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 the discovery of the electron through cathode rays and electrolysis, tracing Faraday and Stoney's charge concepts to Thompson's charge-to-mass ratio and Millikan's electron charge measurement.

Proton, the main particle of atoms, carries a positive charge equal to the electron’s; removing electrons forms ions, while Goldstein and Thomson explored the q/m ratio and proton mass.

Explore how Rutherford proposed a neutral particle in the nucleus to account for atomic mass, and how Chadwick later measured the neutron's mass from nuclear reactions.

Investigate the atom's nucleus, radioactivity and alpha, beta, and gamma radiation, and compare Rutherford's nuclear model with Thomson's plum pudding.

Identify atomic symbols with number Z and mass number A, compute neutrons as A minus Z, note electrons equal protons in neutral atoms, and charges on ions.

Explore the periodic table’s structure, groups, and periods, highlight noble gases and alkaline metals, and explain periodic law, lanthanides, actinides, and metal-nonmetal trends.

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.

Convert six Americium atoms (243 amu per atom) to grams and then to moles, yielding 2.42×10^-21 g, and calculate 0.371 moles and 2.23×10^23 Aluminium atoms from a 10 g sample.

Convert chip mass from mg to g, apply silicon atomic mass and Avogadro's number to find silicon atoms; then use cobalt's atomic mass to convert atom counts to moles.

Calculate molar mass by summing atomic masses to obtain grams per mole, as shown with methane and the C10H6O3 example, yielding 174.1 g/mol and 8.96×10^-5 mol.

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.

Compute mass percent by dividing each element’s mass in one mole by the molar mass, illustrated with C2H6O: carbon about 52.4%, hydrogen about 13%, oxygen about 34.2%.

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%.

Learn how to determine an empirical formula from combustion data by calculating mass percentages, converting to moles using molar masses, and comparing ratios to relate it to the molecular formula.

Practice determining empirical and molecular formulas from mass percent data by converting to moles with proper significant figures, then using molar mass to confirm the formula C2H4Cl2.

Determine empirical and molecular formulas from percent composition and molar mass, using a phosphorus and oxygen example. Conclude empirical formula P2O5 and molecular formula P4O10.

Determine caffeine's molecular formula from its mass percentages and molar mass, convert grams per mole to moles, and derive the formula as C8H10N4O2.

Balance chemical equations by conserving atoms and using coefficients to match reactants and products; start with the most complex molecule and balance by inspection (ethanol combustion to CO2 and H2O).

Balance chemical equations by prioritizing hydrogen, then nitrogen, then oxygen; use fractional coefficients if needed and clear them by multiplying to whole numbers, as shown with ammonia and CH4 reactions.

Apply stoichiometric calculations to balance equations, convert masses to moles, and compute reactant and product masses—like determining oxygen needed and CO2 produced from propane.

Balance the LiOH and CO2 reaction, then calculate the CO2 absorbed from one kilogram of LiOH, noting that it forms lithium carbonate and water, yielding about 0.92 kilograms of CO2.

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.

Balance the reaction, identify the limiting reactant, and calculate the nitrogen gas yield using mole ratios and mass-to-mole conversions.

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.

Balance the methanol synthesis equation from carbon monoxide and hydrogen, identify the limiting reagent, and determine the theoretical yield. Convert the actual yield to grams and compute the percent yield.

Determine the empirical formula CH5N and the molecular formula C3H15N3 for a carbon-hydrogen-nitrogen compound from combustion data using CO2 and H2O masses and a 93.18 g/mol molar mass.

Balance the calcium oxide and carbon reaction to compute the theoretical yield of calcium carbide, identify the limiting reactant, and calculate the percent yield from given masses.

Explore water’s polarity and hydration, which dissolve ionic solids into hydrated cations and anions, and explain solubility differences among salts and nonpolar substances.

Explore electrical conductivity in aqueous solutions by linking ionization to current flow and classifying electrolytes as strong, weak, or non-electrolytes. See examples of soluble salts, strong acids, and strong bases.

Explore how weak electrolytes ionize only slightly in water, exemplified by acetic acid and ammonia, and contrast with non-electrolytes like ethanol and sugar that do not conduct electricity.

Calculate molarity by converting grams of solute to moles using molar mass and dividing by the solution volume in liters (moles per liter).

Learn how to determine the concentration of each iron species in ionic solutions by balancing equations, applying molarity, and using coefficients to relate moles of compounds to ions.

calculate moles from volume and molarity in general chemistry, balance equations to find coefficients, and determine solution volume from mass using molarity, with NaCl and blood serum examples.

