Decide Whether These Proposed Lewis Structures Are Reasonable

Determining the reasonableness of proposed Lewis structures is a cornerstone of understanding molecular structure and chemical behavior. A well-constructed Lewis structure not only visualizes how atoms are connected within a molecule but also predicts molecular geometry, polarity, and reactivity. This comprehensive guide will walk you through the essential rules and principles for evaluating Lewis structures, ensuring you can confidently assess their validity.

Understanding Lewis Structures

Lewis structures, also known as electron dot diagrams, are visual representations of the bonding between atoms in a molecule, as well as any lone pairs of electrons that may exist. They help us understand:

• Connectivity: Which atoms are bonded to each other.
• Valence Electrons: The distribution of valence electrons around each atom.
• Bond Order: Whether bonds are single, double, or triple.
• Formal Charges: The charge assigned to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally between atoms.
• Resonance: The possibility of multiple valid Lewis structures for a single molecule.

The Rules of the Game: Constructing Valid Lewis Structures

Before we delve into evaluating proposed structures, let's revisit the rules for constructing them in the first place:

1. Count Valence Electrons: Sum the valence electrons of all atoms in the molecule or ion. For polyatomic ions, add electrons for negative charges and subtract for positive charges.
2. Draw the Skeletal Structure: Connect the atoms with single bonds. The least electronegative atom usually occupies the central position (except for hydrogen, which always occupies a terminal position).
3. Complete Octets (or Duets): Distribute the remaining electrons as lone pairs around the atoms to satisfy the octet rule (eight electrons) for each atom, or the duet rule (two electrons) for hydrogen.
4. Form Multiple Bonds: If any atom (except hydrogen) lacks an octet, form double or triple bonds to satisfy the octet rule.
5. Calculate Formal Charges: Assign formal charges to each atom. Minimize formal charges whenever possible.
6. Consider Resonance Structures: If multiple arrangements satisfy the octet rule and minimize formal charges, draw all plausible resonance structures.

Key Criteria for Evaluating Lewis Structures

Now, let's focus on how to determine if a proposed Lewis structure is reasonable. Here's a breakdown of the key criteria:

1. Correct Number of Valence Electrons

• The Check: The most fundamental check is ensuring the Lewis structure accurately depicts the total number of valence electrons available in the molecule or ion.
• How to Verify: Count all the electrons in the Lewis structure, including those in bonds (each bond contains two electrons) and lone pairs. Compare this number to the sum of valence electrons you calculated in Step 1 of the construction process.
• Example: Consider carbon dioxide (CO₂). Carbon has 4 valence electrons, and each oxygen has 6, for a total of 4 + 6 + 6 = 16 valence electrons. A reasonable Lewis structure for CO₂ must have 16 electrons represented.
• Error Indication: If the number of electrons in the Lewis structure doesn't match the total number of valence electrons, the structure is incorrect.

2. Octet Rule Satisfaction (and Exceptions)

• The Check: Most atoms strive to achieve an octet (eight electrons) in their valence shell. Hydrogen aims for a duet (two electrons). However, there are notable exceptions to the octet rule.
• How to Verify: Examine each atom in the Lewis structure and count the number of electrons surrounding it. Consider both bonding electrons (shared) and lone pair electrons.
• Common Octet Rule Violations:
  • Incomplete Octets: Certain atoms, like boron (B) and beryllium (Be), are stable with fewer than eight electrons. For example, in boron trifluoride (BF₃), boron has only six electrons around it.
  • Expanded Octets: Elements in the third period and beyond (e.g., P, S, Cl, Br, I) can accommodate more than eight electrons due to the availability of d orbitals. For example, sulfur hexafluoride (SF₆) has sulfur surrounded by twelve electrons.
  • Odd-Electron Species: Molecules with an odd number of valence electrons, called free radicals, cannot satisfy the octet rule for all atoms. For example, nitrogen monoxide (NO) has 11 valence electrons.
• Error Indication: If an atom should have an octet but doesn't, or if an atom is depicted with an expanded octet when it shouldn't, the structure is likely unreasonable (unless it fits one of the exceptions).

