CLEP Chemistry Quick Review

Every atom is defined by its number of protons (atomic number), while its mass comes from protons plus neutrons, and the mass number can vary between atoms of the same element.

Atomic number $Z=$ number of protons (defines the element); mass number $A = $ protons $+$ neutrons; isotope notation $^A_ZX$. Average atomic mass is the weighted average of all naturally occurring isotopes: $m= (isotope mass)(fractional abundance)$, measured experimentally by mass spectrometry (peaks = isotope mass-to-charge ratios; peak height = relative abundance).

Isotopes (same $Z$, different $A$/neutron number, same chemical behavior) vs. ions (same $Z$ and $A$, different electron count, different charge).

Treating the mass number on the periodic table as a whole number for a single atom --- the periodic table lists the weighted AVERAGE atomic mass across isotopes, so it is almost never a whole number, while any individual atom's mass number is always an integer.

$Z=$ protons; $A=$ protons+neutrons; periodic table mass = weighted average, not one atom's mass.

Electrons occupy quantized energy levels, and the pattern in which they fill those levels reproduces the structure of the periodic table.

Aufbau principle (fill lowest energy orbitals first), Pauli exclusion principle (max 2 electrons per orbital, opposite spin), Hund's rule (singly occupy degenerate orbitals before pairing). Filling order: $1s 2s 2p 3s 3p 4s 3d 4p 5s…$ Atomic emission spectra (discrete bright lines) result from electrons dropping from higher to lower energy levels, releasing photons of energy $E=h=/hc$.

Ground-state configuration (lowest-energy, most stable arrangement) vs. excited-state configuration (an electron promoted to a higher, unfilled orbital).

Removing electrons from the $3d$ subshell before $4s$ when forming a transition-metal cation --- despite $4s$ filling first, electrons are always removed from the HIGHEST principal quantum number subshell first, so $Fe arrow Fe^2+$ loses its two $4s$ electrons, not $3d$.

Aufbau + Pauli + Hund fill electrons; cations lose highest-$n$ electrons first.

Atomic radius, ionization energy, and electronegativity all trace back to the tug-of-war between nuclear charge and electron shielding.

Across a period (left$→$right): atomic radius $$, ionization energy $$, electronegativity $$, electron affinity generally $$ (more negative). Down a group: atomic radius $$, ionization energy $$, electronegativity $$. Effective nuclear charge $Z_eff≈ Z - S$ (shielding) drives the across-period trends.

First ionization energy (removing the first, outermost electron) vs. successive ionization energies (a large jump occurs after removing all valence electrons, when a full inner shell must be broken).

Ranking isoelectronic ions by size using charge sign alone --- for a fixed electron count, MORE protons pulls electrons in tighter, so radius shrinks as nuclear charge rises: $Al^3+ < Mg^2+ < Na+ < F- < O^2-$.

Across period: size$$, IE$$, EN$$; down group: size$$, IE$$, EN$$.

Unstable nuclei release particles or energy in predictable ways, and the amount of radioactive material remaining decays exponentially with a characteristic half-life.

Alpha decay: $^A_ZXarrow ^A-4_Z-2Y + ^4_2He$ (mass $-4$, atomic number $-2$). Beta decay: $^A_ZXarrow ^A_Z+1Y + ^0_-1e$ (mass unchanged, atomic number $+1$). Gamma decay releases energy only. Half-life: $N=N_0≤ft(/12)^t/t_1/2$; mass number and charge are always conserved in nuclear equations.

Alpha particles (helium nuclei, high mass, low penetration, stopped by paper) vs. beta particles (electrons, low mass, moderate penetration) vs. gamma rays (pure energy, highest penetration).

Forgetting that mass number and atomic number must each balance on both sides of a nuclear equation --- always solve for the unknown particle by conservation, not by guessing the decay type.

Alpha: $A-4,Z-2$; beta: $A$ same, $Z+1$; half-life $N=N_0(1/2)^t/t_1/2$.

The type of chemical bond that forms is dictated by how the electronegativity difference between atoms determines whether electrons transfer, share, or delocalize.

Ionic bonding: large electronegativity difference, metal + nonmetal, electron transfer forms a lattice of oppositely charged ions. Covalent bonding: nonmetal + nonmetal, electrons shared (polar covalent if EN difference moderate, nonpolar if near zero). Metallic bonding: a lattice of metal cations surrounded by a delocalized ``sea'' of electrons, explaining conductivity and malleability.

