Chapter 4: Covalent Compounds

4.1 Introduction to Covalent Bonding

Covalent bonding is a fundamental concept in chemistry, describing how atoms share electrons to form molecules. This section introduces the nature of covalent bonds and their role in molecular formation.

• Covalent bonds result from the sharing of electrons between two atoms, typically nonmetals.

• A covalent bond is a two-electron bond in which the bonding atoms share valence electrons.

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

• Example: Two hydrogen atoms (H) each contribute one electron to form a covalent bond, resulting in a stable H2 molecule with two electrons shared.

4.1 Unshared Electron Pairs and the Octet Rule

Atoms in covalent compounds may have electrons not involved in bonding, called lone pairs. The octet rule guides how atoms achieve stability.

• Nonbonded electron pairs or lone pairs are valence electrons not shared in bonds.

• Atoms share electrons to attain the electronic configuration of the noble gas closest to them in the periodic table.

• Hydrogen shares 2 electrons to achieve stability (like helium).

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

4.1 Covalent Bonding and the Periodic Table

The periodic table helps predict which elements form covalent bonds and how many bonds they can form.

• Lewis structures are electron-dot diagrams showing the location of all valence electrons in a molecule, including bonding and nonbonding pairs.

• Covalent bonds are typically formed between nonmetals or between a metalloid and a nonmetal.

• The number of covalent bonds an atom forms is related to its number of valence electrons.

• Formula for main group elements:

• Examples:

  • Hydrogen: 1 bond

  • Carbon: 4 bonds

  • Nitrogen: 3 bonds

  • Oxygen: 2 bonds

  • Halogens: 1 bond

4.1 The Modern Periodic Table

The periodic table classifies elements as metals, nonmetals, and metalloids, which influences their bonding behavior.

• Nonmetals (right side of the table) commonly form covalent bonds.

• Metalloids (bordering the staircase line) can form covalent bonds with nonmetals.

• Metals (left and center) typically do not form covalent bonds.

4.2 Lewis Structures

Lewis structures are essential for visualizing the arrangement of atoms and electrons in molecules.

• Molecular formula shows the number and identity of atoms but not their connectivity.

• Lewis structure shows how atoms are connected and the location of all bonding and nonbonding valence electrons.

• General rules for drawing Lewis structures:

  • Draw only valence electrons.

  • Give every main group element (except H) an octet of electrons.

  • Give each hydrogen (H) 2 electrons.

• Place hydrogen and halogens on the periphery, as they form only one bond.

• Count total valence electrons using group numbers from the periodic table.

• Arrange electrons to fill octets, using lone pairs as needed.

4.2 Multiple Bonds

Some molecules require double or triple bonds to satisfy the octet rule.

• A double bond contains four electrons in two shared pairs.

• A triple bond contains six electrons in three shared pairs.

• Example: In ethene (C2H4), a double bond is formed between the two carbon atoms to ensure each has an octet.

4.3 Exceptions to the Octet Rule

While most elements follow the octet rule, there are notable exceptions.

• Hydrogen needs only 2 electrons for stability.

• Elements in group 3A (e.g., boron) may have fewer than 8 electrons in a neutral molecule.

• Elements in the third row (e.g., phosphorus, sulfur) can have more than 8 electrons due to available d orbitals.

4.4 Resonance

Resonance structures represent molecules where electron arrangement can vary without changing atom connectivity.

• Resonance structures are two or more Lewis structures with the same arrangement of atoms but different electron arrangements.

• The true structure is a resonance hybrid, stabilizing the molecule by delocalizing electrons.

• Example: Ozone (O3) has two resonance structures, with the double bond in different positions.

4.5 Naming Covalent Compounds

Covalent compounds are named using specific rules and prefixes to indicate the number of atoms.

• Name the first nonmetal by its element name; the second nonmetal uses the suffix “-ide.”

• Add prefixes to show the number of atoms of each element (see Table below).

• The prefix “mono-” is usually omitted for the first element but retained for the second.

4.6 Molecular Shape/Geometry

The shape of a molecule is determined by the number of groups (atoms or lone pairs) around a central atom, using VSEPR theory.

• Linear: 2 groups, bond angle = 180° (e.g., CO2)

• Trigonal planar: 3 groups, bond angle = 120° (e.g., H2CO)

• Tetrahedral: 4 groups, bond angle = 109.5° (e.g., CH4)

• Trigonal pyramidal: 4 groups (one lone pair), bond angle ≈ 107° (e.g., NH3)

• Bent: 4 groups (two lone pairs), bond angle ≈ 105° (e.g., H2O)

4.7 Electronegativity and Bond Polarity

Electronegativity measures an atom’s ability to attract electrons in a bond, affecting bond polarity.

• If two bonded atoms have similar electronegativities, the bond is nonpolar (electrons shared equally).

• If electronegativities differ, the bond is polar covalent (electrons shared unequally), creating a dipole.

• Symbol indicates partial negative charge; indicates partial positive charge.

4.8 Polarity of Molecules

The overall polarity of a molecule depends on both the polarity of individual bonds and the molecular geometry.

• Nonpolar molecules: No polar bonds or bond dipoles cancel due to symmetry.

• Polar molecules: One or more polar bonds with dipoles that do not cancel, resulting in a net dipole moment.

• Example: CO2 is linear and nonpolar; H2O is bent and polar.

Additional info: These principles are foundational for understanding molecular interactions in biological systems, such as water’s polarity affecting its role as a universal solvent.