Which Is Not A Correct Lewis Structure

Lewis structures, also known as Lewis dot diagrams, are visual representations of the bonding between atoms in a molecule and the lone pairs of electrons that may exist in the molecule. They are essential tools for understanding molecular structure, predicting molecular properties, and rationalizing chemical reactions. However, constructing a correct Lewis structure requires adhering to certain rules and guidelines. This article delves into the common errors that lead to incorrect Lewis structures and how to avoid them. Understanding these pitfalls is crucial for accurately representing molecules and predicting their behavior.

Understanding the Basics of Lewis Structures

Before diving into the errors, it's important to recap the fundamental principles of drawing Lewis structures:

• Counting Valence Electrons: The first step is to determine the total number of valence electrons in the molecule. This is done by summing the valence electrons of each atom in the molecule. Remember to adjust for ionic charges: add electrons for anions and subtract for cations.
• Drawing the Skeletal Structure: The next step is to draw a skeletal structure of the molecule, connecting atoms with single bonds. The least electronegative atom is usually placed in the center, except for hydrogen, which always occupies a terminal position.
• Distributing Electrons: Distribute the remaining valence electrons as lone pairs around the atoms, starting with the most electronegative atoms, until each atom (except hydrogen) has an octet of electrons.
• Forming Multiple Bonds: If any atom lacks an octet, form multiple bonds (double or triple bonds) by sharing lone pairs from adjacent atoms.
• Checking Formal Charges: After drawing the Lewis structure, calculate the formal charge on each atom. The formal charge is the difference between the number of valence electrons an atom has in its neutral state and the number of electrons it "owns" in the Lewis structure. The goal is to minimize formal charges on all atoms.

Common Errors in Drawing Lewis Structures

Many mistakes can lead to an incorrect Lewis structure, even when following the general steps. Here's an in-depth look at some of the most prevalent:

1. Incorrectly Counting Valence Electrons

This is a fundamental error that will cascade through the entire Lewis structure.

• Misidentifying Group Numbers: The number of valence electrons is directly related to the group number of the element in the periodic table. For example, Group 1 elements (alkali metals) have 1 valence electron, Group 2 elements (alkaline earth metals) have 2, Group 16 elements (chalcogens) have 6, and Group 17 elements (halogens) have 7. Errors can occur if the group numbers are misidentified, especially with transition metals, which often have variable valencies.
• Forgetting Ionic Charges: When dealing with polyatomic ions, it's crucial to account for the ionic charge. For example, the sulfate ion (SO₄²⁻) has a 2- charge, meaning it has two extra electrons compared to the neutral molecule.
• Ignoring Polyatomic Ions: In complex molecules, polyatomic ions might be present. Students often forget to treat these ions as a unit and fail to account for their charge and electron count. For example, ammonium nitrate (NH₄NO₃) contains both ammonium (NH₄⁺) and nitrate (NO₃⁻) ions, each requiring a separate consideration of valence electrons.

Example: Consider drawing the Lewis structure for the carbonate ion (CO₃²⁻).

• Carbon (C) has 4 valence electrons.
• Each oxygen (O) has 6 valence electrons, so 3 oxygens have 3 * 6 = 18 valence electrons.
• The 2- charge indicates 2 extra electrons.
• Total valence electrons = 4 + 18 + 2 = 24 electrons.

If you incorrectly count the valence electrons, the resulting Lewis structure will be flawed, with either too many or too few electrons distributed.

2. Incorrect Skeletal Structure

The arrangement of atoms in the skeletal structure is crucial for a correct Lewis structure.

• Misidentifying the Central Atom: The central atom is usually the least electronegative atom (excluding hydrogen). However, there are exceptions. For example, in molecules like hypochlorous acid (HOCl), hydrogen is bonded to oxygen, and oxygen is bonded to chlorine.
• Not Considering Symmetry: Molecules often have symmetrical structures. Failing to recognize this symmetry can lead to incorrect arrangements. For example, sulfur dioxide (SO₂) has a bent structure with sulfur as the central atom bonded to two oxygen atoms.
• Wrong Connectivity: The order in which atoms are connected matters. For example, nitrous acid (HNO₂) has the structure HONO, not OHNO. Knowing the common bonding patterns of elements is helpful in determining the correct connectivity.

Example: Consider drawing the Lewis structure for sulfuric acid (H₂SO₄).

