CHM 101 - Dr. Vic

Chemical Equations and Stoichiometry

• Definition: A chemical equation is a shorthand representation of a chemical reaction using chemical symbols and formulae, encompassing both qualitative and quantitative information about reactants and products.

Writing and Balancing Chemical Equations

• Example Reaction: Methane (CH4) reacts with oxygen (O2) to yield carbon dioxide (CO2) and water (H2O).

  • Reaction: CH4 + 2O2 ⟶ CO2 + 2H2O

  • Mole Ratio: 1:2:1:2 (for every 1 methane, 2 oxygen react to produce 1 carbon dioxide and 2 water molecules).

Balancing Equations

• A balanced equation must reflect equal numbers of each type of atom on both sides, adhering to the law of conservation of matter.

  • Method: Count the number of atoms for each element on both sides and adjust coefficients as necessary.

Fundamental Aspects of Chemical Equations

1. Reactants and Products:

   • Reactants are substances on the left side, while products are on the right.

2. Separation Symbols:

   • Plus signs (+) denote different reactants/products and an arrow (⟶) indicates the direction of the reaction.

3. Coefficients:

   • Numbers before formulas indicate the quantity of reactants and products (e.g., 2Na, 2H2O).

4. Physical States:

   • Indicated by abbreviations (s = solid, l = liquid, g = gas, aq = aqueous solution).

5. Special Conditions:

   • Indicated by symbols above or below the reaction arrow (e.g., Δ for heat).

• Example: 2Na(s) + 2H2O(l) ⟶ 2NaOH(aq) + H2(g)

Reaction Stoichiometry

• Definition: Stoichiometry involves calculating the quantitative relationships of reactants and products in a chemical reaction based on a balanced equation.

• Coefficients represent the mole ratios needed for calculations.

Types of Stoichiometric Calculations

1. Mass-Mass Relationship:

   • Steps:

     • Write a balanced equation.

     • Note the moles of each ingredient.

     • Use unit conversion to find unknowns.

   • Example 1: Calculate grams of oxygen needed to burn octane (C8H18).

2. Mass-Volume Relationship:

   • Utilize molar volume of gas at NTP (22.4 L per mole).

   • Example 2: Calculate the volume of CO2 produced from heating calcium carbonate.

3. Volume-Volume Relationship:

   • Related to Avogadro's and Gay-Lussac's laws pertaining to gas reactions under similar conditions.

Chemical Bonding and Intermolecular Forces

Introduction to Chemical Bonds

• Definition: A chemical bond is a force that holds atoms together in molecules or ions.

• Categories: Intramolecular (ionic, covalent, metallic) and intermolecular (hydrogen bonds, van der Waals forces).

Intramolecular Bonds

1. Ionic Bonding:

   • Electrostatic attraction between cations and anions formed through electron transfer.

   • Metals lose electrons, forming positive ions; non-metals gain electrons, forming negative ions.

   • Example: NaCl formation involves Na losing an electron to Cl.

2. Covalent Bonding:

   • Atoms share electrons to achieve stable configurations.

   • Examples: H2, F2, and HF formations demonstrate electron sharing.

3. Coordinate Bonding:

   • Formed when one atom provides both electrons for bonding.

   • Example: Formation of ammonium ion (NH4+).

Properties of Ionic Compounds

• Crystalline structure, high melting/boiling points, conduct electricity when dissolved or molten.

Properties of Covalent Compounds

• Exist as molecules, lower melting/boiling points than ionic compounds, poor electrical conductors, and variable solubility in solvents.

Metals and Metallic Bonding

• Metallic bonds involve a lattice of metal ions surrounded by delocalized electrons, resulting in electrical conductivity and malleability.

Intermolecular Forces

• Importance: Weaker forces compared to intramolecular; crucial for understanding physical properties.

  • Types:

    1. Hydrogen Bonding: Attraction involving H and highly electronegative atoms (N, O, F).

    2. Van der Waals Forces: Weak attractions between molecules.

  • Applications of Hydrogen Bonds:

    • Critical in water properties and biological molecule interactions.

  • Applications of Van der Waals Forces:

    • Responsible for condensation of gases and hardness differences in solids.

Kinetic Theory of Matter

Overview

• All matter consists of tiny particles in constant random motion, leading to kinetic energy.

Kinetic Theory Assumptions

1. Continuous movement of particles.

2. Negligible particle volume in gases.

3. Minimal intermolecular forces.

4. Elastic collisions between particles.

5. Average kinetic energy is directly proportional to temperature.

States of Matter

• Gases: No fixed shape or volume, can be compressed.

• Liquids: Fixed volume, take the shape of their container.

• Solids: Fixed shape and volume, particles are closely packed.