GENERAL CHEMISTRY · ATOMIC AND MOLECULAR STRUCTURE
Count the particles. Then build the model.
Twenty-four study objectives connect the eight official topics—from subatomic identity and experimental evidence to electron configuration, Lewis structures, bonding, geometry, and polarity.
The ADA lists Atomic and Molecular Structure as one General Chemistry domain. Periodic Properties and Liquids and Solids are separate domains.
The three study objectives under each topic organize learning; they do not predict question counts, weights, or test order.
Official hierarchy → study sequence
Eight branches. Twenty-four outcomes.
Each objective includes an outcome, prerequisite knowledge, must-know relationships, a depth boundary, a misconception cue, and free source links.
AElectron configuration3 objectives
Objective 1
Ground-state electron configurations
Write or select a ground-state electron configuration using Aufbau order, Pauli exclusion, and Hund’s rule.
Must know
Fill lower-energy subshells in the conventional Aufbau sequence before higher ones. Degenerate orbitals receive one electron each with parallel spins before pairing, and an orbital holds at most two opposite-spin electrons.
Depth boundary
Common transition-metal exceptions are assessed only when expected by standard introductory treatment or signaled by choices.
Misconception
Electrons pair in the first p orbital before occupying the other p orbitals. Hund’s rule places one electron in each degenerate orbital before pairing.
Prerequisites
Subshell orbital counts and capacities · Atomic number, mass number, and isotopes
Objective 2
Electron configurations of ions
Derive main-group and transition-metal ion configurations from a neutral atom while removing electrons from the correct shell.
Must know
Anions add and cations remove electrons according to charge. For transition-metal cations, electrons are removed from the highest principal shell before the lower-n d subshell.
Depth boundary
Highly unusual oxidation states and relativistic effects are outside scope.
Misconception
Fe²⁺ loses two 3d electrons because 3d is written after 4s in the neutral configuration. Cation formation removes the higher-n 4s electrons before 3d electrons.
Prerequisites
Ground-state electron configurations
Objective 3
Valence and unpaired-electron inference
Use an electron configuration or orbital diagram to identify valence electrons, unpaired electrons, and qualitative magnetic behavior.
Must know
Unpaired electrons produce paramagnetic behavior; all paired electrons produce diamagnetic behavior. For main-group atoms, valence electrons occupy the highest principal shell.
Depth boundary
Quantitative magnetic moments and ligand-field splitting belong to advanced coordination chemistry.
Misconception
Any atom with an odd atomic number must be diamagnetic. Magnetic behavior depends on the actual orbital occupancy and whether electrons remain unpaired.
Prerequisites
Ground-state electron configurations
BOrbital types3 objectives
Objective 1
Orbital types and qualitative shapes
Identify s, p, and d orbital types and their qualitative spatial features.
Must know
s orbitals are spherically symmetric; p orbitals have two lobes separated by a nodal plane. Orbital pictures show probability regions and phase, not solid containers or electron tracks.
Depth boundary
Detailed radial-node equations and complete d-orbital phase derivations are outside scope.
Misconception
A p orbital is two separate orbitals because it has two lobes. The two lobes belong to one p orbital and are separated by an angular node.
Prerequisites
Quantum description and probability
Objective 2
Subshell orbital counts and capacities
Determine the number of orbitals and maximum electrons in s, p, d, and f subshells.
Must know
s, p, d, and f subshells contain 1, 3, 5, and 7 orbitals. Each orbital holds at most two electrons, giving capacities 2, 6, 10, and 14.
Depth boundary
Only allowed subshells for a principal shell need be recognized; degeneracy splitting requires supplied context.
Misconception
A p subshell holds three electrons because it has three orbitals. Each of the three p orbitals holds up to two electrons, for six total.
Prerequisites
Orbital types and qualitative shapes
Objective 3
Allowed quantum-number sets
Recognize allowed combinations of principal, angular-momentum, magnetic, and spin quantum numbers.
Must know
For principal quantum number n, l ranges from 0 through n−1; mₗ ranges from −l through +l. Electron spin mₛ is +1/2 or −1/2.
Depth boundary
Use quantum-number selection rules only to validate states; spectroscopy transition rules are outside scope unless supplied.
Misconception
For n = 2, l = 2 is allowed. l must be smaller than n, so n = 2 permits l = 0 or 1 only.
Prerequisites
Subshell orbital counts and capacities
CLewis-dot diagrams3 objectives
Objective 1
Valence-electron budgets
Calculate total valence electrons for a molecule or polyatomic ion before drawing a Lewis structure.
Must know
Sum valence electrons from every atom, adding electrons for negative charge and subtracting for positive charge. The electron budget constrains all bonds and lone pairs in the final structure.
Depth boundary
Transition-metal electron counting is outside this objective.
Misconception
A negative charge means subtracting one electron from the Lewis budget. Negative charge adds electron density; a −1 ion adds one electron to the budget.