• Sulfur (S) is the central atom, bonded to four oxygen (O) atoms.
• Each oxygen atom is bonded to a hydrogen (H) atom.
• The correct skeletal structure is H-O-S-O-H with the other two oxygen atoms directly bonded to sulfur.

An incorrect skeletal structure will lead to an incorrect distribution of electrons and inaccurate representation of the molecule.

3. Violating the Octet Rule

The octet rule states that atoms tend to gain, lose, or share electrons to achieve a full outer shell of eight electrons. While this rule is generally followed, there are important exceptions.

• Ignoring Exceptions to the Octet Rule:
  • Incomplete Octets: Some atoms, like beryllium (Be) and boron (B), can be stable with fewer than eight electrons around them. For example, in boron trifluoride (BF₃), boron has only six electrons.
  • Expanded Octets: Elements in the third period and beyond (such as sulfur, phosphorus, and chlorine) 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 Number of Electrons: Molecules with an odd number of valence electrons (free radicals) cannot satisfy the octet rule for all atoms. For example, in nitrogen dioxide (NO₂), one atom will have an unpaired electron.
• Forcing the Octet Rule: Trying to force an atom to obey the octet rule when it is an exception can lead to an incorrect Lewis structure with unnecessary formal charges and inaccurate bonding.

Example: Consider drawing the Lewis structure for phosphorus pentachloride (PCl₅).

• Phosphorus (P) is the central atom, bonded to five chlorine (Cl) atoms.
• Phosphorus has 5 valence electrons, and each chlorine has 7, totaling 5 + (5 * 7) = 40 electrons.
• The correct Lewis structure has phosphorus with ten electrons around it, violating the octet rule but accurately representing the molecule.

4. Incorrect Distribution of Lone Pairs and Bonds

Even when the skeletal structure is correct, incorrect placement of lone pairs and bonds can lead to a flawed Lewis structure.

• Not Satisfying Octets: Failing to distribute enough electrons to satisfy the octets (or duets for hydrogen) of all atoms is a common error.
• Overfilling Octets: Assigning too many electrons to atoms that cannot accommodate an expanded octet will also lead to an incorrect structure.
• Incorrect Placement of Multiple Bonds: Multiple bonds (double or triple bonds) should be placed strategically to minimize formal charges and satisfy the octet rule where possible.

Example: Consider drawing the Lewis structure for carbon dioxide (CO₂).

• Carbon (C) is the central atom, bonded to two oxygen (O) atoms.
• Carbon has 4 valence electrons, and each oxygen has 6, totaling 4 + (2 * 6) = 16 electrons.
• The correct Lewis structure has double bonds between carbon and each oxygen atom (O=C=O), with two lone pairs on each oxygen atom.

If you incorrectly place the double bonds, such as having a single bond on one side and a triple bond on the other, the Lewis structure will be incorrect and have non-zero formal charges.

5. Incorrect Calculation and Interpretation of Formal Charges

Formal charges help determine the most stable and correct Lewis structure among several possibilities.

• Miscalculating Formal Charges: The formal charge is calculated as:

  Formal Charge = (Valence Electrons) - (Non-bonding Electrons) - (1/2 Bonding Electrons)

  Errors can occur if any of these components are miscounted.

• Ignoring Formal Charges: Students sometimes draw a Lewis structure without considering the formal charges, which can lead to a less stable and incorrect representation.

• Not Minimizing Formal Charges: The most stable Lewis structure is usually the one with the smallest formal charges on all atoms. Ideally, the formal charges should be zero. If formal charges are unavoidable, they should be placed on the most electronegative atoms.

Example: Consider drawing the Lewis structure for the cyanate ion (OCN⁻). There are multiple possible resonance structures:

• O=C=N⁻ (Formal charges: O = 0, C = 0, N = -1)
• ⁻O-C≡N (Formal charges: O = -1, C = 0, N = 0)
• O≡C-N⁻ (Formal charges: O = +1, C = 0, N = -2)

The first structure (O=C=N⁻) is the most stable because it minimizes the formal charges on all atoms.

6. Not Recognizing Resonance Structures

Resonance occurs when a molecule or ion can be represented by two or more Lewis structures that differ only in the arrangement of electrons.

