Atomic Structure, Bonding, and Stoichiometry
The foundation of AP Chemistry rests on understanding atoms, periodic trends, bonding models, and quantitative relationships in chemical formulas and reactions.
Atomic Structure and Electrons
Atoms have a dense nucleus of protons and neutrons surrounded by electrons. Protons carry a +1 charge, neutrons are neutral, and electrons carry a -1 charge. Atomic number equals the number of protons. Mass number equals protons plus neutrons. Isotopes differ in neutron count but have the same atomic number.
Electron configuration follows three rules: the Aufbau principle (fill orbitals from low to high energy), Hund's rule (one electron per orbital before pairing), and Pauli exclusion (maximum two electrons per orbital with opposite spin). Example: iron (Fe) = [Ar] 4s2 3d6. When forming ions, remove s electrons before d electrons.
Photoelectron spectroscopy (PES) measures the ionization energy of electrons in specific orbitals. Peaks at higher binding energy correspond to electrons closer to the nucleus. Peak intensity tells you how many electrons occupy that orbital. This is a major AP exam topic.
Periodic Trends and Periodic Law
Atomic radius decreases left to right across a period due to higher effective nuclear charge. It increases down a group because new electron shells are added. Ionization energy increases left to right and decreases down a group. Electronegativity follows a similar trend, with fluorine being most electronegative.
Coulomb's Law (F = kq1q2/r2) explains many periodic trends. Higher charge or smaller distance means stronger attraction or repulsion. Use this to explain ionization energies, lattice energies, and bond strengths on the AP exam.
Bonding and Molecular Structure
Ionic bonding involves electron transfer between a metal and nonmetal. Ionic compounds form crystal lattices with high melting points and conduct electricity when dissolved or molten. Covalent bonding involves electron sharing between nonmetals. The strength and polarity depend on electronegativity differences.
Electronegativity differences determine bond type: 0 to 0.4 is nonpolar covalent, 0.4 to 1.7 is polar covalent, and greater than 1.7 is typically ionic. Polar bonds create dipole moments. Whether a whole molecule is polar depends on its geometry.
VSEPR Theory (Valence Shell Electron Pair Repulsion) predicts molecular geometry. Electron domains repel and arrange to minimize repulsion. Common geometries include: linear (2 domains, 180 degrees), trigonal planar (3 domains, 120 degrees), tetrahedral (4 domains, 109.5 degrees), trigonal bipyramidal (5 domains), and octahedral (6 domains, 90 degrees).
Intermolecular Forces and Hydrogen Bonding
Intermolecular forces (IMFs) from weakest to strongest are: London dispersion (all molecules, stronger with more electrons), dipole-dipole (polar molecules), hydrogen bonding (H bonded to N, O, or F), and ion-dipole (ions in polar solvents). IMFs explain boiling points, vapor pressure, and solubility.
Hydrogen bonding occurs when hydrogen is bonded to N, O, or F and interacts with another N, O, or F lone pair. This strong dipole-dipole force is responsible for water's unusually high boiling point, DNA base pairing, and protein secondary structure.
Chemical Formulas and Reactions
Empirical formula shows the simplest whole-number ratio of atoms. Molecular formula shows the actual number of atoms. For example, glucose has an empirical formula of CH2O but a molecular formula of C6H12O6. Calculate empirical formula from percent composition and molecular formula from the molar mass.
Limiting reactant is the reactant that runs out first, limiting product formation. Find it by calculating moles of product from each reactant. The smaller yield identifies the limiting reactant. Percent yield equals (actual yield divided by theoretical yield) times 100.
Reaction types include synthesis (A + B yields AB), decomposition (AB yields A + B), single replacement (A + BC yields AC + B), double replacement (AB + CD yields AD + CB), combustion, acid-base neutralization, and redox reactions.
Net ionic equations show only species directly involved in a reaction. Remove spectator ions (present on both sides unchanged). This technique is essential for precipitation, acid-base, and redox reactions in aqueous solution.
Solubility rules determine which ionic compounds dissolve in water. Alkali metals, ammonium, and nitrates are always soluble. Most chlorides, bromides, and iodides are soluble except Ag+, Pb2+, and Hg22+. Most sulfates are soluble except Ba2+, Pb2+, and partially Ca2+. Most carbonates, phosphates, and hydroxides are insoluble except with alkali metals and ammonium.