Weigh precise solid potassium dichromate, dissolve it to prepare a standard solution, and calculate moles and molarity; apply dilution using m1v1 = m2v2 to reach a target volume.

Explore precipitation reactions by forming insoluble solids from mixing aqueous solutions, using ion exchange to predict solid products and remaining ions.

Practice predicting outcomes of mixing electrolyte solutions by writing separate ions, performing ion interchange, and identifying soluble versus insoluble products to balance precipitation reactions.

Describe the three reaction descriptions in solution: formula equation, complete ionic equation, and net ionic equation. Distinguish the overall stoichiometry from the actual aqueous ion forms and identify spectator ions.

Apply precipitation stoichiometry to form AgCl, using net ionic equations and spectator ion concepts; compute moles from 1.5 L of 0.01 M AgNO3 and convert to NaCl mass (58.44 g/mol).

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.

Examine acid-base reactions via Brønsted–Lowry definitions, proton transfer, and net ionic equations. Learn about water formation, spectator ions, and neutralisation when base matches acid.

Calculate the needed volume of 0.01 M FCl to neutralize 25 mL of 0.35 M acid. Use a 1:1 mole ratio and derive the net ionic equation for water formation.

Compute moles from the given volumes and concentrations, identify the limiting reactant, determine the water formed, and calculate the concentration of excess OH− ions.

Conduct a titration by delivering a measured volume of a known concentration solution to reach the equivalence point, with the endpoint indicated by phenolphthalein's color change from colorless to pink.

Calculate the NaOH concentration by titrating 1.309 g of an acid with one acidic hydrogen (204.22 g/mol) using 41.20 ml NaOH to the phenolphthalein endpoint.

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.

Explore oxidation and reduction through electron transfer and oxidation numbers, identifying oxidizing and reducing agents, with methane combustion and a sodium chloride example.

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 equations using the oxidation numbers method or the half-reaction approach. Assign oxidation states to copper and silver to track electron transfer and balance atoms and charge.

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.

Balance a redox equation in a basic solution using the half-reaction method, assign oxidation states, balance atoms, add oh minus ions, and verify the final equation.

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.

Balance a redox reaction in basic solution using the half-reaction method, identify oxidized and reduced species, balance with water and hydroxide, and obtain coefficients 1 and 1.

Explore thermochemistry by defining energy, heat, and work, apply the first law of thermodynamics and state functions, and distinguish exothermic from endothermic heat transfer.

Calculate ΔE for an endothermic process with Q = 15.6 kJ and W = +1.4 kJ, yielding ΔE = 17.0 kJ.

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.

Mix equal volumes of 1 molar Ba(NO3)2 and Na2SO4 in a calorimeter; calculate the enthalpy change per mole of BaSO4 formed from the temperature rise and the specific heat, exothermic.

Compare the energy of combustion of hydrogen and methane using a calorimeter with a 11.3 kJ/°C heat capacity. Hydrogen yields about 141 kJ/g and methane about 55 kJ/g, roughly 2.5 times greater.

Explore hess's law by showing that enthalpy is a state function, so the overall enthalpy equals the sum of steps, whether one step or multiple.

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.

Apply Hess's law to synthesize B2H6 from elements using a sequence of reactions A–D. Determine the net enthalpy change, which is +36 kJ.

Explain how standard enthalpy of formation defines the enthalpy change to form one mole of a compound from elements in their standard states, using Hess's law and formation data.

Calculate standard enthalpy change using formation enthalpies for combustion of ammonia in air to form nitrogen dioxide and water by balancing and applying multipliers, yielding minus 1396 kJ.

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

Compare the standard enthalpy of combustion of methanol and octane by balancing equations and calculating per-gram energy, noting gasoline’s higher energy per gram and methanol’s smoother burn.

Examine energy sources such as petroleum, natural gas, and coal, and their environmental effects, including sulfur dioxide and carbon dioxide emissions, plus the greenhouse effect guiding future energy needs.

Explore future energy sources such as the sun, solar energy, hydrogen, fusion, biomass, coal gasification, and ethanol, while examining fossil fuel conservation and the challenges of production, storage, and transport.

Apply Hess's law to build the target reaction from A, B, and C by halving A and C, reversing B, and canceling intermediates; derive ΔH for the desired reaction.

