Chemistry study notes
Chemistry for the NMAT
Key points
- Electron configurations follow the aufbau principle, Hund's rule, and the Pauli exclusion principle; periodic trends (radius, ionization energy, electronegativity) follow directly from effective nuclear charge and shell number.
- Ionic bonds involve electron transfer (metal + nonmetal); covalent bonds involve electron sharing (nonmetal + nonmetal), which can be polar or nonpolar depending on electronegativity difference.
- VSEPR theory predicts molecular geometry from the number of bonding and lone-pair electron domains around a central atom; lone pairs compress ideal bond angles.
- Molecular polarity requires both polar bonds and a geometry that does not cancel their dipoles — symmetric molecules like CO₂ can be nonpolar despite polar bonds.
- Hydrogen bonding (H bonded to N, O, or F) is the strongest common intermolecular force and explains anomalies like water's high boiling point.
- Balanced equations give mole ratios; the limiting reagent determines theoretical yield, and percent yield compares actual to theoretical yield.
- Empirical formulas come from percent composition or combustion analysis; multiplying by the correct whole-number factor (using molar mass) gives the molecular formula.
- Redox reactions require tracking oxidation number changes: the reducing agent is oxidized (loses electrons), and the oxidizing agent is reduced (gains electrons).
- Alkenes undergo addition reactions; Markovnikov's rule predicts the major product forms via the more stable (more substituted) carbocation intermediate.
- Alcohol oxidation depends on class: primary → aldehyde → carboxylic acid; secondary → ketone only; tertiary → resistant to oxidation.
- Strong acids and bases dissociate completely; weak acids and bases establish an equilibrium governed by Ka or Kb, and pH + pOH = 14 at 25 °C.
- Buffers (weak acid/conjugate base pairs) resist pH change and follow pH = pKa + log([A⁻]/[HA]).
Atomic Structure and Periodicity
Electrons occupy orbitals — probability regions described by quantum numbers — rather than fixed Bohr-model orbits. Ground-state electron configurations are built using the aufbau principle (fill lowest-energy orbitals first), Hund's rule (spread electrons across degenerate orbitals before pairing), and the Pauli exclusion principle (no two electrons share all four quantum numbers). The periodic table itself is a map of these configurations: periods correspond to principal quantum number, and groups share the same valence electron count, which is why elements in a group behave similarly. Moving across a period, increasing effective nuclear charge pulls electrons in tighter, so atomic radius shrinks while ionization energy and electronegativity rise. Moving down a group, each new row adds an entire outer shell, so radius grows even though nuclear charge also increases. Watch for exceptions: chromium and copper gain extra stability from half-filled or fully filled d subshells, and photoelectron spectroscopy (PES) data — peak position for binding energy, peak height for electron count per subshell — is a common way the NMAT tests whether you truly understand orbital filling rather than just memorizing the table.
Chemical Bonding and Molecular Shape
Bonds form to lower a system's energy. Ionic bonds arise from electron transfer between a metal and a nonmetal, generating an electrostatic lattice of ions; covalent bonds arise from electron sharing between nonmetals, categorized as nonpolar or polar depending on the electronegativity difference between the bonded atoms. Lewis structures track valence electrons as dots and lines, and formal charge — calculated as valence electrons minus nonbonding electrons minus half the bonding electrons — helps pick the most reasonable structure when several are possible (as in resonance). VSEPR theory predicts molecular geometry from the number of electron domains (bonding groups plus lone pairs) around a central atom: two domains give linear geometry, three give trigonal planar, and four give tetrahedral, with lone pairs compressing bond angles below the ideal values (109.5° in CH₄ becomes about 107° in NH₃ and 104.5° in H₂O). Hybridization (sp, sp², sp³, and beyond for expanded octets) explains how atomic orbitals mix to produce these geometries. Whether a molecule is overall polar depends on both individual bond polarities and molecular symmetry — CO₂ has polar bonds but is nonpolar overall because its linear shape cancels the dipoles.
