Common mistake: hydrogen bonding does not happen just because a molecule contains hydrogen and nitrogen/oxygen/fluorine somewhere. The hydrogen must be directly bonded to N, O, or F (donor), and there must be a lone pair on N, O, or F (acceptor) nearby.

Ion–dipole forces

Ion–dipole forces occur between an ion and a polar molecule. For example, when sodium chloride dissolves in water, Na plus interacts with the oxygen end of water, and Cl minus interacts with the hydrogen end.

Ion–dipole attractions help explain why many ionic compounds dissolve in polar solvents and why dissolution can be energetically favorable. Ion–dipole interactions are generally stronger than hydrogen bonding (though actual strength depends on ion charge and size).

Comparing IMF strength: a practical hierarchy

For AP-style predictions, a useful simplified ordering is:

• Ion–dipole (often strongest in solutions)
• Hydrogen bonding (strong dipole–dipole subtype)
• Dipole–dipole
• London dispersion (always present; can become very strong for large, polarizable molecules)

That last point is crucial: a large nonpolar molecule can have stronger dispersion forces than a small polar molecule. For instance, iodine (I2) is nonpolar but has a much higher boiling point than many small polar molecules because it is large and highly polarizable.

Also keep the “force type” separate from bonding type in solids. Ionic lattices are held together by ionic bonding (electrostatic attraction in a crystal), and network covalent solids are held together by extended covalent bonds. Melting those solids typically requires far more energy than overcoming IMFs in molecular substances. Metallic bonding can also be very strong, especially in many transition metals, leading to high melting points.

Bonding and phases (connecting forces to states of matter)

In general, the stronger the attractions holding particles together, the more tightly packed the particles tend to be in condensed phases. Solids have particles in a tight, ordered arrangement; gases have particles far apart with minimal attractive influence most of the time.

Substances with weak IMFs are often gases at room temperature (for example, nitrogen gas). Many substances capable of hydrogen bonding are liquids at room temperature (for example, water) because hydrogen bonding provides stronger cohesion.

Ionic substances are not described as experiencing IMFs between molecules; their phase is determined by strong ion–ion attractions in the lattice, so ionic compounds are usually solids at room temperature.

Worked example: identifying IMFs in a substance

Example: For each substance, identify the most significant intermolecular forces.

1. CH4
2. CH3OH
3. NH3
4. Na plus in water

Reasoning:

• CH4 is nonpolar (tetrahedral symmetry, C–H bonds nearly nonpolar). It has only London dispersion forces.
• CH3OH (methanol) is polar and has an O–H bond. It has dispersion, dipole–dipole, and hydrogen bonding.
• NH3 is polar (trigonal pyramidal) and has N–H bonds plus a lone pair on N, so it can hydrogen bond. It has dispersion, dipole–dipole, and hydrogen bonding.
• Na plus in water forms ion–dipole attractions with polar water molecules.

Exam Focus

• Typical question patterns:
  • Given structures or formulas, determine which IMFs are present and justify using polarity and bonding.
  • Compare two or more substances and predict which has higher boiling point or stronger IMFs.
  • Explain an anomalous trend (for example, water vs H2S) using hydrogen bonding.
• Common mistakes:
  • Declaring a molecule polar based only on polar bonds, without checking geometry and dipole cancellation.
  • Saying hydrogen bonding exists if H is anywhere in the molecule with O/N/F (instead of H bonded to O/N/F).
  • Forgetting that dispersion forces exist in all substances, including polar ones.

How Intermolecular Forces Control Physical Properties

Once you can identify IMFs, you use them to explain macroscopic properties. This is a core AP skill: connect particle-level attractions to lab-scale observations.

Boiling point, melting point, and phase at room temperature

A substance boils when particles in the liquid have enough energy to escape into the gas phase. Stronger IMFs mean particles are held more tightly, so you need a higher temperature to separate them, leading to a higher boiling point.

