Which One Of The Following Lewis Structures Is Definitely Incorrect

The ability to correctly represent molecules using Lewis structures is fundamental to understanding chemical bonding and reactivity. However, not all Lewis structures are created equal; some may violate fundamental rules and principles. Identifying an incorrect Lewis structure involves checking for several key aspects, including the octet rule, formal charges, and overall structure validity. This article delves into the criteria for evaluating Lewis structures and provides a step-by-step guide to determining which Lewis structure is definitely incorrect.

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. These structures help to predict the geometry of molecules and understand their chemical properties.

Key Components of a Lewis Structure:

• Atomic Symbols: Represent the atoms in the molecule.
• Bonds: Lines connecting atoms represent shared pairs of electrons (covalent bonds). A single line represents a single bond (2 electrons), a double line represents a double bond (4 electrons), and a triple line represents a triple bond (6 electrons).
• Lone Pairs: Dots around an atom represent non-bonding pairs of electrons. Each pair consists of two electrons.

Rules for Drawing Lewis Structures:

1. Determine the Total Number of Valence Electrons: Add up the valence electrons of all atoms in the molecule or ion.
2. Write the Skeletal Structure: Arrange the atoms in the molecule. The least electronegative atom usually goes in the center (except for hydrogen, which is always on the periphery).
3. Distribute Electrons: Place electron pairs between atoms to form chemical bonds.
4. Complete Octets: Distribute the remaining electrons as lone pairs to satisfy the octet rule (8 electrons around each atom), starting with the most electronegative atoms. Hydrogen is an exception and only needs 2 electrons.
5. Minimize Formal Charges: If necessary, form multiple bonds to satisfy the octet rule and minimize formal charges.

Criteria for Evaluating Lewis Structures

To determine if a Lewis structure is incorrect, consider the following criteria:

1. Octet Rule Violations

The octet rule states that atoms in a molecule should have eight electrons in their valence shell. There are exceptions, but it's a good starting point:

• Hydrogen (H): Only needs two electrons.
• Beryllium (Be): Can be stable with four electrons.
• Boron (B): Can be stable with six electrons.
• Third-row and heavier elements: Can have more than eight electrons (expanded octet) due to the availability of d orbitals. Examples include sulfur (S), phosphorus (P), chlorine (Cl), and xenon (Xe).

Violation: A Lewis structure is incorrect if atoms, other than the exceptions listed above, do not have a complete octet.

2. Incorrect Number of Valence Electrons

A Lewis structure must accurately represent the total number of valence electrons available in the molecule or ion.

Violation: If the number of electrons depicted in the Lewis structure does not match the calculated total number of valence electrons, the structure is incorrect.

3. Unreasonable Formal Charges

Formal charge is the charge assigned to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally between atoms, regardless of relative electronegativity. The formal charge of an atom in a Lewis structure can be calculated using the formula:

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

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 charge of the ion.
• Lewis structures with smaller formal charges are generally more stable.
• Negative formal charges should be placed on more electronegative atoms.
• Avoid placing large formal charges on atoms.

Violation: A Lewis structure is likely incorrect if it has:

• Large formal charges on atoms.
• Positive formal charges on highly electronegative atoms.
• The sum of formal charges does not match the overall charge of the molecule or ion.

4. Incorrect Atom Connectivity

The arrangement of atoms in a molecule (connectivity) must be correct. This is often determined by the chemical formula and known bonding patterns.

Violation: A Lewis structure is incorrect if the atoms are connected in an improbable or chemically unreasonable manner. For example, oxygen atoms are unlikely to be directly bonded to each other unless the molecule is a peroxide.

5. Violation of Electronegativity Rules

In general, the more electronegative atoms should have more lone pairs and/or be bonded to fewer atoms than the less electronegative atoms.

Violation: If a less electronegative atom has more electron density (more lone pairs and fewer bonds) than a more electronegative atom, the Lewis structure might be incorrect.

Step-by-Step Guide to Identifying an Incorrect Lewis Structure

1. Write Down the Chemical Formula: Clearly identify the molecule or ion for which the Lewis structure is provided.
2. Calculate the Total Number of Valence Electrons: Sum up the valence electrons contributed by each atom in the molecule or ion. For ions, add electrons for negative charges and subtract for positive charges.
3. Examine Each Lewis Structure:
   • Count the Total Number of Electrons: Ensure that the number of electrons depicted in the Lewis structure matches the total number of valence electrons calculated.
   • Check Octets: Verify that each atom (except for hydrogen and exceptions like Be and B) has an octet of electrons. Identify any octet rule violations.
   • Calculate Formal Charges: Determine the formal charge on each atom in the structure. Assess whether the formal charges are minimized and appropriately distributed.
   • Assess Atom Connectivity: Evaluate the arrangement of atoms to ensure it is chemically reasonable.
   • Consider Electronegativity: Check that the distribution of electron density aligns with electronegativity trends.
4. Identify the Incorrect Structure: Based on the above criteria, determine which Lewis structure violates one or more of the fundamental rules and principles.

