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 how atoms bond in molecules and ions.

• 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, representing the nucleus and inner electrons.

• Place dots around the symbol to represent valence electrons (one dot per electron).

• Pair electrons after each side has one electron (maximum of two per side).

• For ions, add or remove electrons as needed and indicate the charge.

Example: Draw the Lewis Dot Symbol for Te.

1. Identify if the element is a main group or transition metal.

2. Place one valence electron on each side before pairing.

3. For ions, adjust the number of electrons and indicate the charge.

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

Chemical Bonding

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 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.

• Metals tend to lose electrons (form cations).

• Nonmetals tend to gain electrons (form anions).

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

Example: Na + Cl → Na+ + Cl− → NaCl

Practice: Identify which compound has the most ionic character.

Covalent Bonding

Covalent bonding involves the sharing of valence electrons between nonmetal atoms to achieve a stable electron configuration (octet).

• Each shared pair of electrons forms a covalent bond.

• Multiple pairs can be shared (single, double, triple bonds).

Example: H2O: Each hydrogen shares one electron with oxygen.

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.

• Responsible for conductivity, malleability, ductility, and luster.

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

Practice: Identify which property is not attributed to 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.

• Fluorine is the most electronegative element.

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

• Represented by a dipole arrow pointing towards the more electronegative atom.

• The greater the difference in EN, the larger the dipole moment.

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:

Example: Which bond is most polar? (Choices: S–Se, S–H, Cl–F, S–F, S–O)

Practice: Identify the bond type between carbon and oxygen.

The Octet Rule and Exceptions

The Octet Rule states that atoms tend to gain, lose, or share electrons to achieve eight valence electrons, similar to noble gases.

• Valence Electrons: Electrons available for bonding.

• Shared Electrons: Electrons shared through covalent bonds.

Exceptions:

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

• Expanded Octet: Elements in period 3 or higher can have more than eight electrons (e.g., P, S, Cl).

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

Practice: Identify compounds 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 for the neutral atom (from periodic table).

• Nonbonding Electrons: Lone pairs on the atom.

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

Net Charge: The sum of all formal charges in a molecule or ion equals the overall charge.

Example: Determine 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

Key Formulas and Concepts

• Formal Charge:

• Electronegativity Difference:

• Octet Rule: Most main group elements aim for 8 valence electrons.

Additional info: These notes expand on the provided slides by including definitions, examples, and summary tables for clarity and completeness.