Melting involves disrupting the ordered structure of a solid. Melting point depends not only on IMF strength, but also on how well particles pack into a solid lattice. That’s why melting point trends can be trickier than boiling point trends: a molecule that packs efficiently may have a higher melting point even if its IMFs are not dramatically stronger.

A frequent comparison involves structural isomers:

• Less branched, more “linear” molecules often have higher boiling points (more surface area, stronger dispersion).
• More symmetrical molecules may have higher melting points (pack more efficiently in a crystal).

Vapor pressure (and evaporation vs boiling)

Vapor pressure is the pressure exerted by vapor molecules above a liquid in a closed container at a given temperature. It reflects how readily molecules escape from the liquid.

• Strong IMFs hold molecules in the liquid more strongly.
• Fewer molecules escape into the gas.
• Vapor pressure is lower.

So at the same temperature:

• Higher vapor pressure implies weaker IMFs.
• Lower vapor pressure implies stronger IMFs.

It also helps to separate two related ideas:

• Evaporation (a form of vaporization) can occur at temperatures below the boiling point, without heating the entire liquid to the boiling point. Even then, molecules still must gain enough energy (often from the surroundings) to overcome IMFs at the surface.
• Boiling occurs when vapor bubbles form throughout the liquid at the boiling point; at this point, energy input goes heavily into overcoming IMFs during the phase change.

As temperature increases, average kinetic energy increases, and a larger fraction of molecules can overcome IMFs, increasing evaporation rate and vapor pressure.

A common confusion is mixing up vapor pressure with “pressure of the container.” Vapor pressure is about equilibrium between evaporation and condensation; it is a characteristic of the substance at that temperature.

Enthalpy of vaporization and enthalpy of fusion

When a liquid becomes a gas, energy is required to overcome IMFs. The enthalpy of vaporization is the energy needed to vaporize a given amount of liquid at its boiling point. Stronger IMFs generally mean a larger enthalpy of vaporization.

Similarly, the enthalpy of fusion is energy needed to melt a solid at its melting point.

You don’t usually memorize values in AP Chemistry here; the key skill is qualitative reasoning: stronger attractions require more energy to separate.

Viscosity and surface tension

Viscosity is a liquid’s resistance to flow. Strong IMFs increase viscosity because molecules resist sliding past one another.

Surface tension is the tendency of a liquid surface to minimize area, caused by cohesive forces pulling surface molecules inward. Stronger IMFs increase surface tension.

Real-world connections:

• Honey flows slowly partly because of strong intermolecular attractions and large molecules.
• Water beads up on waxy surfaces due to strong cohesion within water (hydrogen bonding) and weak adhesion to nonpolar wax.

Heating curves and phase changes (energy vs temperature)

A heating curve plots temperature versus heat added. It visualizes two different uses of energy:

1. Increasing kinetic energy (temperature rises within a single phase)
2. Breaking or loosening intermolecular attractions during a phase change (temperature stays constant during melting or boiling)

During melting/boiling plateaus, added energy does not increase temperature because it is being used to overcome IMFs and separate particles.

If you’re asked to compare two substances’ heating curves, the substance with stronger IMFs typically shows:

• Higher melting and boiling temperatures
• Longer plateaus for phase changes if the same amount of heat is added (more energy required to change phase)

Phase diagrams (qualitative understanding)

A phase diagram shows which phase (solid, liquid, gas) is stable at different temperatures and pressures.

Important features:

• The triple point, where solid, liquid, and gas coexist at equilibrium.
• The critical point, beyond which the liquid and gas phases are indistinguishable (supercritical fluid).

A key qualitative link to IMFs: stronger IMFs tend to raise boiling temperatures, which often shifts the liquid–gas equilibrium boundary.

Water has a famous “exception” behavior: its solid (ice) is less dense than its liquid, giving the solid–liquid boundary a negative slope. On AP-style questions, this is usually explained by the open hydrogen-bonding network in ice.

Worked example: boiling point and vapor pressure

Example: At the same temperature, Substance A has a higher vapor pressure than Substance B. What can you conclude about IMFs and boiling points?

