Each Pictured Lewis Structure Is Invalid

Lewis structures, visual representations of molecules illustrating the bonding between atoms and lone pairs of electrons, are fundamental tools in chemistry. However, not every depiction that appears to follow the rules is actually valid. Identifying invalid Lewis structures requires a keen understanding of the underlying principles, including valence electron counts, octet rule exceptions, formal charge calculations, and resonance structures. Mastering these concepts enables you to critically evaluate and construct accurate representations of molecular structure.

Common Reasons for Invalid Lewis Structures

Invalid Lewis structures typically arise from violations of core principles:

1. Incorrect Valence Electron Count: The cardinal rule is to accurately account for the total number of valence electrons in the molecule or ion. This is derived from the group number of each element in the periodic table. Errors in this count will inevitably lead to an incorrect structure.

2. Octet Rule Violations: While most atoms strive to achieve an octet (eight) of electrons in their valence shell, there are exceptions. Hydrogen only needs two electrons, beryllium often has four, and boron often has six. Furthermore, elements in the third row and beyond can accommodate more than eight electrons due to the availability of d orbitals.

3. Incorrect Atom Connectivity: The central atom is usually the least electronegative element (excluding hydrogen). The skeletal structure, representing how atoms are connected, must be correct before placing electrons.

4. Unnecessary Formal Charges: While formal charges are sometimes unavoidable, minimizing them generally leads to a more stable and accurate Lewis structure. Large formal charges or adjacent atoms with the same formal charge usually indicate an incorrect structure.

5. Failure to Depict Resonance: When multiple valid Lewis structures can be drawn for a molecule, differing only in the arrangement of electrons, resonance must be considered. The actual structure is a hybrid of all contributing resonance structures.

Step-by-Step Analysis of Potential Lewis Structures

To determine if a pictured Lewis structure is invalid, follow these steps:

1. Count the Total Valence Electrons: Sum the valence electrons of all atoms in the molecule or ion. Remember to add electrons for negative charges and subtract electrons for positive charges.

2. Draw the Skeletal Structure: Place the least electronegative atom in the center (excluding hydrogen). Connect the atoms with single bonds.

3. Distribute Electrons as Lone Pairs: Complete the octets of the surrounding atoms first. Then, place any remaining electrons on the central atom.

4. Form Multiple Bonds: If the central atom does not have an octet, form multiple bonds (double or triple bonds) by moving lone pairs from the surrounding atoms.

5. Calculate Formal Charges: Calculate the formal charge on each atom using the formula: Formal Charge = (Valence Electrons) - (Non-bonding Electrons) - (1/2 Bonding Electrons).

6. Evaluate the Structure:

   • Does the structure adhere to the octet rule (with exceptions)?
   • Are the formal charges minimized?
   • Is the connectivity correct?
   • Are resonance structures considered if applicable?

In-Depth Examples of Invalid Lewis Structures

Let's examine several examples to illustrate common errors and how to identify them:

Example 1: Carbon Dioxide (CO₂)

Incorrect Lewis Structure: O=C=O with each oxygen having three lone pairs.

Analysis:

• Valence Electron Count: Carbon (4) + 2 * Oxygen (6) = 16 valence electrons
• The incorrect structure shows 20 electrons (4 bonding + 16 non-bonding).
• Correct Lewis Structure: O=C=O with each oxygen having two lone pairs.

Reason for Invalidity: Incorrect valence electron count. The incorrect structure depicts more electrons than are available in the molecule.

Example 2: Ozone (O₃)

Incorrect Lewis Structure: A linear structure with single bonds between the oxygen atoms, with each terminal oxygen having three lone pairs and the central oxygen having one lone pair.

Analysis:

• Valence Electron Count: 3 * Oxygen (6) = 18 valence electrons
• The incorrect structure shows 20 electrons (4 bonding + 16 non-bonding).
• Correct Lewis Structure: A bent structure with one single bond and one double bond, exhibiting resonance. One resonance structure shows O=O-O, and the other shows O-O=O. The central oxygen has one lone pair, one terminal oxygen has two lone pairs, and the other terminal oxygen has three lone pairs.

