AP + College General Chemistry | Full Course & Exam Practice

Explore subatomic particles, atomic number, mass number, and key atomic models—Thomson's plum pudding model, Millikan oil drop, and Rutherford's nuclear model—to understand how protons, neutrons, and electrons define an atom.

Explain isotopes as same-element atoms with different neutrons, and compute average atomic mass as a weighted average of isotope masses using decimal abundances, with magnesium and carbon examples.

Master the mole, molar mass, and conversions between grams, moles, and particles using Avogadro's number and the formula moles = mass over molar mass.

Master percent composition, empirical formulas, and molecular formulas by converting percent data to grams and moles, then scaling to final formulas for AP chemistry success.

Explore the electronic structure of atoms, including shells, sublevels, and orbitals, and learn how electron configurations follow Aufbau, Pauli, and Hund's rules, noble gas notation, to explain trends and bonding.

Explore how photoelectron spectroscopy uses high-energy photons to eject electrons, measure binding energies, and reveal electron configurations through PES graphs showing peak heights and electron counts.

Understand how effective nuclear charge drives periodic trends in atomic radius, ionization energy, and electronegativity to predict element behavior.

Compare ionic, covalent, and metallic bonding, explaining electron transfer, sharing, and delocalization, and relate their structures, properties, and conductivity to metals, nonmetals, and periodic table positions for AP exam prep.

Learn how ions form by gaining or losing electrons to reach noble gas configurations, making metals cations and nonmetals anions, and predict charges from the periodic table for ionic bonding.

Balance charges to form neutral ionic compounds, using the criss cross method and polyatomic ions, then name them, including transition metals with roman numerals.

Draw Lewis structures for covalent molecules by counting valence electrons, building a skeleton around the central atom, and distributing lone pairs to satisfy the octet with single or multiple bonds.

Explore resonance and formal charge to evaluate multiple Lewis structures and form the resonance hybrid, using the nitrate ion NO3- as a guiding example and ranking formal charges.

Explore bond polarity and molecular polarity using electronegativity, dipole moments, and VSEPR geometry. Draw Lewis structures, assess symmetry, and predict how polarity affects interactions and properties in AP chemistry.

Explore hybridization and the roles of sigma and pi bonds in shaping molecular geometry, with sp, sp2, and sp3 examples, and how bond types relate to structure and reactivity.

Explore how solid structure at the particle level determines melting point, hardness, conductivity, and solubility. Identify solid types: ionic, metallic, molecular, and network covalent and predict properties by interparticle forces.

Explore intermolecular forces (imfs) and their impact on boiling points, melting points, solubility, viscosity, and state of matter, highlighting London dispersion forces, dipole-dipole interactions, and hydrogen bonding.

Explore how intermolecular and interparticle forces—from London dispersion forces and dipole–dipole interactions to hydrogen bonding—shape boiling points, phase changes, and solubility.

Visualize how potential energy curves illustrate attraction and repulsion between particles, revealing the equilibrium distance and how intermolecular forces shape bonding, phase changes, and properties like boiling and melting points.

Explore how intermolecular forces (IMFs) shape the three states of matter—solids, liquids, and gases—linking particle behavior to properties like shape, volume, compressibility, diffusion, and boiling points.

Linking particle motion to gas behavior, the ideal gas law PV=nRT connects pressure, volume, temperature, and moles under KMT assumptions; heat increases pressure, while deviations later explore non-ideal behavior.

Examine why real gases deviate from the ideal gas law PV=NRT by particle volume and intermolecular forces, especially at high pressure or low temperature, and predict deviations.

Explore solutions as homogeneous mixtures, including how solvents dissolve solutes, the difference between ionic and molecular solutes, and how polarity governs solubility and conductivity.

Learn to describe and calculate concentration using molarity for the AP exam, perform dilution calculations with m1v1 = m2v2, and apply molarity to stoichiometry and titrations.

