Which Of The Following Lewis Structures Is Correct

Which of the following Lewis Structures is Correct? A Deep Dive into Resonance, Formal Charge, and Octet Rule Exceptions

Choosing the "correct" Lewis structure isn't always straightforward. While the goal is to represent a molecule's valence electrons and bonding in a way that adheres to the octet rule (or its exceptions), multiple seemingly valid structures can often be drawn. This article explores the key principles – resonance, formal charge, and octet rule exceptions – to determine the most accurate and stable Lewis structure representation for a given molecule.

Understanding Lewis Structures: A Recap

Lewis structures, also known as electron dot diagrams, visually represent the arrangement of valence electrons in a molecule. They depict bonding pairs (shared electrons) as lines and lone pairs (unshared electrons) as dots. The fundamental principles guiding their construction include:

• Valence Electrons: The number of valence electrons dictates the number of bonds and lone pairs a central atom can form.
• Octet Rule: Atoms (except hydrogen and some exceptions discussed later) tend to gain, lose, or share electrons to achieve a stable configuration of eight valence electrons.
• Formal Charge: A tool used to evaluate the distribution of electrons and identify the most likely structure.
• Resonance: A phenomenon where a molecule's actual structure is a hybrid of multiple Lewis structures.

The Importance of Formal Charge in Determining the Correct Lewis Structure

Formal charge helps determine the most plausible Lewis structure when multiple options exist. It's calculated as:

Formal Charge = (Valence electrons) - (Non-bonding electrons) - ½(Bonding electrons)

A lower formal charge on each atom is generally preferred, and a structure with formal charges closest to zero is considered more stable. Furthermore, negative formal charges should ideally reside on the more electronegative atoms.

Example: Let's consider the nitrate ion (NO₃⁻). Three plausible Lewis structures can be drawn, each differing in the placement of the double bond. Each structure shows one N=O double bond and two N-O single bonds. However, calculating formal charges reveals that all three structures are equally valid resonance contributors.

• Structure 1: N=O double bond on top, with single bonds on the other two oxygens
• Structure 2: N=O double bond in the middle, with single bonds on the other two oxygens
• Structure 3: N=O double bond at the bottom, with single bonds on the other two oxygens

In each structure, the formal charge on the nitrogen atom is zero. But for the singly bonded oxygen atoms, the formal charge will be -1, whereas for the doubly bonded oxygen, the formal charge will be zero.

The actual structure of the nitrate ion is a resonance hybrid where the electron density is delocalized across all three N-O bonds, resulting in bond lengths intermediate between single and double bonds.

The Role of Resonance in Lewis Structures

Resonance occurs when a molecule can be represented by two or more Lewis structures that differ only in the placement of electrons (not atoms). The actual molecule is a hybrid of these resonance structures, exhibiting properties intermediate between the contributing structures. The true structure is a blend of characteristics represented in resonance structures. Each individual resonance structure does not accurately represent the molecule; rather, they contribute to the overall description.

The significance of resonance lies in its ability to explain certain molecular properties:

• Bond Lengths: In molecules with resonance, bond lengths between atoms are often equal and intermediate between single and double bonds (e.g., benzene).
• Stability: Resonance structures usually stabilize a molecule, making it less reactive than predicted based on a single Lewis structure.
• Reactivity: Resonance affects where a molecule is most likely to react.

Exceptions to the Octet Rule: Expanding and Incomplete Octets

While the octet rule is a useful guideline, some molecules deviate from it:

1. Expanded Octets: Elements in the third period and beyond (e.g., phosphorus, sulfur, chlorine) can accommodate more than eight valence electrons in their valence shell due to the availability of d-orbitals. This is commonly seen in compounds like phosphorus pentachloride (PCl₅) and sulfur hexafluoride (SF₆).

Example: In PCl₅, phosphorus has five bonds (10 electrons) surrounding it, exceeding the octet rule.

2. Incomplete Octets: Certain molecules, particularly those containing boron or beryllium, may have fewer than eight valence electrons. These elements are sometimes content with six electrons (as in BF₃) or even only four (as in BeCl₂). These molecules are electron-deficient and often exhibit greater reactivity compared to molecules with complete octets.

Example: Boron trifluoride (BF₃) has only six electrons around the boron atom.

Practical Application: Identifying the Correct Lewis Structure Step-by-Step

Let's analyze a specific example to illustrate how to determine the most appropriate Lewis structure: Consider the molecule sulfur dioxide (SO₂).

Step 1: Count Valence Electrons:

• Sulfur (S): 6 valence electrons
• Oxygen (O): 6 valence electrons each (2 oxygen atoms = 12 electrons)
• Total: 18 valence electrons

Step 2: Draw a Skeletal Structure:

Start by arranging the atoms. Typically, the less electronegative atom (sulfur in this case) is placed in the center. This gives us: O-S-O

Step 3: Place Single Bonds:

Connect each atom with a single bond (2 electrons per bond). This uses 4 electrons, leaving 14 electrons.

Step 4: Distribute Remaining Electrons:

Add remaining electrons to atoms as lone pairs, starting with the outer atoms (oxygen). Each oxygen gets 3 lone pairs (6 electrons), using 12 electrons. This leaves 2 electrons.

Step 5: Satisfy the Octet Rule (if possible):

The sulfur atom currently only has 6 electrons around it. To complete the octet, we move a lone pair from one of the oxygen atoms to form a double bond with sulfur.

Step 6: Calculate Formal Charges:

Calculate the formal charges for each atom in this structure. You will find that one oxygen atom has a formal charge of -1, the sulfur atom has a formal charge of 0, and the other oxygen atom has a formal charge of +1.

Step 7: Consider Resonance:

Since we could have placed the double bond with either oxygen atom, we have two resonance structures. The actual molecule is a resonance hybrid of these two structures, exhibiting partial double bond character in both S-O bonds.

Step 8: Choose the "Best" Structure:

While both resonance structures are valid, the one that contributes more to the hybrid structure is generally considered the "best" representation. This choice often involves looking at other factors, like electronegativity or experimental data. In the case of SO₂, the structure minimizing the formal charges would be given more weight.

Conclusion: A Holistic Approach to Lewis Structures

Determining the "correct" Lewis structure is not about finding a single, absolute answer but rather about systematically applying principles like formal charge, the octet rule (and its exceptions), and the concept of resonance to arrive at the most accurate and informative representation of a molecule's bonding and electron distribution. By carefully considering these factors, you can confidently build Lewis structures that accurately reflect the properties and behavior of molecules. Remember, practicing with various examples and understanding the underlying principles is key to mastering this fundamental concept in chemistry.