Reason for Invalidity: Incorrect electron count and failure to depict resonance. The incorrect structure does not account for the delocalization of electrons in ozone.

Example 3: Sulfur Trioxide (SO₃)

Incorrect Lewis Structure: Sulfur single-bonded to three oxygen atoms, each oxygen with three lone pairs, and sulfur with one lone pair.

Analysis:

• Valence Electron Count: Sulfur (6) + 3 * Oxygen (6) = 24 valence electrons
• The incorrect structure obeys the octet rule but results in large formal charges. Sulfur has a +3 formal charge, and each oxygen has a -1 formal charge.
• Correct Lewis Structure: Sulfur double-bonded to one oxygen and single-bonded to two oxygens, with resonance. Alternatively, a structure with three double bonds to sulfur is also valid because sulfur can exceed the octet.

Reason for Invalidity: High formal charges and failure to consider resonance and octet expansion. Minimizing formal charges and acknowledging resonance leads to a more accurate representation.

Example 4: Nitrogen Dioxide (NO₂)

Incorrect Lewis Structure: Nitrogen double-bonded to both oxygen atoms, each oxygen with two lone pairs.

Analysis:

• Valence Electron Count: Nitrogen (5) + 2 * Oxygen (6) = 17 valence electrons
• The incorrect structure depicts an even number of electrons, which is impossible for NO₂. Nitrogen would have a positive charge, and each oxygen would have no charge.
• Correct Lewis Structure: Nitrogen is double-bonded to one oxygen and single-bonded to the other, with a lone electron on the nitrogen. Resonance is possible between the two N-O bonds.

Reason for Invalidity: NO₂ has an odd number of valence electrons, making it a radical. The incorrect structure attempts to force an even number of electrons.

Example 5: Formaldehyde (CH₂O)

Incorrect Lewis Structure: Carbon single-bonded to two hydrogens and one oxygen, with the oxygen having three lone pairs.

Analysis:

• Valence Electron Count: Carbon (4) + 2 * Hydrogen (1) + Oxygen (6) = 12 valence electrons
• The incorrect structure shows all atoms single-bonded, leaving the carbon with only six electrons, violating the octet rule.
• Correct Lewis Structure: Carbon double-bonded to oxygen and single-bonded to the two hydrogens. The oxygen has two lone pairs.

Reason for Invalidity: Carbon does not have a full octet. Forming a double bond between carbon and oxygen satisfies the octet rule for both atoms.

Example 6: The Cyanide Ion (CN⁻)

Incorrect Lewis Structure: Carbon single-bonded to nitrogen, with each atom having three lone pairs.

Analysis:

• Valence Electron Count: Carbon (4) + Nitrogen (5) + 1 (from the negative charge) = 10 valence electrons
• The incorrect structure has 14 electrons (2 bonding + 12 non-bonding).
• Correct Lewis Structure: Carbon triple-bonded to nitrogen. Carbon has one lone pair, and nitrogen has one lone pair.

Reason for Invalidity: Incorrect valence electron count. Forming a triple bond satisfies the octet rule for both carbon and nitrogen while using the correct number of electrons.

Example 7: Boron Trifluoride (BF₃)

Incorrect Lewis Structure: Boron single-bonded to three fluorine atoms, each fluorine having three lone pairs, giving boron a full octet through dative bonding.

Analysis:

• Valence Electron Count: Boron (3) + 3 * Fluorine (7) = 24 valence electrons
• The incorrect structure forces an octet on boron through dative bonds, leading to a high formal charge on Boron (+3) and some Fluorine atoms.
• Correct Lewis Structure: Boron single-bonded to three fluorine atoms, each fluorine having three lone pairs. Boron only has six electrons.

Reason for Invalidity: While dative bonds are possible, the most stable structure minimizes formal charges. Boron is an exception to the octet rule and can be stable with only six valence electrons.

Example 8: Phosphorus Pentachloride (PCl₅)

Incorrect Lewis Structure: Phosphorus single-bonded to five chlorine atoms, with each chlorine having three lone pairs, but incorrectly placing lone pairs on phosphorus to complete an octet.

