Covalent Compounds Study Notes

Covalent Compounds

4.1 Introduction to Covalent Bonding

• Covalent bonds: Result from the sharing of electrons between two atoms.

• A covalent bond is a two-electron bond where bonding atoms share valence electrons.

• A molecule: A discrete group of atoms held together by covalent bonds.

• Nonbonded electron pairs: Also known as lone pairs.

• Atoms share electrons to achieve the electronic configuration of the nearest noble gas.

  • Hydrogen shares 2 electrons.

  • Other main group elements share electrons until attaining an octet (8 electrons in their outer shell).

4.1A Covalent Bonding and the Periodic Table

• Lewis structures: Electron-dot structures that illustrate the location of all valence electrons.

• Covalent bonds typically form when nonmetals combine or when a metalloid bonds to a nonmetal.

Covalent Bonds Formed by Various Atoms:

• Hydrogen: 1 bond (1 valence electron).

• Atoms with 1-3 valence electrons form the same number of bonds as valence electrons.

• Atoms with 4 or more valence electrons form bonds to achieve an octet:
  $ext{Predicted number of bonds} = 8 - ext{number of valence electrons}$

General Bonding Rule:

$ext{Number of bonds} + ext{Number of lone pairs} = 4$

4.2 Lewis Structures

Importance of Lewis Structures:

• A molecular formula identifies the number and types of atoms but not their bonding.

• A Lewis structure shows both atom connectivity and electron placement (bonding and nonbonding valence electrons).

Steps to Draw Lewis Structures:

1. Arrange atoms: Position bonded atoms close to each other; place H and halogens on the periphery.

2. Count valence electrons: Sum equals the total for the structure.

3. Distribute electrons: Place 1 bond (2 electrons) between each pair, filling octets (8 for main group elements, 2 for H).

   • Example: For CH3Cl, total valence electrons = 14.

4. Use multiple bonds (double/triple) as necessary to achieve octets.

4.2B Multiple Bonds

• Double bonds: 4 electrons (two 2-electron bonds).

• Triple bonds: 6 electrons (three 2-electron bonds).

Example: Lewis Structure for C2H4

• Step 1: Arrange atoms.

• Step 2: Total valence electrons = 12.

• Step 3: Complete structure.

• Step 4: Adjust for octet if needed, forming a double bond if necessary. T

4.3 Exceptions to the Octet Rule

• Hydrogen requires only 2 electrons (an exception).

• Elements in Group 3A may not form an octet due to low valence electrons.

• Third-row elements (like P and S) can exceed 8 electrons due to empty d orbitals.

4.4 Resonance

• Resonance structures: Different Lewis structures showing the same arrangement of atoms but different electron placements.

• Resonance stabilizes a molecule by spreading out electrons and lone pairs over a larger area.

• Example: Ozone (O3) has two resonance structures.

4.5 Naming Covalent Compounds

Steps for Naming:

1. Name the first nonmetal by element name, the second with the suffix “-ide”.

   • Example: NO2 = nitrogen dioxide, N2O4 = dinitrogen tetroxide.

2. Add prefixes to indicate the number of atoms (from Table 4.1):

   • Mono (1), Di (2), Tri (3), Tetra (4), Penta (5), Hexa (6), Hepta (7), Octa (8), Nona (9), Deca (10).

4.6 Molecular Shape

• VSEPR Theory: Shapes are determined by the numbe t r of electron groups (bonding and lone pairs).

  • An atom surrounded by 2 groups is linear (180°).

  • An atom surrounded by 3 groups is trigonal planar (120°).

  • An atom surrounded by 4 groups is tetrahedral (109.5°).

• Examples:

  • CO2: Linear

  • BF3: Trigonal planar

  • CH4: Tetrahedral

  • NH3: Trigonal pyramid (107°)

  • H2O: Bent (105°)

4.7 Electronegativity and Bond Polarity

• Electronegativity: Measures an atom's attraction for electrons in a bond.

Bond Type Based on Electronegativity Difference:

• Less than 0.5: Nonpolar

• 0.5 to 2.0: Polar covalent

• Greater than 2.0: Ionic

4.8 Polarity of Molecules

• Molecular polarity: Depends on individual bond polarities and overall shape.

  • If individual bond dipoles cancel, the molecule is nonpolar; if they do not, it is polar.

Steps to Determine Polarity:

1. Identify polar bonds based on electronegativity differences.

2. Assess the shape around atoms based on groups.

3. Determine if dipoles cancel or reinforce each other.