General inorganic chemistry : 14H course + solved problems

Get a concise overview of general inorganic chemistry through a promotional video that highlights the course in a glance and introduces solved problems from the 14h course.

Learn foundational general inorganic chemistry concepts, including quantities and units, atomic structure, periodic table, formulas, balancing chemical equations, gas laws, thermodynamics, and electrochemistry.

Apply Hess's law to compute the total heat required for warming ice, melting, vaporizing, and heating steam from -10 to 120 degrees Celsius.

Learn to communicate scientific findings by performing and reporting measurements in standardized procedures, expressing quantities with numbers and units, and converting values between measurement systems.

Explore the international system of units (si), including base units like meter, kilogram, and second, and how prefixes and multipliers form compound units and use symbols and division conventions.

Define temperature as the property guiding heat flow, and compare the scales. Learn Kelvin, Celsius, and Fahrenheit definitions, absolute zero, triple point, normal freezing and boiling points, and scale conversions.

Learn to handle units and dimensions correctly, keep units with numbers, convert where needed, and apply multiplication and division rules to obtain valid results such as cubic metre.

Practice dimensional analysis and the unit factor method to solve problems by tracking unit changes, canceling units, and converting values such as 10 inches to 25.4 cm.

Solve unit conversion problems across length, volume, and mass, converting centimeters, inches, meters, and cubic meters, then compute cubic centimeters, liters, and density.

Dalton's atomic theory explains that atoms share protons, isotopes differ in neutrons, as carbon-12, carbon-13, and carbon-14 show, and the periodic table assigns their average atomic masses.

The nucleus holds most of an atom’s mass, built from protons and neutrons; protons carry positive charge, and isotopes differ in neutron number described by Z and A.

Explore the periodic table’s structure, including groups and periods, and labeling schemes. Learn how outer electrons determine main-group numbers and identify alkali metals, halogens, and noble gases.

Define the atomic mass unit as exactly 1/12 the mass of a carbon-12 atom and explain how average atomic mass arises from isotopic compositions through weight percentages.

Defines the mole as a fixed number of entities and links it to Avogadro's number using carbon-12, then explains converting mass to moles with molar mass.

Explore symbols and formulas for elements and compounds, learn how to read molecular and empirical formulas, and calculate molar masses from atomic masses, using CO2 and Avogadro's number as references.

Explore atomic, molecular, and molar masses through solved problems. Calculate silicon and carbon isotopic abundances, and determine molar masses for cadmium and hydrogen sulfide, using isotope data.

Compute the empirical formula from percent composition by converting to moles, dividing by the smallest mole value to obtain the lowest whole-number ratio, yielding MgAl2O4 in the example.

Explore how a chemical formula reveals fixed mass ratios between elements, compute masses from atomic masses, and determine the fractional weights and percent composition of a compound.

Compute nuclidic molecular masses by summing average atomic masses, using isotopic abundances; a mass spectrometer separates fragments to help deduce exact molecular formulas, illustrated with CO2.

Develop methods to derive molecular and empirical formulas from percent composition and molar masses, including hydrates and combustion data, using mass, moles, and ratio analysis.

Learn to classify compounds and correctly name binary nonmetal compounds and hydrates, writing accurate formulas to prevent dangerous errors from naming differences.

Learn to name binary nonmetal compounds by placing the left or below element first, adding -ide to the second, and using prefixes like mono and di to show atom counts.

Learn how to name ionic compounds by identifying the cation first, then the anion, following a fixed naming order that does not depend on the compound’s nature.

Name monoatomic cations by using roman numerals for metals with multiple charges, such as copper(II) and cobalt(III). Metals with fixed charges require no numeral, like magnesium or zinc.

Learn how to name polyatomic cations by their charge, using -ic for higher charge and -ous for lower charge, with Latin names and ferric/ferrous iron examples.

Discover how to name monoatomic anions by applying the group number minus eight rule, with examples such as hydroxide, cyanide, sulfide, nitride, and chloride.

Learn to name oxyanions by applying prefixes and suffixes—hypo-, -ite, -ate, and per-—as shown with chlorine oxoanions such as hypochlorite, chlorite, chlorate, and perchlorate.

Explore the naming of special anions, focusing on chromate and permanganate, their analogous formulas, and how prefixes indicate one more oxygen in these oxygen-containing anions.

Learn how to name ionic compounds by pairing a cation first with an anion, balancing charges to form a neutral compound, with examples calcium nitrate and chromium(III) chlorate.

Balance ionic formulas by using known charges to achieve zero net charge, memorizing ion charges and writing the smallest whole-number ratios, with examples like NaCl, AlCl3, (NH4)2SO4, and Al2O3.

