Covalent bonding

Chapter 8: Covalent Bonding

I. Covalent Bonds

  • Definition: A covalent bond is a chemical bond resulting from the sharing of valence electrons, typically between nonmetals.

A. Types of Covalent Bonds

  1. Single Bond:

    • Comprises one shared pair of electrons between atoms.

    • Example:

  2. Double Bond:

    • Comprises two shared pairs of electrons between atoms.

    • Example:

  3. Triple Bond:

    • Comprises three shared pairs of electrons between atoms.

    • Example:

B. Molecules

  • Definition: A molecule is a neutral group of covalently bonded atoms.

  • Examples:

    • H2, H2O, H2SO4.

  • Diatomic Molecules: Seven elements occur as diatomic molecules in nature; students should memorize this list.

C. Chemical Formula

  • Also known as the molecular formula, it uses element symbols and subscripts to indicate the number of atoms in a molecule.

D. Polyatomic Ions

  • Defined as a charged group of covalently bonded atoms.

    • Example: OH- in NaOH.

  • Important Note: Polyatomic ions differ from molecules and are named distinctly; polyatomic nomenclature should not be used for molecules.

II. Properties of Molecules

  1. Usually exist as liquids or gases at room temperature.

  2. Generally have low melting points and boiling points.

  3. Electrolytic Properties:

    • Strong acids/bases are strong electrolytes (strong conductors of electricity).

    • Weak acids/bases are weak electrolytes (weak conductivity).

    • All other molecules are non-electrolytes (do not conduct electricity).

A. Nomenclature and Formulas of Molecular Compounds

  1. Organic Compounds:

    • Compounds containing carbon bonded to hydrogen.

    • Examples include millions of organic compounds due to carbon’s ability to bond with itself, forming long chains or rings.

    • Focus on straight-chain alkanes and cycloalkanes.

    • Straight-Chain Alkane: Hydrocarbon with only single bonds; example: C2H6 is ethane.

    • Naming of Straight-Chain Alkanes:

      • Base Names:

        • 1 Carbon: meth

        • 2 Carbons: eth

        • 3 Carbons: prop

        • 4 Carbons: but

        • 5 Carbons: pent

        • 6 Carbons: hex

        • 7 Carbons: hept

        • 8 Carbons: oct

        • 9 Carbons: non

        • 10 Carbons: dec

      • Suffix: Indicates bonding. If single bonds only, the suffix is -ane.

III. Naming and Structural Formulas

A. Chemical or Molecular Formula for Straight-Chain Alkanes

  • General formula: CnH2n+2 (where n = number of carbons).

  • Examples:

    • Propane: C3H8

    • Decane: C10H22

B. Structural Formulas

  • Depicts which atoms are covalently bonded using symbols and lines.

  • To draw: Connect carbon symbols with lines representing single bonds and attach hydrogen symbols to achieve four shared electron pairs per carbon.

  • Structural formula of Methane:

  • Example Drawings: Propane and Pentane.

IV. Cycloalkanes

  1. Definition: Cycloalkanes are hydrocarbons with only single bonds forming a ring.

    • Example: C4H8 is cyclobutane.

    • Nomenclature: Same as straight-chain alkanes but with the prefix "cyclo."

    • Example Naming: C3H6.

  2. Molecular Formula: For cycloalkanes, CnH2n.

    • Example: What is the molecular formula for cyclohexane?

  3. Structural Formulas: Can include full or condensed structural formulas.

    • Example Drawings: Cyclopropane and Cyclopentane.

V. Inorganic Compounds

  1. Definition: Compounds that do not contain carbon bonded to hydrogen.

    • Examples: H2O, CO, NaBr, SF6, HCl.

  2. Acids: Produce hydrogen ions (H+) in water; formulas often start with hydrogen.

    • Examples: HCl (hydrochloric acid), HNO3 (nitric acid); note that water does not fit this definition.

VI. Nomenclature of Acids

  1. Oxyacids: Contain oxyanions.

    • If the oxyanion ends in "ate," the acid name is "ic acid."

      • Examples: H2SO4 (sulfuric acid), HNO3 (nitric acid).

    • If the oxyanion ends in "ite," the name is "ous acid."

      • Examples: H2SO3 (sulfurous acid), HNO2 (nitrous acid).

  2. Non-Oxy Acids: Name starts with "hydro" and ends with "ic acid."

    • Examples: HCl (hydrochloric acid), HI (hydroiodic acid).

  3. Naming Guidelines for Acids:

    • They should be named when in aqueous solution; naming differs when not in water.

VII. Chemical Formulas for Acids

  • Formulas for acids utilize oxidation numbers, adjusting hydrogen count to sum oxidation numbers to zero.

  • Examples:

    • Hydrosulfuric acid d

    • Nitrous acid

    • Perchloric acid

  • Structural formulas exist but are not covered in this course.

VIII. Binary Molecular Compounds

  1. Definition: Compounds composed of two elements.

    • Examples: SF6, H2O, CO.

  2. Naming Rules:

    • Use common names

    • First element retains its name; the second alters to end with "ide."

    • Apply prefixes indicating atom quantity.

      • Examples of Prefixes: mono, di, tri, tetra, penta, etc.

  • Note: Drop "o" or "a" from prefixes if followed by a vowel; maintain "i" in di and tri.

IX. Examples of Naming and Formulas

  • Provide practice for naming and writing formulas for various compounds: CO2, CO, P2O5, NO2.

  • Learn to derive chemical formulas from names appropriately.

  • Structural formulas covered minimally in class (e.g., structural formula for water).

X. Acids and Bases

  1. Acid Characteristics: Produces H+ ions, corrosive, sour taste, electrolytes.

  2. Base Characteristics: Produces OH- ions, corrosive, slippery feel, electrolytes (often referred to as alkali when involving metals bonded to hydroxide).

  3. pH Scale: Numeric scale (0-14) indicating acidity (0-7), neutral (7), and basicity (7-14).

    • Example pH measurements: stomach acid ~2, water 7, household ammonia 11.2.

  4. Indicators: Substances that change color based on pH; includes

    • Litmus paper (indicates acid/base but not exact pH).

    • pH paper (color changes corresponding to pH) and liquid indicators (change with pH).