Determine the standard enthalpy of formation for slf3 from the given reaction using the products minus reactants with their coefficients and provided values, yielding minus 169 kilojoules.

Explore how electromagnetic radiation travels as waves at the speed of light, linking wavelength and frequency via lambda and nu, from gamma rays to visible, infrared, and radio waves.

Calculate the frequency of red light from a 650 nm emission using c over lambda and unit conversions, yielding about 4.61e14 Hz.

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.

The photoelectric effect shows electrons emitted from a metal surface when light exceeds a threshold frequency; below threshold no emission, while emission and kinetic energy rise with frequency and intensity.

Compare the electron's and ball's wavelengths using lambda = h/(mv), showing the electron's about 7.3e-11 m and the ball's about 1.9e-4 m, highlighting matter's particulate and wave properties.

Explore the hydrogen emission spectrum and line spectrum, showing how excited hydrogen atoms release light at discrete wavelengths, revealing quantized energy levels and Planck's postulates.

Bohr's 1913 quantum model explains the hydrogen atom with quantized, allowed circular orbits and a corresponding energy-level expression, E = -2.178×10^-18 Z^2/n^2, showing that binding energy increases as orbits tighten.

Compute the energy to excite hydrogen from ground state to a higher level and determine the corresponding absorption wavelength using delta E and lambda = h c over delta E.

Compute the ionization energy of hydrogen in its ground state by taking the final energy as zero at infinite separation, yielding Delta E = +2.178e-18.

Trace the shift from the ball model to the quantum mechanical model, introducing wave mechanics, Schrodinger's equation, orbitals as wavefunctions, and Heisenberg's uncertainty principle.

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.

Trace the history of periodic table organization from triads to Mendeleev's predictions and see how hydrogen-like orbitals and Pauli exclusion principle shape electron configurations and atomic-number order.

This lecture explains orbital filling order with sodium and neon, defines valence and core electrons, and explains trends across groups, transition metals, and the lanthanide and actinide series.

Investigate periodic trends in ionization energy, electron affinity, and atomic size, using aluminium’s successive ionizations to explain core versus valence electrons and the large fourth ionization jump.

Compare electron configurations to identify the atom with the largest first ionization energy and the one with the smallest second ionization energy, noting nucleus proximity and core shielding.

Explore electron affinity as the energy change when adding an electron to a gaseous atom, with exothermic processes yielding negative values, and atomic radii vary across periods and down groups.

The ionic radii of B2+, Mg2+, Ca2+, and Sr2+ increase down the group, with B2+ smallest and Sr2+ largest, mirroring the neutral-atom trend.

Learn how to write electron configurations for sulfur, cadmium, and radium and identify the number of electrons in their last orbital using noble gas guidance.

Explore chemical bonds, bond energy, and the difference between ionic and covalent bonding, illustrated by sodium chloride ion formation and hydrogen molecule covalent bonding, plus bond length concepts.

Explore how bonding ranges from ionic to covalent, highlighting polar covalent bonds and dipole moments in HF, H2O, CO2, SO3, and CH4.

Explore electronegativity, the attraction of shared electrons, and how differences create polar covalent and ionic bonds, with bond energies and trend patterns across periods and groups.

Explain how electron configurations seek noble gas arrangements to form covalent and ionic bonds; predict ionic formulas like CaO, note electronegativity differences, group trends, and transition metal exceptions.

Ionic radii are inferred from ion-center distances; cations shrink and anions expand relative to their parent atoms, with isoelectronic ions like F−, Mg2+, and Al3+ ordered from largest to smallest.

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.

Learn how valence electrons form Lewis structures, obey the duet and octet rules, and predict bonding and lone pairs in molecules like H2, F2, H2O, CO2, and CN−.

Draw lewis structures by counting valence electrons, forming single, double, or triple bonds, and distributing lone pairs to satisfy the octet rule for hf, n2, nh3, ch4, cf4, and ions.

Demonstrates resonance by presenting three valid Lewis structures for nitrate and explains that the actual structure is the average of these resonance forms, shown with double-headed arrows.

Explore drawing the Lewis structures for the nitrite ion, assign valence electrons, compare single and double bonds, illustrate the two resonance structures, and note the bent geometry.

Apply the formal charge method to select the best Lewis structure, ensuring charges near zero and negatives on the most electronegative atoms, using lone pairs and shared electrons.

Compute formal charges for sulfate ion by assigning valence electrons to oxygen and sulfur, compare resonance structures, and conclude the two double-bond form is preferred due to lower formal charges.