Learn how solubility defines the maximum amount dissolved at a given temperature and pressure, and how temperature affects solids and gases for AP exam, including unsaturated, saturated, and supersaturated solutions.

Identify separation techniques that isolate mixture components without breaking bonds, using boiling point, polarity, and solubility, with filtration, distillation, chromatography, and evaporation.

Explore spectroscopy's role in analyzing solutions by measuring light absorption across the electromagnetic spectrum, especially UV-visible spectroscopy, to identify substances and determine concentration via Beer-Lambert law.

Explain how chemical reactions rearrange atoms by breaking and forming bonds, conserve mass, and form substances, with color changes, gas evolution, a precipitate, and distinguish chemical from physical changes.

Classify and predict products for the five main reaction types—synthesis, decomposition, single replacement, double replacement, and combustion—and relate acid-base, precipitation, gas evolution, and redox processes.

Write net ionic equations by eliminating spectator ions to reveal changing species, using three steps: balance, break strong electrolytes into ions, and cancel spectators, for precipitation, acid-base, and redox reactions.

Study acid-base reactions and neutralization, focusing on proton transfer and salt formation. Learn strong versus weak acids and bases, write molecular, ionic, and net ionic equations, and understand titration.

Explore redox reactions by identifying oxidation numbers, balancing half-reactions in acidic or basic solutions, and writing net ionic equations to trace electron transfer and distinguish oxidizing and reducing agents.

Balance ordinary chemical reactions using coefficients, not subscripts, to honor the law of conservation of mass, with step-by-step examples and practice toward AP-style stoichiometry.

Master stoichiometry by using a balanced equation to relate amounts through mole ratios. Convert grams to moles with molar mass and predict product amounts for AP chemistry.

Identify the limiting reactant using a balanced equation and mole ratios, then calculate percent yield from actual versus theoretical yields, with stoichiometry guiding the process.

Define reaction rates as the change in concentration over time, with reactants decreasing and products forming. Identify how concentration, temperature, surface area, catalysts, and nature of reactants influence rate.

learn to express reaction speed with rate laws and concentrations, recognize zero, first, and second orders, and determine overall order from experimental data.

Determine rate laws from experimental data by analyzing data tables and trials, holding one reactant constant to reveal zero, first, or second order, and deriving the rate law.

Examine concentration versus time graphs to identify reaction order by curve shape and linearity, recognizing zero, first, and second order plots, their slopes, and the integrated rate law.

Explore how half-life governs first-order decay, with t1/2 = 0.693/k, independent of concentration, and use integrated rate laws and ln plots to predict remaining reactant.

Collision theory explains when reactions occur via energy, orientation, and collisions, affected by temperature, concentration, and surface area; reaction mechanisms break reactions into steps with intermediates, catalysts, and rate-determining step.

Explore how catalysts speed up chemical reactions by providing a lower activation energy and a new, multistep pathway, without changing products, enthalpy, or equilibrium, illustrated by energy diagrams.

Explore endothermic and exothermic reactions, where energy flows between system and surroundings, with enthalpy changes and the first law of thermodynamics illustrated by methane combustion and ice melting.

Analyze how heat, temperature, and heat capacity relate, using q=mcΔT to quantify energy changes, and distinguish specific heat capacity from heat capacity in endothermic and exothermic processes.

Explore the enthalpy of reactions and how delta H indicates heat exchange, exothermic or endothermic processes, using standard states, thermochemical equations, and calorimetry to quantify energy changes.

Hess's law lets you determine a reaction's enthalpy by combining known steps, flipping or scaling reactions, and canceling terms, since enthalpy is a state function and path independent.

Learn to estimate reaction enthalpies using bond enthalpies by adding bonds broken and subtracting bonds formed, with endothermic and exothermic energy flow in kilojoules per mole.

Calculate reaction enthalpy using standard enthalpies of formation by summing product and reactant values per mole and applying the products-minus-reactants method, with elements in standard states having zero enthalpy.

