Compound Structure & Properties: every key term you need (+ practice quiz)
89 flashcard terms for AP Chemistry Unit 2, written to match the course framework. Read them here, drill them as flashcards, or take the 49-question quiz. Free, no account needed.
An attraction between atoms that holds them together. Bonding lowers potential energy, making the bonded state more stable than separate atoms.
Three Types of Bonds
Ionic (metal + nonmetal, electron transfer), covalent (nonmetal + nonmetal, electron sharing), and metallic (metal atoms, shared 'sea' of electrons).
Valence Electrons in Bonding
Only outer-shell (valence) electrons participate in bonding. Atoms bond to achieve a stable, typically full, valence shell (octet).
Octet Rule
Atoms tend to gain, lose, or share electrons to achieve eight valence electrons (a noble-gas configuration). Hydrogen aims for two.
Ionic Bond
Forms when a metal transfers electrons to a nonmetal, creating oppositely charged ions (cation and anion) that attract via electrostatic force.
Cation and Anion
A cation is a positive ion (lost electrons, usually a metal); an anion is a negative ion (gained electrons, usually a nonmetal).
Ionic Compound (Salt)
A crystalline lattice of alternating cations and anions held by strong electrostatic attractions in a repeating 3-D pattern.
Lattice Energy
The energy released when gaseous ions form an ionic solid — a measure of ionic bond strength. Higher lattice energy means a stronger, more stable ionic solid.
Lattice Energy Trends (Coulomb's Law)
Lattice energy increases with larger ionic charges and smaller ionic radii (F ∝ q₁q₂/r²). MgO (2+/2−) has far higher lattice energy than NaCl (1+/1−).
Properties of Ionic Compounds
High melting points, brittle, conduct electricity when molten or dissolved (mobile ions) but not as solids, and often soluble in water.
Why Ionic Solids Are Brittle
Shifting the lattice aligns like charges, and their repulsion shatters the crystal — a structure-property relationship.
Covalent Bond
A bond formed when two nonmetal atoms share one or more pairs of electrons, each attracted to both nuclei.
Single, Double, Triple Bonds
One, two, or three shared electron pairs. More shared pairs (higher bond order) means a shorter, stronger bond.
Bond Order
The number of shared electron pairs between two atoms. Higher bond order → shorter bond length and greater bond energy.
Bond Length and Strength
As bond order increases, bond length decreases and bond energy (strength) increases: triple > double > single.
Molecular (Covalent) Compound
Discrete molecules held internally by strong covalent bonds; between molecules, weaker intermolecular forces give lower melting points than ionic solids.
Nonpolar Covalent Bond
Electrons shared equally between atoms of equal (or very similar) electronegativity (e.g., H–H, C–H). No partial charges.
Polar Covalent Bond
Electrons shared unequally because of an electronegativity difference; the more electronegative atom gains a partial negative charge (δ−).
Bonding is a continuum: small ΔEN = nonpolar covalent; moderate ΔEN = polar covalent; large ΔEN = ionic (electrons effectively transferred).
Metallic Bond
Metal cations sit in a 'sea' of delocalized valence electrons that move freely, bonding the whole structure.
Properties of Metals
Metallic bonding explains conductivity (mobile electrons), malleability and ductility (cations slide without breaking bonds), and luster.
Alloys
Mixtures of metals (or metal + nonmetal). Substitutional alloys replace atoms with similar-sized ones; interstitial alloys fit small atoms into gaps (e.g., steel).
Lewis Dot Structure
A diagram showing an atom's or molecule's valence electrons as dots and lines, used to predict bonding and shape.
Drawing Lewis Structures
Count total valence electrons, place the least electronegative atom in the center, form bonds, then distribute remaining electrons to satisfy octets.
Lone Pairs vs. Bonding Pairs
Bonding pairs are shared between atoms; lone (nonbonding) pairs belong to one atom. Both affect molecular shape.
Exceptions to the Octet Rule
Some atoms have fewer than 8 (H, Be, B) or expanded octets (P, S, Cl and beyond, using d orbitals), and odd-electron species (radicals) exist.
Resonance
When one Lewis structure can't describe a molecule, several equivalent structures (resonance forms) are drawn; the true structure is an average (hybrid).
Resonance Example
In ozone (O₃) or carbonate (CO₃²⁻), the double bond is delocalized, so all bonds are identical and of intermediate length — not one single and one double.
Formal Charge
A bookkeeping charge assigned to each atom in a Lewis structure: valence electrons − (lone-pair electrons + ½ bonding electrons). The best structure minimizes formal charges.
VSEPR Theory
Valence Shell Electron Pair Repulsion: electron domains (bonds and lone pairs) around a central atom spread out as far as possible, determining molecular shape.
Electron Domains
Regions of electron density around a central atom — each single/double/triple bond and each lone pair counts as one domain.
Linear Geometry
Two electron domains, 180° bond angle (e.g., CO₂, BeCl₂).
