Which Of The Following Lewis Structures Is Correct

The ability to accurately depict molecular structures is fundamental to understanding chemical properties and reactions. Lewis structures, also known as electron dot diagrams, are a simple yet powerful tool used to visualize the arrangement of atoms and electrons within a molecule. Determining which Lewis structure is correct among several possibilities involves adhering to specific rules and guidelines, considering factors such as valence electrons, formal charges, and resonance.

Understanding Lewis Structures

Lewis structures illustrate how atoms are connected within a molecule and where the valence electrons (electrons in the outermost shell of an atom) are distributed. These diagrams help predict molecular geometry, polarity, and reactivity.

Key Components of a Lewis Structure:

Chemical Symbols: Represent the atoms in the molecule.
Lines: Indicate covalent bonds between atoms, where each line represents a shared pair of electrons.
Dots: Represent non-bonding valence electrons, also known as lone pairs, which are not involved in bonding.

Rules for Drawing Lewis Structures:

Determine the Total Number of Valence Electrons: Sum the valence electrons of all atoms in the molecule or ion.
Write the Skeletal Structure: Connect the atoms with single bonds. The least electronegative atom usually occupies the central position (hydrogen is always terminal).
Distribute Electrons to Outer Atoms: Complete the octets (or duet for hydrogen) of the atoms bonded to the central atom.
Place Remaining Electrons on the Central Atom: If any valence electrons remain after satisfying the octets of the outer atoms, place them as lone pairs on the central atom.
Form Multiple Bonds if Necessary: If the central atom does not have an octet, form double or triple bonds by sharing lone pairs from the outer atoms.

Criteria for Evaluating Lewis Structures

When presented with multiple possible Lewis structures, several criteria can help determine which one is the most correct or plausible. These include adherence to the octet rule, minimization of formal charges, and consideration of resonance structures.

1. The Octet Rule

The octet rule states that atoms tend to gain, lose, or share electrons in order to achieve a full outer electron shell with eight electrons, similar to noble gases. Hydrogen is an exception as it only requires two electrons (duet rule).

Satisfying the Octet Rule: A correct Lewis structure should ensure that most atoms have an octet of electrons. Exceptions include:
- Hydrogen (H): Forms one bond and has two electrons.
- Beryllium (Be): Can be stable with four electrons.
- Boron (B): Can be stable with six electrons.
- Elements beyond the Second Period: Can have expanded octets, accommodating more than eight electrons due to available d orbitals.

2. Formal Charges

Formal charge is the hypothetical charge on an atom in a molecule if all bonding electrons were shared equally between atoms. Calculating formal charges helps to evaluate the distribution of electrons in a Lewis structure.

Formula for Formal Charge:

Formal Charge = (Valence Electrons) - (Non-bonding Electrons) - (Number of Bonds)

Principles for Formal Charges:
- Minimize Formal Charges: The best Lewis structure is typically the one with the smallest formal charges on each atom.
- Negative Formal Charges on More Electronegative Atoms: If formal charges cannot be completely avoided, negative formal charges should be placed on the more electronegative atoms and positive formal charges on the less electronegative atoms.
- Sum of Formal Charges Equals Overall Charge: The sum of the formal charges in a molecule must equal zero, and in an ion, it must equal the ion's charge.

3. Electronegativity

Electronegativity is a measure of an atom's ability to attract electrons in a chemical bond. Differences in electronegativity between atoms in a molecule can influence the distribution of electrons and the polarity of bonds.

Role of Electronegativity: In a correct Lewis structure, more electronegative atoms should have a greater share of electrons, reflected in lower (or more negative) formal charges compared to less electronegative atoms.
Trends in Electronegativity: Electronegativity generally increases across a period (from left to right) and decreases down a group (from top to bottom) in the periodic table.

4. Resonance Structures

Resonance occurs when more than one valid Lewis structure can be drawn for a molecule or ion, differing only in the arrangement of electrons (not atoms). The actual electronic structure is a resonance hybrid, an average of all resonance structures.

Identifying Resonance: Look for molecules with multiple bonds and lone pairs that can be rearranged without changing the positions of the atoms.
Evaluating Resonance Structures:
- Equivalent Resonance Structures: If resonance structures are equivalent, they contribute equally to the resonance hybrid.
- Non-equivalent Resonance Structures: Structures that minimize formal charges and place negative charges on more electronegative atoms are more significant contributors to the resonance hybrid.

Step-by-Step Approach to Determining the Correct Lewis Structure

To determine the correct Lewis structure among several possibilities, follow these steps:

1. Draw All Possible Lewis Structures:

Start by calculating the total number of valence electrons.
Arrange the atoms in different bonding patterns, keeping in mind that the least electronegative atom is typically the central atom.
Distribute electrons to satisfy the octet rule.

2. Calculate Formal Charges for Each Atom in Each Structure:

Use the formal charge formula to calculate the charge on each atom in each possible Lewis structure.