Prerequisites
Valence and unpaired-electron inference
Objective 2
Lewis structures and formal charge
Select a valid Lewis structure that satisfies the electron budget and minimizes defensible formal charges.
Must know
Count every bonding pair and lone pair against the total electron budget. Formal charge equals valence electrons minus nonbonding electrons minus one-half of bonding electrons.
Depth boundary
Electron-deficient and expanded-octet cases are included only for standard introductory examples.
Misconception
A structure is correct whenever every atom has an octet, even if the total electron budget is wrong. The total budget, connectivity, octet exceptions, and formal charges must all be checked.
Prerequisites
Valence-electron budgets
Objective 3
Resonance and octet exceptions
Recognize resonance contributors and bounded octet-rule exceptions without treating contributors as rapidly switching molecules.
Must know
Resonance contributors keep atom connectivity fixed while electron placement changes. The real electron distribution is a resonance hybrid, not an alternation between separate structures.
Depth boundary
Detailed resonance-energy calculations and molecular-orbital derivations are outside scope.
Misconception
Resonance structures differ by moving atoms or changing connectivity. Only electron placement changes among valid contributors; nuclei remain in the same connectivity.
Prerequisites
Lewis structures and formal charge
DAtomic theory3 objectives
Objective 1
Evidence and atomic models
Match landmark observations to the atomic-model claim they supported or rejected.
Must know
Cathode-ray evidence supported negatively charged particles within atoms. Alpha-particle scattering supported a small, dense, positively charged nucleus and mostly empty atomic volume.
Depth boundary
Memorize evidence-to-model logic rather than biographical chronology beyond what a prompt supplies.
Misconception
Rutherford’s scattering showed positive charge spread uniformly through the atom. Rare large deflections contradicted diffuse positive charge and supported a concentrated nucleus.
Prerequisites
Start here
Objective 2
Atomic number, mass number, and isotopes
Relate atomic number, mass number, nuclide symbols, and isotope identity without substituting average atomic mass.
Must know
Atomic number Z is the proton count; mass number A is protons plus neutrons for one nuclide. Isotopes share Z but differ in neutron count and therefore A.
Depth boundary
Mass defect and binding energy belong to Nuclear Reactions.
Misconception
The decimal atomic mass on a periodic table is the neutron count of one atom. It is an abundance-weighted average across naturally occurring isotopes.
Prerequisites
Particle inventory from nuclide notation
Objective 3
Average atomic mass
Calculate or interpret an abundance-weighted average atomic mass from isotope masses and fractional abundances.
Must know
Convert percentages to fractions before multiplying each isotope mass by its abundance. The weighted average lies between the contributing isotope masses and nearer the more abundant isotope.
Depth boundary
Exact isotope masses or rounded mass numbers are used as provided; do not introduce unlisted isotopes.
Misconception
Average atomic mass is the simple mean of isotope masses regardless of abundance. Each isotope’s contribution is weighted by its fractional abundance.
Prerequisites
Atomic number, mass number, and isotopes
EQuantum theory3 objectives
Objective 1
Photon wavelength, frequency, and energy
Relate wavelength, frequency, and photon energy qualitatively or with supplied constants.
Must know
For electromagnetic radiation, c = λν, so wavelength and frequency are inversely related. Photon energy E = hν, so higher frequency means higher photon energy.
Depth boundary
Constants and unit conversions are supplied when numerical precision matters.
Misconception
Longer wavelength means greater photon energy. Longer wavelength means lower frequency and therefore lower photon energy.
Prerequisites
Start here
Objective 2
Quantized energy and line spectra
Interpret absorption or emission lines as transitions between allowed energy levels.
Must know
An emitted photon carries the energy lost in a downward electronic transition. Discrete spectral lines reflect discrete allowed energy differences rather than every possible energy.
Depth boundary
Hydrogen-like calculations use equations supplied in the prompt; many-electron spectra are interpreted qualitatively.
Misconception
An electron emits a photon while moving to a higher energy level. Moving upward requires absorption; a downward transition can emit a photon.
Prerequisites
Photon wavelength, frequency, and energy
Objective 3
Quantum description and probability
Distinguish orbitals as probability distributions from fixed classical electron paths.
Must know
An orbital describes a spatial probability distribution for an electron in an atom. Quantum numbers label allowed states; they do not specify a miniature planetary trajectory.
Depth boundary
Wavefunction derivation and operator algebra are outside scope.
Misconception
An orbital is the exact circular path traveled by an electron. An orbital is a quantum state represented by a probability distribution, not a classical path.
Prerequisites
Quantized energy and line spectra
FMolecular geometry3 objectives
Objective 1
Electron-group and molecular geometry
Use a valid Lewis structure to distinguish electron-group geometry from molecular geometry.
Must know
VSEPR counts each bond—single or multiple—as one electron group and each lone pair as one group. Electron-group geometry includes lone pairs; molecular geometry names atom positions.
Depth boundary
Focus on standard two- through six-group VSEPR cases represented in introductory chemistry.