Trigonal Planar Geometry
Three electron domains, 120° angles, flat triangle (e.g., BF₃). With one lone pair it becomes bent.
Tetrahedral Geometry
Four bonding domains, 109.5° angles (e.g., CH₄) — the classic shape when the central atom has four bonds and no lone pairs.
Trigonal Pyramidal Geometry
Four domains with one lone pair (e.g., NH₃); the lone pair pushes bonds down, giving ~107° angles.
Bent (Angular) Geometry
Four domains with two lone pairs (e.g., H₂O, ~104.5°) or three domains with one lone pair — lone pairs bend the molecule.
Lone pairs repel more strongly than bonding pairs, compressing bond angles below the ideal (e.g., water's 104.5° vs. 109.5°).
Trigonal Bipyramidal & Octahedral
Five domains give trigonal bipyramidal (90°/120°); six give octahedral (90°) — expanded-octet geometries (e.g., PCl₅, SF₆).
Hybridization
The mixing of atomic orbitals to form new equivalent hybrid orbitals that match a molecule's geometry: sp (linear), sp² (trigonal planar), sp³ (tetrahedral).
sp³ Hybridization
Four hybrid orbitals from one s and three p orbitals; matches tetrahedral geometry (four electron domains), as in methane.
sp² Hybridization
Three hybrid orbitals (one s + two p); matches trigonal planar geometry (three domains); the leftover p orbital allows a pi bond.
sp Hybridization
Two hybrid orbitals (one s + one p); matches linear geometry (two domains); leftover p orbitals allow two pi bonds (triple bond).
Sigma (σ) Bond
A covalent bond formed by head-on orbital overlap along the axis between nuclei. Every single bond is a sigma bond.
Pi (π) Bond
A bond from side-by-side overlap of p orbitals, above and below the axis. Double bonds have 1 σ + 1 π; triple bonds have 1 σ + 2 π.
Bond Polarity (Dipole)
A polar bond has a dipole — a separation of partial charge — pointing toward the more electronegative atom.
Molecular Polarity
A molecule is polar if it has polar bonds AND an asymmetrical shape so the bond dipoles don't cancel.
Symmetry Cancels Dipoles
Symmetrical molecules (CO₂ linear, CCl₄ tetrahedral) can have polar bonds yet be nonpolar overall because the dipoles cancel.
Water Is Polar
H₂O has polar O–H bonds AND a bent shape, so the dipoles don't cancel — giving water a strong net dipole and its solvent power.
Predicting Polarity
Draw the Lewis structure, find the geometry (VSEPR), then check whether the bond dipoles cancel by symmetry. Lone pairs on the central atom usually make a molecule polar.
Ionic vs. Covalent Properties
Ionic: high melting point, conducts when molten/dissolved, brittle. Molecular covalent: lower melting point, usually doesn't conduct, softer.
Network Covalent Solids
Atoms bonded covalently in a continuous 3-D network (diamond, quartz, graphite). Extremely hard with very high melting points.
Coulomb's Law in Bonding
The strength of ionic and covalent attractions traces back to Coulomb's law — the balance of charge magnitude and distance between attracting particles.
Bond Energy
The energy required to break one mole of a bond in the gas phase. Stronger (shorter, higher-order) bonds have greater bond energy.
Delocalized Electrons
Electrons spread over multiple atoms rather than fixed between two — as in resonance, metallic bonding, and graphite — often conferring conductivity or stability.
Charged groups of covalently bonded atoms (e.g., NO₃⁻, SO₄²⁻, NH₄⁺) that act as a unit in ionic compounds.
Structure Determines Properties
A substance's macroscopic properties (melting point, conductivity, hardness, solubility) follow directly from its bonding type and structure.
Lattice Energy Factors
Lattice energy grows with higher ionic charges and smaller ionic radii (Coulomb's law). MgO has a far larger lattice energy than NaCl because both ions carry charge 2 and are small.
Bond Order and Bond Length
Higher bond order (single < double < triple) means shorter, stronger bonds with higher bond energy. Resonance gives fractional bond orders (e.g., 4/3 in NO₃⁻).
Formal Charge
FC = valence electrons − nonbonding electrons − ½ bonding electrons. The best Lewis structure minimizes formal charges and places negative FC on the more electronegative atom.
Expanded Octets
Period 3 and heavier central atoms (S, P, Xe, Cl) can hold more than 8 electrons using d-orbital availability, giving species like SF₆, PCl₅, and XeF₄.
Incomplete Octets
Be and B commonly have fewer than 8 electrons (BeCl₂, BF₃). Radicals like NO have an odd electron count and cannot satisfy every octet.
Electron Domain vs. Molecular Geometry
Electron geometry counts all domains (bonds and lone pairs); molecular geometry names only atom positions. Four domains with one lone pair: tetrahedral electron geometry, trigonal pyramidal shape.
