Chem 109 Unit II Study Guide: Bonding, Reactions, Gases, Solutions, Acids/Bases, and Redox

Bonding & Molecular Properties

Naming Compounds and Polyatomic Ions

Understanding chemical nomenclature is essential for identifying and communicating about compounds. Compounds may be classified as molecules (nonmetal + nonmetal), formula units (metal + nonmetal), or those containing polyatomic ions.

• Molecules: Composed of nonmetals; named using prefixes (e.g., carbon dioxide, CO2).

• Formula Units: Ionic compounds formed from metals and nonmetals (e.g., sodium chloride, NaCl).

• Polyatomic Ions: Charged groups of atoms acting as a unit (e.g., sulfate, SO42-).

• Naming: Use handouts or reference tables for naming compounds with polyatomic ions.

Example: Na2SO4 is named sodium sulfate.

Lewis Structures

Lewis structures represent the arrangement of electrons in molecules, showing bonds and lone pairs.

• Steps: Count valence electrons, arrange atoms, distribute electrons to satisfy octet rule.

• Key Point: Lone pairs and bonds determine molecular shape.

Example: Water (H2O) has two bonds and two lone pairs on oxygen.

Electronegativity and Bond Dipoles

Electronegativity is the tendency of an atom to attract electrons. Differences in electronegativity create bond dipoles, indicated by δ+ (partial positive) and δ- (partial negative).

• Bond Dipole: The more electronegative atom attracts electrons, becoming δ-.

Example: In HCl, Cl is δ-, H is δ+.

VSEPR Shapes

The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes based on the number of bonds and lone pairs.

• Tetrahedral: Four bonds, no lone pairs (e.g., CH4).

• Pyramidal: Three bonds, one lone pair (e.g., NH3).

• Bent: Two bonds, two lone pairs (e.g., H2O).

• Linear: Two bonds, no lone pairs (e.g., CO2).

Phase Changes vs. Chemical Changes

Phase changes are physical changes (e.g., evaporation, condensation, sublimation) that do not alter chemical identity, while chemical changes involve the formation of new substances.

• Evaporation: Liquid to gas.

• Condensation: Gas to liquid.

• Sublimation: Solid to gas.

Example: Melting ice is a phase change; burning hydrogen is a chemical change.

Chemical Reactions and Stoichiometry

The Mole Concept

The mole is a fundamental unit in chemistry, analogous to a "dozen" for counting atoms or molecules.

• 1 mole = $6.022 \times 10^{23}$ particles (Avogadro's number).

• Calculating: Number of particles = moles × Avogadro's number.

Example: 2 moles of H2O contains $2 \times 6.022 \times 10^{23}$ molecules.

Molar Mass Calculation

Molar mass is the mass of one mole of a compound, calculated from its chemical formula.

• Formula: Sum atomic masses of all atoms in the compound.

Example: Molar mass of H2O = $2 \times 1.01 + 16.00 = 18.02$ g/mol.

Balancing Chemical Reactions

Balancing reactions ensures the same number of atoms of each element on both sides of the equation.

• Steps: Adjust coefficients, not subscripts.

Example: $2H_2 + O_2 \rightarrow 2H_2O$

Mole Maps and Stoichiometry

Mole maps help determine quantities of reactants and products in reactions.

• Use: Convert between grams, moles, particles, and volume.

Example: Given moles of reactant, use balanced equation to find moles of product.

Gases

Kinetic Molecular Theory

This theory explains gas behavior based on the motion of particles.

• Assumptions: Gas particles are in constant, random motion; collisions are elastic.

• Properties: Pressure, volume, temperature, and number of moles affect gas behavior.

Ideal Gas Law

The ideal gas law relates pressure, volume, temperature, and moles of gas.

• Equation: $PV = nRT$

• Variables: P = pressure (atm), V = volume (L), n = moles, R = gas constant ($0.0821$ L·atm/mol·K), T = temperature (K).

• Note: Always use Kelvin for temperature.

Law of Combining Volumes

At constant temperature and pressure, equal volumes of gases contain equal numbers of moles.

• Application: Useful for stoichiometry involving gases.

Example: 1 L of H2 and 1 L of O2 at same T and P contain equal moles.

Solutions

Definition and Components

A solution is a homogeneous mixture of two or more substances. The solute is dissolved in the solvent.

