Decide Whether These Proposed Lewis Structures Are Reasonable.

Lewis structures are fundamental tools in chemistry, providing a visual representation of the bonding between atoms in a molecule, as well as any lone pairs of electrons that may exist. However, not all proposed Lewis structures are created equal. Some adhere to the established rules and principles of chemical bonding, while others violate them, leading to unreasonable or incorrect depictions of the molecule. Deciding whether a proposed Lewis structure is reasonable involves a careful evaluation of several key factors, including the octet rule, formal charges, electronegativity, and resonance. This article delves into the criteria and procedures for assessing the reasonableness of proposed Lewis structures, providing detailed examples and guidelines to help you navigate this crucial aspect of chemical understanding.

Understanding Lewis Structures

Before we can evaluate the reasonableness of proposed Lewis structures, it's essential to have a solid grasp of what they are and the principles they represent. A Lewis structure is a diagram that shows the bonding between atoms of a molecule and the lone pairs of electrons that may exist in the molecule.

Here are the key components and principles of Lewis structures:

• Valence Electrons: Only valence electrons (electrons in the outermost shell) are shown in Lewis structures.
• 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. Hydrogen is an exception, as it only needs two electrons to achieve a full outer shell (duet rule).
• Bonds: Covalent bonds are represented by lines between atoms, with each line representing a pair of shared electrons.
• Lone Pairs: Non-bonding pairs of electrons, also known as lone pairs, are represented by dots around an atom.
• Formal Charge: Formal charge is the hypothetical charge an atom would have if all bonding electrons were shared equally between atoms. It's calculated using the formula: Formal Charge = (Valence Electrons) - (Non-bonding Electrons) - (1/2 Bonding Electrons).
• Resonance: When multiple valid Lewis structures can be drawn for a molecule, the actual structure is a resonance hybrid, an average of all the possible structures.

Criteria for Evaluating Lewis Structures

To determine whether a proposed Lewis structure is reasonable, consider the following criteria:

1. Correct Number of Valence Electrons: The total number of valence electrons shown in the Lewis structure must match the sum of the valence electrons of all the atoms in the molecule.
2. Octet Rule Compliance: Each atom (except hydrogen) should have an octet of electrons surrounding it. Hydrogen should have a duet.
3. Formal Charges: The formal charges on atoms should be minimized. A Lewis structure with smaller formal charges is generally more reasonable than one with larger formal charges.
4. Electronegativity: Negative formal charges should be placed on more electronegative atoms, and positive formal charges should be placed on less electronegative atoms.
5. Resonance Structures: If multiple Lewis structures can be drawn that satisfy the octet rule and other criteria, consider resonance. The resonance hybrid will represent the actual structure of the molecule.
6. Molecular Geometry: Consider the three-dimensional arrangement of atoms in the molecule, which can be predicted using VSEPR (Valence Shell Electron Pair Repulsion) theory. The Lewis structure should be consistent with the predicted molecular geometry.

Step-by-Step Procedure for Assessing Lewis Structures

Here's a step-by-step procedure for assessing the reasonableness of proposed Lewis structures:

• Step 1: Count Valence Electrons: Determine the total number of valence electrons for all atoms in the molecule.
• Step 2: Draw a Skeletal Structure: Draw a basic structure of the molecule, connecting atoms with single bonds. Place the least electronegative atom in the center (except for hydrogen).
• Step 3: Complete Octets: Add lone pairs to the surrounding atoms to satisfy the octet rule (or duet rule for hydrogen).
• Step 4: Place Remaining Electrons: If there are remaining valence electrons, place them on the central atom.
• Step 5: Minimize Formal Charges: If the central atom does not have an octet, form multiple bonds by moving lone pairs from the surrounding atoms.
• Step 6: Calculate Formal Charges: Calculate the formal charge on each atom in the Lewis structure.
• Step 7: Evaluate Reasonableness: Evaluate the reasonableness of the Lewis structure based on the criteria outlined above.
• Step 8: Consider Resonance: If multiple reasonable Lewis structures can be drawn, consider resonance.
• Step 9: Refine the Structure: If the Lewis structure is unreasonable, refine it by adjusting the arrangement of electrons or bonds.

Detailed Examples

Let's illustrate these principles with several examples:

Carbon Dioxide (CO2)

Suppose we are presented with the following proposed Lewis structure for carbon dioxide:

O-C-O

1. Count Valence Electrons: Carbon (C) has 4 valence electrons, and each oxygen (O) has 6 valence electrons. Total = 4 + 2(6) = 16 valence electrons.

