Chemical Bonding and Lewis Structures

Lewis Dot Symbols

Lewis Dot Symbols (also called 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 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 (one dot per electron).

• Start from the right of the symbol and add dots clockwise, pairing them after each side has one dot.

• For ions, add or remove dots to reflect the gain or loss of electrons, 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 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 to another, resulting in oppositely charged ions that attract each other.

• Cations (usually metals) lose electrons.

• Anions (usually nonmetals) gain electrons.

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

Example: Na• + Cl•••••• → Na+ + Cl−

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

Covalent Bonding

Covalent bonds involve the sharing of valence electrons between nonmetal atoms, resulting in the formation of molecules.

• Each shared pair of electrons constitutes a single covalent bond.

• Multiple pairs can be shared, forming double or triple bonds.

Example: Which element is unlikely to form covalent bonds? (Choices: S, H, K, Ar, Si)

Metallic Bonding

Metallic bonding is the attractive force between free-flowing (delocalized) electrons and positively charged metal ions in a metal lattice. This bonding gives metals their unique properties.

• Delocalized electrons can move freely throughout the metal structure.

• Properties: ductility, luster, malleability, and high electrical conductivity.

Example: Which best describes the free-flowing electrons in metallic bonding?

Practice: Identify physical properties and bonding types in various compounds.

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.

• Polarity: Unequal sharing of electrons creates a dipole (partial positive and negative ends).

• Difference in Electronegativity (ΔEN): ΔEN = |ENatom1 − ENatom2|

• The dipole arrow points towards the more electronegative atom.

Example: Calculate the ΔEN 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? (Choices: S–Se, S–H, Cl–F, S–F, S–O)

Practice: Identify bond types between given atoms.

The Octet Rule and Exceptions

Most main group elements tend to achieve an octet (8 valence electrons) through bonding, similar to noble gases.

• Valence Electrons: Electrons an element possesses based on its group number.

• Shared Electrons: Electrons shared through chemical bonds.

Exceptions:

• Incomplete Octet: Some elements (e.g., Be, B) are stable with fewer than 8 electrons.

• Expanded Octet: Elements in period 3 or higher can have 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 to determine the distribution of electrons in a molecule or ion. It helps identify the most stable Lewis structure.

• Formula:

• Valence Electrons: Number of electrons in the neutral atom.

• Nonbonding Electrons: Electrons not involved in bonding (lone pairs).

• Bonding Electrons: Electrons shared in bonds (count all electrons in bonds to the atom).

Example: Determine 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 or dissolved.

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

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

Key Formulas and Concepts

• Electronegativity Difference:

• Formal Charge:

Additional info: The notes above expand on the provided slides by including definitions, periodic trends, and examples of exceptions to the octet rule, as well as a summary table for quick reference.