Eighteen focused outcomes connect molecular structure, interparticle attraction, phase behavior, vapor pressure, solid organization, and observable properties across the six official topics.
Structure suggests behavior only when the comparison is controlled.
The ADA names six Liquids and Solids topics but does not publish topic-level weights or question quotas. These outcomes organize study; they do not predict how many questions will appear.
Polarity, size, shape, particle type, temperature, and pressure can compete. The map states when a trend is safe and when supplied measurements or a fuller model are required. Solution polarity and miscibility remain in the separate Solutions domain.
Evidence sequence
Three layers before a property claim.
01
Name the particles
Atoms, ions, discrete molecules, or an extended network? Crystalline or amorphous?
02
Inventory attractions
Include dispersion for every molecular substance, then add supported permanent-dipole, hydrogen-bond, ionic, metallic, or network interactions.
03
Control the comparison
Check temperature, pressure, size, shape, and measured evidence before linking attraction to vapor pressure, boiling, flow, or solid behavior.
Official hierarchy → learning sequence
Six branches. Eighteen outcomes.
Open a topic to see its outcome, must-know relationships, model boundary, prerequisite sequence, misconception correction, and free references.
AIntermolecular forces3 objectives
Objective 1
Interparticle versus intraparticle forces
Distinguish attractions between particles from the ionic, covalent, or metallic bonding that determines the identity of each particle or extended solid.
Must know
Intermolecular forces act between distinct molecules or condensed-phase particles; intramolecular covalent bonds hold the atoms of one molecule together. Melting or boiling a molecular substance primarily overcomes attractions between molecules rather than breaking the covalent bonds within each molecule.
Model boundary
For ionic, metallic, and covalent-network solids, name the extended bonding model rather than forcing every attraction into a molecular-force label.
Misconception
Boiling water breaks each O–H covalent bond. Vaporization separates intact H₂O molecules by overcoming intermolecular attractions; it does not decompose water.
Dispersion, dipole, hydrogen-bond, and ion–dipole recognition
Identify the attractions available from particle charge, molecular polarity, polarizability, and the presence of a valid hydrogen-bond donor and partner.
Must know
London dispersion attractions occur between all atoms and molecules; permanent dipole–dipole attractions additionally occur between polar molecules. A standard hydrogen-bond donor has H covalently bonded to N, O, or F; ion–dipole attraction acts between an ion and a polar molecule.
Model boundary
A molecule may experience several attractions at once. Classify all supported interactions instead of assigning one exclusive label; detailed solution energetics belongs to the Solutions domain.
Misconception
A nonpolar molecule has no intermolecular attractions. Every atom and molecule can form instantaneous dipoles and therefore experiences London dispersion attraction.
Compare intermolecular attraction using charge type, polarity, polarizability, molecular size, and contact shape while controlling competing variables.
Must know
More easily distorted electron clouds are more polarizable and generally support stronger dispersion attractions. For similar substances, greater surface contact can strengthen dispersion; a simple force-name hierarchy is unsafe when size, shape, and multiple attractions differ.
Model boundary
Do not assign exact force energies or a universal boiling-point order without enough structural or measured evidence.
Misconception
Hydrogen bonding always makes a substance’s attractions stronger than those of every larger nonpolar substance. Hydrogen bonding is important, but a much larger or more polarizable species can have substantial dispersion attraction; compare all relevant evidence.
Earlier objectives in this map
Dispersion, dipole, hydrogen-bond, and ion–dipole recognition
BPhase changes3 objectives
Objective 1
Six phase changes and energy direction
Name melting, freezing, vaporization, condensation, sublimation, and deposition and determine whether the substance absorbs or releases energy during each transition.
Must know
Melting, vaporization, and sublimation are endothermic for the substance; freezing, condensation, and deposition are their exothermic reverses. A phase change changes particle arrangement and separation without changing chemical identity.
Model boundary
Energy signs are stated for the substance as system; the surroundings have the opposite heat direction.
Misconception
Condensation absorbs heat because gas particles must lose energy. The condensing substance releases energy as particles enter the lower-energy liquid phase.
Earlier objectives in this map
Interparticle versus intraparticle forces
Objective 2
Heating curves and latent heat
Interpret a heating or cooling curve and calculate sensible or phase-change heat when the needed heat capacity or phase-change enthalpy is supplied.
Must know
Within one phase, added heat changes temperature according to the supplied heat capacity; during a pure-substance transition at fixed pressure, temperature remains constant while phase proportions change. Use q = mcΔT on sloped segments and the supplied molar or mass phase-change enthalpy on a plateau.
Model boundary
Do not apply a plateau rule to an unspecified mixture or changing-pressure path; constants and phase-change quantities must be supplied when calculation is required.
Misconception
Temperature must rise continuously whenever heat enters a sample. During a phase transition, added energy changes particle separation and phase proportion while temperature can remain constant.