• Failing to Draw All Resonance Structures: Some molecules have multiple valid resonance structures. Failing to draw all of them can lead to an incomplete understanding of the molecule's electronic structure.
• Treating Resonance Structures as Different Molecules: Resonance structures are not different molecules; they are different representations of the same molecule. The actual molecule is a hybrid of all resonance structures.
• Incorrectly Drawing Resonance Structures: Resonance structures must have the same skeletal structure; only the placement of electrons can differ. You cannot move atoms to create resonance structures.

Example: Consider drawing the Lewis structure for ozone (O₃).

• Ozone has two major resonance structures: O=O-O and O-O=O.
• The actual ozone molecule is a hybrid of these two structures, with the bond lengths between the oxygen atoms being equal and intermediate between a single and a double bond.

7. Overcomplicating Structures

Sometimes, students try to create excessively complex Lewis structures when a simpler structure is more appropriate.

• Adding Unnecessary Bonds: Adding extra bonds to atoms that already satisfy the octet rule or have a minimized formal charge is a common mistake.
• Creating Unrealistic Structures: Attempting to force molecules into specific shapes or bonding patterns that are not supported by experimental evidence can lead to unrealistic and incorrect Lewis structures.

Example: Consider drawing the Lewis structure for water (H₂O).

• The correct Lewis structure has oxygen as the central atom bonded to two hydrogen atoms, with two lone pairs on the oxygen atom.
• Trying to add a double bond between oxygen and hydrogen or creating a linear structure would be incorrect.

Practical Tips for Drawing Correct Lewis Structures

To avoid these common errors and draw correct Lewis structures, consider the following tips:

• Practice Regularly: The more you practice drawing Lewis structures, the better you will become at recognizing patterns and avoiding mistakes.
• Use a Systematic Approach: Follow a step-by-step approach to ensure that you don't miss any crucial steps.
• Check Your Work: After drawing a Lewis structure, always double-check your work to ensure that you have counted the valence electrons correctly, satisfied the octet rule (where applicable), and minimized formal charges.
• Consult Reliable Resources: Use textbooks, online resources, and expert guidance to clarify any doubts or uncertainties.
• Understand Electronegativity: A good grasp of electronegativity trends helps in predicting the central atom and distribution of electrons.
• Learn Common Bonding Patterns: Familiarize yourself with the typical bonding patterns of common elements to make informed decisions about the skeletal structure.
• Use Molecular Modeling Kits: Physical models can help visualize the three-dimensional structure of molecules and aid in drawing accurate Lewis structures.

Examples of Incorrect Lewis Structures and Their Corrections

Let's examine a few specific examples of incorrect Lewis structures and how to correct them:

1. Carbon Monoxide (CO)

• Incorrect Structure: C=O with two lone pairs on each atom.
• Error: This structure does not minimize formal charges.
• Correct Structure: ⁻C≡O⁺ with one lone pair on carbon and one lone pair on oxygen. This structure, though having formal charges, is more stable due to the electronegativity of oxygen.

2. Nitrogen Dioxide (NO₂)

• Incorrect Structure: O=N=O with complete octets on all atoms.
• Error: This structure requires nitrogen to have an expanded octet and does not account for the odd number of valence electrons.
• Correct Structure: A resonance hybrid with one N=O double bond and one N-O single bond, with the unpaired electron on the nitrogen atom.

3. Sulfur Dioxide (SO₂)

• Incorrect Structure: O-S-O with single bonds only and complete octets on all atoms.
• Error: This structure does not minimize formal charges and does not account for the expanded octet capability of sulfur.
• Correct Structure: A resonance hybrid with one S=O double bond and one S-O single bond, and a lone pair on sulfur.

4. Methane (CH₄)

• Incorrect Structure: Carbon with double bonds to two hydrogen atoms.
• Error: Hydrogen can only form one bond, and carbon's octet is exceeded.
• Correct Structure: Carbon with single bonds to four hydrogen atoms.

Conclusion

Drawing correct Lewis structures is a fundamental skill in chemistry. Avoiding common errors such as incorrectly counting valence electrons, drawing incorrect skeletal structures, violating the octet rule, miscalculating formal charges, and not recognizing resonance structures is crucial for accurately representing molecules and predicting their properties. By understanding these pitfalls and following a systematic approach, you can improve your ability to draw correct Lewis structures and deepen your understanding of chemical bonding. Consistent practice, attention to detail, and a solid grasp of the basic principles of chemistry are key to mastering this essential skill. Remember that Lewis structures are simplified models and may not perfectly represent the electronic structure of all molecules, but they provide a valuable framework for understanding chemical bonding and molecular behavior.