Reasoning: Higher vapor pressure means molecules escape more easily, so IMFs are weaker in A than in B. If A has weaker IMFs, it generally has a lower boiling point than B.

Exam Focus

• Typical question patterns:
  • Rank substances by boiling point, melting point, viscosity, surface tension, or vapor pressure using IMFs and structure.
  • Interpret a heating curve: identify which segments correspond to warming vs phase changes and explain why temperature is constant during plateaus.
  • Interpret a phase diagram: locate triple point, critical point, and infer which phase exists at given conditions.
• Common mistakes:
  • Assuming melting point trends always match boiling point trends (packing effects can complicate melting points).
  • Saying “temperature increases during boiling” on a heating curve plateau (it does not at constant pressure).
  • Reversing vapor pressure logic (stronger IMFs mean lower, not higher, vapor pressure).

Solids: Structure, Types, and Properties

Solids are not all held together the same way. AP Chemistry expects you to connect a solid’s microscopic structure to macroscopic properties like conductivity, malleability, solubility, and melting point.

A helpful starting definition: a solid has particles arranged in a fixed, closely packed structure (often crystalline), so it has definite shape and volume.

Four major categories of solids

Molecular solids

A molecular solid consists of neutral molecules held together by IMFs (dispersion, dipole–dipole, hydrogen bonding).

Typical properties:

• Relatively low melting points compared with ionic or network solids
• Poor electrical conductivity (no mobile charges)
• Often soft

Examples include ice (molecular solid with hydrogen bonding) and solid carbon dioxide (dry ice, held by dispersion).

Ionic solids

An ionic solid is a lattice of cations and anions held together by electrostatic attraction. The “force” here is not an IMF between molecules; it is ionic bonding in an extended crystal lattice.

Typical properties:

• High melting points (strong attractions in the lattice)
• Brittle (shifting layers can bring like charges next to each other, causing repulsion and fracture)
• Conduct electricity when molten or dissolved (ions can move), but not as a solid

A common confusion: ionic solids contain ions, but the ions are not free to move in a rigid lattice, so they do not conduct as solids.

Metallic solids

A metallic solid consists of metal atoms with valence electrons delocalized in a “sea of electrons.” Metallic bonding can be very strong, and many metals (especially many transition metals) have high melting points.

Typical properties:

• Good electrical and thermal conductivity (mobile electrons)
• Malleable and ductile (layers of atoms can slide without breaking a directional bond network)
• Variable melting points

Covalent network solids

A covalent network solid is a giant, continuous network of atoms bonded by covalent bonds (not discrete molecules). Network covalent solids are among the hardest substances to melt because melting requires breaking covalent bonds across the network.

Typical properties:

• Very high melting points
• Very hard (for example, diamond)
• Usually poor conductors (graphite is a notable exception because it has delocalized electrons within layers)

The exam often probes whether you can distinguish a network solid from a molecular substance with a similar formula. For example, carbon in diamond is not “molecules of carbon”; it is a network.

Structure-property reasoning: what to say and why

When AP asks you to “justify” a property, a strong answer names the type of solid and connects it to the kind of particles and forces:

• Molecular solid: molecules + IMFs, so lower melting point, no mobile charges
• Ionic solid: ions + lattice attractions, so high melting point, brittle, conducts when molten/dissolved
• Metallic: delocalized electrons, so conducts, malleable
• Network: covalent bonds throughout, so very high melting point, very hard

Worked example: classifying solids

Example: Predict which has the highest melting point and explain.

• Solid neon (Ne)
• Solid water (H2O)
• Solid sodium chloride (NaCl)
• Solid silicon dioxide (SiO2)

Reasoning:

• Ne is a molecular (atomic) solid held by dispersion only, very low melting.
• H2O is a molecular solid with hydrogen bonding, higher but still limited.
• NaCl is ionic, high melting.
• SiO2 (quartz) is a covalent network solid, extremely high melting because melting requires breaking covalent bonds across the network.

