Chapter 4: Reactions in Aqueous Solution

Solutions and Their Components

A solution is a homogeneous mixture composed of a solvent (the substance present in the greatest amount) and one or more solutes (substances dissolved in the solvent). When water is the solvent, the solution is called aqueous and denoted as (aq).

• Solvent: Substance doing the dissolving (e.g., water in salt water).

• Solute: Substance being dissolved (e.g., NaCl in salt water).

• Aqueous (aq): Indicates the solute is dissolved in water.

Solvation and Dissociation

Solvation is the process where solvent particles surround solute particles. Dissociation refers to the separation of an ionic compound into its constituent ions when dissolved in water.

• Example:

• Molecular compounds (e.g., sugar) may dissolve but do not dissociate into ions.

Electrolytes: Classification and Properties

Electrolytes are substances that produce ions in solution and thus conduct electricity. They are classified as:

• Strong: One-way arrow ()

• Weak: Double arrow ()

How to Classify Substances

• Soluble ionic compound? Strong electrolyte

• One of the 7 strong acids or 8 strong bases? Strong electrolyte

• Other acids or weak bases (e.g., NH3)? Weak electrolyte

• Sugar, alcohol, or molecular compound? Nonelectrolyte

Common trap: All soluble ionic compounds are strong electrolytes, even if not acids or bases.

Strong Acids and Bases to Memorize

• Acids not on the list (e.g., HF, CH3COOH) are weak acids.

Solubility Rules (Simplified)

Dissociation of Ionic Compounds

When ionic compounds dissolve, they separate into ions according to their formula subscripts.

Polyatomic ions (e.g., PO43-) remain intact during dissociation.

Molarity and Dilution

• Molarity (M):

• Moles from molarity:

• Dilution: (moles of solute remain constant)

Example: To prepare 450 mL of 0.10 M solution from 3.0 M stock:

• mL (use 15 mL stock, dilute to 450 mL)

Acid-Base Neutralization

• General reaction: Acid + Base → Salt + Water

• Net ionic:

• Use mole ratios from balanced equations for stoichiometry.

Worked Examples and Common Traps

• CaBr2: Soluble ionic compound, strong electrolyte, bright conductivity.

• K2S (0.15 M): [K+] = 2 × 0.15 M = 0.30 M

• Na3PO4: To get [Na+] = 0.800 M, need [Na3PO4] = 0.267 M

• AgNO3 + MgCl2: Use stoichiometry and molar mass to find mass of AgCl precipitate.

Chapter 5: Thermochemistry

Electrostatic Potential Energy (Eel)

Electrostatic potential energy is the energy due to interactions between charged particles.

• Like charges close: higher (repulsion)

• Opposite charges close: lower (attraction)

• Bond formation: Releases energy

• Bond breaking: Requires energy

First Law of Thermodynamics

Energy is conserved; it cannot be created or destroyed, only transferred or converted.

• q: Heat; positive if system gains heat, negative if loses

• w: Work; positive if work done on system, negative if by system

• Memory trick: Positive = energy enters system

System and Surroundings

Internal Energy and State Functions

• State function: Depends only on initial and final states (e.g., E, H)

• Heat (q) and work (w) are not state functions

Enthalpy and Heat at Constant Pressure

• At constant pressure:

Thermochemical Calculations (Grams to kJ)

• Convert grams to moles using molar mass

• Use as a conversion factor (match to equation coefficients)

Example: , kJ

• 4.50 g CH4 → 0.281 mol → kJ released

Calorimetry: Measuring Heat Flow

• c (specific heat): For water, J/g°C

Example: 250 g water, °C → J = 78.5 kJ

• If solution warms, positive, negative (and vice versa)

Hess's Law and Enthalpy of Formation

• Standard enthalpy of formation (): Enthalpy change for forming 1 mol of a compound from elements in standard states

• Elements in standard state:

Example:

• Products: kJ

• Reactants: kJ

• kJ

Chapter 6: Electronic Structure of Atoms

Wavelength, Frequency, and the Speed of Light

• Wavelength (): Distance between wave peaks (m, nm)

• Frequency (): Waves per second (Hz or s-1)

• Speed of light (): m/s

• Longer wavelength = lower frequency; shorter wavelength = higher frequency

Example: nm → s-1

Photon Energy

• Photon: Packet of light energy

• or

• Planck's constant (): J·s

• High frequency/short wavelength = high energy

Example: s-1 → J

Electromagnetic Spectrum

• Order (highest to lowest energy): Gamma rays > X-rays > Ultraviolet > Visible > Infrared > Microwave > Radio

• Visible light: 400–700 nm (violet = shorter wavelength, higher energy; red = longer wavelength, lower energy)

Quantized Energy and the Photoelectric Effect

• Energy is quantized (comes in discrete packets)

• Photoelectric effect: Electrons ejected from metal only if photon energy is sufficient ()

• Low-frequency (long-wavelength) light may not eject electrons, regardless of brightness

Line Spectra and the Bohr Model

• Atoms emit light at specific wavelengths (line spectra) due to electron transitions between energy levels

• Electron moves to higher (energy level): absorbs energy ( positive)

• Electron moves to lower : emits energy ( negative)

Atomic Orbitals

• Each orbital holds max 2 electrons

Quantum Numbers

Electron Configurations

• Describes electron arrangement in orbitals

• Order: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p

• 4s fills before 3d

Pauli Exclusion Principle and Hund's Rule

• Pauli exclusion principle: No two electrons in an atom have the same four quantum numbers; max 2 electrons per orbital with opposite spins

• Hund's rule: Electrons fill degenerate (equal-energy) orbitals singly before pairing

• Example: Carbon (1s2 2s2 2p2): 2p electrons occupy separate orbitals first

Periodic Table Blocks

Common Electron Configuration Anomalies

Final Review Checklist

• Define solution, solute, solvent, aqueous

• Classify electrolytes and memorize strong acids/bases

• Write dissociation equations and compare ion concentrations

• Use , for solutions

• Apply , , products reactants

• Use , for light calculations

• Write electron configurations and apply Pauli/Hund rules

• Remember 4s fills before 3d; know Cr and Cu exceptions

Study order: 1) Memorize formulas/lists. 2) Practice setup steps. 3) Redo worked examples. 4) Be able to explain reasoning for each step.

Additional info: This guide includes all essential formulas, definitions, and worked examples for Chapters 4–6, focusing on solution chemistry, thermochemistry, and atomic structure. It is suitable for exam preparation in a general chemistry course.