Analysis:

• Valence Electron Count: Phosphorus (5) + 5 * Chlorine (7) = 40 valence electrons
• The incorrect structure tries to force an octet on phosphorus, but phosphorus can exceed the octet rule due to the availability of d orbitals.
• Correct Lewis Structure: Phosphorus single-bonded to five chlorine atoms, each chlorine having three lone pairs. No lone pairs are placed on phosphorus.

Reason for Invalidity: Phosphorus can have more than eight electrons in its valence shell.

Example 9: Xenon Tetrafluoride (XeF₄)

Incorrect Lewis Structure: Xenon single-bonded to four fluorine atoms, each fluorine having three lone pairs, and incorrectly depicts xenon with fewer than two lone pairs.

Analysis:

• Valence Electron Count: Xenon (8) + 4 * Fluorine (7) = 36 valence electrons
• The incorrect structure does not account for all valence electrons on Xenon.
• Correct Lewis Structure: Xenon single-bonded to four fluorine atoms, each fluorine having three lone pairs, and two lone pairs on Xenon.

Reason for Invalidity: Xenon needs two lone pairs to satisfy the valence electron count.

Example 10: Ammonium Ion (NH₄⁺)

Incorrect Lewis Structure: Nitrogen single-bonded to four hydrogen atoms but fails to remove an electron for the positive charge, and doesn't correctly account for the number of lone pairs.

Analysis:

• Valence Electron Count: Nitrogen (5) + 4 * Hydrogen (1) - 1 (from the positive charge) = 8 valence electrons
• The incorrect structure does not properly account for the positive charge.
• Correct Lewis Structure: Nitrogen single-bonded to four hydrogen atoms, with no lone pairs on the nitrogen, and enclosed in brackets with a "+" charge outside.

Reason for Invalidity: Failure to account for the positive charge.

Advanced Considerations: Resonance and Formal Charge

Resonance is crucial when multiple valid Lewis structures can be drawn for a molecule or ion. The actual structure is a hybrid of all resonance contributors, and the electrons are delocalized across the molecule. Consider nitrate (NO₃⁻) as an example. Three resonance structures can be drawn, each with one nitrogen-oxygen double bond and two nitrogen-oxygen single bonds. The actual structure is an average of these, with each N-O bond having a bond order of 1 1/3.

Formal charge helps assess the relative stability of different Lewis structures. The best Lewis structure typically minimizes formal charges. For instance, consider carbon monoxide (CO). Two possible Lewis structures can be drawn:

• C≡O (Formal charges: C = -1, O = +1)
• C=O (Formal charges: C = 0, O = 0)

While both structures satisfy the octet rule, the structure with zero formal charges is more stable and contributes more to the actual electronic structure of CO.

Common Pitfalls to Avoid

• Forgetting Lone Pairs: Always ensure that all non-bonding electrons are explicitly drawn as lone pairs.
• Ignoring Formal Charges: Calculate and consider formal charges to assess the stability of different structures.
• Overlooking Resonance: Recognize when resonance is possible and draw all significant resonance contributors.
• Misunderstanding Octet Rule Exceptions: Be aware of elements that can have fewer or more than eight electrons in their valence shell.
• Incorrectly Determining Connectivity: Double-check the skeletal structure to ensure atoms are connected correctly.

Tools and Resources for Mastering Lewis Structures

Several resources can aid in mastering Lewis structures:

• Textbooks: General chemistry textbooks provide detailed explanations of Lewis structures and related concepts.
• Online Tutorials: Websites like Khan Academy and Chem LibreTexts offer interactive tutorials and practice problems.
• Molecular Modeling Software: Programs like ChemDraw and GaussView allow you to visualize and manipulate molecular structures.
• Practice Problems: Work through a variety of examples to solidify your understanding.

Conclusion

Constructing and evaluating Lewis structures is a fundamental skill in chemistry. By understanding the underlying principles, recognizing common errors, and practicing diligently, you can confidently determine the validity of Lewis structures and gain a deeper understanding of molecular bonding and structure. Mastering these concepts provides a solid foundation for further exploration of chemical principles and reactions. Always remember to double-check your work, consider resonance and formal charges, and be mindful of octet rule exceptions.