So the highest melting point is SiO2.

Exam Focus

• Typical question patterns:
  • Given a formula or description, identify the solid type (molecular, ionic, metallic, network) and predict properties.
  • Explain conductivity differences: solid vs molten vs aqueous for ionic compounds; graphite vs diamond.
  • Compare melting points using the kind of bonding/structure.
• Common mistakes:
  • Calling ionic lattice attractions “intermolecular forces” (they are not IMFs between molecules).
  • Assuming all covalent substances have low melting points (network solids are the counterexample).
  • Forgetting that ionic solids conduct only when ions can move (molten or dissolved).

Gases: Pressure, Temperature, and the Ideal Gas Model

Gases behave very differently from liquids and solids because gas particles are far apart, moving rapidly, and interacting relatively weakly most of the time. AP Chemistry uses gases to strengthen your particle model thinking: macroscopic variables (pressure, volume, temperature) are explained by molecular motion.

What pressure really is

Gas pressure comes from collisions of gas particles with the walls of a container. Each collision exerts a tiny force; the sum over countless collisions becomes the measurable pressure.

So when something changes pressure, the “why” is usually one of these:

• Particles collide more often (higher number density, smaller volume)
• Particles hit harder (higher speed from higher temperature)

Temperature and kinetic energy

In the particle model, temperature is proportional to the average kinetic energy of particles. For an ideal gas, the average kinetic energy per particle depends only on absolute temperature.

The kinetic energy of a single gas molecule can be written as:

KE = 0.5mv^2

Here, m is the mass of the molecule in kilograms, v is the speed in meters per second, and KE is in joules.

If different gases are present at the same temperature, they have the same average kinetic energy, but not the same average speed. Lighter molecules must move faster on average than heavier molecules to have the same kinetic energy.

The ideal gas law

The ideal gas law relates pressure, volume, amount of gas, and temperature:

PV = nRT

Where P is pressure, V is volume, n is moles of gas, T is temperature in kelvin, and R is the gas constant (value depends on units chosen). A commonly used value is:

R = 0.0821 \frac{L\cdot atm}{mol\cdot K}

Temperature must be in kelvin. A practical conversion is:

T(K) = T(^\circ C) + 273.15

Combined gas law (moles constant)

When the number of moles is constant:

\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}

This is just the ideal gas law applied to two states with constant n.

It’s also helpful to recognize the classic special-case relationships:

• If volume is constant: as pressure increases, temperature increases.
• If temperature is constant: as pressure increases, volume decreases (Boyle’s Law).
• If pressure is constant: as temperature increases, volume increases (Charles’s Law).

Dalton’s law of partial pressures

When multiple gases share a container, each gas contributes to the total pressure as if it were alone:

P_{total} = P_1 + P_2 + P_3 + \dots

This is especially important when gases are collected over water (water vapor contributes to total pressure).

Mole fraction and partial pressure

The mole fraction of a gas component is:

X_i = \frac{n_i}{n_{total}}

For ideal gases:

P_i = X_iP_{total}

A common procedure for partial pressure is: divide moles of that gas by total moles (mole fraction), then multiply by total pressure.

Kinetic molecular theory (KMT)

The kinetic molecular theory is the set of assumptions used to derive ideal gas behavior. Core assumptions (idealized):

• Gas particles have negligible volume compared to the container.
• Particles are in constant random motion.
• Collisions are elastic (no net loss of kinetic energy).
• There are no attractive or repulsive forces between particles (except during collisions).

Maxwell–Boltzmann diagrams (velocity distributions)

A Maxwell–Boltzmann diagram shows the range of velocities for molecules of a gas.

• For the same gas at different temperatures: higher temperature produces a broader distribution with the peak shifted to higher speeds (more variation of velocities, and a higher average speed).
• For different gases at the same temperature: average kinetic energy is the same, but the lighter gas has higher typical speeds while the heavier gas has lower typical speeds.