| Term | Meaning |
|---|---|
| Atomic Structure | Atoms have a dense nucleus of protons (+1 charge, ~1 amu) and neutrons (neutral, ~1 amu) surrounded by electrons (-1 charge, ~1/1836 amu). Atomic number = protons; mass number = protons + neutrons. Isotopes differ in neutron count. |
| Mole Concept | 1 mole = Avogadro's number = 6.022 × 10^23 particles. Molar mass in g/mol equals atomic or molecular mass in amu. Bridges the particle and macroscopic scales: moles = mass / molar mass = particles / 6.022 × 10^23. |
| Electron Configuration | Orbitals filled by Aufbau principle (low to high energy), Hund's rule (one per degenerate orbital before pairing), and Pauli exclusion (max two electrons per orbital, opposite spin). Example: Fe = [Ar] 4s^2 3d^6. Ions: remove s electrons before d. |
| Photoelectron Spectroscopy (PES) | Measures the ionization energy of electrons in specific orbitals. Peaks at higher binding energy correspond to electrons closer to the nucleus (1s). Peak intensity is proportional to the number of electrons in that orbital. Major AP exam topic. |
| Periodic Trends | Atomic radius: decreases left to right (more effective nuclear charge), increases down a group (new shell). Ionization energy: increases left to right, decreases down. Electronegativity: similar trend; F is most electronegative. Metallic character opposes these. |
| Coulomb's Law | F = kq1q2/r^2. Explains many periodic trends: higher charge or smaller distance means stronger attraction/repulsion. Used qualitatively on the AP exam to explain ionization energies, lattice energies, and bond strengths. |
| Ionic vs. Covalent Bonding | Ionic: electron transfer between metal and nonmetal; forms crystal lattices with high melting points and electrical conductivity when dissolved or molten. Covalent: electron sharing between nonmetals; molecular structures with varying polarity. |
| Polar vs. Nonpolar Covalent Bonds | Determined by electronegativity difference. 0-0.4: nonpolar covalent. 0.4-1.7: polar covalent. >1.7: typically ionic. Polar bonds create dipole moments; whether a molecule is polar also depends on geometry. |
| VSEPR Theory | Valence Shell Electron Pair Repulsion: electron domains (bonding pairs and lone pairs) repel and arrange to minimize repulsion. Common geometries: linear (2 domains, 180°), trigonal planar (3, 120°), tetrahedral (4, 109.5°), trigonal bipyramidal (5), octahedral (6, 90°). |
| Intermolecular Forces (IMFs) | From weakest to strongest: London dispersion (all molecules, stronger with more electrons), dipole-dipole (polar molecules), hydrogen bonding (H bonded to N, O, or F with another N/O/F), and ion-dipole (ions in polar solvents). IMFs explain boiling points, vapor pressure, solubility. |
| Hydrogen Bonding | A strong dipole-dipole force occurring when H is bonded to N, O, or F and interacts with another N/O/F lone pair. Responsible for water's unusually high boiling point, DNA base pairing, and protein secondary structure. |
| Empirical vs. Molecular Formula | Empirical formula: simplest whole-number ratio of atoms (e.g., CH2O for glucose). Molecular formula: actual number of atoms (C6H12O6 for glucose). Calculated from percent composition and, for molecular, the molar mass. |
| Limiting Reactant | The reactant that runs out first, limiting the amount of product formed. Found by calculating mol product from each reactant; the smaller yields. Excess reactant is what remains. Percent yield = (actual yield / theoretical yield) × 100. |
| Types of Reactions | Synthesis (A + B → AB), decomposition (AB → A + B), single replacement (A + BC → AC + B), double replacement (AB + CD → AD + CB), combustion (hydrocarbon + O2 → CO2 + H2O), acid-base neutralization, and redox reactions. |
| Net Ionic Equations | Show only species directly involved in a reaction. Spectator ions (present on both sides unchanged) are eliminated. Used for precipitation, acid-base, and redox reactions in aqueous solution. Key tool on the AP free-response section. |
| Solubility Rules | Alkali metals, ammonium, and nitrates are always soluble. Most chlorides, bromides, iodides are soluble except Ag+, Pb2+, Hg2^2+. Most sulfates are soluble except Ba2+, Pb2+, Ca2+ (partially). Most carbonates, phosphates, hydroxides are insoluble except with alkali metals and ammonium. |
Gases, Solutions, and Kinetics
Understanding gas behavior, solution properties, and reaction rates is essential for predicting how matter behaves and changes.
Gas Laws and Kinetic Molecular Theory
Ideal Gas Law is PV = nRT, where R = 0.0821 L·atm/mol·K or 8.314 J/mol·K. At STP (1 atm, 273.15 K), 1 mole of ideal gas occupies 22.4 liters. Deviations occur at high pressure (molecular volume matters) and low temperature (intermolecular forces matter).
Kinetic Molecular Theory explains gas behavior with five postulates: gas particles move in constant random motion, have negligible volume, undergo elastic collisions, have no intermolecular forces, and have average kinetic energy proportional to temperature. This theory explains pressure, temperature, and diffusion.
Maxwell-Boltzmann Distribution shows molecular speeds in a gas. Higher temperature shifts the curve right and flattens it, placing more molecules at higher speeds. The fraction of molecules above the activation energy determines reaction rate, explaining why higher temperature speeds up reactions.
Dalton's Law of Partial Pressures states that total pressure equals the sum of individual partial pressures: P(total) = P1 + P2 + P3... Mole fraction times total pressure equals partial pressure. This law applies to gas collection over water.
