Chemistry Lecture Notes

Key Study Areas in Chemistry

• Chemistry studies matter, its properties, composition, and structure.
• It examines the changes matter undergoes and the energy involved in these processes.

Intermolecular Forces in Different States of Matter

• Solids: Strong intermolecular forces, particles are closely packed.
• Liquids: Moderate intermolecular forces, particles can slide past each other.
• Gases: Weak intermolecular forces, particles are far apart and move freely.

Crystal Structure of Ionic Substances

• Ionic substances form a crystal lattice structure.
• This structure is characterized by a repeating pattern of positive and negative ions.

Polarity of Water Molecules

• Water molecules are polar due to the difference in electronegativity between hydrogen and oxygen.
• This causes a partial negative charge on the oxygen and a partial positive charge on the hydrogens.

Formation of Precipitates in Ionic Reactions

• A precipitate forms when two ionic compounds in solution combine to form an insoluble solid.
• Example: Combining Ba(NO3)2 and Na2SO4 forms BaSO4 as a precipitate.

General Solubility Guidelines

1. Sodium, potassium, and ammonium compounds are generally soluble.
2. Nitrates, acetates, and chlorates are typically soluble.
3. Chlorides are soluble except for those of silver, mercury(I), and lead.
4. Sulfates are soluble except for barium, strontium, and lead.
5. Carbonates, phosphates, and silicates are generally insoluble.
6. Sulfides are mostly insoluble except for calcium, strontium, sodium, potassium, and ammonium.

Molarity Calculations

• Molarity (M) = moles of solute/liters of solution
• Example: Molarity of 0.202 mol KCl in 7.98 L = 0.0253 M
• Example: Molarity of 125 g NaCl in 4.00 L = (mass/molar mass) / volume

Writing Chemical Equations

• Complete reaction equations include all reactants and products.
• Involves balancing to ensure the law of conservation of mass.
• Hydrogen gas is represented as H2 in equations.
• The number in front of a formula is the coefficient.

Law of Conservation of Mass

• States that matter cannot be created or destroyed.
• Mass of reactants equals mass of products in a chemical reaction.

Types of Chemical Reactions

1. Synthesis: Two or more substances combine to form a new compound.
2. Decomposition: A single compound breaks down into two or more simpler substances.
3. Single Replacement: One element replaces another in a compound.
4. Double Replacement: Ions in two compounds exchange places.
5. Combustion: A substance combines with oxygen, releasing energy.

Yield Calculations

• Actual yield: The measured amount of product obtained from a reaction.
• Theoretical yield: The maximum amount of product expected based on stoichiometry.
• Percent yield: (Actual yield/Theoretical yield) x 100

Bonding and Electron Configuration

• Electrons involved in bonds are called valence electrons.
• Atoms become more stable when they bond.
• Metallic Bonds: Sea of electrons, good conductivity.
• Ionic Bonds: Transfer of electrons, strong and rigid.
• Covalent Bonds: Sharing of electrons, can form polar or non-polar molecules.

Lewis Structures

• Nonmetals, except hydrogen, are surrounded by eight electrons to satisfy the octet rule.
• Central atom is usually the least electronegative.
• Includes bonding and non-bonding electrons.

Properties of Matter

• Ductility: Ability to be drawn into wires.
• Malleability: Ability to be hammered or rolled into sheets.
• Conductivity: Ability to conduct electricity or heat.

States of Matter

• Solids: High density, low energy, fixed shape.
• Liquids: Mid-density, flow, take shape of container.
• Gases: Low density, high energy, fill the container.

Measurement Units in Chemistry

• Mass: Kilograms (kg)
• Volume: Liters (L)
• Density: Kilograms per cubic meter (kg/m³)
• Length: Meters (m)

Elements and Isotopes

• Use atomic symbols and numbers to represent elements and isotopes.
• Example:
  • Potassium-40 (K) with atomic number 19 and mass number 40.
  • Helium-4 (He) with atomic number 2 and mass number 4.