Your FREE ACS General Chemistry Flashcards 2026 – 100+ Cards

Electrons fill the lowest-energy orbitals first when building an atom's ground-state configuration — so 4s fills before 3d.

No two electrons in an atom can have the same four quantum numbers; an orbital holds at most 2 electrons, with opposite spins.

Electrons singly occupy each orbital in a subshell, with parallel spins, before any orbital is doubly occupied.

n (energy level), ℓ (orbital shape: s/p/d/f), mℓ (orientation), and mₛ (spin, ±½).

Atoms of the same element (same protons) with different numbers of neutrons, and so different mass numbers.

Atomic number = number of protons (defines the element); mass number = protons + neutrons.

Decreases across a period (left → right) and increases down a group.

Increases across a period and decreases down a group (it takes more energy to remove an electron from a smaller, tightly held atom).

Increases across a period and decreases down a group; fluorine is the most electronegative element.

E = hν = hc/λ — higher frequency (shorter wavelength) means higher energy.

s holds 2, p holds 6, d holds 10, f holds 14 (each orbital holds 2).

Remove electrons from the highest n first — for transition metals the 4s electrons leave before the 3d (e.g., Fe²⁺ is [Ar]3d⁶).

The net positive charge an electron feels after shielding by inner electrons; it rises across a period, pulling electrons in.

Elements with a full valence shell, very stable and unreactive — all are monatomic gases at room temperature.

The abundance-weighted average mass of an element's naturally occurring isotopes.

Electrostatic attraction between oppositely charged ions formed when a metal transfers electrons to a nonmetal (e.g., NaCl).

A bond formed by two atoms sharing one or more pairs of electrons (typically between nonmetals).

Unequal electron sharing (electronegativity difference ~0.4–1.7) is polar covalent; near-equal sharing is nonpolar.

Electron groups around a central atom arrange to minimize repulsion, which sets the molecule's geometry.

Bonded atoms + lone pairs on the central atom; it sets the electron-domain geometry and the hybridization.

2 → linear (180°), 3 → trigonal planar (120°), 4 → tetrahedral (109.5°).

2 → sp, 3 → sp², 4 → sp³, 5 → sp³d, 6 → sp³d².

Oxygen has steric number 4 (2 bonds + 2 lone pairs); the lone pairs push the bonds to a bent ~104.5° shape.

Electron geometry counts all bonding pairs and lone pairs; molecular geometry describes only the arrangement of bonded atoms.

FC = (valence electrons) − (nonbonding electrons) − ½(bonding electrons); the best Lewis structure minimizes formal charges.

Structures that differ only in electron placement, not nuclear positions; the real molecule is a hybrid of all of them.

A sigma bond is head-on orbital overlap along the bond axis; a pi bond is side-to-side overlap above and below the axis.

A single bond is 1 sigma; a double is 1 sigma + 1 pi; a triple is 1 sigma + 2 pi (shorter and stronger).

It is polar only if its individual bond dipoles do NOT cancel — CO₂ is nonpolar (symmetric), H₂O is polar (bent).

Total valence electrons → least-electronegative atom (not H) in center → single bonds → distribute lone pairs to satisfy octets → add multiple bonds if needed.

Atoms tend to gain a full 8-electron valence shell; exceptions include H (2), incomplete octets (B), and expanded octets (P, S).

Points from the less electronegative atom toward the more electronegative atom.

Ion–dipole > hydrogen bonding > dipole–dipole > London dispersion.

A strong dipole–dipole attraction when H is bonded to N, O, or F; it gives water its high boiling point and surface tension.

Weak, temporary-dipole attractions present in ALL molecules; they grow with molar mass and polarizability.

Higher boiling point and melting point, higher viscosity and surface tension, and lower vapor pressure.

PV = nRT — relates pressure, volume, moles, and absolute temperature; R = 0.0821 L·atm/mol·K.

At constant temperature, pressure and volume are inversely related: P ∝ 1/V, so P₁V₁ = P₂V₂.

At constant pressure, volume is directly proportional to absolute temperature: V ∝ T.

P₁V₁/T₁ = P₂V₂/T₂ — temperature must be in kelvin.

The total pressure of a gas mixture equals the sum of the partial pressures of its components.

Lighter gases effuse faster: r₁/r₂ = √(M₂/M₁).

Gas particles are tiny, in constant random motion, with negligible volume and no IMFs; average kinetic energy ∝ absolute temperature.

A solution property that depends on the NUMBER of dissolved particles, not their identity.

A non-volatile solute raises the boiling point: ΔTb = i·Kb·m (m = molality, i = van't Hoff factor).

A solute lowers the freezing point: ΔTf = i·Kf·m — why salt melts ice on roads.

The vapor pressure of a solvent over an ideal solution is lowered: P = X(solvent)·P° (X = mole fraction).

The number of particles a solute produces in solution: NaCl gives i ≈ 2, glucose i = 1.

Polar/ionic solutes dissolve in polar solvents; nonpolar solutes dissolve in nonpolar solvents.

The temperature and pressure above which a distinct liquid phase can no longer exist (a supercritical fluid).

Using the mole ratios in a balanced equation to relate amounts of reactants and products.

The amount of substance containing Avogadro's number (6.022 × 10²³) of particles; it bridges mass and number of particles.

n = mass ÷ molar mass.

The reactant consumed first; it caps the amount of product (theoretical yield).

(actual yield ÷ theoretical yield) × 100%.

Concentration = moles of solute ÷ liters of solution (mol/L, M).

M₁V₁ = M₂V₂ — moles of solute are conserved when you add solvent.

Oxidation Is Loss of electrons; Reduction Is Gain of electrons. They always occur together.

The oxidizing agent is reduced (gains electrons); the reducing agent is oxidized (loses electrons).