Lone Pair Repulsion
Lone pairs occupy more space than bonding pairs and compress bond angles: CH₄ 109.5°, NH₃ 107°, H₂O 104.5°.
Hybridization from Domains
2 domains → sp (linear), 3 → sp² (trigonal planar), 4 → sp³ (tetrahedral). AP Chemistry limits hybridization to sp, sp², and sp³.
Sigma and Pi Bonds
A single bond is one σ bond; a double bond is one σ + one π; a triple is one σ + two π. π bonds prevent rotation about the bond axis.
Bond Polarity vs. Molecular Polarity
Polar bonds cancel in symmetric geometries (CO₂, CCl₄, BF₃) giving nonpolar molecules; asymmetric shapes or lone pairs (SO₂, NH₃, CHCl₃) give net dipoles.
Metallic Bonding
A lattice of cations in a delocalized 'sea' of electrons explains conductivity, malleability, and luster. Alloys modify these properties by disrupting the lattice.
Interstitial vs. Substitutional Alloys
Interstitial: small atoms (C in steel) fill gaps between larger metal atoms, hardening the metal. Substitutional: similar-sized atoms replace host atoms (Cu/Zn in brass).
Ionic Solids and Conductivity
Ionic compounds conduct only when molten or dissolved because ions must be mobile; in the solid the ions are locked in the lattice.
Network Covalent Solids
Diamond, graphite, and SiO₂ are held by continuous covalent bonds, giving very high melting points and (except graphite) poor conductivity.
Resonance Structures
Multiple valid Lewis structures differing only in electron placement; the real molecule is a hybrid with equal bond lengths, as in O₃ and CO₃²⁻.
Bond Enthalpy Reasoning
Bond energy correlates with bond order and inversely with length; comparing C–O (1), C=O (2), and C≡O (3) predicts increasing energy and decreasing length.
Among valid Lewis structures, prefer the one with formal charges nearest zero and any negative formal charge placed on the most electronegative atom.
Resonance and Bond Length Averaging
In nitrate, all three N–O bonds measure identically between a single and double bond because the real structure is a resonance hybrid with bond order 4/3.
Delocalization and Stability
Spreading electron density over several atoms lowers the total energy, so molecules with valid resonance forms are more stable than any single structure suggests.
Bond Order from Resonance Count
Bond order = (total bonding pairs between the two atoms across all forms) ÷ (number of resonance forms). Carbonate's C–O bond order is 4/3.
Predicting Bond Angle Compression
Each lone pair on the central atom compresses the ideal angle by roughly 2–3°: CH₄ is 109.5°, NH₃ is 107°, and H₂O is 104.5°.
Seesaw and T-Shaped Geometries
Five electron domains with one lone pair give a seesaw shape; two lone pairs give T-shaped. Lone pairs always occupy the roomier equatorial positions.
Square Planar vs. Tetrahedral
Six domains with two lone pairs (as in XeF₄) place the lone pairs axially, leaving four bonds in a square plane at 90°, unlike tetrahedral's 109.5°.
Polarity by Vector Cancellation
A molecule is nonpolar only if bond dipoles cancel by symmetry. CO₂ cancels linearly, while SO₂'s bent shape leaves a net dipole even with identical bonds.
sp³d and sp³d² Hybridization
Five domains require sp³d and six require sp³d² hybrid orbitals, which is why expanded octets appear only for period 3 and beyond where d orbitals are energetically accessible.
Pi Bonds and Rotational Rigidity
Sideways p-orbital overlap in a pi bond blocks rotation about the axis, which is why alkenes show cis/trans isomers but alkanes rotate freely.
Lattice Energy and Coulomb's Law
Lattice energy rises with larger ionic charges and smaller ionic radii, so MgO (2+/2−, small ions) far exceeds NaCl in melting point and hardness.
Photon Energy and Bond Breaking
A photon breaks a bond only when its energy exceeds the bond enthalpy; UV photons carry enough energy to break O–O bonds in ozone, while visible light does not.
Estimating ΔH from Bond Enthalpies
ΔH ≈ Σ(bonds broken) − Σ(bonds formed). Positive values mean the bonds formed are weaker than those broken, so the reaction absorbs energy.
Alloys and Malleability
Substitutional alloys such as brass keep the parent lattice with mismatched atom sizes that impede layer slipping, raising hardness while retaining metallic conductivity.
Band Theory Sketch
Overlapping atomic orbitals in a metal form continuous energy bands; partially filled bands let electrons move freely, explaining conductivity and the metal's luster.
Semiconductor Doping
Adding a Group 15 atom to silicon supplies extra electrons (n-type); a Group 13 dopant creates holes (p-type). Both raise conductivity above the pure element.
Covalent Network Melting Points
Silicon dioxide and diamond melt extremely high because melting requires breaking a continuous three-dimensional array of strong covalent bonds, not merely overcoming intermolecular forces.