Earlier objectives in this map
Six phase changes and energy direction
Objective 3
Phase diagrams and special points
Read pressure–temperature phase regions, equilibrium boundaries, the triple point, the critical point, and pressure-dependent melting or boiling from a supplied phase diagram.
Must know
A boundary line represents two phases in equilibrium; the triple point joins three boundaries and the critical point terminates the liquid–gas boundary. Below the triple-point pressure a liquid phase is unavailable, and above the critical temperature pressure alone cannot produce an ordinary liquid–gas boundary.
Model boundary
Infer the stable phase and transition path from the supplied diagram; do not assume every solid–liquid line has water’s unusual slope.
Misconception
Every point inside a phase region contains two phases in equilibrium. A single region represents one stable phase; two phases coexist along a boundary.
Earlier objectives in this map
Six phase changes and energy direction
CVapor pressure3 objectives
Objective 1
Evaporation and dynamic vapor equilibrium
Explain equilibrium vapor pressure as the pressure of vapor above its condensed phase when evaporation and condensation rates are equal in a closed system.
Must know
Evaporation occurs from a liquid surface at temperatures below the boiling point; condensation returns vapor particles to the liquid. In a closed container with both phases present at fixed temperature, dynamic equilibrium has equal forward and reverse rates, not stopped molecular motion.
Model boundary
Equilibrium vapor pressure is a substance-and-temperature property while condensed phase remains; an open container may lose vapor and fail to establish that equilibrium.
Misconception
At vapor-pressure equilibrium, evaporation and condensation both stop. Both processes continue at equal rates, so macroscopic amounts remain constant.
Earlier objectives in this map
Six phase changes and energy direction
Objective 2
Temperature, attraction, and volatility
Predict vapor-pressure and volatility changes from temperature and defensible intermolecular-attraction comparisons.
Must know
A liquid’s equilibrium vapor pressure rises with temperature because a larger fraction of particles can escape the condensed phase. At the same temperature, otherwise comparable liquids with stronger attractions generally have lower vapor pressure and lower volatility.
Model boundary
Do not use one force label to rank structurally dissimilar substances when polarizability, shape, or measured data could reverse the comparison.
Misconception
A liquid with stronger intermolecular attraction has a higher vapor pressure because its molecules pull harder. Stronger attraction holds molecules in the liquid, reducing the equilibrium vapor pressure at the same temperature.
Earlier objectives in this map
Evaporation and dynamic vapor equilibrium · Relative attraction and polarizability
Objective 3
Boiling point and external pressure
Use vapor-pressure curves or supplied data to locate boiling conditions and predict how external pressure changes boiling temperature.
Must know
Boiling occurs when a liquid’s vapor pressure equals the external pressure; normal boiling point is defined at 1 atm. Lower external pressure lowers the boiling temperature, while higher external pressure raises it.
Model boundary
Read quantitative vapor pressures from supplied curves or data; Clausius–Clapeyron calculations require a supplied equation and constants.
Misconception
A liquid boils at one fixed temperature regardless of pressure. The boiling temperature is the point where vapor pressure matches the current external pressure.
Earlier objectives in this map
Temperature, attraction, and volatility
DStructures3 objectives
Objective 1
Crystalline and amorphous solids
Distinguish long-range crystalline order from amorphous structure and connect that difference to melting behavior.
Must know
Crystalline solids have a repeating long-range arrangement; amorphous solids lack that long-range periodic order. A pure crystalline solid has a characteristic melting temperature, while an amorphous solid generally softens across a range.
Model boundary
Amorphous does not mean particles have no local organization, and a sketch alone supports only the structural features it actually shows.
Misconception
Every solid is a perfectly ordered crystal. Glasses and many other solids are amorphous and lack long-range repeating order.
Earlier objectives in this map
Interparticle versus intraparticle forces
Objective 2
Four major crystalline-solid types
Classify ionic, metallic, covalent-network, and molecular solids from their constituent particles, dominant attractions, and structural evidence.
Must know
Ionic solids contain ions, metallic solids contain metal atoms with delocalized electrons, covalent-network solids contain bonded atomic networks, and molecular solids contain discrete molecules. The dominant bonding or attraction model—not appearance alone—supports the classification.
Model boundary
Real materials can contain defects or mixed structural features; classify the idealized evidence supplied in an introductory prompt.
Misconception
Any solid made from covalent bonds is a covalent-network solid. Molecular solids contain discrete covalent molecules held together by intermolecular attractions; network solids have covalent bonding extended through the solid.
Earlier objectives in this map
Crystalline and amorphous solids · Dispersion, dipole, hydrogen-bond, and ion–dipole recognition
Count effective particles and identify basic simple-cubic, body-centered-cubic, or face-centered-cubic packing from a supplied unit-cell representation.
Must know
A corner particle contributes one eighth to one cubic unit cell, a face-centered particle one half, and a body-centered particle one whole. Simple cubic, body-centered cubic, and face-centered cubic cells contain 1, 2, and 4 effective particles respectively in their conventional unit cells.
