Honors Chemistry Unit 5

Page 1

Page 2: Unit 5: Bonding and Nomenclature Content Outline: Chemical Bonds (5.1)

I. Chemical Bonds

  • Definition: Mutual attractions between atom's valence shell electrons resulting in new molecules.

  • Bonds can affect the electrical charge of an atom or molecule.

  • Stability: Promotes atomic or molecular stability, mirroring Noble gas stability.

    • Potential Energy (PE) decreases in a bond; more stable atoms mean less PE.

    • Energy is released when bonds are formed and required for bonds to break.

II. Types of Chemical Bonds

  1. Covalent Bonds

    • Electrons are shared, leading to stability as atoms behave like Noble gases.

    • Molecule Types:

      • Non-polar Molecules: Equal distribution of electrons; electronegativity difference: 0 to 0.3.

      • Polar Molecules: Unequal distribution of electrons; one end positive (δ+) and one end negative (δ-); electronegativity difference: 0.3 to 1.7.

        • Example: Water, crucial for supporting life.

    • Covalent bonds are generally stronger than other bonds but not always.

  2. Ionic Bonds

    • Formed from electrical charge attractions between oppositely charged atoms after electron transfer.

      • The atom losing an electron: oxidation (cation).

      • The atom gaining an electron: reduction (anion).

    • Typically stronger than covalent bonds but can vary.

    • Electronegativity difference: > 1.7.

  3. Metallic Bonds

    • Involves electrically charged metals with mobile electrons, forming a ‘sea of electrons.’

    • Imparts conductivity, luster, malleability, and ductility to metals.

    • Strength depends on nuclear charge and valence electrons.

    • Enthalpy of Vaporization: Energy required to convert a liquid to a gas helps measure bond strength.

Page 3

  • Continued from Page 2 regarding metallic bonds.

Page 4

  • Reiterates properties of metallic bonds, discussing electrons' mobility and their implications for conductivity and malleability.

Page 5: Unit 5: Bonding and Nomenclature Content Outline: Drawing Molecules (5.2)

I. Structural Formula

  • Represents atom types, quantity, and arrangement of atoms, but not all valence shell electrons.

    • Example: H2O as H-O-H.

II. Electron Dot / Lewis Dot Notation

  • Interactive resource available for drawing Lewis structures.

  • Shows all interacting (bonded) and non-interacting (lone pairs) electrons.

    • Paired electrons represented by dots or dashes (X-Y).

    • Double and triple bonds shown by X=Y and X≡Y.

  • Exceptions to the Octet Rule: Hydrogen (max 2), Helium (max 2), Lithium & Beryllium (max 2, 4 respectively); Boron (fills at 6).

III. Resonance Structures

  • Molecules or ions that can't be depicted by a single Lewis Dot Structure (e.g., Ozone).

IV. Polyatomic Ions

  • Groups of atoms bonded with molecular and ionic characteristics, forming ionic compounds.

  • Charges result from electron clumping (negative) or absence (positive).

  • Examples: NH4+, SO4 2-, PO4 3-.

Page 6

  • Further explanation of polyatomic ions and their Lewis Dot structures.

Page 7: Unit 5: Bonding and Nomenclature Content Outline: Molecular Geometry and VSEPR Theory (5.3)

I. Structure = Function

  • 3D shapes determine chemical properties like polarity and reactivity.

II. VSEPR Theory

  • Valence-Shell, Electron-Pair Repulsion: Electron pairs maximally separate around the central atom.

    • Lone pairs exert more repulsion than shared pairs.

  • Bond Types:

    • Sigma (σ) Bonds: Primary bonds from overlapped orbitals.

    • Pi (π) Bonds: Secondary bonds from overlapping p orbitals.

III. Molecular Shapes

  1. Linear: Bond angle 180° (AB2).

  2. Trigonal-Planar: Bond angle 120° (AB3).

  3. Bent/Angular: Bond angle > 120° (AB2E).

  4. Tetrahedral: Bond angle 109.5° (AB4).

  5. Trigonal-Pyramidal: Bond angle 107° (AB3E).

  6. Trigonal-Bipyramidal: Bond angles 90° and 120° (AB5).

  7. Octahedral: Bond angles 90° (AB6).

Page 8

  • Contains additional details and visual aids on molecular shapes and hybridization.

Page 9: Unit 5: Bonding and Nomenclature Content Outline: Intermolecular Forces (5.4)

I. Intermolecular Forces

  • Attractions between molecules in solids, liquids, and gases; generally weaker than chemical bonds.

  • Strength influences boiling points: higher interactions lead to higher boiling points.

II. Types of Intermolecular Forces

  1. Network Solid: 3D structures, e.g., diamond and graphite.

  2. Hydrogen Bonds: Attractions in polar molecules involving H, O, N, or F.

  • Essential for DNA structure and water properties.

  1. Dipole-Dipole Forces: Occurs between polar molecules.

  2. Induced Dipole Moment: Non-polar molecules behaving like polar temporarily, e.g., O2 in water.

  3. London Dispersion Forces: Temporary attractions existing even in noble gases due to fluctuating electron density.