Explore how acid salts form by retaining hydrogen atoms and balancing charges to match the parent acid, using dihydrogen phosphate and bicarbonate naming conventions (hydrogen prefixes and bi-).

Learn how hydrates are named by pairing a stable ionic compound with the exact number of water molecules, illustrated by a hydrate containing five water molecules.

Learn to write and balance chemical equations with coefficients, guided by the law of conservation of matter that mass is conserved. Deduce molecular and mass relationships between reactants and products.

Balance chemical equations by adjusting coefficients to conserve atoms, read reactions as sentences, and use stoichiometry to predict ammonia and oxygen forming nitrogen and water.

Demonstrates mass relations and the conservation of mass in chemical equations, showing how molar masses, molecular formulas, and polymers align across reactions with ammonia, oxygen, nitrogen, and water.

Identify the limiting reactant by comparing the amounts of reactants, and predict product amounts based on the limiting one while noting excess reactants remain unused.

Balance chemical equations with a clear action plan: balance from left to right, apply consistent coefficients, check each element, and avoid fractions by doubling when needed.

Identify and balance combustion, single replacement, and double displacement (metathesis) reactions. Explore acid-base reactions producing water and salt, precipitation, and composition reactions transformed by electricity or heat.

Explore how gas volumes change with temperature and pressure, and define pressure as force per area in newtons, square meters, and pascals, linking to standard atmospheres and millimeters of mercury.

Explore the gas laws, including Boyle's, Charles's, and Gay-Lussac's, and learn how volume, temperature, and pressure govern gas behavior, while noting deviations due to intermolecular forces.

Explore Boyle's law, showing the inverse relationship between pressure and volume at constant temperature, and apply the relation P1V1 = P2V2 to connect initial and final conditions.

Explore Charles' law, where at constant pressure the gas volume varies directly with temperature in kelvin. Learn the proportional relationship V1/T1 = V2/T2 and how kelvin temperature governs gas behavior.

Explore Gay-Lussac's law at constant volume, showing how gas pressure varies directly with temperature in Kelvin and applying P1/T1 = P2/T2 to relate initial and final conditions.

Derive the combined gas law from the individual gas laws, relating pressure, volume, and temperature for a fixed amount of gas. Use kelvin for temperature, with P1V1/T1 = P2V2/T2.

Explore how the density of an ideal gas varies inversely with volume while mass stays constant, and derive volume relations under changing temperature and pressure via the combined gas law.

Explore Dalton's law for partial pressures, showing how the total pressure of a four-gas mixture equals the sum of individual partial pressures, with emphasis on ideal gas behavior.

Collect a gas over a liquid by subtracting the liquid's vapor pressure from the total pressure to get the dry gas, accounting for water vapor or other volatile liquids.

Explore how real gases deviate from ideal behavior; compression and low temperatures can liquefy a gas, signaling not ideal behavior, while predictions are accurate at low pressure and high temperature.

Apply pressure calculations for water and mercury using P = ρ g h, then use Boyle’s, Gay‑Lussac’s, and Dalton’s laws to solve gas problems and partial pressures.

Avogadro's hypothesis states that equal volumes of gas at the same temperature and pressure contain the same number of molecules, enabling relative molar masses to be inferred from volumes.

Learn about molar volume under standard conditions, where one mole of any gas occupies about 22.4 liters, assuming ideal gas behavior, with real gases not always ideal.

derive the ideal gas law from standard conditions, linking pressure, volume, temperature, and moles via PV = nRT with standard molar volume of 22.4 L per mole and 1 atm.

Explain how balanced chemical equations link reactant and product volumes under standard temperature and pressure, using the ammonia–oxygen system and molar volume 22.4 L.

Convert gas volume to mass via the ideal gas law, using volume, temperature, and pressure, then apply balanced stoichiometry to compute the lithium hydroxide mass needed to remove CO2.

Apply the ideal gas law to solve problems at standard temperature and pressure, calculating molar masses, gas volumes, and reaction volumes using Avogadro's hypothesis.

Explore how energy in all forms transforms with no net loss during heat exchanges, and how heat capacity links temperature change to energy transfer via calorimetry.

Explore how energy partitions into work and heat during chemical processes, and how internal energy and enthalpy relate to heat at constant pressure in open versus closed systems.

Calculate and interpret enthalpy changes for heating, cooling, and phase transitions using specific heat, latent heats of fusion, vaporization, sublimation, and standard heats of formation.

Explore how enthalpy depends only on initial and final states, and apply Hess's law to sum fusion, vaporization, sublimation, and formation enthalpies for reactions.

Solve thermochemistry problems through heat transfer calculations, using heat capacity, temperature changes, phase changes, and Hess's law to analyze copper, water, coal, steam, and calcium carbonate.