Chapter 4 – Types of Reactions and Solution Stoichiometry

Water shape

Bent or V-shaped molecule

Water bond angle

Approximately 104.5°

Water polarity

Polar; oxygen is partially negative, hydrogen is partially positive

Reason water is polar

Unequal sharing of electrons between O and H atoms

Why water is a good solvent

Its polarity allows it to dissolve ionic and polar compounds

Solution

A homogeneous mixture of solute(s) and solvent

Solvent

The substance present in the largest amount; does the dissolving

Solute

The substance being dissolved

Electrolyte

A substance that produces ions and conducts electricity when dissolved in water

Nonelectrolyte

A substance that does not produce ions in solution and does not conduct electricity

Acid

Substance that produces H⁺ ions in aqueous solution

Base

Substance that produces OH⁻ ions in aqueous solution

Strong acid

Completely ionizes in water (examples: HCl, HNO₃, H₂SO₄)

Weak acid

Partially ionizes in water (examples: HF, CH₃COOH)

Strong base

Completely dissociates in water (examples: NaOH, KOH, Ba(OH)₂)

Weak base

Partially dissociates or reacts weakly with water (example: NH₃)

Titration

A lab method used to determine unknown concentration by reacting it with a known one

Precipitate

An insoluble solid formed from two aqueous solutions

Redox reaction

A chemical reaction involving transfer of electrons between species

Oxidation

Loss of electrons (OIL = Oxidation Is Loss)

Reduction

Gain of electrons (RIG = Reduction Is Gain)

Oxidizing agent

The substance that gets reduced and causes oxidation of another

Reducing agent

The substance that gets oxidized and causes reduction of another

Soluble

Capable of dissolving significantly in water

Insoluble

Does not dissolve significantly in water

Half reaction

A reaction showing only oxidation or reduction separately with electrons included

Molecular equation

Shows all reactants and products in their neutral, undissociated forms

Complete ionic equation

Shows all strong electrolytes as ions in solution

Net ionic equation

Shows only species directly involved in the chemical change

Six basic reaction types

Synthesis, Decomposition, Combustion, Single Replacement, Double Replacement, Acid-Base

Three main types of solution reactions

Precipitation, Acid-Base, and Redox

Precipitation reaction

Two aqueous solutions form an insoluble product (solid)

Acid-base reaction

Acid reacts with base to produce water and a salt

Redox reaction

Electrons are transferred between substances; oxidation and reduction both occur

Difference between precipitation and acid-base

Precipitation forms an insoluble solid; acid-base forms water

What is unique about redox reactions

They involve electron transfer and oxidation state changes

Example of precipitation reaction

AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)

Example of acid-base reaction

HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

Example of redox reaction

Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)

Stoichiometry in neutralization reactions

Use balanced equation to find mole ratios between acid and base

Oxidation state

The apparent charge of an atom based on electron gain/loss rules

How to determine oxidation states

Apply rules (oxygen = –2, hydrogen = +1, sum = compound charge)

OIL RIG meaning

Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons)

Half reaction type identification

Oxidation loses electrons; Reduction gains electrons

Balancing redox in acidic solution

Add H₂O, H⁺, and e⁻ as needed to balance atoms and charge

Balancing redox in basic solution

Add H₂O, OH⁻, and e⁻ to balance atoms and charge

💨 Chapter 5 – Gases

Boyle’s Law

P₁V₁ = P₂V₂ (pressure × volume = constant; temperature constant)

Boyle’s Law relationship

Inverse: as pressure ↑, volume ↓

Boyle’s Law graph

Downward-curved (hyperbola) for P vs. V

Charles’s Law

V₁/T₁ = V₂/T₂ (volume ÷ temperature = constant; pressure constant)

Charles’s Law relationship

Direct: as temperature ↑, volume ↑

Charles’s Law graph

Straight line through origin (V vs. T in Kelvin)

Gay-Lussac’s Law

P₁/T₁ = P₂/T₂ (pressure ÷ temperature = constant; volume constant)

Gay-Lussac’s Law relationship

Direct: as temperature ↑, pressure ↑

Gay-Lussac’s Law graph

Straight line through origin (P vs. T in Kelvin)

Avogadro’s Law

V₁/n₁ = V₂/n₂ (volume ÷ moles = constant; P and T constant)

Avogadro’s Law relationship

Direct: as moles ↑, volume ↑

Combined Gas Law

(P₁V₁)/T₁ = (P₂V₂)/T₂ (combines Boyle’s, Charles’s, and Gay-Lussac’s)

Ideal Gas Law

PV = nRT (relationship between pressure, volume, moles, and temperature)

Graham’s Law of Effusion

Rate₁/Rate₂ = √(M₂/M₁); lighter gases effuse faster

Dalton’s Law of Partial Pressures

P(total) = P₁ + P₂ + P₃ + …

Van der Waals Equation

Corrects ideal gas law for real gases (accounts for molecular volume and attractions)

Pressure conversions

1 atm = 101.325 kPa = 101,325 Pa = 760 torr = 760 mmHg

STP meaning

Standard Temperature and Pressure

STP numeric values

0°C (273 K) and 1 atm (101.3 kPa)

Kinetic Molecular Theory

Gases are in constant motion; collisions are elastic; average kinetic energy ∝ temperature

Universal gas constant (R) values

0.0821 L·atm/(mol·K) and 8.314 J/(mol·K)

Why two R values

They depend on the pressure unit used (atm vs. Pa)

Ideal vs. real gases

Real gases deviate from ideal behavior at high pressure and low temperature

Why real gases deviate

Intermolecular forces and molecular volume become significant

Effusion

Movement of gas through a tiny hole without collisions

Diffusion

Spreading or mixing of gases due to random motion

Absolute zero

0 K (–273.15°C); temperature at which molecular motion theoretically stops

Why absolute zero matters

It’s the lowest possible temperature; defines Kelvin scale

When 1 mol = 22.4 L

At STP, one mole of an ideal gas occupies 22.4 L

Gas stoichiometry use

Can determine gas volume or moles using PV = nRT and balanced equations

Using gas stoichiometry with density

Can calculate molar mass of an unknown gas using density and gas laws