Use the valence shell electron pair repulsion model to predict molecular geometries by drawing Lewis structures and minimizing electron-pair repulsions. Covers linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral shapes.

Apply vsepr model 2 to predict electron-pair geometries for four and five pairs, tetrahedral, trigonal pyramidal, bent, seesaw, and linear shapes. Lone pairs alter bond angles in CH4, NH3, H2O.

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.

Use bond energy values to estimate delta h for two gas-phase reactions by applying the bond-energy formula (bonds broken minus bonds formed), yielding delta a over delta b = 1.68.

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.

Calculate the moles of hydrogen gas using the ideal gas law PV = nRT, converting Celsius to Kelvin and plugging in P, V, R, and T.

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.

The lecture demonstrates solving a two-state ideal gas problem using PV/T = P2V2/T2, with constant moles and Kelvin temperature conversion to find V2.

Under standard temperature and pressure, use the ideal gas law to relate pressure, volume, and temperature; at STP, one mole occupies 22.42 liters.

Explain the thermal decomposition of calcium carbonate (CaCO3) to CO2, and calculate the CO2 volume from 152 g CaCO3 (1.52 mol) at STP, about 34.1 L.

Master gas stoichiometry with the ideal gas law to balance methane combustion and identify the limiting reactant. Calculate the CO2 volume at the given pressure and temperature.

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.

Apply Dalton's law of partial pressures to calculate the total pressure of gas mixtures by summing individual partial pressures, using the ideal gas law and mole fractions.

Apply the ideal gas law to calculate helium and oxygen partial pressures in a five-liter tank at 25 C, yielding PHe 9.3 atm, PO2 2.4 atm, Ptot 11.7 atm.

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.

Explore how real gases deviate from the ideal gas law by analyzing P, V, T, and n. Use the Vanderwaals equation with a and b corrections to fit observed behavior.

Explore liquids and solids as condensed states driven by intermolecular forces, including covalent and ionic bonding, and how state changes arise from these forces; examine surface tension and capillary action.

Classify solids as crystalline or amorphous, explore lattice and unit cell concepts, and compare ionic, molecular, and atomic solids, including dipole-dipole attractions, hydrogen bonding, and London dispersion forces.

Explore how evaporation and vaporization drive a liquid to gas, reach equilibrium with condensation, and determine vapor pressure through the enthalpy of vaporization and its temperature dependence.

Calculate water’s vapor pressure at 50 degrees Celsius using the vapor pressure equation with ΔHvap and R, converting temperatures to kelvin and obtaining P2.

Explore how heating drives state changes from solid to liquid to vapor, via melting and boiling points, a heating curve, enthalpy of fusion, and phase diagrams.

Apply stoichiometry to find the mass of aluminium that reacts with two liters of oxygen at STP, using 26.98 g/mol and the balanced Al + O2 -> Al2O3 reaction.

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.

Learn rate laws, identify the rate constant k and the order n, and relate forward rates to reactant concentrations using NO2 consumption and O2 production.

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.

Calculate activation energy using the Arrhenius equation from two rate constants at different temperatures, converting to kelvin and solving for Ea.

Explore how catalysts speed up reactions by lowering activation energy, not being consumed, using enzymes as biological catalysts and industrial catalysts, without changing the energy difference between reactants and products.

Explore chemical equilibrium as a dynamic state where reactant and product concentrations stay constant in a closed system, and learn to calculate equilibrium concentrations as forward and reverse reactions balance.

Explain how to write the equilibrium expression for reactions using concentrations and coefficients. Relate the equilibrium constant to equilibrium positions and partial pressures via the ideal gas law.

Predict the shift toward equilibrium by comparing Q to K. Include initial concentrations in Q and exclude pure solids or liquids from the expression, as in ammonia synthesis.

Explore ammonia synthesis at 500 degrees, comparing reaction quotient Q to the equilibrium constant K to predict how the system shifts left or right under three initial condition scenarios.

Compute equilibrium concentrations for the PCl5 ⇌ PCl3 + Cl2 system in a 1 liter flask, using initial moles and stoichiometry, then determine the expression and value of Kc.

Analyze water-gas shift CO + H2O ⇌ CO2 + H2 at 700 K (K = 5.1) with 1 M, yielding CO and H2O 0.613 M, CO2 and H2 1.387 M.

