Comprehensive Study Guide: Core Concepts in General Chemistry

Chemical Equations and Stoichiometry

Writing and Balancing Chemical Equations

Chemical equations represent chemical reactions, showing the reactants and products with their respective quantities. Balancing ensures the law of conservation of mass is obeyed.

• Balanced Equation: The number of atoms of each element is the same on both sides of the equation.

• Steps: Identify reactants and products, write their formulas, and adjust coefficients to balance atoms.

• Example: $2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$

Types of Chemical Reactions

• Synthesis (Combination): Two or more substances combine to form one product.

• Decomposition: A single compound breaks down into two or more simpler substances.

• Single Replacement: One element replaces another in a compound.

• Double Replacement: Exchange of ions between two compounds.

• Combustion: A substance reacts with oxygen, releasing energy, usually as heat and light.

The Mole and Avogadro’s Number

• Mole: The amount of substance containing $6.022 \times 10^{23}$ entities (Avogadro’s number).

• Avogadro’s Number: $6.022 \times 10^{23}$ particles/mol.

Mass, Moles, and Molecular Weight

• Molecular Weight (Molar Mass): The mass of one mole of a substance, in grams per mole (g/mol).

• Conversions:

  • Grams to moles: $\text{moles} = \frac{\text{grams}}{\text{molar mass}}$

  • Moles to grams: $\text{grams} = \text{moles} \times \text{molar mass}$

• Stoichiometry: Use balanced equations to convert between moles of reactants and products.

• Example: Given $4$g of$\text{H}_2$, how many grams of$\text{H}_2\text{O}$ are produced?

Percent Yield

• Percent Yield: $\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$

Energy, Enthalpy, and Reaction Rates

Energy and Its Units

• Energy: The capacity to do work or produce heat. Units: Joules (J), calories (cal).

Bond Dissociation Energy and Enthalpy

• Bond Dissociation Energy: Energy required to break a bond in a molecule.

• Enthalpy ($\Delta H$): The heat change at constant pressure.

• Signs of Enthalpy:

  • Positive $\Delta H$ (endothermic): Heat absorbed.

  • Negative $\Delta H$ (exothermic): Heat released.

Reaction Rates and Catalysis

• Reaction Rate: The speed at which reactants are converted to products.

• Catalyst: Substance that increases reaction rate without being consumed.

Chemical Equilibrium

Equilibrium Constant

• Expression: For $aA + bB \rightleftharpoons cC + dD$, $K = \frac{[C]^c[D]^d}{[A]^a[B]^b}$

• Interpretation:

  • $K \gg 1$: Products favored.

  • $K \ll 1$: Reactants favored.

States of Matter and Gas Laws

States of Matter

• Solid: Definite shape and volume.

• Liquid: Definite volume, takes shape of container.

• Gas: No definite shape or volume.

Gas Laws

• Boyle’s Law: $P_1V_1 = P_2V_2$ (constant T, n)

• Charles’s Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ (constant P, n)

• Combined Gas Law: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$

• Ideal Gas Law: $PV = nRT$

Intermolecular Forces and Properties of Liquids/Solids

Types of Intermolecular Forces

• London Dispersion: Weak, present in all molecules.

• Dipole-Dipole: Between polar molecules.

• Hydrogen Bonding: Strong, between H and N, O, or F.

Effects on Physical Properties

• Stronger forces increase boiling/melting points, viscosity, and surface tension; decrease vapor pressure.

Heating Curves and Specific Heat

• Specific Heat ($c$): Amount of heat to raise 1 g of substance by 1°C.

• Heat Calculation: $q = mc\Delta T$

• Heating Curve: Shows temperature changes and phase transitions as heat is added.

Solutions and Their Properties

Definitions

• Solution: Homogeneous mixture of solute and solvent.

• Solute: Substance dissolved.

• Solvent: Substance doing the dissolving.

• Mixture: Physical blend of two or more substances.

Electrolytes and Equivalents

• Electrolyte: Substance that conducts electricity when dissolved in water.

• Equivalents: Amount of ion that carries 1 mole of charge.

Solubility Principles

• Solubility: Maximum amount of solute that dissolves in a solvent at a given temperature.

• Factors Affecting Solubility: Temperature, pressure, nature of solute/solvent.

Concentration Calculations

• Weight/Volume % (w/v): $\frac{\text{grams solute}}{\text{mL solution}} \times 100\%$

• Volume/Volume % (v/v): $\frac{\text{mL solute}}{\text{mL solution}} \times 100\%$

• Molarity (M): $\frac{\text{moles solute}}{\text{L solution}}$

• Dilution: $M_1V_1 = M_2V_2$

Colligative Properties, Osmosis, and Diffusion

• Colligative Properties: Depend on number of particles, not identity (e.g., boiling point elevation, freezing point depression).

• Osmosis: Movement of solvent through a semipermeable membrane from low to high solute concentration.

• Diffusion: Movement of particles from high to low concentration.

Acids, Bases, and Buffers

Definitions

• Acid: Donates H+ ions in solution.

• Base: Accepts H+ ions or donates OH- ions.

Conjugate Acid-Base Pairs

• Conjugate Acid: Formed when a base gains H+.

• Conjugate Base: Formed when an acid loses H+.

Neutralization Reactions

• General Form: Acid + Base → Salt + Water

• Example: $\text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O}$

Acid/Base Strength and pH

• Strength: Strong acids/bases dissociate completely; weak only partially.

• pH Scale: Measures acidity/basicity; $\text{pH} = -\log[\text{H}^+]$

• Calculations:

  • $[\text{H}^+] = 10^{-\text{pH}}$

  • $[\text{OH}^-] = 10^{-\text{pOH}}$

  • $\text{pOH} = 14 - \text{pH}$

Buffers

• Buffer: Solution of a weak acid and its conjugate base (or weak base and conjugate acid) that resists changes in pH.

• Function: Maintains pH upon addition of small amounts of acid or base.