Chapter 9 Chemical Bonding I: Basic Concepts

Chapter 9 Chemical Bonding I: Basic Concepts

Valence Electrons

  • Valence electrons are the outer shell electrons of an atom and play a critical role in chemical bonding.

  • The number of valence electrons determines how an atom can bond with other atoms to form molecules.

Lewis Dot Symbols

  • Lewis dot symbols are visual representations used to illustrate the distribution of valence electrons around an atom.

  • For representative elements, each dot around the element symbol represents one valence electron, which helps to visualize how atoms bond with each other, particularly in covalent and ionic bonding.

Ionic Bonding

Ionic Bond

  • An ionic bond is the electrostatic force that holds ions together in an ionic compound, resulting from the attraction between positively charged cations and negatively charged anions.

  • Typically formed by the transfer of electrons from a metal (which loses electrons and becomes a cation) to a nonmetal (which gains electrons and becomes an anion).

Example of Ionic Bond Formation

  • Aluminum Oxide Formation:

  • In the formation of aluminum oxide (Al₂O₃), aluminum (Al) transfers its three valence electrons to oxygen (O). This transfers forms two Al³⁺ ions and three O²⁻ ions, resulting in the neutral compound Al₂O₃.

  • This follows the electroneutrality principle, where the total positive charge equals the total negative charge in the compound.

Lattice Energy

  • Lattice energy (U) is the energy required to separate one mole of a solid ionic compound into its gaseous ions.

  • Factors affecting lattice energy include the charges on the ions (higher charges result in higher lattice energy) and the distance between them (shorter distances result in higher lattice energy due to stronger attractions).

Born-Haber Cycle

  • The Born-Haber cycle is a thermodynamic cycle that helps analyze the formation of ionic compounds and calculate lattice energy through various energy changes, which includes ionization energy, electron affinity, and enthalpy changes.

  • An example of this cycle is seen with lithium fluoride (LiF), where each step includes enthalpy changes that collectively help evaluate lattice energy.

Lattice Energies and Properties of Compounds

  • Compound Lattice Energy (kJ/mol) | Melting Point (°C)

    • LiF: 1017 | 845

    • NaCl: 788 | 801

    • MgO: 3890 | 2800

Covalent Bonding

Covalent Bond

  • A covalent bond is formed when two or more electrons are shared between two atoms, allowing them to achieve a more stable electronic configuration.

  • Types of covalent bonds include:

    • Single bonds (1 pair of shared electrons)

    • Double bonds (2 pairs of shared electrons)

    • Triple bonds (3 pairs of shared electrons)

Average Bond Lengths of Common Bonds

  • Bond Type | Bond Length (pm)

    • C−H: 107

    • C=C: 133

    • C≡C: 120

Comparison: NaCl vs. CCl₄

  • Property | NaCl | CCl₄

    • Appearance: White solid | Colorless liquid

    • Melting Point (°C): 801 | −23

    • Solubility in water: High | Very low

Polar Covalent Bonds

  • Polar covalent bonds occur when electrons are shared unequally between two atoms, resulting in partial charges (dipoles).

  • The more electronegative atom pulls electrons closer, creating a shift in electron density.

Electronegativity

  • Electronegativity is a measure of an atom's ability to attract electrons in a bond; this property varies across the periodic table. For example, fluorine (F) is the most electronegative element, while elements further down a group exhibit lower electronegativities.

  • An example of electronegativity values illustrates this trend among various elements.

Electronegativity and Bond Classification

  • Bond types can be classified based on the difference in electronegativity between the bonded atoms:

    • 0: Covalent

    • 2: Ionic

    • 0 < difference < 2: Polar Covalent

Writing Lewis Structures

Steps to Write Lewis Structures

  1. Sketch the skeletal structure of the compound.

  2. Count the total number of valence electrons, adjusting for any charges from ions.

  3. Complete the octets for main-group elements, noting that hydrogen (H) only requires 2 electrons to fill its valence shell.

  4. Use double or triple bonds if necessary to satisfy the octet rule for elements that require more than 8 electrons.

Examples of Lewis Structures

  • Lewis structures for molecular nitrogen trifluoride (NF₃) and nitric acid (HNO₃) are illustrated, demonstrating how the octet rule and formal charge considerations come into play in correct bonding pictures.

Exceptions to the Octet Rule

  • Types of Exceptions:

    • Incomplete Octet Example: Beryllium Hydride (BeH₂)

    • Odd-Electron Molecules Example: Nitric Oxide (NO)

    • Expanded Octets Example: Phosphorus Pentafluoride (PF₅)

Formal Charge Calculation

  • The formal charge for an atom in a Lewis structure is calculated using the formula:

    Formal Charge = (Valence Electrons) - (Nonbonding Electrons + 0.5 imes Bonding Electrons)

  • Examples elucidate the approach to calculating formal charges and their significance.

Resonance Structures

  • A resonance structure is one of several possible Lewis structures that collectively represent a molecule that cannot be accurately depicted by a single formula. Nitrous oxide (N₂O) is used to illustrate different resonance forms and their associated formal charges.

Bond Enthalpy

  • Bond enthalpy is defined as the energy change required to break a specific bond in one mole of gaseous molecules and provides insight into the stability of bonds.

  • Example calculations illustrate how understanding bond enthalpies plays a vital role in determining the overall enthalpy of a given reaction.

Notes compiled using a variety of examples, tables, and visuals to reinforce core concepts of chemical bonding.