The Complete High School and College Chemistry

Explore chemistry as a branch of science studying the composition, properties, and structure of matter, and apply scientific methodology from observation to theory.

Explore how chemistry opens careers in medicine, pharmacy, teaching, geology, mining, petrochemicals, manufacturing, research, and international health organizations.

Learn how to separate mixtures using filtration, decantation, and sieving, then apply simple and fractional distillation to differences in boiling points.

Explore evaporation, sublimation, precipitation, magnetic separation, and chromatography as methods to recover solids from solutions. Learn how heat, freezing, and solvent movement separate mixtures like salt, sugar, and dyes.

Explore crystallisation as a purification technique that uses heating and cooling of a supersaturated solution to form crystals, and compare it with evaporation, fractional crystallisation, flotation, and separating funnel methods.

Balance equations to derive mole ratios between reactants and products, and use coefficients to predict how reactants convert to products in reactions such as ammonia formation and sulfur dioxide production.

Explore the mole concept with Avogadro's number, linking mass, moles, and particle count to concentration, and examine valency, radicals, and exchange to form compounds like sodium chloride and aluminium chloride.

Balance equations by counting atoms on both sides and applying coefficients, not changing subscripts; examples include h2 + cl2 → 2 hcl.

Balance chemical equations by placing coefficients in front of reactants and products to equalize carbon, hydrogen, and oxygen on both sides, using trial-and-error, and never adjusting subscripts.

Explore various reaction types in chemistry, including combination, decomposition, displacement, and redox reactions, and learn how the electrochemical series governs which elements displace others.

Explore double decomposition reactions where salts exchange ions to form a precipitate, with examples like silver nitrate and sodium chloride, and examine homogeneous and heterogeneous catalysis and catalyst properties.

Explain reversible reactions with a double arrow and redox reactions, including oxidation and reduction, oxidising and reducing agents, and dissociation driven by heat with ammonium chloride and nitrogen oxide examples.

Explore the parametric relationships among volume, mass, concentration, and particles using Avogadro's number, and learn to calculate molar mass (grams per mole) from atomic masses.

this lecture teaches calculating masses in the lead nitrate and sodium chloride reaction, using balanced equations and molar masses to show 10.7 g lead nitrate yields 9 g lead chloride.

Balance the chemical equation, identify the reactant and carbon dioxide, then use molar masses and a cross-multiplication to calculate the mass of CO2 produced from a given reactant mass.

Master stoichiometry by balancing equations and converting masses to moles to calculate sodium hydroxide needed for 100 g of HCl, and calcium chloride from limestone.

Convert grams to moles using molar mass, as 5.3 g of sodium carbonate yields about 0.05 mol, and use stoichiometry to predict ethanol from glucose.

Balance the reaction of S02 with oxygen to form S03, then perform stoichiometric mass calculations using cross-multiplication, and apply these steps to a copper nitrate problem.

Master mass–volume–mole calculations to determine the volume of dry oxygen produced from decomposition, balancing the equation and using 22.4 L per mole.

Cross-multiply to obtain 0.96 g of oxygen from 2.45 g of potassium chloride, then convert to ~0.03 mol O2 and a stp volume of about 0.672 L.

Compute the mass of a reactant to yield a specific chlorine gas volume using a balanced equation; convert volume to moles (22.4 L/mol) and apply stoichiometry.

Explore limiting reagents and excess reagents in a reaction, and compute percent yield by comparing actual yield to the theoretical yield using a potassium chloride decomposition example.

Learn to identify limiting reagents and compute theoretical yields from balanced equations using molar masses, then compare with actual yields to assess percent yield in chemistry reactions.

Tackle chemistry test problems on methanol synthesis and theoretical yield, compute percent yield for salicylic acid to aspirin with acetic acid byproduct, and survey past MCQs for study.

Learn to distinguish empirical formulas from molecular formulas and calculate empirical formulas from percent composition using atomic masses and relative molecular mass, with step-by-step examples.

Compute empirical formula from percent composition by dividing masses by atomic masses and normalizing to the smallest value, then derive molecular formula from the empirical formula and molar mass.

Calculate the percentage composition by mass of a compound by dividing each element's mass by the total mass and multiplying by 100, with examples like water and ammonium sulfate.

Demonstrate the law of conservation of mass and the law of definite proportion, showing mass before and after reactions remains equal and elements combine in fixed ratios.

Demonstrate the law of definite proportion through copper oxide and calcium oxide experiments, showing constant mass ratios and similar percentage composition across methods.

The lecture explains the law of multiple proportions: when two elements form several compounds, the masses of one element combine with a fixed mass of the other in simple ratios.