Model boundary
Coordination numbers or packing efficiencies may be interpreted when supplied or explicitly represented; Miller indices, diffraction derivations, and complex ionic lattices are outside this objective.
Misconception
Every particle drawn on a unit-cell boundary belongs entirely to that one cell. Boundary particles are shared with neighboring cells and contribute fractional counts.
Earlier objectives in this map
Four major crystalline-solid types
EPolarity3 objectives
Objective 1
Bond dipoles and net molecular polarity
Determine molecular polarity by combining bond-dipole directions with three-dimensional molecular geometry and symmetry.
Must know
Polar bonds do not guarantee a polar molecule because vector dipoles can cancel in a symmetric geometry. A net molecular dipole requires both bond polarity and a geometry whose dipole vectors do not cancel.
Model boundary
Use a valid structure and geometry; do not infer polarity from formula order or electronegativity difference alone.
Misconception
Any molecule containing polar bonds must be polar. Symmetric arrangements such as linear CO₂ can cancel equal bond dipoles and produce no net molecular dipole.
Use molecular polarity to determine which permanent-dipole attractions are possible without forgetting universal dispersion attraction or hydrogen-bond structural requirements.
Must know
Polar molecules can align permanent dipoles; nonpolar molecules lack a permanent molecular dipole but still experience dispersion. Hydrogen bonding is not a synonym for polarity: it requires the appropriate H–N, H–O, or H–F donor structure and a suitable partner.
Model boundary
Molecular polarity alone does not quantify total attraction; polarizability, size, shape, and ionic charge may also matter.
Misconception
Every polar molecule hydrogen-bonds with itself. Many polar molecules have dipole–dipole attraction but lack an H bonded to N, O, or F and therefore are not hydrogen-bond donors.
Earlier objectives in this map
Bond dipoles and net molecular polarity · Dispersion, dipole, hydrogen-bond, and ion–dipole recognition
Objective 3
Evidence-controlled polarity comparisons
Use polarity together with controlled size, shape, and phase-property evidence to support a bounded comparison between condensed substances.
Must know
For closely matched molecular size and shape, a permanent dipole can help explain stronger attraction, higher boiling point, or lower vapor pressure. When several structural variables change, measured data or a fuller interaction analysis is needed before claiming polarity is the sole cause.
Model boundary
Solvent–solute miscibility and quantitative solution behavior belong to the separately published Solutions domain.
Misconception
The more polar molecule always has the higher boiling point, regardless of size or shape. Polarity is one contributor; dispersion and molecular contact can dominate when structures differ substantially.
Earlier objectives in this map
Polarity and available attractions · Relative attraction and polarizability
FProperties3 objectives
Objective 1
Viscosity and molecular motion
Predict qualitative viscosity changes from temperature, intermolecular attraction, and molecular size or shape when the comparison controls competing variables.
Must know
Viscosity is resistance to flow; stronger attractions and structures that hinder sliding generally increase it. For a given liquid, viscosity generally decreases as temperature rises because molecular motion more readily overcomes attractions.
Model boundary
Do not compute viscosity or rank unrelated liquids without supplied measurements or enough structural evidence.
Misconception
Heating a liquid makes it more viscous because its particles move faster. For most liquids, higher temperature lets particles move past one another more readily, so viscosity decreases.
Earlier objectives in this map
Relative attraction and polarizability
Objective 2
Cohesion, adhesion, and surface behavior
Explain surface tension, wetting, meniscus shape, and capillary rise or depression by comparing cohesive and adhesive attractions.
Must know
Cohesion acts between like particles and contributes to surface tension; adhesion acts between a liquid and a different surface. A concave or convex meniscus and capillary rise or depression depend on the relative strengths of liquid–surface adhesion and liquid–liquid cohesion.
Model boundary
Capillary-height calculations require the relevant equation and quantities; density alone does not determine the meniscus direction.
Misconception
Capillary rise occurs only because the tube pushes the liquid upward. Liquid–surface adhesion and liquid cohesion produce the rise; tube radius and other supplied properties affect its extent.
Earlier objectives in this map
Dispersion, dipole, hydrogen-bond, and ion–dipole recognition
Objective 3
Structure–property patterns in solids
Relate solid type and particle mobility to qualitative melting behavior, hardness, brittleness or malleability, and electrical conductivity.
Must know
Ionic solids are typically brittle and conduct when ions become mobile, metallic solids conduct and are malleable, network solids are often hard with high melting points, and molecular solids are generally nonconducting with lower melting points. Conductivity depends on mobile charged particles or delocalized electrons, not simply the presence of charged nuclei.
Model boundary
These are introductory patterns with exceptions such as conductive graphite; use supplied structural evidence instead of treating every descriptor as universal.
Misconception
A solid ionic compound conducts electricity because it contains ions. Its ions are fixed in the solid lattice; conductivity appears when the ions can move in the molten or dissolved state.