2. Skeletal Structure: The basic structure is O-C-O.

3. Complete Octets: Add lone pairs to oxygen atoms: :O-C-O: This structure uses 4 electrons for the single bonds and needs more electrons for octets.

4. Place Remaining Electrons: Add lone pairs to oxygen atoms to satisfy the octet rule: :O=C=O: Each oxygen has 6 electrons (2 bonding, 4 non-bonding).

5. Minimize Formal Charges: Calculate formal charges:

   • Carbon: 4 (valence) - 0 (non-bonding) - 1/2(8 bonding) = 0
   • Each Oxygen: 6 (valence) - 4 (non-bonding) - 1/2(4 bonding) = 0

6. Evaluate Reasonableness: The Lewis structure :O=C=O: is reasonable because each atom has an octet, and the formal charges are minimized (all are zero).

Ozone (O3)

Consider a proposed Lewis structure for ozone:

O-O-O

1. Count Valence Electrons: Each oxygen atom has 6 valence electrons. Total = 3(6) = 18 valence electrons.

2. Skeletal Structure: The basic structure is O-O-O.

3. Complete Octets: Add lone pairs to oxygen atoms: O-O-O

4. Place Remaining Electrons: To satisfy the octet rule, we form a double bond between two oxygen atoms and add lone pairs: O=O-O

5. Minimize Formal Charges: Calculate formal charges:

   • Double-bonded Oxygen: 6 (valence) - 4 (non-bonding) - 1/2(4 bonding) = 0
   • Central Oxygen: 6 (valence) - 2 (non-bonding) - 1/2(6 bonding) = +1
   • Single-bonded Oxygen: 6 (valence) - 6 (non-bonding) - 1/2(2 bonding) = -1

6. Evaluate Reasonableness: The Lewis structure O=O-O has formal charges, but they are minimized. However, we can also draw another valid Lewis structure with the double bond on the other side: O-O=O. This indicates resonance.

7. Consider Resonance: Ozone exhibits resonance, and the actual structure is a resonance hybrid of the two Lewis structures.

8. Refine the Structure: The resonance hybrid is a more accurate representation of ozone, with bond order of 1.5 between each oxygen atom.

Nitrate Ion (NO3-)

Given a proposed Lewis structure for the nitrate ion:

O-N-O | O

1. Count Valence Electrons: Nitrogen (N) has 5 valence electrons, each oxygen (O) has 6 valence electrons, and there is an additional electron due to the negative charge. Total = 5 + 3(6) + 1 = 24 valence electrons.

2. Skeletal Structure: The basic structure is:

   O | O-N-O

3. Complete Octets: Add lone pairs to oxygen atoms:

   :O: | :O-N-O:

4. Place Remaining Electrons: To satisfy the octet rule, form a double bond between nitrogen and one of the oxygen atoms:

   :O: || :O-N-O:

5. Minimize Formal Charges: Calculate formal charges:

   • Nitrogen: 5 (valence) - 0 (non-bonding) - 1/2(8 bonding) = +1
   • Double-bonded Oxygen: 6 (valence) - 4 (non-bonding) - 1/2(4 bonding) = 0
   • Single-bonded Oxygen (2): 6 (valence) - 6 (non-bonding) - 1/2(2 bonding) = -1

6. Evaluate Reasonableness: The Lewis structure has formal charges. However, we can draw resonance structures with the double bond on different oxygen atoms.

7. Consider Resonance: Nitrate exhibits resonance, and the actual structure is a resonance hybrid of the three Lewis structures. Each oxygen atom has a partial negative charge, and the nitrogen atom has a partial positive charge.

Sulfur Hexafluoride (SF6)

Consider a proposed Lewis structure for sulfur hexafluoride:

F | F-S-F | F

| F-F

1. Count Valence Electrons: Sulfur (S) has 6 valence electrons, and each fluorine (F) has 7 valence electrons. Total = 6 + 6(7) = 48 valence electrons.

2. Skeletal Structure: The basic structure is:

   F | F-S-F | F

   | F-F

3. Complete Octets: Sulfur has 12 electrons around it, which is an expanded octet. Fluorine has an octet when six fluorine atoms are bonded to sulfur with single bonds.

4. Evaluate Reasonableness: Sulfur can have an expanded octet because it is in the third period and has available d-orbitals. The Lewis structure is reasonable.