Explore reversible reactions and dynamic equilibrium, where forward and reverse rates balance in a closed system, keeping concentrations constant and guiding optimization in lab and industry.

Learn to write equilibrium constant expressions using Kc for aqueous and gaseous species, interpret k values to predict whether a reaction favors products or reactants, and note temperature effects.

Explore how reaction quotient Q, a snapshot, compares with equilibrium constant K to predict direction toward equilibrium, with Q<K shifting right, Q>K left, Q=K at equilibrium, Le Chatelier's principle.

Le Chatelier's principle explains how chemical systems at equilibrium respond to disturbances by shifting to counteract changes in concentration, temperature, or pressure, forming a new equilibrium.

Apply the equilibrium constant k to calculate concentrations at equilibrium and to determine k from known data, using ice tables and the k expression for aqueous and gaseous species.

Manipulate the equilibrium constant k when reversing, multiplying, or adding reactions using Hess's law logic. Reversing yields the reciprocal; multiplying raises k to a power; adding reactions multiplies k values.

Explore solubility equilibrium for slightly soluble salts, apply KSP and Le Chatelier's principle to the common-ion effect, and determine precipitation using ICE tables.

Predict precipitate formation by comparing the reaction quotient Q to KSP, accounting for dilution. Apply unsaturated, saturated, and supersaturated rules with a PbI2 example.

Learn Arrhenius and Bronsted-Lowry acid-base definitions, identify conjugate pairs and amphoteric species, and apply proton transfer to predict equilibria and prep for pH and pOH concepts.

Calculate the pH and pOH for strong acids and bases that fully dissociate in water. Use initial concentration to find [H+] or [OH-] and apply the pH and pOH formulas.

Distinguish strong and weak acids and bases by ionization in water, equilibrium behavior, and conductivity, and apply Ka, Kb, and ICE for pH and titration problems.

Explore weak acids and bases in water, analyze equilibrium with ICE tables, compare acid strength via Ka and base strength via Kb, and calculate pH or pOH using HCN.

Explore how Ka, Kb, and Kw relate conjugate acid-base strengths, enabling prediction of acid-base behavior and solving equilibrium problems, illustrated by NH4+/NH3 using Kw equals Ka times Kb.

Learn to calculate the pH of weak acids, bases, and buffers using equilibrium principles and the Henderson–Hasselbalch equation, including buffer capacity and practical examples.

Explore how salts dissociate in water and hydrolyze to determine the solution’s pH, depending on the parent acid and base, with practical examples.

Explore how acid strength arises from molecular structure, including bond polarity and strength, electronegativity, resonance, and oxygen content in oxyacids.

Learn how titrations determine acid or base concentrations, how buffer regions stabilize pH, and how pH curves reveal equivalence points and pKa values.

Explore how Gibbs free energy determines spontaneity by linking enthalpy and entropy through temperature. Learn to predict spontaneity using delta G, delta H, delta S, T, and standard free energy changes.

Explore galvanic (voltaic) and electrolytic cells, showing how redox reactions convert chemical energy to electrical energy and back, with anode, cathode, salt bridge, electrodes, electrolytes, and external wire.

See how electrochemical cell potential relates to Gibbs free energy and spontaneity. Calculate standard cell potential and delta G to predict spontaneous redox reactions.

Explore how concentration and temperature shift non-standard cell potentials in electrochemical cells, contrasting standard versus non-standard conditions and examining the role of the reaction quotient Q and Le Chatelier's principle.

Explore Faraday's law in electrochemistry by calculating total charge, moles of electrons, and moles of substance formed from current and time, with worked examples and concentration cells introduced.

Explore concentration cells, a galvanic setup with identical electrodes at different ion concentrations, where electrons flow from the dilute to the concentrated side as they reach equilibrium.

Explore how non-spontaneous redox reactions are driven by external energy in electrolysis, using electrolytic cells to couple energy to chemical change in applications like electroplating and battery charging.