Intermolecular Forces and States of Matter
While covalent and ionic bonds hold atoms together within a molecule or lattice, intermolecular forces (IMFs) act between separate molecules and govern bulk properties like boiling point, melting point, and solubility. From weakest to strongest for comparably sized molecules: London dispersion forces (temporary, induced dipoles present in all molecules, growing stronger with more electrons), dipole-dipole interactions (between molecules with permanent, unequal dipoles), and hydrogen bonding (a special, strong dipole interaction when hydrogen is bonded directly to nitrogen, oxygen, or fluorine). Metallic bonding — a lattice of cations surrounded by a delocalized "sea" of valence electrons — explains why metals conduct electricity and heat, and why they are malleable and ductile rather than brittle. On the NMAT, IMF questions typically ask you to rank boiling points or explain an anomaly (such as water's unusually high boiling point relative to other Group 16 hydrides), so identifying which IMF dominates for a given molecule is the key skill.
Stoichiometry and Chemical Reactions
The mole is chemistry's bridge between the atomic and macroscopic worlds: one mole of any substance contains Avogadro's number (6.022 × 10²³) of particles, and molar mass converts between grams and moles. Balanced chemical equations express the law of conservation of mass and give the mole ratios needed for stoichiometric calculations — converting from grams or moles of one substance to grams or moles of another via the coefficients in the balanced equation. When two reactants are given, the limiting reagent (the one that runs out first) caps how much product can form; percent yield then compares the actual amount recovered to this theoretical maximum. Other core skills include deriving empirical and molecular formulas from percent composition or combustion data, applying the ideal gas law (PV = nRT) to gas-phase reactions, and handling solution stoichiometry (molarity, dilution with M₁V₁ = M₂V₂, and acid-base titrations). Redox reactions, where oxidation numbers change, require tracking electron transfer explicitly — the reducing agent loses electrons (is oxidized) and the oxidizing agent gains them (is reduced).
Organic Chemistry Basics
Carbon's ability to form four stable covalent bonds, including to itself, creates the vast diversity of organic compounds. Functional groups define reactivity: alkanes (C–C, C–H only) are relatively unreactive; alkenes (C=C) and alkynes (C≡C) readily undergo addition reactions, such as hydrogenation (addition of H₂) or hydrohalogenation (addition of HX), with Markovnikov's rule predicting that H adds to the carbon with more existing hydrogens while the other substituent adds to the more substituted carbon, via the more stable carbocation intermediate. Alcohols oxidize in a stepwise fashion — primary alcohols to aldehydes and then to carboxylic acids, secondary alcohols to ketones only, and tertiary alcohols resist oxidation entirely because they lack a hydrogen on the carbinol carbon. Carboxylic acids and alcohols undergo acid-catalyzed condensation (Fischer esterification) to form esters and water; esters can be hydrolyzed back, and base-catalyzed hydrolysis of fats (saponification) yields glycerol plus fatty acid salts (soap). Structural isomers share a molecular formula but differ in atom connectivity, a distinction the NMAT tests frequently.
Solutions, Acids, and Bases
Concentration can be expressed as molarity (mol solute/L solution), molality (mol solute/kg solvent), or percent by mass; molality and mass percent, unlike molarity, do not change with temperature because they are mass-based rather than volume-based. Colligative properties — boiling-point elevation, freezing-point depression, vapor-pressure lowering, and osmotic pressure — depend only on the number of dissolved particles, not their identity, so electrolytes that dissociate into multiple ions (tracked by the van 't Hoff factor) produce a proportionally larger effect than nonelectrolytes at the same molal concentration. In Brønsted-Lowry theory, acids donate protons and bases accept them, forming conjugate acid-base pairs; strong acids and bases dissociate completely, while weak acids and bases establish an equilibrium described by Ka or Kb. pH = −log[H⁺] and pH + pOH = 14 at 25 °C. Buffers, made of a weak acid and its conjugate base (or vice versa), resist pH changes and are described by the Henderson-Hasselbalch equation, pH = pKa + log([A⁻]/[HA]) — a relationship worth memorizing cold, since it underlies both textbook buffer problems and physiological ones like blood pH regulation.
References
- Atomic Structure and Periodicity — OpenStax
- Ionic Bonding — OpenStax
- Covalent Bonding and Molecular Geometry — OpenStax
- Stoichiometry of Chemical Reactions — OpenStax
- Solutions and Colloids — OpenStax
- Acids and Bases — Khan Academy
- Electronic Structure of Atoms — Khan Academy
- Introduction to Organic Chemistry — LibreTexts
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