Honors Chemistry Unit 5 Test Review

Section 1: Covalent and Ionic Bonds

  • Covalent Bond:
    • Definition: A covalent bond is a type of chemical bond where nonmetals or metalloids share electron pairs (e's) between them.
    • Comparison with Ionic Bond: Ionic bonds involve the transfer of electrons, typically between metals and nonmetals, resulting in the formation of cations and anions.

Section 2: Molecules and Molecular Compounds

  • Molecule:

    • Definition: A neutral group of atoms joined by covalent bonds.

  • Molecular Compound:

    • Definition: A compound that is composed of molecules, showcasing a specific chemical formula that indicates the types and numbers of atoms present in the compound.

  • Molecular Formula:

    • Definition: The chemical formula of a molecular compound that displays the number of each kind of atom within a molecule.

Section 3: Structural vs. Molecular Formula

  • Structural Formula (Lewis Structure):

    • Definition: Represents the bonding and general shape of the molecule. It reflects how the atoms within the molecule are connected and the arrangement of the electrons participating in bonding.

  • Difference from Molecular Formula:

    • A molecular formula shows only the kinds and amounts of atoms present, while a structural formula elucidates the bonding and spatial arrangement of those atoms.

Section 4: Representative Units in Compounds

  • Representative Units of a Covalent Compound:

    • Molecule (specific to covalent compounds).

  • Representative Units of an Ionic Compound:

    • Formula unit (specific to ionic compounds).

Section 5: Diatomic Molecules

  • Diatomic Molecule:
    • Definition: A molecule that contains two atoms that are typically found together in nature.
    • Examples: H₂, O₂, N₂, F₂, Cl₂, Br₂, I₂.

Section 6: Characteristics of Compounds

  • Ionic Compounds:

    • Generally very specific in structure and composition.

  • Covalent Compounds:

    • Exhibit properties that are often the opposite of ionic compounds, such as lower melting and boiling points, and they can exist in all three states of matter (solid, liquid, gas).

Section 7: Octet Rule

  • Octet Rule:
    • Principle stating that atoms tend to bond in a manner that allows them to achieve the electron configuration of a noble gas, often resulting in eight valence electrons.

Section 8: Unshared Electron Pairs

  • Definition:
    • Unshared electron pairs are pairs of electrons in the outer shell of an atom that are not involved in bonding, but are part of the molecular orbitals.
    • Example: F₂ (since each fluorine atom has three unshared pairs and one shared pair in the bond).

Section 9: Types of Bonds

  • Single Bond:

    • Description: Involves the sharing of 2 electrons.
    • Example: Breaking down CH₄ (methane).

  • Double Bond:

    • Description: Involves sharing of 4 electrons.
    • Example: O₂ (oxygen gas).

  • Triple Bond:

    • Description: Involves sharing of 6 electrons.
    • Example: N₂ (nitrogen gas).

Section 10: Lewis Structures

  • Drawing Lewis Structures:

    • Ability to draw Lewis structures for various molecules is essential.

    • Examples:

    • CH₄:

      • H
      • |
      • H - C - H
      • |
      • H

    • NH₃:

      • H
      • |
      • H - N - H

    • CO₂:

      • O = C = O

Section 11: Coordinate Covalent Bonds

  • Definition:
    • A bond where one atom provides both bonding electrons.
    • Example: In ammonium ion (NH₄⁺), nitrogen donates a pair of electrons to form a bond with hydrogen.

Section 12: Exceptions to the Octet Rule

  • Examples:
    • Nitric oxide (NO) has an odd number of electrons.
    • Sulfur tetrafluoride (SF₄) features an expanded octet.

Section 13: Bond Dissociation Energy

  • Definition:
    • The energy required to break a bond, measured in KJ/mol.
    • Higher bond dissociation energy indicates a stronger bond.

Section 14: Resonance Structures

  • Definition:
    • Different ways to represent the same molecule with the same arrangement of atoms, displaying delocalized electrons.
    • Examples: O₃, NO₂.