Intramolecular forces (ionic, covalent, metallic bonds --- hold atoms/ions together within a compound) vs. intermolecular forces (London dispersion, dipole-dipole, hydrogen bonding --- act between separate molecules).

Assuming a bond is purely ionic or purely covalent from one electronegativity cutoff --- CLEP expects a continuum: EN difference $≈ 0$ = nonpolar covalent, small-to-moderate = polar covalent, large ($ 1.7$) = ionic.

Metal+nonmetal $→$ ionic (transfer); nonmetal+nonmetal $→$ covalent (share); metal+metal $→$ metallic (delocalize).

A correctly drawn Lewis structure predicts a molecule's three-dimensional shape because electron domains around the central atom arrange themselves to minimize repulsion.

VSEPR shapes by electron domain count: 2 domains = linear ($180^$); 3 = trigonal planar ($120^$); 4 = tetrahedral ($109.5^$); with lone pairs, 4 domains with 1 lone pair = trigonal pyramidal ($≈ 107^$, e.g. $NH3$), 4 domains with 2 lone pairs = bent ($≈ 104.5^$, e.g. $H2O$); 5 domains = trigonal bipyramidal; 6 = octahedral. Lone pairs repel more strongly than bonding pairs, compressing bond angles.

Electron-domain geometry (counts lone pairs and bonding pairs) vs. molecular geometry (describes only the positions of atoms, ignoring lone pairs) --- e.g., water is tetrahedral electron-domain geometry but bent molecular geometry.

Forgetting to count lone pairs on the central atom as electron domains --- omitting them gives the wrong steric number and therefore the wrong predicted shape and bond angle.

Count all domains (bonds + lone pairs) first; lone pairs squeeze bond angles smaller.

Hybridization mixes atomic orbitals into new orbitals whose geometry matches the number of electron domains, and every bond is built from a sigma framework with pi bonds added for extra bond order.

$sp$ (2 domains, linear), $sp^2$ (3 domains, trigonal planar), $sp^3$ (4 domains, tetrahedral), $sp^3d$ (5, trigonal bipyramidal), $sp^3d^2$ (6, octahedral). A single bond = 1 sigma bond (head-on overlap); a double bond = 1 sigma + 1 pi bond (side-on overlap); a triple bond = 1 sigma + 2 pi bonds.

Sigma bonds (allow free rotation around the bond axis) vs. pi bonds (restrict rotation, which is why double bonds create cis/trans geometric isomers).

Counting a double or triple bond as more than one electron domain when assigning hybridization --- a multiple bond still occupies only ONE electron domain/one direction in space.

Domains = hybridization letters; single=$$; double=$+$; triple=$+2$.

Whether a molecule is polar depends on both individual bond polarities and molecular symmetry, since symmetric arrangements of polar bonds can cancel to give a nonpolar molecule.

A molecule is polar only if it has polar bonds AND an asymmetric geometry so bond dipoles do not cancel (e.g., $CO2$ is linear and nonpolar despite polar C=O bonds; $H2O$ is bent and polar). Resonance structures are different valid Lewis structures for the same molecule that differ only in electron placement; the real structure is a weighted hybrid/average, not an alternating mix.

A polar bond (unequal electron sharing within one bond, judged by EN difference) vs. a polar molecule (net dipole moment of the whole 3-D structure, judged by geometry).

Assuming any molecule with polar bonds must itself be polar --- always check molecular symmetry; symmetric molecules like $CO2$, $BF3$, and $CCl4$ are nonpolar overall.

Polar bonds + asymmetric shape = polar molecule; polar bonds + symmetric shape = nonpolar.

Transition metals form coordination complexes with ligands that donate electron pairs, and the same molecular formula can correspond to structurally distinct isomers.

A coordination complex consists of a central metal ion bonded to surrounding ligands (Lewis bases) through coordinate covalent (dative) bonds, where the ligand supplies both electrons. Structural isomers have the same molecular formula but different atom connectivity; stereoisomers (e.g., cis/trans) have identical connectivity but different spatial arrangement.

Structural (constitutional) isomers (atoms bonded in a different order) vs. stereoisomers (same bonding order, different 3-D arrangement in space, such as cis/trans about a double bond or ring).

Assuming two molecules with the same molecular formula must be identical --- always check connectivity and spatial arrangement before concluding two structures are the same compound.

Same formula: different connectivity = structural isomer; different 3-D arrangement = stereoisomer.

All the individual gas laws (Boyle's, Charles's, Gay-Lussac's) are special cases of one master equation relating pressure, volume, moles, and temperature.