Misconception
A double bond counts as two electron groups in VSEPR. All electron density connecting the same two atoms counts as one group.
Prerequisites
Lewis structures and formal charge
Objective 2
Lone pairs and bond-angle trends
Predict common molecular shapes and qualitative bond-angle compression from bonding and lone-pair domains.
Must know
Lone-pair electron density generally repels more strongly than bonding-pair density. Lone pairs can change the molecular shape and compress adjacent bond angles relative to the ideal electron-group angle.
Depth boundary
Exact experimental angles are required only when supplied or conventionally emphasized.
Misconception
CH₄, NH₃, and H₂O all have the same molecular geometry because each has four electron groups. They share tetrahedral electron-group geometry but have different molecular geometries because lone pairs hide vertices.
Prerequisites
Electron-group and molecular geometry
Objective 3
Bond dipoles and molecular polarity
Combine bond polarity with molecular geometry to determine whether dipoles cancel or produce a net molecular dipole.
Must know
Polar bonds do not guarantee a polar molecule; vector direction and geometry determine cancellation. Symmetry cancels equivalent bond dipoles only when the surrounding atoms and geometry support it.
Depth boundary
Quantitative dipole moments are outside scope unless supplied.
Misconception
Any molecule containing polar bonds must be polar. Equivalent bond dipoles can cancel in a symmetric geometry, as in CO₂.
Prerequisites
Lone pairs and bond-angle trends · Ionic, covalent, and metallic bonding
GBond types3 objectives
Objective 1
Ionic, covalent, and metallic bonding
Classify a dominant bonding model from composition, electron behavior, and characteristic structure.
Must know
Ionic bonding is electrostatic attraction among oppositely charged ions; covalent bonding shares electron density between atoms. Metallic bonding involves delocalized valence electrons across a metal lattice.
Depth boundary
Bonding is a model continuum; classify the dominant introductory model from supplied evidence.
Misconception
An ionic bond is a single shared electron pair between two atoms. Ionic bonding is the collective electrostatic attraction among ions in a lattice.
Prerequisites
Valence and unpaired-electron inference
Objective 2
Bond order, length, strength, and multiplicity
Relate single, double, and triple covalent bonds to sigma/pi composition, bond order, length, and typical strength.
Must know
A single bond is one sigma bond; double and triple bonds add one and two pi bonds. For the same bonded atoms, greater bond order is generally shorter and stronger.
Depth boundary
Quantitative bond energies are used only when supplied; resonance can produce intermediate bond orders.
Misconception
A double bond contains two pi bonds. A double bond contains one sigma bond and one pi bond.
Prerequisites
Ionic, covalent, and metallic bonding · Lewis structures and formal charge
Objective 3
Orbital overlap and hybridization
Associate standard sp, sp², and sp³ hybridization with electron-group geometry and sigma-bond frameworks.
Must know
Two, three, and four electron groups correspond to sp, sp², and sp³ hybrid sets in the standard valence-bond model. Unhybridized p orbitals form pi bonds after the sigma framework is established.
Depth boundary
d-orbital participation and alternative hypervalent bonding models are not inferred unless a prompt specifies a model.
Misconception
A carbon in a double bond is sp³ because it forms four shared electron pairs overall. Three electron groups around that carbon give sp² hybridization, with an unhybridized p orbital forming the pi bond.
Prerequisites
Orbital types and qualitative shapes · Lewis structures and formal charge
HSub-atomic particles3 objectives
Objective 1
Subatomic charge, mass, and location
Compare protons, neutrons, and electrons by relative charge, relative mass, and location.
Must know
Protons are +1 and neutrons are neutral in the nucleus; electrons are −1 outside it. Protons and neutrons each have about one atomic mass unit, while electron mass is much smaller.
Depth boundary
Quark structure and particle-physics interactions are outside this objective.
Misconception
Neutrons are negatively charged because their name begins with neutral. Neutrons have zero net charge; electrons carry the negative charge.
Prerequisites
Start here
Objective 2
Particle inventory from nuclide notation
Determine proton, neutron, and electron counts from atomic number, mass number, and ionic charge.
Must know
Protons equal atomic number and neutrons equal mass number minus atomic number. For charge q in elementary-charge units, electron count equals protons minus q.
Depth boundary
Use whole-number mass number for a nuclide, not the periodic-table average atomic mass.
Misconception
A +2 ion has two extra electrons. Positive charge means electrons have been removed; a +2 ion has two fewer electrons than protons.
Prerequisites
Subatomic charge, mass, and location
Objective 3
Isotope, ion, and element identity
Classify particle-count changes as isotope formation, ion formation, or a change of element.
Must know
Changing neutron count changes isotope but not element. Changing electron count changes charge; changing proton count changes element identity.
Depth boundary
Nuclear reaction mechanisms belong to the Nuclear Reactions domain.
Misconception
Adding a neutron creates a different element. Element identity is fixed by proton count, not neutron count.