Effusion

Effusion is the rate at which a gas escapes from a container through microscopic holes, moving from higher pressure to lower pressure. A balloon slowly losing gas over time is a common everyday example.

Effusion rate depends on:

• Speed of the molecules (faster molecules collide more often with container walls, increasing chances of finding a hole)
• Temperature (higher temperature increases molecular speeds, increasing effusion rate)
• Molar mass (lower molar mass gases move faster at the same temperature and effuse more quickly)

Gas density and molar mass from density

Density of a gas can be calculated using:

D = \frac{m}{V}

To find molar mass (MM) from density, temperature, and pressure:

MM = \frac{DRT}{P}

Worked problem: ideal gas law

Example: A sample contains 0.500 mol of gas in a 10.0 L container at 298 K. If the pressure is needed in atmospheres, use the ideal gas law with an appropriate R.

Use:

PV = nRT

Rearrange:

P = \frac{nRT}{V}

Substitute values using:

R = 0.08206 \frac{L\cdot atm}{mol\cdot K}

Then:

P = \frac{(0.500)(0.08206)(298)}{10.0}

Compute (conceptually): multiply numerator then divide by 10. This gives approximately 1.22 atm.

Common checking step: pressure near 1 atm is reasonable for about half a mole in 10 L near room temperature.

Worked problem: partial pressures

Example: A gas mixture contains 2.0 mol nitrogen and 1.0 mol oxygen at a total pressure of 3.0 atm. Find the partial pressure of oxygen.

First find mole fraction of oxygen:

X_{O_2} = \frac{1.0}{2.0 + 1.0} = \frac{1.0}{3.0}

Then:

P_{O_2} = X_{O_2}P_{total} = \frac{1}{3}(3.0\ atm) = 1.0\ atm

Exam Focus

• Typical question patterns:
  • Use the ideal gas law to solve for P, V, n, or T; often combined with stoichiometry (moles from a reaction).
  • Use Dalton’s law and mole fractions to find partial pressures in mixtures.
  • Explain pressure changes using collision reasoning (not just formula manipulation).
• Common mistakes:
  • Using Celsius instead of kelvin.
  • Mixing incompatible units for R (for example, using liters with an R in different volume units).
  • Forgetting that collecting gas over water requires subtracting water vapor pressure when determining the gas’s partial pressure.

Deviations from Ideal Gas Behavior (Real Gases)

The ideal gas model works best when gas particles are far apart and moving fast—conditions typically achieved at high temperature and low pressure. Under other conditions, real gases deviate because two ideal assumptions fail:

1. Particles do have volume.
2. Particles do attract each other.

Why attractions matter

At moderate to high pressures, particles are closer together, so intermolecular attractions become more relevant. Attractions pull particles inward and reduce the force of collisions with container walls. That means the observed pressure can be lower than the ideal prediction.

This effect is most noticeable:

• At lower temperatures (particles move slower, so attractions matter more)
• For gases with stronger IMFs (polar molecules, larger molecules)

Why particle volume matters

At high pressures, the volume taken up by gas particles is not negligible compared to the container volume. That means there is less free space for motion than the container volume suggests. This tends to increase the observed pressure relative to ideal predictions at very high pressures.

Recognizing when ideal assumptions break

AP questions commonly ask you to reason qualitatively:

• A gas is more ideal at high temperature, low pressure.
• A gas deviates more at low temperature, high pressure.
• Larger, more polarizable gases (stronger dispersion) deviate more than small, nonpolar gases.

Gases under normal conditions can deviate slightly if they have strong IMFs.

A misconception to avoid: “Real gases always have lower pressure than ideal.” Not always. Attractions can lower pressure, but finite particle volume can raise pressure at sufficiently high pressure. The key is which non-ideal effect dominates under given conditions.

Worked example: which gas is most ideal?

Example: At the same temperature and pressure, which behaves most ideally: He, CO2, or NH3?

