Chemical Bonding and Lewis Structures

Lewis Dot Symbols

Lewis Dot Symbols (also known as Electron Dot Diagrams) are visual representations of the valence electrons in an atom or ion. They are essential for understanding how atoms bond and form molecules.

• Valence Electrons: Electrons in the outermost shell of an atom, involved in chemical bonding.

• For Main Group Elements, the number of valence electrons equals the group number (for Groups 1A-8A).

• For Transition Metals, the number of valence electrons can vary and is often less straightforward.

Example: Which element will possess the most valence electrons? (Choices: S, Al, Ca, Cl, Br)

Drawing Lewis Dot Symbols

• Write the element symbol to represent the nucleus and inner electrons.

• Place dots around the symbol to represent valence electrons.

• Start by placing one electron on each side (top, right, bottom, left), then pair up as needed.

• For ions, add (anions) or remove (cations) electrons as appropriate, and indicate the charge.

Example: Draw the Lewis Dot Symbol for Te.

Practice: Draw the Lewis Dot symbol for Co+, Cd2+, and P3−.

Chemical Bonding Overview

Chemical bonds are the attractive forces that hold atoms or ions together in compounds. Atoms bond to achieve a more stable electron configuration, often resembling that of noble gases.

Ionic Bonding

Ionic bonding occurs between metals and nonmetals, involving the transfer of electrons from one atom (typically a metal) to another (typically a nonmetal).

• Cation: Atom that loses electrons (usually a metal).

• Anion: Atom that gains electrons (usually a nonmetal).

• Ionic bond formation is exothermic, lowering the energy of the system.

Example: Na + Cl → Na+ + Cl− (with electron transfer shown by arrows).

Practice: Identify which species has the most ionic character and the main source of ionic bond strength.

Covalent Bonding

Covalent bonding involves the sharing of valence electrons between nonmetal atoms. Each shared pair of electrons constitutes a covalent bond.

• Atoms achieve stability by sharing electrons to complete their valence shells.

• Single, double, or triple bonds can form depending on the number of shared electron pairs.

Example: H2O: Each hydrogen shares one electron with oxygen, forming two single covalent bonds.

Practice: Identify which element is unlikely to form covalent bonds.

Metallic Bonding

Metallic bonding is the attractive force between free-flowing valence electrons and positively charged metal ions in a metal lattice. This 'sea of electrons' gives metals their unique properties.

• Electrons are delocalized and can move freely throughout the metal structure.

• Responsible for properties such as conductivity, malleability, ductility, and luster.

Example: In metallic bonding diagrams, positive metal ions are surrounded by a sea of delocalized electrons.

Practice: Identify properties and characteristics of metallic bonding.

Electronegativity and Dipole Moment

Electronegativity (EN) is a measure of an atom's ability to attract electrons in a chemical bond.

• Periodic Trend: Electronegativity increases from left to right across a period and from bottom to top within a group.

• Dipole Moment: Occurs when there is a significant difference in electronegativity between two bonded atoms, resulting in a polar bond.

• The dipole arrow points towards the more electronegative atom.

Example: Calculate the difference in EN between C and F.

Practice: Arrange molecules in order of decreasing dipole moment.

Chemical Bond Classifications

The type of chemical bond depends on the difference in electronegativity between the two atoms:

• Nonpolar Covalent: ΔEN ≈ 0 (equal sharing of electrons)

• Polar Covalent: 0 < ΔEN < 1.7 (unequal sharing)

• Ionic: ΔEN > 1.7 (electron transfer)

Example: Which bond is most polar? (Compare ΔEN values)

Practice: Identify bond types between given atoms.

Octet Rule and Exceptions

The Octet Rule states that atoms tend to form bonds to achieve eight electrons in their valence shell, similar to noble gases.

• Valence Electrons: Electrons available for bonding.

• Shared Electrons: Electrons shared between atoms in a bond.

• Exceptions: Some elements have incomplete or expanded octets.

• Incomplete Octet: Stable with fewer than 8 electrons (e.g., Be, B).

• Expanded Octet: Stable with more than 8 electrons (e.g., P, S, Xe).

Example: How many octet electrons are around phosphorus in PH4+?

Practice: Identify molecules with incomplete octets.

Formal Charge

Formal Charge is a bookkeeping tool used to determine the distribution of electrons in a molecule or ion. It helps identify the most stable Lewis structure.

• Formula:

• Sum of formal charges in a molecule equals the overall charge.

• Structures with formal charges closest to zero are generally more stable.

Example: Calculate the formal charge of nitrogen in NH3.

Practice: Calculate formal charges for atoms in CO, NO2−, and SCN−.

Summary Table: Bond Types and Properties

• Ionic Bonds: Metal + Nonmetal, electron transfer, high melting point, conducts electricity when molten.

• Covalent Bonds: Nonmetal + Nonmetal, electron sharing, low to moderate melting point, poor conductor.

• Metallic Bonds: Metal atoms, delocalized electrons, high conductivity, malleable and ductile.

Key Formulas and Concepts

• Electronegativity Difference:

• Formal Charge:

Additional info: The notes above are based on standard General Chemistry curriculum and expand on the provided slides and images for clarity and completeness.