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.

• 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 on each side before pairing).

• For ions, add or remove dots to reflect the gain or loss of electrons, and indicate the charge.

Steps:

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

2. Place one valence electron as a dot on each of the four sides of the element symbol, then pair as needed.

3. For ions, add (anions) or remove (cations) 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

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 (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 species has the most ionic character.

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 (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 valence electrons and positively charged metal ions in a metal lattice.

• Electrons are delocalized and can move freely, giving rise to properties such as conductivity, malleability, and luster.

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

Electronegativity and Bond Polarity

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.

• A dipole arrow points towards the more electronegative atom.

• The greater the difference in EN, the more polar the bond.

Example: Calculate the EN difference between C and F.

Bond Classifications by Electronegativity Difference

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

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 8 electrons.

• Expanded Octet: Elements in period 3 or higher can have more than 8 electrons (e.g., P, S, Cl).

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

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 (each bond counts as one for 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 the formal charge for each oxygen atom in NO2−.

• Calculate the formal charge of carbon in CO.

• Determine the net charge of a given compound based on formal charges.

• Identify which element in the thiocyanate ion (NCS−) has a negative charge.

Summary Table: Bond Types and Properties

Key Terms

• Valence Electrons: Outermost electrons involved in bonding.

• Lewis Dot Symbol: Diagram showing valence electrons as dots around the element symbol.

• Ionic Bond: Electrostatic attraction between cations and anions.

• Covalent Bond: Sharing of electron pairs between atoms.

• Metallic Bond: Attraction between free electrons and metal ions.

• Electronegativity: Atom's ability to attract electrons in a bond.

• Dipole Moment: Measure of bond polarity due to electronegativity difference.

• Octet Rule: Tendency to have eight electrons in the valence shell.

• Formal Charge: Calculated charge on an atom in a molecule, assuming equal sharing of electrons.

Additional info: Some explanations and examples have been expanded for clarity and completeness, including the summary tables and stepwise procedures for drawing Lewis structures and calculating formal charge.