Solutions and Colligative Properties
Molarity (M) equals moles of solute divided by liters of solution. Use the dilution equation M1V1 = M2V2 to solve dilution problems. Molality (m) equals moles of solute divided by kilograms of solvent and does not change with temperature. The AP exam typically uses molarity for aqueous solutions.
Colligative properties depend on solute particle concentration, not identity. These include boiling point elevation (ΔTb = iKbm), freezing point depression (ΔTf = iKfm), vapor pressure lowering, and osmotic pressure (π = iMRT). The variable i is the van't Hoff factor (particles per formula unit).
Reaction Rate and Rate Laws
Reaction rate equals change in concentration divided by change in time. Measure it by monitoring disappearance of reactant or appearance of product. Rate factors include concentration, temperature, surface area, catalyst presence, and the nature of reactants.
Rate law is expressed as: Rate = k[A]^m[B]^n. The exponents m and n are determined experimentally, not from stoichiometric coefficients. Overall order equals m plus n. The units of k depend on overall order. Zero-order reactions proceed at constant rate. First-order reactions show linear behavior when plotting ln[A] versus time. Second-order reactions show linear behavior when plotting 1/A versus time.
Integrated rate laws connect concentration and time. For zero-order: [A] = [A]0 - kt. For first-order: ln[A] = ln[A]0 - kt. The half-life for first-order is ln2/k (constant). For second-order: 1/A = 1/A0 + kt. The half-life for second-order is 1/(k[A]0) and depends on initial concentration.
Activation Energy and Catalysts
Arrhenius Equation is k = Ae(-Ea/RT). Plotting ln k versus 1/T yields a straight line with slope -Ea/R. This shows the exponential dependence of rate on temperature. A 10 degree Celsius rise often roughly doubles the rate. Catalysts lower activation energy without being consumed. They provide an alternative pathway with lower Ea.
Collision Theory states that reactions occur when molecules collide with sufficient energy (at least Ea) and proper orientation. This explains why rate depends on concentration (more collisions) and temperature (more energetic collisions). Low steric factor means only specific orientations produce a reaction.
Elementary reactions are single-step reactions whose rate laws can be written from stoichiometric coefficients. Unimolecular reactions (A yields products, rate = k[A]) and bimolecular reactions (A + B yields products, rate = k[A][B]) are common. Termolecular reactions are rare.
Reaction mechanism is a sequence of elementary steps whose sum gives the overall reaction. The slowest step, called the rate-determining step, governs the overall rate law. Intermediates appear in the mechanism but not in the overall equation. Catalysts appear on both sides or in mechanism steps only.
| Term | Meaning |
|---|---|
| Ideal Gas Law | PV = nRT, where R = 0.0821 L·atm/mol·K or 8.314 J/mol·K. At STP (1 atm, 273.15 K), 1 mol of ideal gas occupies 22.4 L. Deviations occur at high pressure (molecular volume matters) and low temperature (IMFs matter). |
| Kinetic Molecular Theory | Postulates: gas particles are in constant random motion, have negligible volume, undergo elastic collisions, have no IMFs, and have average kinetic energy proportional to temperature (KE = 3/2 RT). Explains pressure, temperature, and diffusion. |
| Maxwell-Boltzmann Distribution | Distribution of molecular speeds in a gas. Higher temperature shifts the curve right and flattens it, more molecules at higher speeds. Fraction above activation energy determines reaction rate, explaining why reactions speed up with temperature. |
| Dalton's Law of Partial Pressures | In a mixture of non-reacting gases, total pressure equals the sum of individual partial pressures: P_total = P1 + P2 + P3 + ... Mole fraction (Xi) times total pressure equals partial pressure. Used in gas collection over water. |
| Molarity and Dilution | Molarity (M) = mol solute / L solution. Dilution equation: M1V1 = M2V2. Molality (m) = mol solute / kg solvent (not affected by temperature). The AP exam favors molarity for aqueous solutions. |
| Colligative Properties | Properties that depend on solute particle concentration, not identity: boiling point elevation (ΔTb = iKbm), freezing point depression (ΔTf = iKfm), vapor pressure lowering, and osmotic pressure (π = iMRT). i = van't Hoff factor (particles per formula unit). |
| Reaction Rate | Rate = change in concentration / change in time. Can be measured by monitoring disappearance of reactant or appearance of product. Rate factors: concentration, temperature, surface area, catalyst, and nature of reactants. |
| Rate Law | Rate = k[A]^m[B]^n, where m and n are reaction orders determined experimentally (not from coefficients). Overall order = m + n. Units of k depend on overall order. Zero-order: rate independent of [A]. First-order: ln[A] vs. t is linear. Second-order: 1/[A] vs. t is linear. |
| Integrated Rate Laws | Zero-order: [A] = [A]0 - kt. First-order: ln[A] = ln[A]0 - kt; half-life = ln2/k (constant). Second-order: 1/[A] = 1/[A]0 + kt; half-life = 1/(k[A]0) (depends on initial concentration). |