Ideal gas law: $PV=nRT$, with $R=0.0821\ /L·atmmol·K$ (or $8.314\ /Jmol·K$). Combined gas law: $/P_1V_1T_1=/P_2V_2T_2$. At STP ($0^C$, 1 atm) 1 mole of ideal gas occupies 22.4 L. Dalton's law of partial pressures: $P_total= P_i$, with $P_i = X_i P_total$.

STP (standard temperature and pressure, used for molar volume: $0^$C, 1 atm) vs. standard conditions for thermodynamics/electrochemistry ($25^$C, 1 atm, 1 M).

Plugging a Celsius temperature directly into $PV=nRT$ --- temperature MUST be converted to Kelvin ($K=^C+273$) first, or every calculated value will be wrong.

$PV=nRT$; always convert to Kelvin; 22.4 L/mol only at STP.

The ideal gas model assumes molecules have no volume and no intermolecular attraction, so real gases deviate most where those assumptions break down: high pressure and low temperature.

Kinetic molecular theory: gas particles are in constant random motion; average kinetic energy is proportional to Kelvin temperature ($KE_avg=/32RT$ per mole); collisions are elastic. Graham's law of effusion: $/rate_1rate_2=sqrt/M_2M_1$ (lighter gases effuse faster). Real gases deviate from ideal behavior at high pressure (molecular volume becomes significant) and low temperature (intermolecular attractions become significant).

Effusion (gas escaping through a tiny hole into a vacuum) vs. diffusion (gas spreading through another gas already present) --- both follow the same inverse-square-root molar mass relationship.

Assuming heavier gases always effuse faster --- Graham's law shows the LIGHTER gas (smaller molar mass) effuses faster, since rate is inversely proportional to $sqrtM$.

$KE_avg T$ only; light gases effuse fast; real gases misbehave at high $P$, low $T$.

A phase diagram maps which state of matter is stable at every combination of temperature and pressure, with special points marking where phases coexist or become indistinguishable.

Phase diagram regions are separated by curves representing fusion (solid$≤ftrightarrow$liquid), vaporization (liquid$≤ftrightarrow$gas), and sublimation (solid$≤ftrightarrow$gas) equilibria. The triple point is the unique temperature/pressure where all three phases coexist; the critical point is where the liquid-gas boundary ends and the distinction between liquid and gas disappears.

Melting/boiling point (temperature at 1 atm where phase change occurs) vs. the triple point (a single fixed pressure-temperature combination unique to each substance, not adjustable by choosing 1 atm).

Reading a phase diagram's axes backwards --- pressure is virtually always the y-axis and temperature the x-axis; misreading which axis is which flips every phase-boundary interpretation.

Triple point = all 3 phases coexist; critical point = liquid/gas distinction vanishes.

The same solution can be described by several different concentration units, and choosing the right one (and converting correctly between them) is essential for solving stoichiometry and colligative-property problems.

Molarity $M=/mol soluteL solution$; molality $m=/mol solutekg solvent$; mole fraction $X_i=/n_in_total$; dilution equation $M_1V_1=M_2V_2$. ``Like dissolves like'': polar solvents dissolve ionic/polar solutes; nonpolar solvents dissolve nonpolar solutes.

Molarity (moles per liter of total SOLUTION, temperature-dependent since volume expands/contracts) vs. molality (moles per kilogram of SOLVENT only, temperature-independent) --- molality is used for colligative-property calculations.

Using the volume of solvent added instead of the final total solution volume when calculating molarity --- molarity always uses the volume of the completed solution, not just the water poured in.

Molarity: mol/L solution; molality: mol/kg solvent (temperature-independent).

Colligative properties depend only on the NUMBER of dissolved solute particles, not on their chemical identity, so ionic compounds that dissociate have a magnified effect.

Freezing point depression: $ T_f = iK_fm$; boiling point elevation: $ T_b=iK_bm$; osmotic pressure: $=iMRT$; Raoult's law for vapor pressure lowering: $P_solution=X_solventP^_solvent$. The van't Hoff factor $i$ = number of particles a formula unit produces in solution (e.g., $i=2$ for $NaCl$, $i=3$ for $CaCl2$, $i=1$ for a nonelectrolyte like glucose).

A nonelectrolyte (does not dissociate, $i=1$, e.g. sucrose) vs. a strong electrolyte (fully dissociates into ions, $i>1$, e.g. $MgCl2$ has $i=3$) --- forgetting to multiply by $i$ is the single most common colligative-property error.