3. Appropriate Formal Charges

• The Check: Formal charge helps assess the distribution of electrons in a Lewis structure. While a structure with zero formal charges on all atoms is ideal, this is not always possible. Minimizing formal charges and placing negative formal charges on more electronegative atoms contributes to a more reasonable structure.
• How to Calculate Formal Charge: Formal Charge = (Valence Electrons) - (Non-bonding Electrons) - (1/2 Bonding Electrons)
• Guidelines for Formal Charges:
  • The sum of the formal charges in a neutral molecule should be zero.
  • The sum of the formal charges in an ion should equal the ion's charge.
  • Minimize the magnitude of formal charges on all atoms.
  • Place negative formal charges on more electronegative atoms and positive formal charges on less electronegative atoms.
• Example: Consider the ozone molecule (O₃). We can draw two resonance structures. One structure has formal charges of 0, +1, and -1. The other is equivalent.
• Error Indication: A Lewis structure with large formal charges, or with positive formal charges on highly electronegative atoms, is likely less reasonable than a structure with minimal formal charges and appropriate placement.

4. Molecular Geometry Predictions (VSEPR Theory)

• The Check: Valence Shell Electron Pair Repulsion (VSEPR) theory states that electron pairs (both bonding and non-bonding) around a central atom will arrange themselves to minimize repulsion. The Lewis structure should be consistent with the molecular geometry predicted by VSEPR theory.
• How to Apply VSEPR Theory:
  1. Determine the central atom.
  2. Count the number of electron groups (bonding pairs and lone pairs) around the central atom.
  3. Predict the electron-pair geometry based on the number of electron groups:
     • 2 electron groups: Linear
     • 3 electron groups: Trigonal Planar
     • 4 electron groups: Tetrahedral
     • 5 electron groups: Trigonal Bipyramidal
     • 6 electron groups: Octahedral
  4. Determine the molecular geometry by considering the number of bonding pairs and lone pairs. Lone pairs exert a greater repulsive force, distorting the bond angles.
• Example: Methane (CH₄) has four bonding pairs and no lone pairs around the central carbon atom. This leads to a tetrahedral electron-pair geometry and a tetrahedral molecular geometry, with bond angles of approximately 109.5°.
• Error Indication: A Lewis structure that implies a drastically different molecular geometry than what VSEPR theory predicts raises suspicion about its validity.

5. Resonance Structures

• The Check: Resonance occurs when multiple valid Lewis structures can be drawn for a molecule or ion. These structures differ only in the arrangement of electrons, not in the position of the atoms. The actual structure is a hybrid of all resonance contributors.
• How to Identify Resonance: Look for molecules with multiple bonds and lone pairs adjacent to each other. These arrangements allow for electron delocalization.
• Rules for Resonance Structures:
  • Only electrons can be moved; the positions of the atoms must remain the same.
  • All resonance structures must be valid Lewis structures.
  • Resonance structures should have the same number of unpaired electrons.
  • Resonance structures with lower formal charges are more stable and contribute more to the resonance hybrid.
• Example: Benzene (C₆H₆) has two major resonance structures, each with alternating single and double bonds in the ring. The actual structure is a hybrid, with all carbon-carbon bonds having equal length and strength.
• Error Indication: Failing to recognize resonance when it's possible, or drawing resonance structures that violate the rules above, suggests an incomplete or incorrect understanding of the molecule's electronic structure.

6. Bond Lengths and Bond Strengths

• The Check: While Lewis structures don't directly provide quantitative bond lengths or strengths, they can indicate relative differences. Triple bonds are shorter and stronger than double bonds, which are shorter and stronger than single bonds.
• How to Interpret: If experimental data on bond lengths or strengths is available, compare it to the predictions based on the Lewis structure. A Lewis structure implying a bond order inconsistent with experimental data is suspect.
• Example: In carbon monoxide (CO), the Lewis structure shows a triple bond between carbon and oxygen. This correlates with the experimentally determined short bond length and high bond dissociation energy of CO.
• Error Indication: A significant discrepancy between the predicted relative bond orders (from the Lewis structure) and experimental bond lengths or strengths suggests the Lewis structure may not be an accurate representation of the molecule.

7. Electronegativity Considerations

• The Check: Electronegativity differences between atoms influence bond polarity and the distribution of electron density. A reasonable Lewis structure should reflect the expected polarity based on electronegativity differences.
• How to Apply: Remember that electronegativity increases across a period and up a group in the periodic table. The greater the electronegativity difference between two bonded atoms, the more polar the bond.
• Example: In water (H₂O), oxygen is more electronegative than hydrogen. The Lewis structure correctly shows that the oxygen atom has a partial negative charge (δ-) and the hydrogen atoms have partial positive charges (δ+).
• Error Indication: A Lewis structure that contradicts expected bond polarities based on electronegativity differences may be incorrect, especially when considering formal charge placement.

