AP Chemistry Unit 2: Compound Structure & Properties

Vocabulary you must be able to define

• Electronegativity (EN): a measure of how strongly an atom pulls on the shared electrons in a bond. Higher EN means a stronger pull. It rises across a period (more protons, same shell) and falls down a group (the valence shell is farther from the nucleus and better shielded).
• Polarity (of a bond): an uneven distribution of electron density in a bond, caused by a difference in electronegativity (ΔEN) between the two atoms. A bond with an uneven distribution is polar; one with an even distribution is nonpolar.
• Partial charge (δ+ / δ−): the fraction of a charge that builds up on an atom in a polar bond. The more electronegative atom holds electron density longer, so it is δ−, and the atom it pulled away from is δ+. These are less than one full electron of charge, which is why they are called partial.
• Net (full, formal) charge: what forms when an electron is transferred outright rather than shared. Na loses one electron entirely and becomes Na⁺; Cl gains it and becomes Cl^-. Ions carry whole charges; atoms in covalent bonds carry partial ones.

Vocabulary you must be able to define

• Internuclear separation: the distance between the nuclei (the centers) of two atoms. It is the r in Coulomb’s law and the x-axis of every potential energy curve. At the energy minimum, the internuclear separation is called the bond length.
• Potential energy: stored energy that depends on position — here, on how far apart the charged particles are. Two particles infinitely far apart are defined as having zero potential energy. Attraction pulls the energy below zero (negative, more stable); repulsion pushes it above zero (positive, less stable).
• Bond energy (bond enthalpy): the energy required to break one mole of a bond in the gas phase, in kJ/mol. It equals the depth of the well below zero. Breaking bonds always costs energy (endothermic, +); forming bonds always releases it (exothermic, −).
• Lattice energy: the Coulombic version of bond energy for an ionic solid — the energy change when one mole of gaseous ions comes together into a crystal lattice (see Section 3).

How to read the diagram, and why it matters

• The axis is energy, not time. Higher on the page means higher potential energy and less stability. The two horizontal lines are the two states being compared: loose gaseous ions on top, the assembled crystal on the bottom.
• The drop is the lattice energy. Na⁺(g) + Cl^-(g) → NaCl(s) releases 787 kJ/mol, so ΔH = −787 kJ/mol. The ions attract, fall to lower energy, and the surroundings warm up. This is just the deep well of a potential energy curve redrawn as levels.
• The dashed arrow up is the same number reversed. To rip the solid apart into free gaseous ions you must supply +787 kJ/mol. That is why the size of the drop is the strength of the ionic bonding.
• Why it is important: a bigger drop predicts a higher melting point, a harder and less soluble solid, and a smaller heat of solution. Every “rank these by melting point” ionic question is really asking you to rank the depth of this drop, and Coulomb’s law tells you the depth from charge and radius alone.
• Sign warning: some sources define lattice energy as the positive energy needed to break the solid apart. Both describe the same 787 kJ/mol. State the process you mean (“energy released when the lattice forms”) and your sign can never be marked wrong.

Vocabulary you must be able to define

• Malleable: able to be hammered, pressed, or rolled into thin sheets without shattering (gold leaf, aluminum foil). The cations can slide to new positions because the electron sea is nondirectional and simply flows with them.
• Ductile: able to be drawn out into a wire. Same cause as malleability, just pulled in one direction instead of flattened (copper wiring).
• Conductivity: the ability to let electric charge or heat flow through a material. Metals conduct because the delocalized valence electrons are free to move through the entire solid, carrying charge and transferring kinetic energy.
• Luster: the shiny, reflective appearance of a metal surface. The mobile electrons absorb and immediately re-emit visible light of all wavelengths instead of letting it pass through.

Contrast with an ionic solid, which is brittle: shifting a layer lines up like charges, and the repulsion cracks the crystal. Metals are malleable for exactly the reason ionic solids are not.

Six steps that always work

1. Count total valence electrons. Add one per negative charge and subtract one per positive charge.
2. Pick the central atom: the least electronegative atom that is not hydrogen (carbon is central whenever it is present).
3. Connect every outer atom to the center with a single bond.
4. Complete the octets of the outer atoms with lone pairs, giving hydrogen nothing extra.
5. Place any electrons still left over on the central atom.
6. If the central atom is short of an octet, pull a lone pair from an outer atom in to make a double or triple bond.

Vocabulary: isomer and isomeric

• Isomer: one of two or more compounds that have the same molecular formula but a different arrangement of atoms. Same atoms, same count, different connectivity or different geometry — and therefore different properties.
• Isomeric: the adjective. Two substances are “isomeric” (or are “isomers of each other”) when they stand in that relationship, as in “ethanol and dimethyl ether are isomeric C₂H₆O compounds.”
• Structural (constitutional) isomers: the atoms are connected in a different order, as in ethanol (C–C–O–H) versus dimethyl ether (C–O–C).
• Geometric (cis/trans) isomers: the connectivity is the same but the groups are locked on the same side (cis) or opposite sides (trans) of a double bond, which cannot rotate. Cis-1,2-dichloroethene is polar; the trans form is nonpolar.

Do not confuse isomers with resonance structures (Section 6). Isomers are genuinely different substances you could bottle separately, because the atoms moved. Resonance structures are two drawings of one substance, because only the electrons moved.

Vocabulary you must be able to define

• Dipole: a separation of positive and negative charge across a bond or a molecule, with a δ+ end and a δ− end.
• Dipole moment (μ): the size of that separation, measured as μ = q × d — the amount of separated charge times the distance between the centers of charge. It is a vector, drawn as an arrow pointing toward the δ− end with a plus sign on the tail, and is reported in debyes (D). A larger ΔEN or a longer bond gives a larger dipole moment. Bond dipole moments add as vectors, so a molecule’s net (molecular) dipole moment is zero when the shape makes them cancel: μ = 0 for CO₂, 1.85 D for H₂O.
• Enthalpy (H): the heat content of a system at constant pressure. Only the change, ΔH, is measurable. ΔH < 0 is exothermic (heat released, products more stable); ΔH > 0 is endothermic (heat absorbed).
• Bond enthalpy (bond dissociation energy): the enthalpy change to break one mole of a specific bond in the gas phase, always positive because breaking bonds costs energy. Forming the same bond releases the same magnitude. This is the number tabulated as “bond energy” and it is the well depth from Section 2.
• Estimating a reaction enthalpy from bond enthalpies: ΔH_rxn ≈ Σ(bonds broken) − Σ(bonds formed). If the bonds in the products are stronger, the reaction is exothermic.

1. Classify the bonding: metal + nonmetal, nonmetal + nonmetal, or metal + metal
2. For an ionic question, go straight to Coulomb’s law: compare charges first, then radii
3. Count total valence electrons and draw the Lewis diagram
4. Check for resonance; if it exists, calculate formal charges and bond order
5. Count electron domains on the central atom, remembering that multiple bonds count once
6. Name the electron geometry, then subtract lone pairs to get the molecular geometry and bond angle
7. Assign hybridization from the domain count: sp, sp², or sp³
8. Judge polarity from bond polarity plus the symmetry of the shape