Forgetting the van't Hoff factor $i$ for ionic solutes --- a 1 m $NaCl$ solution depresses the freezing point roughly TWICE as much as a 1 m glucose solution because it produces 2 particles per formula unit, not 1.

$ T=iKm$; always multiply by $i$ = particles produced per formula unit.

Three progressively broader definitions of acids and bases (Arrhenius, Br nsted-Lowry, Lewis) let chemists classify acid-base behavior even when no water or protons are involved.

Arrhenius: acids produce $H+$ in water, bases produce $OH-$. Br nsted-Lowry: acids are proton ($H+$) donors, bases are proton acceptors (broader --- explains conjugate acid-base pairs). Lewis: acids are electron-pair acceptors, bases are electron-pair donors (broadest --- explains reactions with no protons at all, like $BF3$ accepting a lone pair). Amphoteric species (e.g., $H2O$, $HCO3-$, $Al(OH)3$) can act as either acid or base depending on the other reactant.

Amphoteric (can act as either an acid or a base, chemically) vs. amphiprotic (specifically can either donate or accept a proton) --- amphiprotic is a subset of amphoteric.

Assuming every acid-base reaction must involve $H+$ transfer --- Lewis acid-base reactions (like metal-ligand coordination) involve no proton transfer at all, only electron-pair donation.

Arrhenius $$ Br nsted-Lowry $$ Lewis (broadest = electron pairs, no protons required).

Water self-ionizes to a tiny but fixed extent at a given temperature, which links the $H+$ and $OH-$ concentrations of every aqueous solution, acidic or basic.

$K_w=[H+][OH-]=1.0×10^-14$ at $25^$C; $pH=-$, $pOH=-$, and $pH+pOH=14$ at $25^$C. Strong acids/bases dissociate completely (memorize: $HCl, HBr, HI, HNO3, H2SO4, HClO4$ are strong acids; Group 1 hydroxides and $Ca(OH)2, Sr(OH)2, Ba(OH)2$ are strong bases); weak acids/bases only partially ionize, governed by $K_a$ or $K_b$.

A strong acid/base (completely dissociates --- a statement about the EXTENT of ionization) vs. a concentrated acid/base (a statement about the AMOUNT of acid per volume) --- a dilute strong acid can still have low pH-lowering power in total moles, and a concentrated weak acid is still a weak acid.

Treating pH like a linear scale --- pH is logarithmic, so each whole-number drop in pH means a TENFOLD increase in $[H+]$; a solution of pH 3 is 100 times more acidic than pH 5, not twice as acidic.

pH $=-$; log scale $⇒$ each unit = $10×$ change in $[H+]$.

When two ionic solutions are mixed, a precipitate forms only if the combination produces an insoluble salt, and the net ionic equation isolates just the species that actually react.

Solubility rules (memorize the big ones): nitrates, alkali metal salts, and ammonium salts are always soluble; most chlorides/bromides/iodides are soluble EXCEPT with $Ag+, Pb^2+, Hg2^2+$; most sulfates are soluble EXCEPT $BaSO4, PbSO4, CaSO4$; most carbonates, phosphates, and hydroxides are insoluble (exceptions with Group 1/$NH4+$). Spectator ions (unchanged, present on both sides) are canceled to write the net ionic equation.

A complete ionic equation (shows ALL strong electrolytes as separated ions, including spectators) vs. a net ionic equation (spectators removed, showing only the species that actually form the precipitate/react).

Forgetting to cancel spectator ions, or canceling an ion that is NOT actually a spectator because its physical state differs on each side (e.g., an ion that becomes part of a solid precipitate is not a true spectator).

Soluble = stays as ions; insoluble = precipitates (s); net ionic = cancel true spectators only.

Redox reactions transfer electrons between species, and arranging that transfer through an external circuit in a galvanic cell converts chemical energy directly into electrical energy.

Oxidation = loss of electrons (oxidation number increases); reduction = gain of electrons (oxidation number decreases) --- ``OIL RIG.'' Standard cell potential: $E^_cell=E^_cathode(red)-E^_anode(red)$; a positive $E^_cell$ means the reaction is spontaneous. In a galvanic (voltaic) cell, oxidation occurs at the anode and reduction at the cathode (``AN OX, RED CAT''); electrons flow anode$→$cathode through the wire. Electrolytic cells use an external power source to force a nonspontaneous redox reaction.