| Arrhenius Equation | k = Ae^(-Ea/RT). Plotting ln k vs. 1/T yields a straight line with slope -Ea/R. Shows exponential dependence of rate on temperature: a 10°C rise often roughly doubles the rate. Catalysts lower Ea without being consumed. |
| Reaction Mechanism and Rate-Determining Step | A sequence of elementary steps whose sum gives the overall reaction. The slowest step (rate-determining step) governs the rate law. Intermediates appear in mechanism but not in overall equation; catalysts appear on both sides (or in mechanism steps only). |
| Catalysts | Substances that increase reaction rate by providing an alternative pathway with lower activation energy. Not consumed in the overall reaction. Homogeneous (same phase as reactants) vs. heterogeneous (different phase, e.g., solid catalyst with gas reactants). |
| Collision Theory | Reactions occur when molecules collide with sufficient energy (≥ Ea) and proper orientation. Explains why rate depends on concentration (more collisions) and temperature (more energetic collisions). Low steric factor means only specific orientations produce reaction. |
| Elementary Reactions | Single-step reactions whose rate laws can be written from stoichiometric coefficients. Unimolecular: A → products, rate = k[A]. Bimolecular: A + B → products, rate = k[A][B]. Termolecular is rare. |
| Activation Energy Diagrams | Plot energy vs. reaction progress. Peak represents the transition state (activated complex). Height of peak = Ea. Difference between reactants and products = ΔH (exothermic if products lower, endothermic if higher). Catalyst lowers the peak. |
| Beer-Lambert Law | A = εbc, where A is absorbance, ε is molar absorptivity, b is path length, and c is concentration. Used in spectrophotometry to determine concentration of colored solutions. Linear relationship makes calibration curves useful. |
Thermodynamics and Equilibrium
Energy in chemical processes and the final state reached by reversible reactions drive the highest-value conceptual free-response questions on the AP exam.
Energy and Enthalpy
First Law of Thermodynamics states that energy is conserved: ΔE = q + w. A system gains energy from absorbed heat (q greater than 0) and work done on it (w greater than 0). Internal energy change depends only on initial and final states, not the path taken.
Enthalpy (H) represents heat content at constant pressure: ΔH = qp. Exothermic reactions have ΔH less than 0 (release heat). Endothermic reactions have ΔH greater than 0 (absorb heat). Standard enthalpy of formation (ΔHf degrees) is the enthalpy change to form 1 mole of a compound from elements in standard states.
Hess's Law states that the enthalpy change of a reaction equals the sum of enthalpy changes of its steps, regardless of the path. This lets you calculate ΔH from known reactions. Use the formula: ΔH°(rxn) = Sum of ΔHf degrees (products) minus Sum of ΔHf degrees (reactants).
Calorimetry measures heat using q = mcΔT for a substance or q = CΔT for a calorimeter. A coffee cup calorimeter (constant pressure) measures ΔH. A bomb calorimeter (constant volume) measures ΔE. The specific heat of water is 4.18 J/g degrees C.
Entropy and Spontaneity
Entropy (S) measures disorder or microstate multiplicity. Entropy increases with temperature, volume, number of particles, and phase changes (solid to liquid to gas). Entropy increases when dissolving substances. The Second Law of Thermodynamics states that entropy of the universe always increases in spontaneous processes.
Gibbs Free Energy is calculated as ΔG = ΔH - TΔS. Reactions are spontaneous if ΔG is less than 0. The effect of temperature varies based on the signs of ΔH and ΔS: negative ΔH and positive ΔS are always spontaneous; positive ΔH and positive ΔS are spontaneous at high temperature; negative ΔH and negative ΔS are spontaneous at low temperature; positive ΔH and negative ΔS are never spontaneous.
Equilibrium and the Equilibrium Constant
Equilibrium Constant (K) relates concentrations at equilibrium. For aA + bB yields cC + dD, Kc = [C]^c[D]^d divided by [A]^a[B]^b. Solids and pure liquids are excluded. K greater than 1 favors products. K less than 1 favors reactants. Only temperature changes K.
Le Chatelier's Principle states that a disturbed equilibrium system shifts to counteract the stress. Adding reactant shifts right. Adding product shifts left. Increasing pressure shifts toward fewer gas moles. Raising temperature shifts in the endothermic direction. Lowering temperature shifts in the exothermic direction.
Reaction Quotient (Q) is calculated like K but with current concentrations. If Q is less than K, the reaction proceeds forward. If Q is greater than K, it reverses. If Q equals K, the system is at equilibrium. This predicts the direction of shift before equilibrium is reached.
ΔG and Equilibrium are connected by: ΔG = ΔG degrees + RT ln Q. At equilibrium, ΔG equals 0 and Q equals K, so ΔG degrees = -RT ln K. A large negative ΔG degrees gives a large K (products favored). A large positive ΔG degrees gives a small K (reactants favored).
