Complete General Chemistry 20+ hours of lectures & examples

Trace the origin of the atom from ancient Greece to Dalton's atomic theory, outlining atoms, elements, chemical reactions, conservation of mass, definite proportions, and the law of multiple proportions.

Explore isotopes as atoms of the same element with the same atomic number but different mass numbers from varying neutrons. Identify their similar chemical properties with a mass spectrometer.

Explain how modern atomic masses use carbon-12 as a standard, measured by mass spectrometry, and how natural carbon's isotopic mix yields an average atomic mass of about 12.1 amu.

Calculate average atomic mass from isotopes by multiplying each isotope’s abundance by its mass, then apply method to natural copper using isotopes 63 and 65 and a mass spectrometer.

Explore the mole concept and Avogadro's number, where carbon-12 defines one mole as 6.022×10^23 atoms in 12 grams, and relate grams to AMMU.

Compute the molar mass of calcium carbonate (CaCO3) by treating it as Ca2+ and CO3 2−, then find the mass for 4.86 moles and the CO3 2− mass.

Calculate the mass percent of each element in C10H14O by dividing each element’s mass by the molar mass and multiplying by 100, confirming sums to 100%.

Balance and compare baking soda and magnesium hydroxide as antacids, show reaction balancing, and demonstrate magnesium hydroxide neutralizes more acid per gram than sodium bicarbonate.

Identify the limiting reactant by comparing mole ratios from the balanced equation and checking for a stoichiometric mixture. In the ammonia example, hydrogen is limiting and nitrogen is in excess.

Determine the theoretical yield, the maximum product from the limiting reactant, and compute percent yield with the equation: percent yield = (actual yield / theoretical yield) times 100.

Calculate the mass of PbSO4 precipitate formed when mixing the two solutions, using the net ionic equation Pb2+ + SO4 2- → PbSO4(s) to identify the limiting reactant, yielding 15.2 g of PbSO4.

Explore redox reactions and how oxidation states track electron transfer in covalent bonds and energy production.

Master the assignment of oxidation states in neutral and charged species, ensuring sums equal zero or the overall charge, and use average oxidation numbers for complex molecules.

Practice identifying oxidation states and distinguishing oxidizing and reducing agents in example reactions for each case, clarifying which species are oxidized or reduced.

Balance redox reactions with the half-reaction method by separating oxidation and reduction, balancing atoms, oxygen, hydrogen, and charges, then equalizing electrons and combining the half-reactions, noting acidic or basic conditions.

Identify the net ionic reaction Al3+ + 3 OH- → Al(OH)3(s), determine the limiting reactant from the given volumes and molarities, and calculate the precipitated Al(OH)3 mass as 5.2 g.

Explain how gas expansion and compression perform work on surroundings using a piston, with W = P ΔV and P = F/A, and note expansion implies negative work.

Apply W = -P ΔV under pressure to calculate work for ΔV = 18 L at P = 15 atm, then relate Q and W to ΔE ≈ 8×10^7 J.

Enthalpy equals E + P V and is a state function. At constant pressure, enthalpy change equals heat transferred; note calorimetry, heat capacity, and extensive and intensive properties.

Apply Hess's law to convert graphite to diamond by combining graphite and diamond combustion reactions, yielding a net endothermic delta h of +2 kilojoules per mole.

Calculate the standard enthalpy change of the thermite reaction Fe2O3 + 2 Al → Al2O3 + 2 Fe from formation enthalpies, using balanced coefficients.

Explore how 19th-century views treated matter as particles and energy as waves, then show how Planck and Einstein revealed energy quantization and photons.

Compute the energy of a blue photon emitted from copper(I) chloride by converting a 450 nm wavelength to frequency and applying E = h nu, illustrating photon energy calculation.

Explore quantum numbers that define orbitals: the principal n and angular l. Identify how magnetic m_l orients orbitals and how n and l determine size, energy, and shape.

Explore orbitals as probability distributions and nodal surfaces, noting spherical s and two-lobed p orbitals, and how energy and spin differ in hydrogen and poly electronic atoms.

Explains bond energy across single, double, and triple bonds and how environment affects bond strength, using methane decomposition to illustrate reaction enthalpy from bond energies.

Use bond energies to calculate delta edge for the methane–chlorine–fluorine reaction forming Freon 12 and CO2, balancing and applying coefficients to count bonds broken and formed.

the lecture explains iso electronic ions with 36 electrons; selenium is the largest among the four ions, followed by bromide, rubidium plus, and strontium two plus.

Explore gas pressure and measurement, from the barometer and torricelli to mmHg, torr, atm, and pascals, and use a manometer to relate gas pressure to atmosphere.

Examine the gas laws, including Boyle's, Charles's, and Avogadro's, and derive the ideal gas law PV = NRT, outlining behavior at low pressure.

Apply ideal gas law and Boyle's law to calculate final pressure of ammonia gas when volume changes from 7 ml to 2.7 ml at constant temperature using P1V1 = P2V2.

Calculate molar mass from density, pressure, and temperature using M = DRT/P, shown with 1.95 g/L at 1.5 atm and 27 °C.

Explore the kinetic molecular theory of ideal gases, with negligible particle volume, motion, no intermolecular forces, and kinetic energy proportional to Kelvin temperature, linking to Boyle's law and Charles' law.

Determine the partial pressures of three gases in a 1.0 liter flask at 0 °C by converting data to moles and applying the ideal gas law pv = nrt.

Explore chemical kinetics by defining reaction rates, distinguishing consumption and production, and applying stoichiometric relationships in the decomposition of nitrogen dioxide.

Investigate how temperature boosts reaction rates using the Arrhenius equation, defining activation energy and the transition state, and deriving the relationship ln K = -Ea/R(1/T) + ln A.

Apply Le Chatelier's principle to predict how changes in concentration affect equilibrium, using the ammonia synthesis example to show how adding reactants shifts the system to the right.

determine the pH of a 1 M weak acid by solving Ka = [H+][F-]/[HF], set initial [HF] = 1 M, let x be the change, and use x ≈ sqrt(Ka).

Compute the pH of a 0.1 M hypochlorous acid solution using Ka = 3.5e-8, showing [H+] = [OCl-] ≈ 5.9e-5 and pH ≈ 4.23 due to weak acid dissociation.

Calculate percent dissociation by dividing the dissociated moles per liter by the initial concentration and multiplying by 100; for a 1.0 M solution, 0.027 M dissociates, giving 2.7%.

Calculate the percent dissociation of acetic acid using ka = 1.8×10^-5 for 1.0 m and 0.01 m solutions, showing increased dissociation with dilution to 0.42% and 1.3%.

Explore polyprotic acids such as sulfuric and phosphoric acids, which dissociate in steps to release multiple protons. Ka1 is greater than Ka2 and Ka3, so each successive proton is weaker.

Explore the Lewis acid-base model, where a Lewis acid accepts an electron pair and a Lewis base donates one, illustrated by ammonia and boron trifluoride forming adducts beyond Brønsted-Lowry acids.

Identify the lowest acid and base in each reaction, showing ammonia donates its lone pair as a Lewis base, while protons act as Lewis acids and water forms hydronium.

Demonstrate how adding 0.01 mole of base to a buffer changes pH by about 0.02, while the same addition to water shifts pH by roughly 5, highlighting buffer capacity.