Examples of Identifying Incorrect Lewis Structures

Let's illustrate how to identify incorrect Lewis structures with some examples:

Example 1: Carbon Dioxide (CO₂)

Suppose you are given three possible Lewis structures for CO₂:

(a) O-C-O (each single bonded, with appropriate lone pairs)

(b) O=C=O (each double bonded, with appropriate lone pairs)

(c) O≡C-O (one triple bond, one single bond, with appropriate lone pairs)

Analysis:

1. Chemical Formula: CO₂
2. Valence Electrons: C (4) + 2 O (6 each) = 4 + 12 = 16 valence electrons
3. Evaluation:
   • (a) O-C-O: Contains 16 electrons. Oxygen and Carbon don't achieve octets. Incorrect.
   • (b) O=C=O: Contains 16 electrons. Each atom has a formal charge of zero. Carbon and Oxygen achieve octets. This structure is plausible.
   • (c) O≡C-O: Contains 16 electrons. Carbon and Oxygen achieve octets, but the formal charges are +1 on the right oxygen, -1 on the left oxygen, and 0 on the carbon. This structure is less plausible due to the formal charges.

Conclusion:

Structure (a) is definitely incorrect due to octet violations. Structure (b) is the most correct as it fulfills the octet rule and minimizes formal charges. Structure (c) is technically valid, but less likely due to the non-zero formal charges.

Example 2: Ozone (O₃)

Consider these Lewis structures for ozone:

(a) O=O-O (with appropriate lone pairs and formal charges)

(b) O-O-O (each single bonded, with appropriate lone pairs)

(c) O=O=O (each double bonded, with appropriate lone pairs)

Analysis:

1. Chemical Formula: O₃
2. Valence Electrons: 3 x O (6 each) = 18 valence electrons
3. Evaluation:
   • (a) O=O-O: Contains 18 electrons, fulfills the octet rule with one double bond and one single bond. This is the correct structure.
   • (b) O-O-O: Contains 20 electrons. Incorrect number of valence electrons.
   • (c) O=O=O: Contains 16 electrons. Incorrect number of valence electrons.

Conclusion:

Structures (b) and (c) are definitely incorrect because they do not have the correct number of valence electrons.

Example 3: Nitrogen Dioxide (NO₂)

Suppose you are given the following Lewis structures for NO₂:

(a) O=N-O (with appropriate lone pairs and one odd electron on N)

(b) O-N-O (each single bonded, with appropriate lone pairs and one odd electron on N)

(c) O=N=O (with appropriate lone pairs and one odd electron on N)

Analysis:

1. Chemical Formula: NO₂
2. Valence Electrons: N (5) + 2 x O (6 each) = 5 + 12 = 17 valence electrons
3. Evaluation:
   • (a) O=N-O: 17 electrons. Satisfies octet for two O atoms, but N has an odd number of electrons, implying a radical structure. Plausible.
   • (b) O-N-O: Number of electrons might be incorrect. Octets are not satisfied. Likely Incorrect.
   • (c) O=N=O: Contains 18 electrons. Incorrect number of valence electrons.

Conclusion:

Structure (c) is incorrect since it does not contain the correct number of valence electrons. The correct structure (a) shows one resonance form of nitrogen dioxide, acknowledging the odd electron count which makes NO₂ a radical.

Common Mistakes to Avoid

• Miscounting Valence Electrons: Ensure you are using the correct number of valence electrons for each atom.
• Forgetting Formal Charges: Always calculate formal charges to assess the stability of the structure.
• Ignoring Exceptions to the Octet Rule: Be aware of atoms that can have fewer or more than eight electrons.
• Overlooking Connectivity: Make sure the arrangement of atoms makes chemical sense.
• Assuming Symmetry: Not all molecules are symmetrical; draw the structure based on rules, not assumptions.

Advanced Considerations

Resonance Structures

Sometimes, multiple valid Lewis structures can be drawn for a single molecule. These are called resonance structures. The true structure is a hybrid of all resonance forms, which contributes to the overall stability of the molecule. If you encounter resonance, look for structures that minimize formal charges and fulfill the octet rule in the most balanced manner.

Hypervalency

Elements in the third row and beyond (e.g., sulfur, phosphorus) can sometimes accommodate more than eight electrons in their valence shell. This phenomenon is known as hypervalency and is attributed to the availability of d orbitals. When assessing Lewis structures of molecules with central atoms from these elements, consider the possibility of expanded octets.

Conclusion

Identifying incorrect Lewis structures requires a systematic approach and a strong understanding of fundamental principles. By carefully checking the number of valence electrons, octet rule compliance, formal charges, and atom connectivity, one can effectively evaluate and distinguish between correct and incorrect Lewis structures. This skill is crucial for predicting molecular properties, understanding chemical reactions, and advancing in the field of chemistry. Remember to practice regularly and consult reliable resources to reinforce your understanding.