Solving Equilibrium Problems
ICE Tables (Initial, Change, Equilibrium) systematically solve equilibrium problems. Set up initial concentrations, express changes as plus or minus x using stoichiometry, and substitute equilibrium expressions into K. A small K often allows the approximation that A approximately equals A.
Solubility Product (Ksp) is the equilibrium constant for dissolution of sparingly soluble salts. For AgCl(s) yields Ag+ plus Cl-, Ksp = [Ag+][Cl-]. The common ion effect occurs when you add a common ion, shifting equilibrium toward solid and decreasing solubility.
Acids, Bases, and Buffers
Acids (Brønsted-Lowry definition) are proton donors. Bases are proton acceptors. Strong acids (HCl, HBr, HI, HNO3, H2SO4, HClO4) and strong bases (group 1 hydroxides, Ca/Sr/Ba hydroxides) fully dissociate in water. Weak acids and bases partially dissociate with Ka or Kb values.
pH equals negative log of [H+]. pOH equals negative log of [OH-]. At 25 degrees C, pH plus pOH equals 14 and Kw = [H+][OH-] = 1.0 times 10^-14. Neutral pH is 7. Acidic is below 7. Basic is above 7. Each pH unit represents a 10-fold change in [H+].
Buffers contain a weak acid and its conjugate base, or a weak base and its conjugate acid. They resist pH change when small amounts of acid or base are added. Use the Henderson-Hasselbalch equation: pH = pKa + log([A-]/[HA]). Buffer capacity is highest when [HA] approximately equals [A-].
Titration involves measuring equivalence points. Strong acid plus strong base gives equivalence pH of 7. Weak acid plus strong base gives equivalence pH greater than 7 (conjugate base is basic). Weak base plus strong acid gives equivalence pH less than 7. The half-equivalence point occurs at pH = pKa.
| Term | Meaning |
|---|---|
| First Law of Thermodynamics | Energy is conserved: ΔE = q + w, where q is heat and w is work. System gains energy from heat absorbed (q > 0) and work done on it (w > 0). Internal energy change depends on initial and final states, not the path. |
| Enthalpy (H) | Heat content at constant pressure: ΔH = qp. Exothermic: ΔH < 0 (releases heat). Endothermic: ΔH > 0 (absorbs heat). Standard enthalpy of formation (ΔHf°) is the enthalpy change to form 1 mol of compound from elements in standard states. |
| Hess's Law | The enthalpy change of a reaction equals the sum of the enthalpy changes of its steps, regardless of path. Enables calculation of ΔH from known reactions. Formula: ΔH°rxn = ΣΔHf°(products) - ΣΔHf°(reactants). |
| Calorimetry | Measuring heat: q = mcΔT (for a substance) or q = CΔT (for a calorimeter with heat capacity C). Coffee cup (constant pressure): measures ΔH. Bomb calorimeter (constant volume): measures ΔE. Specific heat of water is 4.18 J/g°C. |
| Entropy (S) | Measure of disorder or microstate multiplicity. Entropy increases with temperature, volume, number of particles, phase changes (s → l → g), and dissolving. Second law: the entropy of the universe always increases in spontaneous processes. |
| Gibbs Free Energy | ΔG = ΔH - TΔS. Spontaneous if ΔG < 0. Temperature-dependent: negative ΔH and positive ΔS always spontaneous; positive ΔH and positive ΔS spontaneous at high T; negative ΔH and negative ΔS spontaneous at low T; positive ΔH and negative ΔS never spontaneous. |
| ΔG and Equilibrium | ΔG = ΔG° + RT ln Q. At equilibrium, ΔG = 0 and Q = K, so ΔG° = -RT ln K. Large negative ΔG° → large K (products favored). Large positive ΔG° → small K (reactants favored). K is unitless thermodynamically. |
| Equilibrium Constant (K) | For aA + bB ⇌ cC + dD, Kc = [C]^c[D]^d / [A]^a[B]^b. Solids and pure liquids excluded. K > 1 favors products; K < 1 favors reactants. Temperature is the only factor that changes K. |
| Le Chatelier's Principle | A system at equilibrium disturbed by a stress shifts to partially counteract the stress. Adding reactant shifts right; adding product shifts left. Increasing pressure shifts toward fewer gas moles. Raising T shifts endothermic direction; lowering T shifts exothermic direction. |
| Reaction Quotient (Q) vs. K | Q is calculated like K but with current concentrations. If Q < K: reaction proceeds forward. If Q > K: reverses. If Q = K: at equilibrium. Useful for predicting direction of shift before a system reaches equilibrium. |
| ICE Tables | Initial, Change, Equilibrium, a systematic way to solve equilibrium problems. Set up initial concentrations, express changes as +x or -x using stoichiometry, and substitute equilibrium expressions into K. Small K often allows the approximation [A]initial ≈ [A]equilibrium. |
