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 useful for predicting bonding behavior and the formation of molecules.

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

• Main Group Elements: Number of valence electrons equals the group number (for groups 1A-8A).

• Transition Metals: Number of valence electrons varies; often includes electrons from the (n-1)d and ns orbitals.

Example: Which element will possess the most valence electrons? (S, I, Ca, H, Br)

• Drawing Lewis Dot Symbols:

  1. Identify if the element is a Main Group or Transition Metal.

  2. Place one valence electron at a time on the four sides of the element symbol, starting from the right and moving clockwise.

  3. Continue adding electrons, pairing them up only after each side has one electron.

  4. If the element is an ion, indicate the charge in the upper right corner and adjust the number of electrons accordingly.

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. The main types are ionic, covalent, and metallic bonds.

Ionic Bonding

Ionic bonds form between metals and nonmetals through the transfer of electrons, resulting in oppositely charged ions that attract each other.

• Metals: Tend to lose electrons and form cations.

• Nonmetals: Tend to gain electrons and form anions.

• Ionic bond formation: Lowers the potential energy of the system.

Example: Na+ and Cl− combine to form NaCl.

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 molecules.

• Shared electrons: Each atom contributes one or more electrons to the bond.

• Bonding pairs: Electrons shared between atoms.

Example: H2O, where O shares electrons with two H atoms.

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

Metallic Bonding

Metallic bonds occur between metal atoms, where valence electrons are delocalized and free to move throughout the metal lattice.

• Delocalized electrons: Responsible for properties such as conductivity, malleability, and luster.

Example: Free-flowing electrons in a metal lattice.

Practice: Identify physical properties attributed to metallic bonding and classify compounds by bonding type.

Electronegativity and Dipole Moment

Electronegativity (EN): A measure of an atom's ability to attract electrons in a chemical bond. It increases from left to right across a period and decreases down a group.

• Difference in Electronegativity (ΔEN): Determines bond polarity.

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

Example: Calculate 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 atoms involved.

Practice: Identify the most polar bond and the bond type between C and O.

Octet Rule and Shared Electrons

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

• Valence electrons: Electrons available for bonding.

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

Example: Analyze the number of shared electrons in a compound.

Practice: Count shared electrons around oxygen in CO2 and identify atoms with the most valence electrons.

Incomplete Octet vs. Expanded Octet

Some elements can be stable with fewer or more than eight electrons in their valence shell.

• Incomplete Octet: Elements stable with less than eight electrons (e.g., Be, B).

• Expanded Octet: Elements stable with more than eight electrons (e.g., P, S, Cl, Xe).

• Group Number: Number of valence electrons for main group elements equals their group number.

Example: Count octet electrons around phosphorus in a compound.

Practice: Identify compounds with incomplete octets.

Formal Charge

Formal charge is a bookkeeping tool used to determine the distribution of electrons in molecules and ions.

• 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 (counted as pairs).

• Net Charge: Sum of formal charges in a molecule or ion.

Example: Calculate the formal charge of nitrogen in NH3.

Practice: Calculate formal charges for atoms in NO2−, CO, and the thiocyanate ion (NCS−).

Summary Table: Bond Types and Properties

Additional info:

• Practice questions throughout the notes reinforce key concepts and help students apply their understanding.

• Periodic table diagrams are used to illustrate trends in valence electrons and electronegativity.

• Bond classification tables help distinguish between pure covalent, polar covalent, and ionic bonds based on electronegativity differences.