Section 15: Orbital Theory in Bonding

  • Types of Orbitals:
    • Atomic Orbital, Molecular Orbital, Bonding Orbital.

Section 16: Sigma and Pi Bonds

  • Sigma Bonds:

    • Formed by the end-to-end overlap of atomic orbitals.

  • Pi Bonds:

    • Occur when orbitals overlap side-by-side, creating a bond above and below the axis of the bonded atoms.

Section 17: VSEPR Theory

  • Definition:
    • Valence Shell Electron Repulsion (VSEPR) theory states that electron pairs around a central atom will arrange themselves as far apart from each other as possible to minimize repulsion.

Section 18: Hybridization

  • Definition:
    • The blending of atomic orbitals to create new hybrid orbitals.
    • Types include: s, sp, sp², sp³, dsp³, d²sp³.

Section 19: Molecular Geometry

  • Main Shapes and Bond Angles:
    • Tetrahedral: 109.5° - Example: CH₄.
    • Trigonal Pyramidal: 107° - Example: NH₃.
    • Trigonal Planar: 120° - Example: CO₂.
    • Linear: 180° - Example: O=O.
    • Bent: 104.5° - Example: H₂O.

Section 20: Polar vs. Nonpolar Bonds

  • Nonpolar Covalent Bond:

    • Definition: A bond where the electrons are shared equally between the two atoms.

  • Polar Covalent Bond:

    • Definition: A bond where the electrons are shared unequally, creating slight charges (dipoles).

  • Determining Bond Type by Electronegativity:

    • Calculate difference between the electronegativity (Pauling scale) of involved elements.
    • 0 to 0.4 = Nonpolar Covalent Bond
    • 0.5 to 1.8 = Polar Covalent Bond
    • Presence of metal = Ionic Bond.

Section 21: Dipole Moment and Polar Molecules

  • Polar Molecule (Dipole):
    • Definition: A molecule that has regions of positive and negative charges due to unequal sharing of electrons.
    • Example: CH₄ is nonpolar, while CH₃Cl is polar due to its molecular shape and electronegativities involved.

Section 22: Hybridization and Molecular Shape

  • Determining Hybridization and Bond Angle:
    • Able to analyze given molecules’ hybridization, bond angles, and geometric shapes based on electron pairs and bond types.

Section 23: Ionic Bonds and Compounds

  • Ionic Bonds:
    • Definition: Formed through electrostatic attraction between positive (cations) and negative ions (anions).
    • Examples of ionic compounds include:
    1. Na⁺ and Br⁻ → NaBr
    2. Mg²⁺ and Br⁻ → MgBr₂
    3. Mg²⁺ and S²⁻ → MgS.

Section 24: Naming Ionic Compounds

  • Chemical formula writing for ionic compounds formed by pairs of ions:
    • K⁺ and Cl⁻ → KCl
    • Mg²⁺ and Cl⁻ → MgCl₂
    • Mg²⁺ and Se²⁻ → MgSe.

Section 25: Naming Covalent Compounds

  • Naming examples of covalent compounds include:
      • N₂O₂ → Dinitrogen dioxide
    • PCl₅ → Phosphorous pentachloride
    • NO → Nitrogen monoxide
    • CO₂ → Carbon dioxide
    • CO → Carbon monoxide
    • NH₃ → Nitrogen trihydride (ammonia).

Section 26: Valence Electrons

  • Definition:
    • Valence electrons are the electrons in the highest occupied energy level of an atom, significant for bonding.

Section 27: Determining Valence Electrons

  • Method:
    • Ability to determine the valence electrons for any representative element from its group on the periodic table.

Section 28: Properties of Ionic Compounds

  • Characteristics of Ionic Compounds:
    1. Solid state at room temperature with high melting points.
    2. Crystalline structure, typically brittle.
    3. Conduct electricity when melted or dissolved in water.

Section 29: Metallic Bonds

  • Definition:
    • A metallic bond is a type of bond formed through the attraction between free-floating valence electrons and positively charged metal cations (often referred to as a