| Solubility Product (Ksp) | Equilibrium constant for dissolution of a sparingly soluble salt. For AgCl(s) ⇌ Ag+(aq) + Cl-(aq), Ksp = [Ag+][Cl-]. Common ion effect: adding a common ion shifts equilibrium toward solid, decreasing solubility. |
| Acids and Bases (Brønsted-Lowry) | Acid: proton donor. Base: proton acceptor. Strong acids (HCl, HBr, HI, HNO3, H2SO4, HClO4) and strong bases (group 1 hydroxides, Ca/Sr/Ba hydroxides) fully dissociate in water. Weak acids and bases partially dissociate, with Ka or Kb. |
| pH and pOH | pH = -log[H+], pOH = -log[OH-]. At 25°C, pH + pOH = 14, and Kw = [H+][OH-] = 1.0 × 10^-14. Neutral pH = 7; acidic < 7; basic > 7. Each pH unit represents a 10-fold change in [H+]. |
| Buffers and Henderson-Hasselbalch | Buffer: weak acid and its conjugate base (or weak base and conjugate acid). Resists pH change when small amounts of acid or base are added. pH = pKa + log([A-]/[HA]). Buffer capacity is highest when [HA] ≈ [A-] (pH ≈ pKa). |
| Titration and Equivalence Point | Strong acid + strong base: equivalence pH = 7. Weak acid + strong base: equivalence pH > 7 (conjugate base is basic). Weak base + strong acid: equivalence pH < 7. Half-equivalence point: pH = pKa. Indicator chosen to change color at equivalence point. |
Redox, Electrochemistry, and Organic Basics
Electron-transfer reactions, galvanic and electrolytic cells, and carbon-based chemistry form the basis of many organic reactions.
Oxidation-Reduction Reactions
Oxidation is loss of electrons and an increase in oxidation number. Reduction is gain of electrons and a decrease in oxidation number. Remember: OIL RIG (Oxidation Is Loss, Reduction Is Gain). The oxidizing agent is reduced. The reducing agent is oxidized.
Oxidation number rules simplify electron counting. Elements have oxidation number 0. Monatomic ions have oxidation number equal to the charge. Hydrogen is plus 1 (except minus 1 with metals in hydrides). Oxygen is minus 2 (except minus 1 in peroxides and plus 2 with fluorine). Group 1 elements are plus 1. Group 2 elements are plus 2. Fluorine is always minus 1. The sum equals the charge on the species.
Balancing redox in acidic solution uses these steps: split into half-reactions, balance atoms other than O and H, balance O with H2O, balance H with H+, and balance charge with electrons. Multiply half-reactions so electrons cancel. Add and simplify. In basic solution, add OH- to both sides to neutralize H+ and form water.
Galvanic and Electrolytic Cells
Galvanic (Voltaic) cells generate electrical energy from spontaneous redox reactions. Oxidation occurs at the anode (negative terminal). Reduction occurs at the cathode (positive terminal). Electrons flow through the external circuit from anode to cathode. Ions flow through the salt bridge to maintain charge balance.
Standard cell potential is calculated as: E°(cell) = E°(cathode) minus E°(anode). Positive E°(cell) means spontaneous (ΔG degrees = -nFE°(cell)). Standard reduction potentials are measured versus the standard hydrogen electrode (SHE, E degrees = 0). The more positive E°(red), the better the oxidizing agent.
Nernst Equation adjusts cell potential for non-standard conditions: Ecell = E°(cell) - (RT/nF) ln Q. At 25 degrees C: Ecell = E°(cell) - (0.0592/n) log Q. As reactants deplete, Q rises and Ecell falls. At equilibrium, Ecell equals 0 and Q equals K. This explains how battery voltage decreases with use.
Electrolytic cells are driven by external power sources to cause non-spontaneous redox reactions. Uses include electroplating, purifying metals, and producing chemicals like Cl2, NaOH, and Al. Oxidation still occurs at the anode, reduction at the cathode, but now the anode is positive and the cathode is negative.
Faraday's Law of Electrolysis quantifies the relationship between charge and chemical change: Moles = (It) divided by (nF). Here I is current in amperes, t is time in seconds, n is moles of electrons per mole of substance, and F = 96,485 C/mol (Faraday's constant).
Organic Chemistry Basics
Alkanes contain single bonds only with formula CnH(2n+2). They are sp3 hybridized with tetrahedral geometry. Alkenes contain one or more C=C double bonds with formula CnH(2n). They are sp2 hybridized with trigonal planar geometry. Alkynes contain one or more C triple bond C with formula CnH(2n-2). They are sp hybridized with linear geometry.
Functional groups determine reactivity and polarity. Common groups include alcohol (-OH), aldehyde (-CHO), ketone (C=O internal), carboxylic acid (-COOH), ester (-COO-), amine (-NH2), amide (-CONH2), ether (R-O-R'), and halide (R-X).
