Chemistry unit 2 book notes.

Bonding: General Concepts

Types of Chemical Bonds

  • Chemical Bond: Force holding atoms together to form compounds.

    • Covalent Bonds: Atoms share electrons.

    • Ionic Bonds: Formed when an atom donates an electron to another atom.

  • Molecules: A collection of atoms.

    • Representations include:

    • Chemical Formula: e.g., CO2

    • Structural Formula: e.g., H—O—H

    • Space-filling Model: Shows relative sizes/orientation.

    • Ball and Stick Model: 3D representation.

Attributes of Molecules

  • Physical Properties:

    • Melting Point

    • Hardness

    • Electrical & Thermal Conductivity

    • Solubility

    • Electric Charge

  • Bond Energy: Energy required to break a bond.

Types of Chemical Bonding

  • Ionic Bonding:

    • Involves atoms that easily lose electrons (metals) and atoms with a high affinity for electrons (non-metals).

    • Example: formation of ionic compounds like NaCl.

  • Bond Length: Distance between two atoms at minimum energy.

    • As atoms get closer, energy rises due to repulsive forces.

Coulomb’s Law

  • Describes interaction energy between ions:

    • E = k * (Q1 * Q2) / r

    • E: Interaction energy (Joules)

    • r: Distance between ions (nm)

    • Q1 & Q2: Charges of the ions.

Electronegativity

  • Ability of an atom in a molecule to attract shared electrons towards itself.

    • Pauling Scale: Assigns values to elements based on electron-attraction ability.

    • Trends: Increase across a period, decrease down a group in the periodic table.

    • Examples to compare bond polarities:

    • H—H < S—H < Cl—H < O—H < F—H

Polar Covalent Bonds

  • Occur when electrons are unequally shared, resulting in partial charges (e.g., HCl).

    • Partial Positive & Negative Charges arise with unequal sharing.

Lewis Structures

  • Show how valence electrons are arranged.

    • Steps to write:

    1. Sum valence electrons.

    2. Form bonds between atoms with pairs of electrons.

    3. Distribute remaining electrons to complete octets (or duets for H).

    • Octet Rule: Atoms gain stability with full outer electron shells.

Common Exceptions to the Octet Rule

  • Boron Compounds: Often have fewer than 8 electrons around boron.

  • Expanded Octets: Elements in Period 3 or higher can hold more than 8 electrons (e.g., SF6).

  • Odd-Electron Molecules: E.g., NO states that some molecules have odd numbers of total valence electrons.

Resonance Structures

  • Used when multiple valid Lewis structures can represent a molecule (e.g., NO3−).

    • Depicted with double-headed arrows.

Formal Charge

  • Used to assess the distribution of electrons in structures:

    • Defined as: Formal Charge = (Valence Electrons of the Atom) - (Assigned Electrons in the Lewis Structure)

  • Guidelines:

    • Aim for formal charges of zero on all atoms for stability.

    • If not possible, negative charges should reside on the more electronegative elements.

Naming Simple Compounds

  • Binary Ionic Compounds (Type I): Named with a cation followed by an anion (e.g., NaCl -> Sodium Chloride).

  • Type II Compounds: Transition metals with variable charges use Roman numerals to indicate charge (e.g., FeCl2 -> Iron(II) Chloride).

  • Binary Covalent Compounds (Type III): Compounds composed of nonmetals use prefixes to denote the number of atoms (e.g., CO2 -> Carbon Dioxide).

  • Acids: Their naming depends on the anion!

    • Ends in -ide (e.g., Cl− -> HCl = Hydrochloric acid)

    • Ends in -ate or -ite can be modified as well.

This comprehensive overview covers concepts from chemical bonding to naming compounds, providing a strong foundation for further study in chemistry.