A galvanic/voltaic cell (spontaneous, $ G<0$, $E^_cell>0$, generates electricity) vs. an electrolytic cell (nonspontaneous, $ G>0$, $E^_cell<0$, consumes electricity to force the reaction).

Subtracting the half-reaction potentials in the wrong order, or reversing the sign of $E^$ when flipping a reduction half-reaction to an oxidation --- always use $E^_cell=E^_cathode-E^_anode$ with BOTH values taken from the reduction-potential table as given (never flip the tabulated sign).

$E^_cell=E^_cathode-E^_anode$; positive $E^_cell$ = spontaneous = galvanic.

A balanced equation is a statement of conservation --- of both mass (atoms) and, for redox/ionic equations, of charge.

Balance atoms first using coefficients (never subscripts, which would change the identity of the compound); for redox equations balance charge too by ensuring electrons lost = electrons gained. Ionic species must balance in both mass AND net charge on each side of a net ionic equation.

Changing a coefficient (scales the whole formula, keeps the compound's identity, always allowed) vs. changing a subscript (changes the compound into a different substance entirely, never allowed when balancing).

Balancing atoms but forgetting to check that total charge also balances on both sides of an ionic or redox equation --- an equation can look mass-balanced yet still be charge-unbalanced, which is invalid.

Balance atoms with coefficients only; for ionic/redox equations, charge must balance too.

The mole is the bridge that converts between the microscopic world of atoms/molecules and the macroscopic world of grams and liters that can be measured in lab.

$1\ mol=6.022×10^23$ particles (Avogadro's number). Molar mass converts grams$≤ftrightarrow$moles. Stoichiometric mole ratios come directly from balanced-equation coefficients: $mol A×/coefficient Bcoefficient A=mol B$. General stoichiometry path: grams A $→$ moles A $→$ moles B (mole ratio) $→$ grams B.

Empirical formula (simplest whole-number mole ratio of atoms) vs. molecular formula (actual number of atoms in one molecule, an integer multiple of the empirical formula, found using molar mass).

Converting directly from grams of A to grams of B without going through moles and the mole-ratio step --- skipping the mole ratio (the balanced-equation coefficients) is the single most common stoichiometry error on CLEP.

Grams $→$ moles $→$ (mole ratio from coefficients) $→$ moles $→$ grams --- never skip the mole step.

When reactants are not present in the exact stoichiometric ratio, one runs out first and caps how much product can actually form, regardless of how much of the other reactant remains.

The limiting reactant is identified by comparing the mole ratio available to the mole ratio required by the balanced equation --- whichever reactant would produce LESS product is limiting. Percent yield $=/actual yieldtheoretical yield×100%$. For solution stoichiometry, moles of solute $= M× V$ before applying the mole ratio.

Theoretical yield (calculated maximum product from the limiting reactant, assuming 100% efficiency) vs. actual yield (the real, experimentally measured amount, always $≤$ theoretical due to side reactions/losses).

Identifying the limiting reactant as whichever reagent has the smaller starting MASS or fewer starting MOLES --- this is wrong; you must convert each reactant's amount through the mole ratio to see which produces the least product.

Limiting reactant = least product possible, found via mole ratio, not raw mass/moles.

At equilibrium the forward and reverse reaction rates are equal (not zero), and the equilibrium constant is a fixed ratio of products to reactants that a system will shift to restore whenever disturbed.

$K_c=/[C]^c[D]^d[A]^a[B]^b$ for $aA+bBleftharpoons cC+dD$ (pure solids/liquids omitted, activity $=1$). $K_p=K_c(RT)^ n$, where $ n=$(mol gas products)$-$(mol gas reactants). Le Ch\^atelier's principle: a system at equilibrium shifts to relieve an applied stress (added/removed reactant or product, volume/pressure change, temperature change); a catalyst speeds up the approach to equilibrium but does NOT shift its position.

The reaction quotient $Q$ (calculated at any moment using current concentrations) vs. the equilibrium constant $K$ (the fixed value $Q$ equals only at equilibrium) --- if $Q<K$ the reaction shifts forward; if $Q>K$ it shifts in reverse.

Forgetting to raise each concentration to the power of its balanced-equation coefficient in the $K_c$ expression, or including pure solids/liquids in that expression --- both errors give a wrong equilibrium constant.

$K_c$: products over reactants, each raised to its coefficient; solids/liquids omitted; catalyst never shifts $K$.

Weak acids and bases only partially ionize, and mixing a weak acid with its conjugate base creates a buffer that resists pH change by consuming added acid or base.