Structural isomers have different connectivity (n-butane versus isobutane). Stereoisomers have the same connectivity but different spatial arrangement. Cis-trans isomers differ around double bonds. Enantiomers are non-superimposable mirror images of a chiral center.
Hybridization mixes atomic orbitals to form equivalent hybrids. sp3 creates 4 equivalent orbitals with tetrahedral geometry (methane). sp2 creates 3 equivalent orbitals plus an unhybridized p orbital with trigonal planar geometry (ethene). sp creates 2 equivalent orbitals plus two unhybridized p orbitals with linear geometry (ethyne).
Sigma (σ) bonds result from head-on overlap along the internuclear axis. Every single bond is a sigma bond. Pi (π) bonds result from sideways overlap of unhybridized p orbitals. Double bonds contain 1 sigma and 1 pi. Triple bonds contain 1 sigma and 2 pi. Pi bonds restrict rotation.
Nuclear Chemistry
Alpha decay emits a helium-4 nucleus. Mass decreases by 4 and atomic number decreases by 2. Beta decay involves a neutron converting to a proton and electron. Atomic number increases by 1. Gamma emission releases a high-energy photon with no change in mass or charge. Fusion combines light nuclei (occurs in stars). Fission splits heavy nuclei (used in reactors).
Half-life is the time for half of a radioactive sample to decay. This is a first-order process: t(1/2) = ln2/k. Carbon-14 dating uses a half-life of 5,730 years. After n half-lives, (1/2)^n of the original sample remains.
Common constants worth memorizing: R = 0.0821 L·atm/mol·K or 8.314 J/mol·K. F = 96,485 C/mol. Kw = 1.0 times 10^-14 at 25 degrees C. NA = 6.022 times 10^23. Specific heat of water is 4.18 J/g degrees C. 1 atm equals 760 mmHg or 101,325 Pa. Molar volume at STP is 22.4 liters. The AP exam provides a formula sheet but many constants are expected to be known.
| Term | Meaning |
|---|---|
| Oxidation and Reduction | Oxidation: loss of electrons, oxidation number increases. Reduction: gain of electrons, oxidation number decreases. Mnemonic: OIL RIG (Oxidation Is Loss, Reduction Is Gain). Oxidizing agent is reduced; reducing agent is oxidized. |
| Oxidation Number Rules | Elements: 0. Monatomic ions: equal to charge. H: +1 (except -1 with metals in hydrides). O: -2 (except -1 in peroxides, +2 with F). Group 1: +1. Group 2: +2. F: always -1. Sum = charge on species. |
| Balancing Redox in Acidic Solution | Split into half-reactions. Balance atoms other than O and H, then balance O with H2O, H with H+, and charge with electrons. Multiply half-reactions so electrons cancel. Add and simplify. In basic solution, add OH- to both sides to neutralize H+ and form water. |
| Galvanic (Voltaic) Cell | Spontaneous redox reaction generates electrical energy. Oxidation at anode (negative terminal in galvanic), reduction at cathode (positive terminal). Electrons flow through external circuit from anode to cathode; ions flow through salt bridge to maintain charge balance. |
| Standard Cell Potential | E°cell = E°cathode - E°anode. Positive E°cell means spontaneous (ΔG° = -nFE°cell). Standard reduction potentials are measured vs. the standard hydrogen electrode (SHE, E° = 0). The more positive E°red, the better the oxidizing agent. |
| Nernst Equation | Ecell = E°cell - (RT/nF) ln Q, or at 25°C: Ecell = E°cell - (0.0592/n) log Q. As reactants deplete, Q rises and Ecell falls. At equilibrium, Ecell = 0 and Q = K. Explains how battery voltage decreases with use. |
| Electrolytic Cell | Non-spontaneous redox reaction driven by an external power source. Used for electroplating, purifying metals, and producing chemicals (Cl2, NaOH, Al). Oxidation at anode, reduction at cathode, same as galvanic, but anode is positive and cathode is negative. |
| Faraday's Law of Electrolysis | Moles of substance deposited or dissolved = (It)/(nF), where I is current in amperes, t is time in seconds, n is moles of electrons per mole of substance, and F = 96,485 C/mol (Faraday's constant). Links electrical charge to chemical change. |
| Alkanes, Alkenes, and Alkynes | Alkanes: single bonds only, CnH2n+2, sp3 hybridized, tetrahedral geometry. Alkenes: one or more C=C double bonds, CnH2n, sp2, trigonal planar. Alkynes: one or more C≡C triple bonds, CnH2n-2, sp, linear. |
| Functional Groups | Alcohol (-OH), aldehyde (-CHO), ketone (C=O internal), carboxylic acid (-COOH), ester (-COO-), amine (-NH2), amide (-CONH2), ether (R-O-R'), halide (R-X). Each group has characteristic reactivity and polarity. |
| Isomers | Structural (constitutional) isomers: different connectivity (n-butane vs. isobutane). Stereoisomers: same connectivity, different spatial arrangement. Cis-trans (E/Z) around double bonds. Enantiomers: non-superimposable mirror images of a chiral center. |
| Hybridization | Mixing of atomic orbitals to form equivalent hybrids. sp^3: 4 equivalent orbitals, tetrahedral (methane). sp^2: 3 equivalent orbitals plus an unhybridized p, trigonal planar (ethene). sp: 2 equivalent orbitals plus two unhybridized p, linear (ethyne). |