Reasoning: Helium is small and nonpolar with very weak dispersion forces, so it has minimal attractions and minimal particle volume effects compared with larger, more interactive molecules. CO2 is larger and more polarizable; NH3 is polar and has stronger attractions than He. So He is most ideal.

Exam Focus

• Typical question patterns:
  • Explain (qualitatively) why a gas deviates from ideal behavior at certain T and P.
  • Compare gases (size, polarity) to decide which deviates more.
  • Interpret particle-level diagrams: closer particles imply higher pressure or lower temperature conditions and more non-ideal behavior.
• Common mistakes:
  • Claiming ideal behavior improves at high pressure (it’s the opposite).
  • Ignoring molecular size and polarizability when comparing deviations.
  • Treating “non-ideal” as a separate law to memorize instead of a consequence of attractions and finite volume.

Solutions and Mixtures: Forming, Describing, and Predicting

A solution is a homogeneous mixture at the particle level (uniform composition throughout). The component present in the greatest amount is usually the solvent, and the dissolved component is the solute.

Understanding solutions requires blending IMF reasoning with energy ideas: dissolving happens when solute–solvent attractions can replace (or compete successfully with) solute–solute and solvent–solvent attractions.

“Like dissolves like” (and what it really means)

The phrase like dissolves like is shorthand for IMF compatibility:

• Polar solvents (like water) dissolve polar solutes and ionic compounds well because they can form strong attractions (dipole–dipole, hydrogen bonding, ion–dipole).
• Nonpolar solvents dissolve nonpolar solutes well because dispersion forces can be similar in both substances.

Examples:

• NaCl dissolves in water because ion–dipole interactions between ions and water compensate for breaking the ionic lattice and disrupting some water–water hydrogen bonding.
• Oil and water are immiscible because water–water hydrogen bonding is strong, and oil cannot replace those interactions effectively.

When an ionic solute dissociates into ions, the presence of mobile ions can allow the solution to conduct electricity (electrolyte behavior).

Energetics of dissolution (qualitative)

Dissolving can be analyzed in steps:

1. Separate solute particles (endothermic)
2. Separate solvent particles (endothermic)
3. Form solute–solvent attractions (exothermic)

The net enthalpy change (often called enthalpy of solution) depends on the balance of these steps.

Solubility and miscibility

Solubility describes how much solute can dissolve in a given amount of solvent at a specified temperature. When no more solute dissolves, the solution is saturated.

For liquids, the analogous idea is miscibility:

• Miscible liquids mix in all proportions (for example, ethanol and water).
• Immiscible liquids form separate layers (for example, hexane and water).

Temperature effects are sometimes addressed qualitatively: many solid solutes become more soluble in water as temperature increases, but there are exceptions.

Concentration: ways to represent solutions

Molarity

Molarity (M) is moles of solute per liter of solution:

M = \frac{moles\ solute}{liters\ solution}

Common pitfall: using liters of solvent rather than liters of solution.

Molality

Molality (m) is moles of solute per kilogram of solvent:

m = \frac{moles\ solute}{kilograms\ solvent}

Molality is useful when temperature changes matter because mass does not change with temperature the way volume can.

Mole fraction

Mole fraction is:

X_i = \frac{n_i}{n_{total}}

Mass percent and parts per million

Mass percent is:

\%\ by\ mass = \frac{mass\ solute}{mass\ solution}\times 100

Very dilute solutions are sometimes described with ppm (parts per million).

Dilution (particle counting logic)

When you dilute a solution, you add solvent, which changes volume but not moles of solute (assuming no reaction):

M_1V_1 = M_2V_2

Worked problem: preparing a solution

Example: You dissolve 5.85 g of NaCl in enough water to make 0.500 L of solution. Find the molarity.

1. Convert grams to moles using molar mass (NaCl about 58.44 g/mol).

n = \frac{5.85}{58.44}

This is about 0.100 mol.

2. Use molarity definition:

M = \frac{n}{V} = \frac{0.100}{0.500} = 0.200\ M

Worked problem: dilution

Example: How much 6.0 M HCl is needed to make 250. mL of 1.5 M HCl?

