10.1 Lewis (Electron) Dot Formulas

Fundamentals of Covalent Bonding and Lewis Dot Structures

• Covalent Bonding Principles: Nonmetal compounds achieve stability by sharing electrons. This sharing allows each atom to reach a more stable electronic configuration, typically an octet (eight valence electrons).

• Molecular Stability: Covalent bonding results from shared electrons between the $s$-sublevel and $p$-sublevel valence electrons to achieve stability.

• The Chlorine Example ($Cl_2$):

  • Two $Cl$ atoms share a pair of electrons in a covalent bond.

  • By sharing, both $Cl$ atoms effectively have eight valence electrons.

  • In the final structure, each atom retains six electrons of its own on the outside, while two electrons are positioned in the middle, shared between the two atoms.

Terminology of Lewis Structures

• Single Bond: Consists of two electrons being shared between two elements. It is represented by one line (e.g., $-$) or two dots (:\b) between the elements.

• Double Bond: Consists of four electrons being shared between two elements. It is represented by two lines (e.g., $=$) or four dots ($::$) between the elements.

• Triple Bond: Consists of six electrons being shared between two elements. It is represented by three lines or six dots between the elements.

• Unshared Electrons: These are electrons that are not shared between two elements. They are also referred to as lone pairs and are usually located on the outside of the molecular structure.

Step-by-Step Process for Drawing Lewis (Electron) Dot Formulas

Drawing Lewis structures follows a consistent, repeatable process for molecular compounds.

1. Step 1: Placement of Atoms: Place the element that is present in the smallest quantity in the center. Surround it with the remaining atoms.

2. Step 2: Calculate Total Valence Electrons: Use the group numbers from the periodic table to determine the number of valence electrons for each individual atom, then sum them together for the entire molecule.

3. Step 3: Establish Single Bonds: Create a single bond (represented by a dash or two dots) between each of the peripheral atoms and the central atom. Each dash accounts for $2$ electrons.

4. Step 4: Satisfy Outer Octets: Fill the octets of the outside atoms by adding electrons in pairs until they have eight (with the exception of Hydrogen, which only needs two).

5. Step 5: Assign Remaining Electrons to Center: Place any remaining electrons from the total count on the central atom to fill its octet, if possible.

6. Step 6: Formation of Multiple Bonds: If the central atom still lacks an octet and no more electrons are available, move unshared electron pairs from an outside atom to the bond area between that atom and the central atom to form a double or triple bond.

Detailed Example #1: Dichlorine Monoxide ($Cl_2O$)

• Step 1: Oxygen ($O$) is the center atom because there is only one, while there are two Chlorines ($Cl$). Layout: $Cl \quad O \quad Cl$.

• Step 2: Calculate valence electrons:

  • $Cl: 7$ valence electrons

  • $O: 6$ valence electrons

  • $Cl: 7$ valence electrons

  • TOTAL: $7 + 6 + 7 = 20$ valence electrons.

• Step 3: Draw single bonds between the center $O$ and the two $Cl$ atoms. This uses $4$ of the $20$ electrons ($2 \times 2 = 4$).

• Step 4: Fill the octets of the outside $Cl$ atoms. This uses $12$ more electrons (six dots around each $Cl$). Total used so far: $4 + 12 = 16$.

• Step 5: Place the remaining $4$ electrons ($20 - 16 = 4$) on the central Oxygen atom.

• Conclusion: Every atom has eight electrons. Each Chloriner has three unshared pairs and one single bond. The Oxygen has two unshared pairs and two single bonds.

Detailed Example #2: Sulfur Dioxide ($SO_2$)

• Step 1: Sulfur ($S$) is the center atom. Layout: $O \quad S \quad O$.

• Step 2: Calculate valence electrons:

  • $O: 6$ valence electrons

  • $S: 6$ valence electrons

  • $O: 6$ valence electrons

  • TOTAL: $6 + 6 + 6 = 18$ valence electrons.

• Step 3: Make a single bond between each peripheral $O$ and the central $S$. (Uses $4$ electrons).

• Step 4: Fill the octets of the outside $O$ atoms. (Uses $12$ electrons). Current total used: $16$.

• Step 5: Place the remaining $2$ electrons ($18 - 16 = 2$) on the central $S$ atom.

• Step 6: The Sulfur atom currently only has six electrons ($2$ from the unshared pair and $4$ from the two single bonds). To satisfy the octet, move an unshared pair of electrons from one of the Oxygen atoms to form a double bond with the Sulfur.

• Final Structure Details:

  • One Oxygen has three pairs of unshared electrons and one single bond.

  • The other Oxygen has two pairs of unshared electrons and one double bond.

  • The Sulfur atom has one unshared pair, one single bond, and one double bond.

Resonance Structures

• Definition: Resonance structures are used when there are two or more possible, valid Lewis structures for the same molecule.

• Representation: Resonance is represented by drawing all possible structures connected by a double-sided arrow ($\leftrightarrow$).

• The Case of $SO_2$: The double bond could be placed on either the left Oxygen or the right Oxygen. Both structures are valid representations.

• Physical Reality: Experiments indicate that the bonds are not actually one single and one double bond, but rather two equivalent "$1.5$ bonds" off the Sulfur. Resonance is a convention used to represent this averaged electronic state.

Note on Ions and Formal Charge

• Lewis Structures for Ions: The process remains similar for ions, though specific handling of charges is required.

• Formal Charge: While often used in drawing Lewis diagrams, students in CHM130 are advised to disregard the mention of formal charge as it is not covered in this level of study.

Practice Problems

• Task: Draw the Lewis Structures for the following molecules:

  • a. Carbon disulfide ($CS_2$)

  • b. Silicon tetrachloride ($SiCl_4$)

  • c. Carbon tetrafluoride ($CF_4$)