3. Evaluate Each Structure Based on the Octet Rule and Formal Charges:

Determine if each atom has an octet of electrons (or follows the duet rule for hydrogen).
Assess the formal charges:
- Are they minimized?
- Are negative formal charges on more electronegative atoms?
- Does the sum of the formal charges equal the overall charge of the molecule or ion?

4. Consider Resonance (If Applicable):

If multiple structures are possible and differ only in the arrangement of electrons, draw all resonance structures.
Evaluate the relative contribution of each resonance structure based on formal charges and electronegativity.

5. Determine the Correct Lewis Structure:

The correct Lewis structure is the one that best satisfies the octet rule, minimizes formal charges, places negative charges on more electronegative atoms, and represents the most stable resonance hybrid (if resonance exists).

Examples with Detailed Explanations

Let's walk through several examples to illustrate how to determine the correct Lewis structure.

Example 1: Carbon Dioxide (CO₂) - A Classic Case

Carbon dioxide (CO₂) is a linear molecule with one carbon atom and two oxygen atoms.

Total Valence Electrons:
- Carbon (C): 4 valence electrons
- Oxygen (O): 6 valence electrons × 2 = 12 valence electrons
- Total: 4 + 12 = 16 valence electrons
Possible Lewis Structures:
- Structure A: O=C=O
- Structure B: O≡C-O (with appropriate lone pairs)
- Structure C: O-C≡O (with appropriate lone pairs)
Formal Charges:
- Structure A (O=C=O):
  - Oxygen: 6 (valence) - 4 (non-bonding) - 2 (bonds) = 0
  - Carbon: 4 (valence) - 0 (non-bonding) - 4 (bonds) = 0
- Structure B (O≡C-O):
  - Triple Bond Oxygen: 6 (valence) - 2 (non-bonding) - 3 (bonds) = +1
  - Carbon: 4 (valence) - 0 (non-bonding) - 4 (bonds) = 0
  - Single Bond Oxygen: 6 (valence) - 6 (non-bonding) - 1 (bond) = -1
- Structure C (O-C≡O):
  - Single Bond Oxygen: 6 (valence) - 6 (non-bonding) - 1 (bond) = -1
  - Carbon: 4 (valence) - 0 (non-bonding) - 4 (bonds) = 0
  - Triple Bond Oxygen: 6 (valence) - 2 (non-bonding) - 3 (bonds) = +1
Evaluation:
- Structure A has formal charges of zero on all atoms, making it the most stable and correct Lewis structure.
- Structures B and C have non-zero formal charges (+1 and -1), which are less favorable.
Thus, the correct Lewis structure for CO₂ is O=C=O.

Example 2: Ozone (O₃) - Illustrating Resonance

Ozone (O₃) consists of three oxygen atoms.

Total Valence Electrons:
- Oxygen (O): 6 valence electrons × 3 = 18 valence electrons
Possible Lewis Structures:
- Structure A: O=O-O (with appropriate lone pairs)
- Structure B: O-O=O (with appropriate lone pairs)
Formal Charges:
- Structure A (O=O-O):
  - Double Bond Oxygen: 6 (valence) - 4 (non-bonding) - 2 (bonds) = 0
  - Central Oxygen: 6 (valence) - 2 (non-bonding) - 3 (bonds) = +1
  - Single Bond Oxygen: 6 (valence) - 6 (non-bonding) - 1 (bond) = -1
- Structure B (O-O=O):
  - Single Bond Oxygen: 6 (valence) - 6 (non-bonding) - 1 (bond) = -1
  - Central Oxygen: 6 (valence) - 2 (non-bonding) - 3 (bonds) = +1
  - Double Bond Oxygen: 6 (valence) - 4 (non-bonding) - 2 (bonds) = 0
Evaluation:
- Both structures A and B are valid Lewis structures and are resonance structures of each other.
- Each oxygen atom has an octet, and the formal charges are minimized. The central oxygen has a formal charge of +1, and one of the terminal oxygens has a formal charge of -1.
- The actual structure of ozone is a resonance hybrid of these two structures, with the double bond delocalized between the two oxygen-oxygen bonds.
The correct representation of ozone involves both resonance structures: O=O-O ↔ O-O=O.

Example 3: Nitrate Ion (NO₃⁻) - An Ionic Case

The nitrate ion (NO₃⁻) consists of one nitrogen atom and three oxygen atoms, with an overall charge of -1.