$K_a=/[H+][A-][HA]$; for conjugate acid-base pairs $K_a× K_b=K_w$. Henderson-Hasselbalch equation for buffers: $pH=pK_a+/[A-][HA]$. Buffer capacity is greatest when $[A-]=[HA]$, i.e., when $pH=pK_a$.

$K_a$ (equilibrium constant for a weak acid's ionization, describes STRENGTH) vs. p$K_a$ ($- K_a$, a convenient pH-scale number where SMALLER p$K_a$ means a STRONGER acid) --- the two move in opposite directions.

Forgetting that a buffer requires a weak acid/base PLUS a significant amount of its conjugate --- mixing a strong acid with its salt does not buffer, because strong acids dissociate completely and leave no equilibrium to shift.

pH $=$ p$K_a+([A-]/[HA])$; buffer works best when p$K_a≈$ target pH.

$K_sp$ quantifies how far a sparingly soluble salt dissolves at equilibrium, and adding an ion the salt already produces suppresses further dissolution.

For $M_aX_bleftharpoons aM^n++bX^m-$: $K_sp=[M^n+]^a[X^m-]^b$. If the ion product $Q_sp>K_sp$, a precipitate forms; if $Q_sp<K_sp$, the solution is unsaturated and more solid can dissolve. The common ion effect: adding a soluble salt that shares an ion with a sparingly soluble salt shifts its dissolution equilibrium backward, DECREASING its molar solubility (Le Ch\^atelier).

$K_sp$ (a fixed equilibrium constant for a given salt at a given temperature) vs. molar solubility (the actual concentration that dissolves, which changes with common ions present, even though $K_sp$ itself does not change).

Forgetting the stoichiometric coefficients as exponents when writing $K_sp$ --- for $Ca(OH)2$, $K_sp=[Ca^2+][OH-]^2$, not $[Ca^2+][OH-]$.

$K_sp$ fixed; common ion added $⇒$ solubility drops, $K_sp$ itself unchanged.

The rate law for a reaction must be determined from experimental data, not read off the balanced equation, because reaction order reflects mechanism, not stoichiometry.

Rate law: $rate=k[A]^m[B]^n$, where $m$ and $n$ (the reaction orders) are found experimentally, typically via the method of initial rates (comparing how rate changes as one concentration is varied while others are held constant). Overall reaction order $=m+n$. First-order integrated rate law: $=-kt+_0$, with a constant half-life $t_1/2=/0.693k$ independent of concentration.

Reaction order (an experimentally determined exponent in the rate law, can be 0, 1, 2, or even fractional) vs. stoichiometric coefficient (comes from the balanced overall equation) --- these are only equal for elementary (single-step) reactions.

Assuming the rate law's exponents equal the balanced equation's coefficients for an overall (multi-step) reaction --- this shortcut only works for an elementary step; for the overall reaction, orders must come from data.

Rate law exponents = experimental, not stoichiometric (unless the step is elementary).

Reaction rate depends on both how often molecules collide and what fraction of those collisions have enough energy and correct orientation to react, and catalysts work by lowering that energy barrier.

Arrhenius equation: $k=Ae^-E_a/RT$ --- raising temperature increases $k$ primarily by increasing the FRACTION of collisions with energy $≥ E_a$, not mainly by increasing collision frequency. A catalyst provides an alternate mechanism with lower activation energy, is not consumed in the overall reaction, and speeds both forward and reverse rates equally (does not shift $K$). In a multistep mechanism, the rate-determining step (the slowest elementary step) controls the overall rate law.

A catalyst (not consumed, appears in the mechanism but cancels out overall) vs. a reaction intermediate (produced in one step, consumed in a later step, never appears in the overall balanced equation or the observed rate law).

Believing a catalyst changes the equilibrium constant or the amount of product formed --- a catalyst only speeds up how FAST equilibrium is reached, it never changes $K$ or the equilibrium position.

$k=Ae^-E_a/RT$; catalyst lowers $E_a$, unconsumed, never shifts $K$.

Enthalpy change is a state function, so the total heat of a reaction can be calculated by adding up steps in any convenient pathway, as long as the overall start and end states match.

Hess's Law: $ H_rxn= H^_f(products)- H^_f(reactants)$; if a reaction is reversed, the sign of $ H$ flips; if scaled by a factor, $ H$ scales by the same factor. Calorimetry: $q=mc T$, where $c$ is specific heat capacity; heat lost by one substance equals heat gained by another in an isolated system ($q_lost=-q_gained$).

