Chapter 4: Three Major Classes of Chemical Reactions

4.1 Solution Concentration & Water as a Solvent

Understanding solution concentration is fundamental in chemistry, especially when dealing with reactions in aqueous solutions. Molarity is the most common unit of concentration, relating the amount of solute to the total volume of solution.

• Molarity (M):

• Finding moles from molarity and volume:

• Finding solution volume:

• Dilution formula: (moles of solute remain constant during dilution)

• Ion molarity:

• Electrolytes: Strong electrolytes dissociate completely (e.g., NaCl), weak electrolytes partially (e.g., CH3COOH), and nonelectrolytes do not dissociate (e.g., sugar).

• Polyatomic ions remain intact during dissociation (e.g., K2SO4 → 2 K+ + SO42−).

Example: To find the molarity of a solution made by dissolving 76.7 g AgNO3 in 846 mL of solution:

• Convert grams to moles:

• Convert mL to L:

• Calculate molarity:

Additional info: Always use liters of total solution, not just solvent, when calculating molarity.

4.2 Precipitation Reactions

Precipitation reactions occur when two soluble ionic compounds exchange ions and form an insoluble product (precipitate).

• Molecular equation: All compounds written as intact units.

• Total ionic equation: Strong electrolytes are written as ions.

• Spectator ions: Ions that do not participate in the reaction.

• Net ionic equation: Only species that change during the reaction are included.

Example: Net ionic equation for AgNO3(aq) + NaCl(aq):

• Molecular: AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)

• Total ionic: Ag+(aq) + NO3−(aq) + Na+(aq) + Cl−(aq) → AgCl(s) + Na+(aq) + NO3−(aq)

• Net ionic: Ag+(aq) + Cl−(aq) → AgCl(s)

Additional info: Do not split solids, liquids, gases, or weak electrolytes into ions in total ionic equations.

4.3 Acid–Base Reactions

Acid–base reactions involve the transfer of protons (H+). A strong acid and strong base neutralization produces water.

• Strong acids: HCl, HBr, HI, HNO3, H2SO4, HClO4

• Strong bases: Group 1 hydroxides, Ca(OH)2, Sr(OH)2, Ba(OH)2

• Weak acids/bases: e.g., HF, CH3COOH, NH3

• Net ionic for strong acid + strong base:

• Titration: Use stoichiometry to relate titrant volume and concentration to the unknown.

Example: 25.00 mL H2SO4 neutralized by 30.00 mL of 0.2000 M NaOH:

• Equation: H2SO4 + 2 NaOH → Na2SO4 + 2 H2O

• Moles NaOH:

• Moles H2SO4:

• Molarity:

Additional info: Coefficients in the balanced equation are crucial for titration calculations.

4.4 Oxidation–Reduction (Redox) Reactions

Redox reactions involve the transfer of electrons. Oxidation is loss of electrons, reduction is gain (OIL RIG).

• Oxidizing agent: Gains electrons, is reduced.

• Reducing agent: Loses electrons, is oxidized.

• Assigning oxidation numbers:

  • Elemental form: 0

  • Monatomic ion: equal to charge

  • Group 1: +1; Group 2: +2

  • F: −1; O: usually −2 (−1 in peroxides); H: +1 with nonmetals, −1 with metals

  • Sum in neutral compound: 0; in polyatomic ion: equals ion charge

Example: Find S in SO3:

• Each O is −2, so 3 × (−2) = −6

• S + (−6) = 0 ⇒ S = +6

Additional info: The oxidizing agent is reduced, and the reducing agent is oxidized.

4.5 Elements in Redox Reactions

Chemical reactions can be classified by their patterns, especially in redox processes.

• Combination: X + Y → Z (fewer products than reactants)

• Decomposition: Z → X + Y (more products than reactants)

• Single displacement: X + YZ → XZ + Y

• Double displacement: AB + CD → AD + CB

• Combustion: Reaction with O2, often oxidation

Example: 2Mg(s) + N2(g) → Mg3N2(s) is a combination and redox reaction.

4.6 Reversibility & the Equilibrium State

At dynamic equilibrium, the forward and reverse reactions occur at equal rates, so concentrations remain constant (not necessarily equal).

• Reversible reactions: Proceed in both directions, denoted by

• Dynamic equilibrium: Microscopic reactions continue, but macroscopic concentrations are constant

• At equilibrium:

Additional info: Equilibrium does not require equal concentrations of reactants and products.