Total Valence Electrons:
- Nitrogen (N): 5 valence electrons
- Oxygen (O): 6 valence electrons × 3 = 18 valence electrons
- Ion Charge: +1 electron (due to the -1 charge)
- Total: 5 + 18 + 1 = 24 valence electrons
Possible Lewis Structures:
- Structure A: O=N-O-O (with appropriate lone pairs and charge distribution)
- Structure B: O-N=O-O (with appropriate lone pairs and charge distribution)
- Structure C: O-N-O=O (with appropriate lone pairs and charge distribution)
Formal Charges:
- Structure A (O=N-O-O):
  - Double Bond Oxygen: 6 (valence) - 4 (non-bonding) - 2 (bonds) = 0
  - Nitrogen: 5 (valence) - 0 (non-bonding) - 4 (bonds) = +1
  - Single Bond Oxygen (two): 6 (valence) - 6 (non-bonding) - 1 (bond) = -1
- Structure B (O-N=O-O):
  - Single Bond Oxygen: 6 (valence) - 6 (non-bonding) - 1 (bond) = -1
  - Nitrogen: 5 (valence) - 0 (non-bonding) - 4 (bonds) = +1
  - Double Bond Oxygen: 6 (valence) - 4 (non-bonding) - 2 (bonds) = 0
  - Single Bond Oxygen: 6 (valence) - 6 (non-bonding) - 1 (bond) = -1
- Structure C (O-N-O=O):
  - Single Bond Oxygen: 6 (valence) - 6 (non-bonding) - 1 (bond) = -1
  - Nitrogen: 5 (valence) - 0 (non-bonding) - 4 (bonds) = +1
  - Single Bond Oxygen: 6 (valence) - 6 (non-bonding) - 1 (bond) = -1
  - Double Bond Oxygen: 6 (valence) - 4 (non-bonding) - 2 (bonds) = 0
Evaluation:
- All three structures are valid resonance structures, as they differ only in the arrangement of electrons.
- Each structure has one nitrogen atom with a formal charge of +1 and two oxygen atoms with formal charges of -1, summing to the overall -1 charge of the ion.
- The actual structure of the nitrate ion is a resonance hybrid of these three structures, with the double bond delocalized among the three oxygen-nitrogen bonds.
The correct representation of the nitrate ion involves all three resonance structures, each contributing equally to the hybrid structure.

Example 4: Sulfur Dioxide (SO₂) - Expanded Octet Considerations

Sulfur dioxide (SO₂) is a molecule with one sulfur atom and two oxygen atoms. Sulfur can exhibit an expanded octet.

Total Valence Electrons:
- Sulfur (S): 6 valence electrons
- Oxygen (O): 6 valence electrons × 2 = 12 valence electrons
- Total: 6 + 12 = 18 valence electrons
Possible Lewis Structures:
- Structure A: O=S=O (with appropriate lone pairs)
- Structure B: O-S≡O (with appropriate lone pairs)
- Structure C: O≡S-O (with appropriate lone pairs)
Formal Charges:
- Structure A (O=S=O):
  - Oxygen: 6 (valence) - 4 (non-bonding) - 2 (bonds) = 0
  - Sulfur: 6 (valence) - 2 (non-bonding) - 4 (bonds) = 0
- Structure B (O-S≡O):
  - Single Bond Oxygen: 6 (valence) - 6 (non-bonding) - 1 (bond) = -1
  - Sulfur: 6 (valence) - 2 (non-bonding) - 4 (bonds) = +1
  - Triple Bond Oxygen: 6 (valence) - 2 (non-bonding) - 3 (bonds) = 1
- Structure C (O≡S-O):
  - Triple Bond Oxygen: 6 (valence) - 2 (non-bonding) - 3 (bonds) = 1
  - Sulfur: 6 (valence) - 2 (non-bonding) - 4 (bonds) = +1
  - Single Bond Oxygen: 6 (valence) - 6 (non-bonding) - 1 (bond) = -1
Evaluation:
- In Structure A, all atoms have formal charges of zero, which is generally the most stable arrangement. Sulfur has 10 electrons around it, demonstrating an expanded octet, which is permissible for third-period elements.
- Structures B and C have formal charges of +1 and -1, which are less favorable.
- The best Lewis structure is O=S=O, recognizing that sulfur can accommodate more than eight electrons.
- However, one must also consider the resonance structures O=S-O and O-S=O each with S having a lone pair. The central S in each of these would bear a positive charge and the single bonded O would bear a negative charge.

Common Mistakes to Avoid

When drawing and evaluating Lewis structures, keep an eye out for these common mistakes:

Incorrectly Counting Valence Electrons: Always double-check the number of valence electrons for each atom.
Violating the Octet Rule: Ensure that most atoms have an octet unless they are exceptions like hydrogen, beryllium, or boron. Elements in the third period and beyond can have expanded octets.
Ignoring Formal Charges: Failing to calculate formal charges can lead to selecting a less stable structure.
Overlooking Resonance: Recognize when multiple valid structures can be drawn and consider the resonance hybrid.
Incorrectly Placing Negative Charges: Ensure negative formal charges are on more electronegative atoms.

Conclusion

Determining the correct Lewis structure involves a systematic approach that considers valence electrons, the octet rule, formal charges, electronegativity, and resonance. By following a step-by-step process and understanding the underlying principles, one can accurately depict molecular structures and gain valuable insights into chemical bonding and molecular properties. Mastery of Lewis structures is not just an academic exercise but a fundamental skill for anyone studying chemistry.