• Solute: Substance present in lesser amount.

• Solvent: Substance present in greater amount.

Example: Salt (solute) dissolved in water (solvent).

Molarity Calculation

Molarity (M) is the concentration of a solution, defined as moles of solute per liter of solution.

• Equation: $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Rearrangement: Use "mole triangle" to solve for M, moles, or L.

Weight Percent Calculation

Weight percent expresses concentration as the mass of solute per 100 g of solution.

• Equation: $\text{wt\%} = \frac{\text{mass of solute}}{\text{mass of solution}} \times 100\%$

Acids and Bases

Physical and Chemical Properties

Acids and bases have distinct properties.

• Acids: Sour taste, turn litmus red, react with metals to produce hydrogen.

• Bases: Bitter taste, slippery feel, turn litmus blue.

Acid-Base Theories

Two main theories describe acids and bases.

• Arrhenius: Acids produce H+ in water; bases produce OH-.

• Brønsted-Lowry: Acids donate protons (H+); bases accept protons.

Strong vs. Weak Acids and Bases

Strong acids and bases dissociate completely in water; weak ones do not.

• Six Strong Acids: HCl, HBr, HI, HNO3, HClO4, H2SO4

• Difference: Strong acids/bases yield high ion concentration; weak ones yield low.

Chemical Equations for Acid/Base Behavior

Acids and bases react in water to form ions.

• Example: $HCl + H_2O \rightarrow H_3O^+ + Cl^-$

• Acidic/Basic Anhydrides: Nonmetal oxides form acids; metal oxides form bases in water.

Conductivity and Ion Concentration

The conductivity of a solution depends on the concentration of ions present.

• Higher ion concentration: Greater conductivity.

Neutralization Reactions

Neutralization occurs when an acid reacts with a base to form water and a salt.

• Equation: $HCl + NaOH \rightarrow NaCl + H_2O$

• Stoichiometry: Balance coefficients so reaction goes to completion.

pH and Acidity

pH measures the acidity or alkalinity of a solution.

• Equation: $pH = -\log[H_3O^+]$

• Correlation: Lower pH = more acidic; higher pH = more basic.

Calculating pH

Given hydronium ion concentration, pH can be calculated.

• Example: If $[H_3O^+] = 1 \times 10^{-3}$, then $pH = 3$.

Redox Reactions

Assigning Oxidation Numbers

Oxidation numbers indicate the charge an atom would have if electrons were transferred completely.

• Rules: Elements = 0; ions = charge; oxygen = -2; hydrogen = +1.

Identifying Oxidation and Reduction

Oxidation is loss of electrons; reduction is gain of electrons.

• Oxidized: Substance loses electrons.

• Reduced: Substance gains electrons.

Example: $Zn + Cu^{2+} \rightarrow Zn^{2+} + Cu$ (Zn is oxidized, Cu2+ is reduced)

Oxidizing and Reducing Agents

The oxidizing agent causes oxidation (accepts electrons); the reducing agent causes reduction (donates electrons).

• Oxidizing agent: Is reduced.

• Reducing agent: Is oxidized.

Cathode and Anode in Electrochemical Cells

Electrochemical cells have two terminals:

• Cathode: Site of reduction.

• Anode: Site of oxidation.

Balancing Half Reactions

Half reactions must be balanced for mass and charge, ensuring equal electrons gained and lost.

• Steps: Balance atoms, add electrons, balance charges.

Formation of Iron(II) and Iron(III) Hydroxide

Iron(II) hydroxide and iron(III) hydroxide form via redox reactions, involving oxidizing and reducing agents and an electrolyte.

• Electrolyte: Facilitates ion movement.

• Iron(II) hydroxide: Fe2+ + 2OH- → Fe(OH)2

• Iron(III) hydroxide: Fe3+ + 3OH- → Fe(OH)3

Example: In water treatment, iron(III) hydroxide precipitates to remove contaminants.

Process	Type	Key Features
Evaporation	Phase Change	Liquid to gas, no new substance
Combustion	Chemical Change	New substances formed, energy released
Neutralization	Chemical Change	Acid + base → salt + water
Sublimation	Phase Change	Solid to gas, no new substance

Additional info: This guide covers all major objectives for Chem 109 Unit II, including bonding, reactions, gases, solutions, acids/bases, and redox, with definitions, examples, and key equations for exam preparation.