Chapter 7 Preview: Quantum Theory & Atomic Structure

7.1 The Nature of Light

Light exhibits both wave and particle properties. Its energy and behavior are described by several key equations.

• Speed of light: (c = 3.00 × 108 m/s)

• Photon energy (frequency): (h = 6.626 × 10−34 J·s)

• Photon energy (wavelength):

• Wavelength (λ) and frequency (ν) are inversely related.

• Electromagnetic spectrum (increasing energy): radio < microwave < infrared < visible < ultraviolet < X-ray < gamma

Example: For λ = 500 nm, ν = c/λ = (3.00 × 108)/(5.00 × 10−7) = 6.00 × 1014 s−1

7.2 Atomic Spectra

Atoms emit or absorb only specific energies, resulting in line spectra. For hydrogen-like atoms, energy levels are quantized.

• Energy change:

• Photon energy:

• Hydrogen spectral lines: , where R is the Rydberg constant

• Photon is emitted when electron drops to a lower energy level; absorbed when moving to a higher level.

7.3 Wave–Particle Duality of Matter and Energy

Both light and matter exhibit wave–particle duality. The de Broglie equation relates a particle's mass and velocity to its wavelength.

• de Broglie wavelength:

• Heisenberg uncertainty principle:

• Lighter particles or slower speeds yield longer wavelengths.

• Position and momentum cannot both be known with unlimited precision.

7.4 Quantum-Mechanical Model of the Atom

Electrons are described by quantum numbers and occupy orbitals, which are regions of probability rather than fixed paths.

• Number of orbitals in a sublevel:

• Maximum electrons in a shell:

• s orbitals are spherical; p orbitals have two lobes.

Chapter 8 Preview: Electron Configuration & Chemical Periodicity

8.1 Characteristics of Many-Electron Atoms

Electron–electron repulsion and shielding affect orbital energies in many-electron atoms. Penetration and effective nuclear charge (Zeff) are key concepts.

• Penetration: s > p > d > f (within the same n)

• Greater penetration = stronger attraction to nucleus = lower energy

• Zeff increases across a period due to increasing nuclear charge

8.2 Quantum-Mechanical Model & the Periodic Table

Electron configurations are determined by the Aufbau principle, Hund's rule, and the Pauli exclusion principle.

• Aufbau principle: Fill lowest energy orbitals first

• Hund's rule: One electron per degenerate orbital before pairing

• Pauli exclusion principle: Max 2 electrons per orbital, opposite spins

• Electron configuration notation: 1s2 2s2 2p6 ...

• Noble-gas shorthand: e.g., Na = [Ne]3s1

• Main-group valence electrons: highest n s and p orbitals

Additional info: Some transition metals have exceptions to the expected filling order (e.g., Cr, Cu).

8.3 Trends in Three Atomic Properties

• Cations are smaller than their neutral atoms; anions are larger.

• For isoelectronic series, more protons = smaller radius.

• Ionization energy: energy required to remove an electron from a gaseous atom.

• Large jumps in successive ionization energies indicate core electron removal.

8.4 Atomic Properties & Chemical Reactivity

Trends in the periodic table influence chemical reactivity and the types of ions formed.

• Metals lose electrons (form cations); nonmetals gain/share electrons.

• Metallic character increases down a group and to the left across a period.

• Group 1 metals become more reactive down the group; halogen reactivity decreases down the group.

• Main-group ion charges often correspond to achieving noble-gas configurations.

Key Equations and Rules Summary

• Molarity:

• Dilution:

• Ion molarity: Compound M × ion subscript

• Precipitation: Use solubility rules; split only strong aqueous electrolytes; cancel spectators

• Strong acid/base net ionic:

• Titration: volume → moles → mole ratio → unknown

• Redox: OIL RIG; oxidizing agent is reduced; reducing agent is oxidized

• Equilibrium:

• Light: ;

• Matter wave:

• Quantum numbers: n; ; ;

• Sublevels: s:1 orbital/2e; p:3/6e; d:5/10e; f:7/14e

• Electron filling: Aufbau; Hund; Pauli

• Radius: Decreases across, increases down

• Ionization energy: Increases across, decreases down

• Isoelectronic ions: More protons = smaller radius

Best exam habit: Before calculating, identify the conversion path or chemistry rule first. The setup usually indicates whether your answer will make sense.