Exothermic ($ H<0$, system releases heat to surroundings, surroundings warm up) vs. endothermic ($ H>0$, system absorbs heat, surroundings cool down).

Dropping or flipping the sign of $ H$ when a reaction is reversed in a Hess's Law sum, or forgetting to multiply $ H$ by the same factor used to scale a reaction's coefficients.

Reverse reaction $⇒$ flip sign of $ H$; scale coefficients $⇒$ scale $ H$ equally.

Whether a reaction is spontaneous depends on the combined balance of enthalpy and entropy changes at a given temperature, captured in a single quantity: Gibbs free energy.

$ G= H-T S$; a reaction is spontaneous when $ G<0$, nonspontaneous when $ G>0$, and at equilibrium when $ G=0$. Entropy $ S$ generally increases when a solid becomes a liquid or gas, when moles of gas increase, or when a solution forms from a pure solid. $ G^=-RT K$ links free energy to the equilibrium constant, and $ G^=-nFE^_cell$ links it to cell potential.

A reaction that is exothermic ($ H<0$) vs. a reaction that is spontaneous ($ G<0$) --- these are NOT synonyms; an endothermic reaction can still be spontaneous if $ S$ is sufficiently positive and $T$ is large enough.

Assuming every exothermic reaction is automatically spontaneous at all temperatures --- spontaneity depends on the SIGN COMBINATION of $ H$ and $ S$ and on temperature; only $ H<0$ with $ S>0$ is spontaneous at every temperature.

$ G= H-T S<0$ = spontaneous; exothermic $$ spontaneous automatically.

An element's chemical reactivity and typical reaction products can be predicted from its horizontal, vertical, and diagonal position on the periodic table.

Group (vertical) relationships: elements in the same group share similar valence electron configuration and similar reactivity (e.g., alkali metals all react vigorously with water to form $MOH+H2$). Period (horizontal) relationships: metallic character decreases and nonmetallic character increases left to right. Diagonal relationships: an element resembles the element diagonally below-right of it more than its own group neighbor in some cases (e.g., $Li$ resembles $Mg$, $Be$ resembles $Al$) due to similar charge density.

A group trend (elements share a similar number of valence electrons, so similar bonding behavior) vs. a diagonal relationship (similar charge-to-size ratio produces similar reactivity despite different groups).

Assuming reactivity trends within a group are always monotonic without exception --- some properties (like first ionization energy) have well-known dips at specific elements (Be/B, N/O) that must be recognized as legitimate exceptions, not errors.

Groups share valence e$^-$ count; periods show metal$→$nonmetal shift; diagonal pairs share charge density.

Main group elements have predictable oxidation states tied directly to their group number, which lets you predict the formulas of the compounds and ions they form.

Group 1 (alkali metals) form $+1$ ions; Group 2 (alkaline earth metals) form $+2$ ions; Group 17 (halogens) form $-1$ ions and diatomic molecules ($F2, Cl2, Br2, I2$); Group 18 (noble gases) are largely unreactive due to full valence shells. Nonmetal oxides tend to form acidic solutions in water (e.g., $CO2 + H2O -> H2CO3$); metal oxides tend to form basic solutions (e.g., $CaO + H2O -> Ca(OH)2$).

A metal oxide (typically basic/ionic, reacts with water or acid) vs. a nonmetal oxide (typically acidic/molecular, reacts with water or base) --- this acid/base character tracks directly with metallic vs. nonmetallic character.

Assuming all main-group elements only ever show their ``default'' group oxidation state --- several (like sulfur, nitrogen, and the halogens beyond fluorine) commonly show multiple oxidation states depending on the compound.

Metal oxide + water $→$ base; nonmetal oxide + water $→$ acid.

Transition metals commonly show multiple oxidation states and form colored complex ions because their partially filled $d$ orbitals split in energy when surrounded by ligands.

Transition metals (unlike most main group elements) typically exhibit VARIABLE oxidation states (e.g., $Fe^2+$/$Fe^3+$, $Cu+$/$Cu^2+$) because $d$-electrons are lost at similar energies. Coordination compounds form when ligands (Lewis bases) donate electron pairs to a central metal cation; the color of many transition-metal complexes arises from electrons absorbing visible light to jump between split $d$-orbital energy levels (crystal/ligand field splitting).

A transition metal's oxidation state (charge after ``losing'' electrons, a bookkeeping number) vs. its coordination number (the number of ligand donor atoms directly bonded to it, e.g., 4 or 6) --- these are independent quantities.

