Atomic Structure & Chemical Bonding

Atomic Structure & Properties

• Protons: Positively charged particles located in the nucleus; determine the atomic number of an element.
• Neutrons: Uncharged particles found in the nucleus; affect the isotope of the element.
• Electrons: Negatively charged particles orbiting the nucleus; determine an element's charge and bonding behavior.
• Isotopes: Atoms of the same element with different numbers of neutrons.
• Atomic radius: Increases down a group, decreases across a period.
• Ionization energy: Decreases down a group, increases across a period.
• Electronegativity: Increases across a period, decreases down a group.

Chemical Bonding

• Ionic bonds: Occur between metals and nonmetals; involve electron transfer, forming cations (positive ions) and anions (negative ions).
• Covalent bonds: Occur between nonmetals; involve electron sharing.
  • Polar covalent: Unequal sharing of electrons.
  • Nonpolar covalent: Equal sharing of electrons.

Lewis Structures & Molecular Geometry

• Lewis structures: Visual representations of bonding in molecules, showing shared and lone pairs of electrons.
• Resonance: Molecules with multiple possible Lewis structures that contribute to the actual bonding.
• VSEPR theory: Valence Shell Electron Pair Repulsion; predicts molecular shapes by considering electron pair repulsion.
• Hybridization: Combining atomic orbitals of an atom to form hybrid orbitals for bonding. Includes sp, sp², and sp³ hybridization.

Stoichiometry

• Balancing chemical equations: Ensuring equal numbers of each element on both sides of the equation.
• Mole concept: 1 mole = 6.022 × 10²³ particles (Avogadro's number).
• Molar mass: Mass of 1 mole of a substance.
• Limiting reactant: Reactant that is completely consumed in a reaction, determining the amount of product formed.
• Percent yield: Percentage of the theoretical yield that is actually obtained experimentally.
  • Percent Yield = (Actual Yield / Theoretical Yield) × 100

Thermochemistry

• ΔE = q + w: Change in internal energy equals the heat absorbed or released plus the work done.
• Enthalpy (ΔH): Heat change at constant pressure.
• Exothermic: Reaction releases heat (ΔH < 0).
• Endothermic: Reaction absorbs heat (ΔH > 0).
• Hess's Law: Total enthalpy change for a reaction equals the sum of enthalpy changes for individual steps.
• Specific Heat Formula: q = mcΔT (heat = mass × specific heat × temperature change).

Gas Laws

• Boyle's Law: P₁V₁ = P₂V₂ (constant temperature).
• Charles's Law: V₁/T₁ = V₂/T₂ (constant pressure).
• Avogadro's Law: V₁/n₁ = V₂/n₂ (constant pressure and temperature).
• Ideal Gas Law: PV = nRT (P = pressure, V = volume, n = moles, R = ideal gas constant, T = temperature in Kelvin).
• Dalton's Law of Partial Pressures: Total pressure is the sum of the partial pressures of all component gases.
• Kinetic Theory of Gases: Explains gas behavior in terms of the motion of gas particles.
  • Graham's Law of Effusion: The rate of effusion of a gas is inversely proportional to the square root of its molar mass.

Solutions & Concentration

• Molarity (M): Moles of solute per liter of solution.
• Molality (m): Moles of solute per kilogram of solvent.
• Boiling Point Elevation: ΔT_b = K_b ⋅ m (change in boiling point).
• Freezing Point Depression: ΔT_f = K_f ⋅ m (change in freezing point).
• Dilution Equation: M₁V₁ = M₂V₂

Acids and Bases

• Arrhenius definition: Acids release H⁺ ions; bases release OH⁻ ions.
• Brønsted-Lowry definition: Acids donate H⁺ ions; bases accept H⁺ ions.
• Lewis definition: Acids accept electron pairs; bases donate electron pairs.
• pH and pOH: pH = -log[H⁺], pOH = -log[OH⁻].
  • Relationship of pH and pOH: pH + pOH = 14.
• Strong vs. weak acids/bases: Strong completely dissociate; weak only partially dissociate.
• Buffers: Resist changes in pH. - Henderson-Hasselbalch equation: pH = pKa + log([A⁻]/[HA])

Chemical Kinetics

• Reaction rate: Rate = k[A]ᵐ[B]ⁿ (describes how fast a reaction occurs).
• Activation energy: Minimum energy needed for a reaction to occur.
• Arrhenius Equation: k = Ae^(-Eₐ/RT)
• Collision Theory: Explains reaction rates based on collisions between molecules.

Equilibrium

• Dynamic equilibrium: The opposing forward and reverse reactions occur at equal rates.
• Le Chatelier's Principle: A system at equilibrium adjusts to counteract stress (changes in concentration, temperature, or pressure).
• Equilibrium Constant (K): Describes the relative amounts of products and reactants at equilibrium.
  • For the reaction: aA + bB ⇌ cC + dD, K_c = [C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ
• ICE Table: Used to calculate equilibrium concentrations.

Key Formulas

• Ideal Gas Law: PV = nRT
• Mole-to-Mass Conversion: mass = moles × molar mass
• pH Calculation: pH = -log[H⁺]
• Molarity: M = moles of solute / liters of solution
• Energy Formula: q = mcΔT

Ionic Species

• Ionic species with the same number of electrons differ in the number of protons. This affects the overall charge and attraction to the nucleus.

Empirical Formula

• Determine the empirical formula by converting given masses to moles, identifying the mole ratio, and simplifying to the smallest whole number ratio. Using combustion analysis, determine the moles of carbon and hydrogen present in a compound to establish the empirical formula.

Quantum Numbers

• Quantum numbers describe the properties of electrons in an atom.
  • Principal quantum number (n): energy level
  • Angular momentum quantum number (l): subshell
  • Magnetic quantum number (ml): orbital orientation
  • Spin quantum number (ms): electron spin (either +½ or -½).

Test your knowledge on atomic structure, properties, and chemical bonding. This quiz covers concepts such as protons, neutrons, electrons, isotopes, and different types of chemical bonds including ionic and covalent bonds. Challenge yourself and see how well you understand these fundamental chemistry topics!