Covalent Compounds and Bonding

Introduction to Covalent Compounds

• Covalent bonds are formed when atoms share electrons, resulting in the formation of a molecule.

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

Chapter Overview

• Covalent Bonds: Definition of covalent bonds.

• Covalent Compounds: Overview of properties and characteristics.

• Naming Covalent Molecules: Methods for naming and writing chemical formulas.

• Drawing Lewis Structures: Instructions on how to visualize covalent bonding.

• Multiple Bonds: Explanation of double and triple bonds.

• Characteristics of Covalent Bonds: Stability, strength, and types of bonds.

• Molecular Shapes: How the shapes of molecules are determined by covalent bonds.

4.1 Covalent Compounds

• A covalent compound forms when two or more atoms share electrons, leading to the creation of a stable molecule.

• Typically, molecular compounds occur when nonmetal atoms share electrons to achieve a more stable electron configuration.

Compounds: Ionic or Covalent?

• Ionic Compounds: Often contain metals or the polyatomic ion ammonium (NH4⁺). Example: K2O (potassium oxide).

• Covalent Compounds: Formed primarily with nonmetals. Example: N2O (dinitrogen oxide).

Identifying Compound Types

• For classification:

  • A: SO3 - Covalent (sulfur trioxide)

  • B: BaCl2 - Ionic (barium chloride)

  • C: (NH4)3PO3 - Ionic (ammonium phosphite)

  • D: Cu2CO3 - Ionic (copper(I) carbonate)

  • E: N2O4 - Covalent (dinitrogen tetroxide)

Covalent Bonding

• Covalent bonds result from the sharing of electron pairs.

• **Types of Pairs:

  • Bonding Pair (shared electrons).

  • Lone Pair (non-bonded electrons).

Bond Representations

• Single Bond: Represented as a single line or by two dots.

• Multiple Bonds:

  • Double Bond: Two shared pairs of electrons.

  • Triple Bond: Three shared pairs of electrons.

Formation of H2 Molecule

• Covalent bond formation illustrated by:

  • H atoms come close together, sharing electrons.

  • Energy dynamics:

  • Far apart, no attraction.

  • Moving closer, attractions pull atoms together.

• Diagram of bond:

  • H-H diagram shows bond formation into the H2 molecule.

Diatomic Molecules

• Elements forming diatomic molecules include:

  • H2 (hydrogen)

  • N2 (nitrogen)

  • O2 (oxygen)

  • F2 (fluorine)

  • Cl2 (chlorine)

  • Br2 (bromine)

  • I2 (iodine)

Molecules Composed of Covalent Bonds

• Examples:

  • Water (H2O)

  • Ethanol (C2H5OH)

  • Sulfuric acid (H2SO4)

  • Carbon dioxide (CO2)

  • Ammonium ion (NH4⁺)

Unshared Electron Pairs

• Nonbonded pairs are referred to as lone pairs.

• Aim for stable configurations resembling noble gases:

  • Hydrogen (H) reaches Helium's configuration by sharing 2 electrons.

  • Main group elements typically achieve an octet for stability.

Lewis Dot Symbols

• Representation of valence electrons.

  • Hydrogen, Carbon, Nitrogen, Oxygen, and Halogens exhibit specific bonding patterns (e.g., number of bonds and lone pairs).

Covalent Bonding and the Periodic Table

• Lewis structures show valence electrons and bonding.

• Predicting the number of covalent bonds:

  • Atoms with one, two, or three valence electrons tend to form one, two, or three bonds, respectively.

  • General formula:
    $ext{Predicted number of bonds} = 8 - ext{number of valence electrons}$

Naming of Covalent Compounds

1. Common Molecular Compounds: Table summarizing some compounds and their commercial uses.

2. Flowchart for Naming Compounds: Simplifies the process of naming ionic vs. covalent compounds.

3. Specific Naming Examples:

   • NO2 = nitrogen dioxide

   • N2O4 = dinitrogen tetroxide

4. Use of Prefixes: Used to denote the number of atoms in covalent compounds (e.g., mono-, di-, tri-, tetra-).

Writing Molecular Formulas

• Rules for creating molecular formulas from names by converting prefixes:

  • Selenium dioxide is SeO2

  • Carbon tetrachloride is CCl4

  • Follow similar conventions for dinitrogen pentoxide (N2O5) and phosphorus triiodide (PI3).

Lewis Structures: Overview

• Definition: Visual representation of how atoms are bonded and valence electrons arranged.

• General Drawing Steps:

  1. Arrange atoms based on bonding potentials.

  2. Count total valence electrons.

  3. Create bonds and assign lone pairs where required.

  4. Adapt structure for octet completion.

Multiple Bonds

• Characteristics of Multiple Bonds:

  • Double bonds occur by sharing two pairs of electrons.

  • Triple bonds occur by sharing three pairs of electrons.

• Lewis Structures involving multiple bonds demonstrate shared electrons between atoms for octet fulfillment.

Exceptions to the Octet Rule

• Not all elements follow the octet rule.

  • Hydrogen only needs 2 electrons.

  • Elements like phosphorus (P) and sulfur (S) can exceed 8 electrons if necessary.

Resonance Structures

• Definition: Multiple valid Lewis structures can exist for a single molecule due to electron arrangement.

• Implications of resonance influence molecular stability and characteristics.

Electronegativity and Bond Polarity

• Electronegativity: A measure of an atom's attraction for electrons.

• Bond Types Based on Electronegativity Differences:

  • Nonpolar Covalent: Electronegative differences < 0.4

  • Polar Covalent: Differences between 0.5 and 1.8

  • Ionic: Differences > 1.8

Polarity of Molecules

• Classification depends on:

  • The polarity of individual bonds.

  • The overall shape of the molecule.

• Examples demonstrating polar and nonpolar molecules based on bond types and symmetry (e.g., carbon dioxide vs. water).

Conclusion

• Understanding covalent bonds is essential for interpreting molecular structure, behavior, and reactivity. The frameworks of naming, drawing Lewis structures, and determining molecular shapes are fundamental skills in chemistry that highlight the diverse nature of compounds in our environment.