5. Molecular Geometry: SF6 has an octahedral geometry.

Incorrect Lewis Structures

Now let's examine some examples of incorrect Lewis structures:

Incorrect Structure for Water (H2O)

Suppose we are given the following incorrect Lewis structure for water:

H-H-O

1. Count Valence Electrons: Each hydrogen (H) has 1 valence electron, and oxygen (O) has 6 valence electrons. Total = 2(1) + 6 = 8 valence electrons.
2. Skeletal Structure: The proposed structure is H-H-O.
3. Complete Octets: This structure violates the octet rule and duet rule. Hydrogen can only form one bond. The correct structure should be H-O-H, with two lone pairs on oxygen.
4. Evaluate Reasonableness: This Lewis structure is unreasonable because it violates the fundamental principles of chemical bonding.

Incorrect Structure for Carbon Monoxide (CO)

Suppose we are given the following incorrect Lewis structure for carbon monoxide:

C=O

1. Count Valence Electrons: Carbon (C) has 4 valence electrons, and oxygen (O) has 6 valence electrons. Total = 4 + 6 = 10 valence electrons.
2. Skeletal Structure: The proposed structure is C=O.
3. Complete Octets: With a double bond, carbon has 6 electrons, and oxygen has 8 electrons. This structure does not satisfy the octet rule for carbon. The correct structure is C≡O, with a triple bond.
4. Evaluate Reasonableness: This Lewis structure is unreasonable because it does not satisfy the octet rule for carbon. The correct Lewis structure for carbon monoxide is :C≡O:, which satisfies the octet rule for both carbon and oxygen, with formal charges of -1 on carbon and +1 on oxygen.

Common Pitfalls to Avoid

When evaluating Lewis structures, be aware of these common pitfalls:

• Forgetting to Count Valence Electrons: Always start by accurately counting the total number of valence electrons in the molecule or ion.
• Violating the Octet Rule: Ensure that each atom (except hydrogen) has an octet of electrons surrounding it.
• Ignoring Formal Charges: Calculate and consider formal charges when evaluating the reasonableness of a Lewis structure.
• Failing to Consider Resonance: If multiple reasonable Lewis structures can be drawn, consider resonance.
• Misinterpreting Electronegativity: Place negative formal charges on more electronegative atoms and positive formal charges on less electronegative atoms.
• Neglecting Molecular Geometry: The Lewis structure should be consistent with the predicted molecular geometry based on VSEPR theory.

Advanced Considerations

Beyond the basic rules, there are some advanced considerations to keep in mind:

Expanded Octets

Some atoms, particularly those in the third period and beyond, can accommodate more than eight electrons around them. This is known as an expanded octet. Examples include sulfur in SF6 and phosphorus in PCl5. The ability to have an expanded octet is due to the availability of d-orbitals.

Incomplete Octets

Some atoms, such as beryllium (Be) and boron (B), may have fewer than eight electrons around them. For example, in boron trifluoride (BF3), boron has only six electrons.

Radicals

Molecules with an odd number of valence electrons are called radicals and cannot satisfy the octet rule for all atoms. For example, nitric oxide (NO) has 11 valence electrons.

FAQ on Lewis Structures

• Q: What is the octet rule, and why is it important?

  • A: The octet rule states that atoms tend to gain, lose, or share electrons to achieve a full outer electron shell with eight electrons. It is important because it helps predict the stability and reactivity of molecules.

• Q: How do I calculate formal charge?

  • A: Formal Charge = (Valence Electrons) - (Non-bonding Electrons) - (1/2 Bonding Electrons).

• Q: What is resonance, and why is it important?

  • A: Resonance occurs when multiple valid Lewis structures can be drawn for a molecule. The actual structure is a resonance hybrid, an average of all possible structures. Resonance is important because it provides a more accurate representation of the electron distribution in the molecule and often leads to increased stability.

• Q: What is VSEPR theory, and how does it relate to Lewis structures?

  • A: VSEPR (Valence Shell Electron Pair Repulsion) theory is a model used to predict the three-dimensional arrangement of atoms in a molecule based on the repulsion between electron pairs. The Lewis structure provides the basis for predicting the molecular geometry using VSEPR theory.

Conclusion

Deciding whether a proposed Lewis structure is reasonable requires a careful evaluation of several key factors, including the number of valence electrons, the octet rule, formal charges, electronegativity, and resonance. By following a systematic approach and considering these criteria, you can accurately assess the reasonableness of Lewis structures and gain a deeper understanding of chemical bonding and molecular structure. This knowledge is essential for success in chemistry and related fields.