Assuming a transition metal ion always shows the same charge in every compound --- always determine oxidation state from the specific compound's overall charge balance, not from memory of ``the'' charge for that metal.

Transition metals: variable oxidation states, colored complexes from split $d$-orbitals.

A molecule's functional group --- the specific atom arrangement responsible for its characteristic chemistry --- determines its reactivity far more than the size of its carbon skeleton does.

Key functional groups: alcohol ($-OH$), carboxylic acid ($-COOH$), ester ($-COO-$), aldehyde ($-CHO$), ketone ($C=O$ within a chain), ether ($-O-$), amine ($-NH2$). Alkanes are saturated (only single C--C bonds); alkenes contain at least one C=C double bond; alkynes contain at least one triple bond. Isomers with the same molecular formula but different structures can have very different physical/chemical properties.

A ketone (carbonyl group WITHIN a carbon chain, both neighbors are carbon) vs. an aldehyde (carbonyl group at the END of a chain, one neighbor is hydrogen) --- both contain a $C=O$ but in different structural contexts with different reactivity.

Assuming compounds with the same molecular formula must share similar chemical properties --- structural isomers with the same formula (e.g., ethanol vs. dimethyl ether, both $C2H6O$) can have drastically different boiling points and reactivity depending on functional group.

Functional group = reactivity ``address''; same formula can still mean very different chemistry.

Choosing the right piece of glassware and reporting the right number of significant figures both directly affect how precise and accurate an experimental measurement can be.

Volumetric flasks and burets measure volume most precisely; graduated cylinders are moderately precise; beakers/Erlenmeyer flasks are for mixing/holding, not precise measurement. Significant figures: all nonzero digits count; zeros between nonzero digits count; leading zeros never count; trailing zeros count only with a decimal point present. In calculations, multiplication/division results are limited by the LEAST number of sig figs among the inputs; addition/subtraction results are limited by the LEAST number of decimal places.

Precision (how close repeated measurements are to EACH OTHER, reproducibility) vs. accuracy (how close a measurement is to the TRUE/accepted value) --- a measurement can be precise but not accurate (consistently off-target).

Reporting a calculated answer with more significant figures than the least-precise measurement used to calculate it --- a calculator gives many digits, but a chemically honest answer keeps only the sig figs justified by the data.

Multiply/divide $→$ fewest sig figs wins; add/subtract $→$ fewest decimal places wins.

Titration uses a precisely known concentration of one solution to determine the unknown concentration of another, with the equivalence point identified experimentally by a color-change endpoint or a pH-curve inflection.

At the equivalence point of an acid-base titration: $mol acid=mol base$ (accounting for stoichiometric ratio), i.e., $M_aV_a×(acid coefficient ratio)=M_bV_b$. An indicator is chosen so its color-change range brackets the equivalence-point pH as closely as possible; for a strong acid/strong base titration the equivalence point is at pH 7, but for a weak acid/strong base titration it is above pH 7 (basic, due to the conjugate base's hydrolysis).

The equivalence point (the theoretical point where moles of acid and base are exactly stoichiometrically equal) vs. the endpoint (the experimentally observed point where the indicator visibly changes color) --- a well-chosen indicator makes these nearly identical, but they are conceptually distinct.

Assuming every acid-base titration's equivalence point occurs exactly at pH 7 --- this is only true for a strong acid/strong base pair; weak acid/strong base titrations have a basic equivalence point, and weak base/strong acid titrations have an acidic one.

Equivalence point = stoichiometric moles equal (theoretical); endpoint = indicator color change (observed).

One formula sheet: gas law, colligative, pH/buffer, $K_c/K_p$, kinetics, $ G$, $E^_cell$, $q=mc T$.

When two terms sound alike, ask: concentration or amount? Theoretical or observed? Reproducible or correct?

Ask first: does charge/oxidation state change? Yes = redox family; no = acid-base/precipitation family.

Nitrates/alkali/ammonium: always soluble. Everything else: check the named exceptions.

Balance $→$ moles of A $→$ mole ratio $→$ moles of B $→$ convert $→$ check sig figs.

75 Qs / 90 min; calculator + periodic table built into the software; recall + apply + interpret.

Split $→$ balance atoms (not O/H) $→$ balance O/H $→$ balance charge with $e^-$ $→$ match electrons $→$ add.

ICE: Initial, Change ($± x$), Equilibrium --- substitute into $K$, solve for $x$, verify any approximation.

When a statement sounds like a shortcut with no exceptions, CLEP is usually testing the exception.