Common Mistakes to Avoid

• Forgetting to count all valence electrons.
• Violating the octet rule without justification (exceptions).
• Placing positive formal charges on highly electronegative atoms.
• Ignoring resonance when it's possible.
• Drawing structures with unreasonable bond angles (inconsistent with VSEPR theory).
• Misinterpreting the meaning of formal charge. Formal charge is a bookkeeping device, not an actual charge on the atom.
• Confusing resonance structures with isomers. Isomers have different arrangements of atoms, while resonance structures differ only in the arrangement of electrons.

Examples: Assessing Proposed Lewis Structures

Let's examine a few examples to illustrate how to apply these criteria:

Example 1: Nitrogen Dioxide (NO₂)

• Proposed Structure: A structure where nitrogen is singly bonded to each oxygen, each oxygen has three lone pairs, and nitrogen has one lone pair.
• Evaluation:
  • Valence Electrons: N (5) + 2O (6 x 2) = 17 valence electrons. The proposed structure has 2 + 6 + 6 + 2 = 16 electrons, which is incorrect.
  • Octet Rule: Oxygen atoms have octets, but nitrogen has only 5 electrons around it.
  • Formal Charges: Can be calculated, but the electron count error makes it invalid.
• Conclusion: This structure is unreasonable due to an incorrect number of valence electrons and a violation of the octet rule for nitrogen. A correct Lewis structure requires resonance and will have one unpaired electron on the nitrogen.

Example 2: Carbon Monoxide (CO)

• Proposed Structure: A structure with a double bond between carbon and oxygen, each having two lone pairs.
• Evaluation:
  • Valence Electrons: C (4) + O (6) = 10 valence electrons. The proposed structure has 4 + 4 + 4 = 12 electrons, incorrect.
  • Octet Rule: The proposed structure doesn't satisfy the octet rule for either carbon or oxygen.
  • Formal Charges: Significant formal charges.
• Conclusion: This structure is unreasonable due to an incorrect number of valence electrons and violation of the octet rule. The correct structure has a triple bond, with formal charges of -1 on carbon and +1 on oxygen. This structure maximizes the number of bonds while minimizing formal charges.

Example 3: Sulfate Ion (SO₄²⁻)

• Proposed Structure: A structure with single bonds between sulfur and each oxygen, each oxygen having three lone pairs, and sulfur having two lone pairs.
• Evaluation:
  • Valence Electrons: S (6) + 4O (6 x 4) + 2 (charge) = 32 valence electrons. The proposed structure has 8 + 24 = 32 electrons – correct.
  • Octet Rule: All atoms satisfy the octet rule.
  • Formal Charges: Sulfur has a formal charge of +2, and each oxygen has a formal charge of -1. The overall charge is -2, consistent with the ion.
• Alternative Structure: A structure with two double bonds and two single bonds between sulfur and oxygen. In this structure, two oxygen atoms have two lone pairs each (double-bonded), and two oxygen atoms have three lone pairs each (single-bonded). Sulfur has one lone pair.
• Evaluation of Alternative:
  • Valence Electrons: Still 32 – correct.
  • Octet Rule: Still satisfied.
  • Formal Charges: Sulfur has a formal charge of 0, two oxygens have a formal charge of 0, and two oxygens have a formal charge of -1.
• Conclusion: While both structures satisfy the octet rule and have the correct number of valence electrons, the second structure (with the double bonds) is generally considered more reasonable because it minimizes formal charges on the atoms. There would be resonance between different arrangements of the double bonds.

The Importance of Practice

Evaluating Lewis structures is a skill that improves with practice. By meticulously applying these rules and criteria, and by working through numerous examples, you'll develop a strong intuition for identifying reasonable and unreasonable Lewis structures. This skill is crucial for understanding chemical bonding, molecular properties, and reactivity.

Conclusion

Determining the reasonableness of proposed Lewis structures is a multi-faceted process that requires a thorough understanding of valence electron counts, the octet rule (and its exceptions), formal charges, VSEPR theory, resonance, and electronegativity. By systematically applying these principles, you can confidently assess the validity of Lewis structures and gain valuable insights into the electronic structure and behavior of molecules. Remember to always prioritize structures that minimize formal charges and are consistent with experimental data whenever available. Mastering this skill will significantly enhance your understanding of chemistry.