M_1V_1 = M_2V_2

Solve for V1:

V_1 = \frac{M_2V_2}{M_1} = \frac{(1.5)(0.250)}{6.0}

This gives 0.0625 L, which is 62.5 mL.

Exam Focus

• Typical question patterns:
  • Predict whether a solute dissolves in a solvent using polarity and IMFs (“like dissolves like”).
  • Calculate concentration (especially molarity) from mass and volume; perform dilution calculations.
  • Interpret particle diagrams of solutions (ions separated and solvated, molecules dispersed uniformly).
• Common mistakes:
  • Treating “dissolves” as meaning “reacts” (dissolution is typically physical separation and solvation, not chemical change).
  • Mixing up molarity and molality definitions (liters of solution vs kilograms of solvent).
  • Using liters of solvent instead of liters of solution when computing molarity.

Representations of Solutions and Mixtures (Particle Diagrams and Reasoning)

AP Chemistry places real weight on your ability to interpret representations: particulate diagrams, molecular drawings, and symbolic equations. For solutions, that means being fluent in what it looks like when something dissolves.

What it means to dissolve an ionic compound

When an ionic compound dissolves in water, the ions separate and become surrounded by oriented water molecules.

• Around a cation, the oxygen ends of water (partially negative) point toward the ion.
• Around an anion, the hydrogen ends (partially positive) point toward the ion.

This is an ion–dipole interaction story, not a “breaking covalent bonds” story.

A common diagram mistake is drawing ion pairs still stuck together in solution. In a typical aqueous solution of a soluble ionic compound, most ions are separated and solvated.

What it means to dissolve a molecular compound

When a molecular compound dissolves, intact molecules disperse among solvent molecules. Whether they dissolve well depends on whether solvent–solute attractions can form.

Examples:

• Ethanol dissolves in water because it can hydrogen bond (through its O–H) and also interact via dispersion.
• Hexane does not dissolve in water because it cannot form strong interactions to replace water–water hydrogen bonding.

Electrolytes vs nonelectrolytes (conceptual)

A solution conducts electricity if it contains mobile charged particles.

• Strong electrolytes: soluble ionic compounds and strong acids (form many ions)
• Weak electrolytes: weak acids/bases (form some ions)
• Nonelectrolytes: molecular substances that do not form ions (for example, sugar)

In Unit 3 context, the main representational idea is: do you draw ions or intact molecules?

Worked example: interpreting a particulate diagram

Example: A diagram shows water molecules with separated spheres labeled as Na and Cl dispersed throughout. What does that represent and what interactions stabilize it?

Reasoning: It represents an aqueous solution of sodium chloride where the crystal lattice has dissociated into Na plus and Cl minus ions. The stabilizing interactions are ion–dipole attractions between ions and polar water molecules.

Exam Focus

• Typical question patterns:
  • Choose the correct particulate diagram for a dissolved ionic vs molecular solute.
  • Explain why ions are oriented with water molecules in a particular way.
  • Identify whether a solution should conduct based on whether ions are present.
• Common mistakes:
  • Drawing ions but labeling a molecular solute (like glucose) as if it dissociates.
  • Confusing “polar” with “ionic” (polar covalent molecules do not automatically form ions).
  • Forgetting orientation: oxygen toward cations, hydrogen toward anions.

Separation of Solutions and Mixtures (with an IMF Perspective)

Separation techniques work because different substances interact differently with a given environment—often through IMFs. A separation method exploits differences in physical properties such as boiling point, solubility, or affinity for a surface.

Chromatography: the big idea

Chromatography separates components based on different attractions to a stationary phase (fixed material) and a mobile phase (moving solvent or gas).

• If a substance is more attracted to the stationary phase, it moves more slowly.
• If it is more attracted to the mobile phase, it moves more quickly.

These attractions are usually IMFs: polarity-based interactions, hydrogen bonding, and dispersion with the stationary phase.