Demonstrates solving for HF formation from H2 and F2 using an ice table and quadratic, calculating concentrations in a 3 L flask, and determining equilibrium concentrations and the equilibrium constant.

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.

Explore how gas-phase equilibria respond to pressure and volume changes, including inert gases, using Le Chatelier's principle; learn how volume shifts affect molecule counts, the equilibrium constant, and temperature-driven changes.

Compute equilibrium concentrations for H2 + F2 ⇌ 2 HF with K ≈ 115, given 3 moles in 1.5 L, determine Q, and show the rightward shift to HF.

Explore Brønsted–Lowry acid–base concepts, proton transfer, and conjugate acid–base pairs through equilibrium ideas, with examples like HCl in water and the role of Ka.

Explore how acids dissociate in water, focusing on hydrochloric acid, acetic acid, and ammonia. HCl dissociates completely to H3O+ and Cl−; acetic acid forms CH3COO− and H3O+; ammonia forms NH4+.

Explore how strong and weak acids differ by equilibrium position and dissociation, with Ka values guiding ionization. See proton transfer from HCl to NH3 and the resulting conjugate bases.

Compare acid strengths using Ka values for HF, NO2-, and CN-, then deduce the conjugate bases (F-, NO2-, CN-) and their relative base strengths.

Explore amphoteric water behavior through water autoprotolysis, derive the ion-product of water, Kw = [H3O+][OH-] = 1.0e-14 at 25 °C, and identify neutral, acidic, and basic solutions.

Calculate concentrations using Kw at 25 °C to determine whether solutions are basic, neutral, or acidic; case A basic, case B neutral, case C acidic, based on H+ and OH-.

Describe how heating to 60 °C raises Kw to 1e-13, signaling an endothermic shift per Le Chatelier, and compute neutral solution [H+] = [OH−] = sqrt(Kw) ≈ 3×10^−7 M.

Explore pH scale and how pH equals the negative log of hydrogen ion concentration, with examples, significant figures, and the pH plus pOH equals 14 at twenty five degrees centigrade.

Calculate pH, pOH, and pKw for two solutions at 25 degrees centigrade using Kw = 1.0×10^-14, with given [edg+] and [oh−], showing pH plus pOH equals 14.

Compute pH and ion concentrations from a blood sample with pH 7.41 at 25 °C using Kw, yielding [H+] ≈ 3.9×10^-8 M and pOH ≈ 6.59.

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%.

Calculate the Ka for lactic acid from a 0.10 M solution using percent dissociation, determine x = 3.7×10^-3 M, and find Ka ≈ 1.4×10^-4 by neglecting small terms.

Explore how strong bases dissociate in water, such as hydroxides of group 1 elements and alkaline earth hydroxides, and how ammonia acts as a base by accepting a proton.

Identify a strong base that dissociates completely to OH- and its counterions. Use Kw to relate H+ and OH-, compute [X+] = Kw/[H+], and find pOH ≈ 12.70.

Apply the ammonia–water equilibrium with Kb = 1.8×10−5 to a 15 M NH3 solution. Solve for [NH4+] and [OH−], then determine pH 12.20 via Pathway A or Pathway B.

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.

Analyze the three-step dissociation of phosphoric acid using Ka1, Ka2, and Ka3 to determine the equilibrium concentrations of H+, H2PO4-, HPO4^2-, and PO4^3-.

Salt dissociates in water into ions that can act as acids or bases. Apply Ka times Kb equals Kw to relate a weak acid and its conjugate base.

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.

Identify the common ion effect: the fluoride ion from NaF shifts the weak acid’s dissociation to the left, reducing acidity.

Using a 1 m HF solution with 1 m NaF as a common ion, the lecture derives [H+]=7.2×10^-4 M from Ka=7.2×10^-4 and shows HF’s percent dissociation is 0.072%.

Examine buffering in acid-base solutions and how buffers resist changes when protons or hydroxide ions are added. Blood illustrates buffering with weak acids or bases and salts using equilibrium calculations.

Analyze a buffer made from 0.5 M acetic acid and 0.5 M sodium acetate, applying Ka = 1.8×10−5 to determine pH. Assume small dissociation and compute pH approximately 4.74.

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.

Explore how buffering capacity stabilizes pH in a weak acid–conjugate base system, showing how added OH- or H+ shifts concentrations without large pH changes, governed by [A-]/[HA].

Apply the Henderson-Hasselbalch relation to calculate pH changes in acetic acid buffers when a strong acid is added, using the given pKa and buffer concentrations.