Chemistry Lecture Notes

Introduction to Chemistry

• Focus on basic topics needed for a chemistry course

Periodic Table Overview

Elements and Groups

• Group 1A (Alkali Metals): Hydrogen (H), Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cesium (Cs)
  • H is a non-metal; others are metals
  • Highly reactive; react violently with water (e.g., Na + H2O)
  • Have 1 valence electron; form ions with +1 charge
• Group 2A (Alkaline Earth Metals): Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba)
  • Reacts but less violently than alkali metals
  • Have 2 valence electrons; form ions with +2 charge

Transition Metals

• Includes Scandium (Sc), Titanium (Ti), Vanadium (V), etc.
  • Variable charges (e.g., Iron can be +2 or +3)
  • Some have consistent charges (e.g., Zinc is +2)

Representative Elements (Groups 13-18)

• Group 3A (Boron Family): Boron (B), Aluminum (Al), Gallium (Ga), etc.
  • 3 valence electrons; tend to form +3 ions
• Group 4A (Carbon Family): Carbon (C), Silicon (Si), etc.
  • 4 valence electrons; can form +4 or +2 ions
• Group 5A (Nitrogen Family): Nitrogen (N), Phosphorus (P), etc.
  • 5 valence electrons; tend to form -3 ions
• Group 6A (Chalcogens): Oxygen (O), Sulfur (S), etc.
  • 6 valence electrons; form -2 ions
• Group 7A (Halogens): Fluorine (F), Chlorine (Cl), etc.
  • 7 valence electrons; form -1 ions
• Group 8A (Noble Gases): Helium (He), Neon (Ne), etc.
  • Chemically inert, stable with 8 valence electrons (except He with 2)

Metals, Nonmetals, and Metalloids

• Metals are on the left side, nonmetals on the right, metalloids in between (e.g., Silicon and Germanium)
  • Metals are conductors, malleable, and ductile
  • Nonmetals are insulators; do not conduct electricity
  • Metalloids have properties between metals and nonmetals

Diatomic Elements

• H2, N2, O2, F2, Cl2, Br2, I2 (exist naturally as diatomic molecules)

Chemical Bonds

Ionic Bonds

• Formed through the transfer of electrons from metals to nonmetals
• Example: Sodium (Na) loses an electron to Chlorine (Cl)
  • Na+ and Cl- attract to form NaCl

Covalent Bonds

• Formed through sharing of electrons between nonmetals
• Example: H2 (hydrogen molecules)
  • Polar vs. Non-Polar bonds based on electronegativity differences

Atomic Structure

• Atomic Number: Number of protons in an atom
• Mass Number: Total number of protons and neutrons
• Electrons in neutral atoms equal the number of protons

Isotopes and Average Atomic Mass

• Isotopes have the same number of protons but different numbers of neutrons
• Average atomic mass is weighted based on isotopes
• Example: Carbon isotopes (C-12, C-13)

Chemical Reactions

Types of Reactions

• Combustion Reactions: Hydrocarbon + O2 → CO2 + H2O
• Single Replacement Reactions: A + BC → AC + B
• Double Replacement Reactions: AB + CD → AD + CB

Balancing Chemical Equations

• Ensure the number of atoms on reactants equals number on products
• Adjust coefficients, not subscripts, to balance

Redox Reactions

• Reduction and oxidation occur, transfer of electrons
• Reducing agent is oxidized, oxidizing agent is reduced

Stoichiometry

• Relationships between reactants and products in a chemical reaction
• Conversions between moles, grams, atoms/molecules
• Molar mass used to convert grams to moles and vice versa
  • Example: 1 mole of Carbon (C) = 12 g

Solutions and Concentrations

• Molarity (M) = moles of solute / liters of solution

Acid-Base Reactions

• Neutralization: Acid + Base → Salt + Water
  • Example: HCl + NaOH → NaCl + H2O

Summary

• Review periodic table, bonding, atomic structure, and types of reactions
• Importance of balancing equations, understanding stoichiometry, and identifying reaction types.

Recommended Study Strategies

• Memorize periodic table elements and their charges
• Practice balancing chemical equations
• Understand and apply concepts of molar mass and stoichiometry
• Familiarize with common polyatomic ions and their charges.