Paper chromatography (common AP context)

In paper chromatography:

• The stationary phase is the paper (cellulose, which is polar and can hydrogen bond).
• The mobile phase is a solvent that rises up the paper via capillary action.

Operationally, paper is suspended with one end submerged in the solvent; the solvent climbs up the paper and can pull components of a mixture apart as it rises (a classic example is separating black ink into different dye colors).

More polar components often interact more strongly with the polar paper and travel less far, though actual outcomes depend strongly on the solvent choice.

Many courses use the retention factor:

R_f = \frac{distance\ traveled\ by\ solute}{distance\ traveled\ by\ solvent\ front}

Conceptually:

• Stronger attraction to the mobile phase (solvent front) means the solute spends more time moving with the solvent and tends to have a larger Rf.
• Stronger attraction to the stationary phase means the solute is retained more and tends to have a smaller Rf.

Column chromatography

In column chromatography, the same stationary vs mobile phase idea applies, but the stationary phase is packed into a column. Components travel different distances (or elute at different times) largely due to polarity-dependent attractions to the stationary material versus solubility in the mobile solvent.

Distillation (boiling point differences)

Distillation separates liquids based on differences in boiling point, which reflect IMF differences.

• The component with lower boiling point (weaker IMFs) vaporizes more readily and can be condensed separately.

Extraction (polarity-based layer separations)

In liquid–liquid extraction, a solute partitions between two immiscible solvents (often water and an organic solvent). The solute prefers the layer where it has stronger solute–solvent attractions—“like dissolves like” used as a separation tool.

Worked example: chromatography reasoning

Example: Two dyes are separated by paper chromatography using a moderately polar solvent. Dye A travels much farther than Dye B. What does that suggest?

Reasoning: Dye A has greater affinity for the mobile phase relative to the stationary phase, so it spends more time dissolved in the solvent moving upward. Dye B likely has stronger interactions with the paper (stationary phase), often meaning Dye B is more polar or better able to hydrogen bond with cellulose under these conditions.

Exam Focus

• Typical question patterns:
  • Given a chromatogram, compare travel distances and infer relative affinity for stationary vs mobile phases.
  • Explain how changing solvent polarity could change separation outcomes.
  • Connect distillation feasibility to boiling point differences using IMF reasoning.
• Common mistakes:
  • Saying “the faster one is lighter” (mass is not the primary driver in chromatography; intermolecular attractions are).
  • Assuming “more polar always travels farther” without considering that a polar stationary phase can retain polar solutes.
  • Calculating Rf using inconsistent reference points (distances must be from the initial spot).

Spectroscopy Connections: Electromagnetic Radiation and Beer’s Law

Several lab techniques discussed alongside solutions rely on electromagnetic radiation and absorption. While not an IMF topic by itself, it commonly appears in chemistry courses in connection with measuring solution concentration.

Electromagnetic radiation relationships

When an electron changes energy levels, the energy absorbed or emitted is related to the frequency of electromagnetic radiation:

E = hv

Here, h is Planck’s constant (6.626 x 10^-34 joule seconds), and the frequency is in per second.

Wavelength can be determined from frequency using:

c = \lambda \nu

Here, c is the speed of light, lambda is wavelength, and nu is frequency.

Beer’s Law (absorbance and concentration)

A spectrometer can be used to measure how the concentration of a solution changes over time by tracking absorbance. Beer’s Law can be written as:

A = abc

In this form, A is absorbance, a is molar absorptivity of the solution, b is pathlength (how far light travels through the solution), and c is concentration.

Exam Focus

• Typical question patterns:
  • Use the relationship between energy and frequency to reason about absorbed light.
  • Use Beer’s Law to connect absorbance to concentration (often conceptually: higher concentration leads to higher absorbance if pathlength and molar absorptivity are constant).
• Common mistakes:
  • Confusing frequency and wavelength trends (they are inversely related through the speed of light).
  • Treating Beer’s Law variables as unitless without keeping track of what each symbol represents (especially pathlength and concentration).