Demonstrate the law of reciprocal proportion by showing fixed-mass elements form simple multiples, as carbon with hydrogen and oxygen, and in water and oxides.

Explore volumetric analysis as a method to measure volumes to determine the exact volume and concentration needed to react completely. Identify burette, pipette, conical flasks, and end-point indicators.

Explore hydration in volumetric analysis to determine exact reaction volumes, prepare 0.1 M and 1.0 M standard and molar solutions using molar mass, and identify endpoints with indicators.

Learn how indicators change color with pH to guide titration in volumetric analysis, with examples like phenolphthalein, and master eye-level readings to avoid parallax and air bubbles.

Apply c1v1 = c2v2 to calculate new acid concentrations during dilution, exemplified by 50 ml of 2 M HCl to 250 ml (c2 = 0.4 M).

Master titration calculations by using volume and concentration to determine the amount of solute, compare molarity and mass concentration, and solve acid-base reaction problems step by step.

Learn to perform titration calculations using a balanced equation, endpoint volumes, and the m1v1=m2v2 relationship to determine the concentration of sodium hydroxide and its mass concentration.

Discover the four states of matter—solids, liquids, gases, and plasma—and how temperature, kinetic energy, and particle attractions drive phase changes like melting, evaporation, boiling, condensation, sublimation, and deposition.

Explain latent heat during phase changes, including fusion and vaporization, and illustrate ice melting at 0 °C and water boiling at 100 °C, while linking Brownian motion, diffusion, osmosis, and kinetic theory.

The kinetic theory of gases treats ideal gases as molecules moving in straight lines with perfectly elastic collisions, negligible molecular volume, and negligible intermolecular forces, where temperature reflects kinetic energy.

Explore the periodic table, arranged by increasing atomic numbers, its history from Mendeleev to the modern periodic law, and how groups, periods, and outermost shells of electrons shape element properties.

Explain group characteristics and block elements, focusing on s-block elements and how to write electron configurations using the s, p, d, and f orbitals. Use magnesium, calcium, potassium as examples.

Explore electron configurations of p-block elements such as nitrogen, phosphorus, and arsenic, and learn how transition metals fit into the d-block with the diagonal arrow method.

Learn to write electron configurations for lanthanides and actinides using noble gas shorthand, long and short forms, and the diagonal-arrow method for building orbitals.

Explore seven periods and groups in the periodic table, learning how each period shares electron shells and how groups share valence electrons across s, p, d, and f blocks.

Atomic size increases down the group and decreases across the period; the across-period decrease arises from more electrons being added and the increasing nuclear charge attracting electrons.

Explain how atomic size increases down a group as more electron shells are added, while it decreases across a period due to stronger nuclear attraction.

Explore how losing electrons forms smaller cations with contracted electron clouds and smaller ionic size, and how gaining electrons expands the cloud to create larger anions such as Cl-.

Explore how melting and boiling points rise across a period with increasing metallic bond strength, and observe conductivity decreasing across a period but increasing down a group.

Explore ionisation energy, the energy to remove an electron, and how it rises across a period and falls down a group due to effective nuclear charge, screening, and electron configuration.

Explore how ionisation energy rises across a period due to increasing nuclear charge and shrinking atomic size, and how electron affinity and electronegativity vary across periods and down groups.

Define oxygen as a gaseous element in the oxygen family, note its abundance in nature and human body, and outline production methods from potassium chloride and hydrogen peroxide, fractional distillation.

Summarize oxygen’s physical properties, covalent bonding with nonmetals, reaction with metals to form basic oxides, and key uses from welding flames to space life support.

Explore acidic, basic, neutral, higher oxides, and peroxides, including hydrogen peroxide, their reactions with acids and bases, and laboratory and industrial preparation methods.

Examine the properties of hydrogen peroxide, including its nearly colorless liquid form and its boiling and freezing points, and its uses as a bleaching agent, antiseptic, and environmental treatment.

Explore hydrogen as the lightest, colorless gas found in the atmosphere and sun, its dual nature, flammability, presence in water and hydrocarbons, and laboratory and industrial production methods.

Explore methane steam reforming to produce synthesis gas of hydrogen and carbon dioxide with iron oxide or chromium oxide catalysts, and compare bus process and electrolysis methods for hydrogen production.

Examine hydrogen’s properties and chemistry, from colorless, odorless, and the lightest element to covalent and coordinate bonds, halogen reactions, and flame tests.

Explore the physical properties of hydrogen and practice questions on its uses, reactions with oxygen to form water, and isotopes.