| Sigma vs. Pi Bonds | Sigma (σ): head-on overlap along the internuclear axis; every single bond is a σ. Pi (π): sideways overlap of unhybridized p orbitals. Double bonds = 1 σ + 1 π. Triple bonds = 1 σ + 2 π. π bonds restrict rotation. |
| Nuclear Reactions | Alpha decay: emits He-4 nucleus; mass -4, atomic number -2. Beta decay: neutron → proton + electron; atomic number +1. Gamma emission: high-energy photon, no change in mass or charge. Fusion: light nuclei combine (stars). Fission: heavy nuclei split (reactors). |
| Half-Life | Time for half of a radioactive sample to decay. First-order process: t(1/2) = ln2/k. Used in carbon-14 dating (t(1/2) = 5730 years) and medical imaging. After n half-lives, (1/2)^n of the original remains. |
| Common Constants to Memorize | R = 0.0821 L·atm/mol·K or 8.314 J/mol·K. F = 96,485 C/mol. Kw = 1.0 × 10^-14 at 25°C. NA = 6.022 × 10^23. Specific heat of water = 4.18 J/g°C. 1 atm = 760 mmHg = 101,325 Pa. Molar volume at STP = 22.4 L. AP exam provides a formula sheet but many are expected to be known. |
How to Study ap chemistry Effectively
Mastering AP Chemistry requires the right study approach, not just more hours. Research in cognitive science consistently shows that three techniques produce the best learning outcomes: active recall (testing yourself rather than re-reading), spaced repetition (reviewing at scientifically-optimized intervals), and interleaving (mixing related topics rather than studying one in isolation). FluentFlash is built around all three.
Active Recall and Spaced Repetition
When you study AP Chemistry with our FSRS algorithm, every term is scheduled for review at exactly the moment you are about to forget it. This maximizes retention while minimizing study time. The most common mistake students make is relying on passive review methods. Re-reading notes, highlighting textbook passages, or watching lecture videos feels productive, but studies show these produce only 10-20 percent of the retention that active recall achieves.
Flashcards force your brain to retrieve information, which strengthens memory pathways far more than recognition alone. Pair this with spaced repetition scheduling, and you can learn in 20 minutes a day what would take hours of passive review.
Your Practical Study Plan
Start by creating 15-25 flashcards covering the highest-priority concepts. Review them daily for the first week using our FSRS scheduling. As cards become easier, intervals automatically expand from minutes to days to weeks. You are always working on material at the edge of your knowledge.
After 2-3 weeks of consistent practice, AP Chemistry concepts become automatic rather than effortful to recall.
- Generate flashcards using FluentFlash AI or create them manually from your notes
- Study 15-20 new cards per day, plus scheduled reviews
- Use multiple study modes (flip, multiple choice, written) to strengthen recall
- Track your progress and identify weak topics for focused review
- Review consistently, daily practice beats marathon sessions
- 1
Generate flashcards using FluentFlash AI or create them manually from your notes
- 2
Study 15-20 new cards per day, plus scheduled reviews
- 3
Use multiple study modes (flip, multiple choice, written) to strengthen recall
- 4
Track your progress and identify weak topics for focused review
- 5
Review consistently, daily practice beats marathon sessions
Why Flashcards Work Better Than Other Study Methods for ap chemistry
Flashcards are one of the most research-backed study tools for any subject, including AP Chemistry. The reason comes down to how memory works. When you read a textbook passage, your brain stores that information in short-term memory, but without retrieval practice, it fades within hours.
The Testing Effect
Flashcards force retrieval, which transfers information from short-term to long-term memory. The "testing effect," documented in hundreds of peer-reviewed studies, shows that students who study with flashcards consistently outperform those who re-read by 30-60 percent on delayed tests. This advantage does not come from flashcards containing more information. It comes because retrieval strengthens neural pathways in a way that passive exposure cannot.
Every time you successfully recall an AP Chemistry concept from a flashcard, you make that concept easier to recall next time. This repeated retrieval builds stronger and longer-lasting memory traces.
FSRS-Based Learning Systems
FluentFlash amplifies this effect with the FSRS algorithm, a modern spaced repetition system that schedules reviews at mathematically-optimal intervals based on your actual performance. Cards you find easy get pushed further into the future. Cards you struggle with come back sooner. Over time, this builds remarkable retention with minimal time investment.
Students using FSRS-based systems typically retain 85-95 percent of material after 30 days. Compare this to roughly 20 percent retention from passive review alone. The difference is dramatic and compounds over weeks and months of study.