A hydrocarbon + O₂ → CO₂ + H₂O, releasing heat (exothermic).

Two soluble salts react to form an insoluble solid (precipitate) — predicted with solubility rules.

An acid + a base → a salt + water (H⁺ + OH⁻ → H₂O).

6.022 × 10²³ — the number of particles in one mole.

Heat exchanged at constant pressure; ΔH < 0 is exothermic (releases heat), ΔH > 0 is endothermic (absorbs heat).

q = mcΔT — heat = mass × specific heat × temperature change.

The overall ΔH of a reaction equals the sum of the ΔH of its steps, because enthalpy is a state function.

ΔH°rxn = ΣΔHf°(products) − ΣΔHf°(reactants).

Zero — for an element in its standard state (e.g., O₂(g), N₂(g)).

A measure of disorder / dispersal of energy and matter; gases have higher entropy than liquids or solids.

The total entropy of the universe (system + surroundings) increases for any spontaneous process.

ΔG = ΔH − TΔS; a process is spontaneous when ΔG < 0, at equilibrium when ΔG = 0.

When ΔH < 0 (exothermic) and ΔS > 0 (more disorder) — both favor a negative ΔG.

When ΔH > 0 (endothermic) and ΔS < 0 (less disorder) — ΔG is positive at every temperature.

Negative — converting more moles of gas to fewer (e.g., N₂ + 3H₂ → 2NH₃) decreases disorder.

Energy is conserved: ΔU = q + w (internal energy change = heat added + work done on the system).

ΔG° = −RT ln K — a negative ΔG° means K > 1 (products favored).

An electrochemical cell that uses a spontaneous redox reaction to generate electricity (ΔG < 0, E°cell > 0).

Oxidation occurs at the ANode; reduction occurs at the CAThode.

E°cell = E°cathode − E°anode; a positive E°cell means a spontaneous reaction.

E = E° − (0.0592/n)·log Q at 25 °C — gives the cell potential under non-standard conditions.

ΔG° = −nFE°cell (F = 96,485 C/mol); a positive E° gives a negative ΔG° (spontaneous).

The reference half-cell, assigned E° = 0 V; all other standard reduction potentials are measured against it.

rate = k[A]ᵐ[B]ⁿ; the orders m and n are found experimentally, NOT from the balanced coefficients.

The exponent on a reactant in the rate law; the overall order is the sum of the exponents.

The minimum collision energy needed for a reaction to occur — the barrier on the energy profile.

It provides an alternate pathway with lower Eₐ, speeding both directions; it is not consumed and does not change ΔH or K.

t½ = 0.693/k — constant, independent of the starting concentration.

k = A·exp(−Eₐ/RT); raising temperature or lowering the activation energy Eₐ increases the rate constant k.

A collision leads to reaction only if the molecules have enough energy (≥ Eₐ) AND the correct orientation.

The slowest elementary step in a mechanism; it controls the overall reaction rate.

Higher temperature increases the fraction of molecules with energy ≥ Eₐ, so the rate rises sharply.

Higher reactant concentration or surface area increases collision frequency, raising the rate.

ln[A] = ln[A]₀ − kt — a plot of ln[A] vs. time is linear for a first-order reaction.

The state where forward and reverse reaction rates are equal, so concentrations stay constant (dynamic, not static).

Products over reactants, each raised to its coefficient: K > 1 favors products, K < 1 favors reactants.

The same ratio as K but at any moment: Q < K shifts forward, Q > K shifts reverse, Q = K is equilibrium.

A disturbed equilibrium shifts to partly counteract the change and restore balance.

Increasing pressure (decreasing volume) shifts toward the side with fewer moles of gas.

Treat heat as a reactant/product: raising T raises K for endothermic reactions and lowers K for exothermic ones.

No — a catalyst speeds the approach to equilibrium but does not change the equilibrium position or K.

pH = −log[H⁺]; below 7 acidic, 7 neutral, above 7 basic. At 25 °C, pH + pOH = 14.

A strong acid ionizes completely (HCl); a weak acid only partially ionizes, described by its Ka.

The ion-product of water, Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C; Ka·Kb = Kw for a conjugate pair.

A solution of a weak acid + its conjugate base that resists pH change when small amounts of acid/base are added.

pH = pKa + log([A⁻]/[HA]); a buffer works best when pH ≈ pKa.

Where moles of added titrant exactly neutralize the analyte; an indicator marks it by changing color.

The equilibrium constant for dissolving a slightly soluble salt; a smaller Ksp means a less soluble compound.

Adding an ion already present in the equilibrium shifts it toward the solid, decreasing solubility.

Two species differing by one H⁺ (e.g., NH₄⁺/NH₃); the stronger the acid, the weaker its conjugate base.

The answer keeps the fewest significant figures of any factor used.

The answer keeps the fewest decimal places of any value used.

Accuracy is closeness to the true value; precision is reproducibility — they are independent.

A systematic error (uncalibrated balance) biases every reading the same way; random error scatters readings unpredictably.

A technique to find an unknown concentration by adding a measured titrant until the equivalence point (indicator color change).

Separates liquids by boiling point — the lower-boiling component vaporizes and is collected first.

Quantifies an analyte by the mass of a precipitate that is filtered, dried, and weighed.

Particulate (atoms/molecules), macroscopic (what you observe), and symbolic (formulas, equations, graphs).

A titration run with all reagents except the analyte, used to correct for impurities in the reagents.

To check linearity — calibrating at two known pH values verifies the meter's response across a range.

Heating with a vertical condenser so vapor condenses and returns, allowing a long reaction without losing volatile reactants.

Exact (counted or defined) numbers have unlimited significant figures; measured numbers carry uncertainty in the last digit.