Decide Whether The Proposed Lewis Structure Below Is Reasonable

Let's delve into the process of evaluating the reasonableness of a proposed Lewis structure. It’s a crucial skill in understanding chemical bonding and molecular behavior, allowing us to predict properties and reactivity. This comprehensive guide will provide you with the tools and knowledge to confidently assess Lewis structures.

The Fundamentals of Lewis Structures

A Lewis structure, also known as an electron dot structure, is a visual representation of the bonding between atoms in a molecule, as well as any lone pairs of electrons that may exist. It's a simplified model, but it provides valuable insight into how atoms share electrons to achieve a stable electron configuration, usually resembling that of a noble gas (octet rule).

Before we can decide if a proposed Lewis structure is reasonable, we need to grasp the core principles involved in creating them:

• Valence Electrons: The number of valence electrons (electrons in the outermost shell) determines how many bonds an atom can form. This number corresponds to the group number of the element in the periodic table (for main group elements).
• Octet Rule (and Exceptions): Atoms "want" to achieve a full outer shell of eight electrons (an octet), similar to the noble gases. They achieve this by sharing electrons through covalent bonds. Hydrogen is an exception, only needing two electrons to fill its outer shell (duet rule). Other exceptions include elements like Boron (often stable with 6 electrons) and elements that can expand their octet (like Sulfur and Phosphorus).
• Formal Charge: Formal charge helps determine the "best" Lewis structure when multiple possibilities exist. It's the difference between the number of valence electrons an atom should have and the number it appears to have in the Lewis structure (number of lone pair electrons + half the number of bonding electrons). The ideal Lewis structure minimizes formal charges.
• Electronegativity: Electronegativity is the measure of an atom's ability to attract electrons in a chemical bond. When assigning formal charges, it's preferable to place negative formal charges on the more electronegative atoms.

Steps to Determine the Reasonableness of a Proposed Lewis Structure

Here’s a systematic approach to evaluating whether a given Lewis structure is reasonable:

1. Verify the Total Number of Valence Electrons:

• This is the absolute first step. If the total number of electrons in the structure doesn't match the sum of valence electrons from all the atoms, the structure is incorrect.
• How to do it: Sum the number of valence electrons for each atom in the molecule or ion. Remember to add electrons for negative charges (anions) and subtract electrons for positive charges (cations).
• Example: For CO2, Carbon has 4 valence electrons, and each Oxygen has 6. Total: 4 + 6 + 6 = 16 valence electrons. The Lewis structure must have exactly 16 electrons represented.

2. Check for Correct Connectivity:

• Is the proposed arrangement of atoms logical? While this can be challenging without prior knowledge of the molecule, some rules of thumb can help:
  • Central Atom: The least electronegative atom is often the central atom (except for Hydrogen, which is always terminal).
  • Symmetry: Many molecules exhibit some degree of symmetry. Look for repeating units or patterns.
  • Common Bonding Patterns: Be aware of common bonding patterns for certain elements (e.g., Carbon often forms four bonds).
• If the proposed connectivity is highly unusual or contradicts established chemical principles, it’s a red flag.

3. Confirm the Octet Rule is Satisfied (Where Possible):

• Do all atoms (except Hydrogen) have eight electrons around them? (Remember Hydrogen only needs two).
• Count the Electrons: Count the number of electrons directly around each atom in the Lewis structure. This includes both bonding electrons (electrons shared in bonds) and lone pair electrons.
• Exceptions: Be mindful of the exceptions to the octet rule:
  • Hydrogen: Forms one bond and has 2 electrons.
  • Boron: Often stable with 6 electrons.
  • Beryllium: Often stable with 4 electrons.
  • Elements in Period 3 and below: Can expand their octet to accommodate more than 8 electrons. This is because they have available d orbitals that can participate in bonding. Examples include S, P, Cl, and Xe.

4. Evaluate Formal Charges:

• Calculate the formal charge on each atom in the Lewis structure.
• Formal Charge Calculation: Formal Charge = (Valence Electrons) - (Non-bonding Electrons) - (1/2 Bonding Electrons)
• Interpreting 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.
  • Minimize formal charges whenever possible. A Lewis structure with all atoms having a formal charge of zero is generally the most stable.
  • Negative formal charges should reside on the more electronegative atoms.

5. Assess Resonance Structures (If Applicable):

• If multiple Lewis structures can be drawn for a molecule that only differ in the arrangement of electrons (not atoms), these are resonance structures.
• Identify Resonance: Look for molecules with multiple bonds (double or triple bonds) and lone pairs that can be moved around to create different arrangements.
• Evaluate Resonance Structures:
  • All resonance structures must be valid Lewis structures (satisfy the octet rule, have the correct number of valence electrons).
  • The "best" resonance structure is the one that minimizes formal charges and places negative formal charges on the most electronegative atoms.
  • The actual molecule is a hybrid of all resonance structures, with the most stable resonance structure contributing more to the hybrid.

6. Consider Bond Lengths and Bond Energies (Advanced):

• While not always readily apparent from the Lewis structure alone, the type of bond (single, double, triple) influences bond length and bond energy.
• Bond Order: The number of chemical bonds between a pair of atoms. Single bond = bond order of 1, double bond = bond order of 2, triple bond = bond order of 3.
• Bond Length: Shorter bond lengths indicate stronger bonds. Triple bonds are shorter than double bonds, which are shorter than single bonds.
• Bond Energy: Higher bond energies indicate stronger bonds. Triple bonds have higher bond energies than double bonds, which have higher bond energies than single bonds.
• If the proposed Lewis structure suggests unusual bond lengths or energies compared to known values, it might be unreasonable. This usually requires comparing the structure to experimental data or using computational chemistry methods.

Examples of Evaluating Lewis Structures

Let's apply these steps to some examples:

Example 1: Carbon Dioxide (CO2)

• Proposed Lewis Structure: O=C=O
• Step 1: Valence Electrons: C (4) + 2 * O (6) = 16 valence electrons. The proposed structure has 16 electrons (4 bonds * 2 electrons/bond + 4 lone pairs * 2 electrons/lone pair), so it's correct.
• Step 2: Connectivity: Oxygen atoms are bonded to the central Carbon atom, which is reasonable.
• Step 3: Octet Rule: Each atom has an octet (8 electrons around each atom).
• Step 4: Formal Charges:
  • Carbon: 4 (valence) - 0 (non-bonding) - 1/2 * 8 (bonding) = 0
  • Oxygen: 6 (valence) - 4 (non-bonding) - 1/2 * 4 (bonding) = 0
• Conclusion: The proposed Lewis structure for CO2 (O=C=O) is reasonable because it satisfies the octet rule, has the correct number of valence electrons, has reasonable connectivity, and minimizes formal charges.

Example 2: Ozone (O3)

• Proposed Lewis Structure: O=O-O (one double bond, one single bond) with appropriate lone pairs.
• Step 1: Valence Electrons: 3 * O (6) = 18 valence electrons. The structure should have 18 electrons.
• Step 2: Connectivity: A chain of three Oxygen atoms is possible.
• Step 3: Octet Rule: Check if all Oxygen atoms have an octet. By placing lone pairs strategically, this can be achieved.
• Step 4: Formal Charges: Calculating formal charges reveals that the central oxygen has a +1 formal charge, one terminal oxygen has a -1 formal charge, and the other terminal oxygen has a 0 formal charge.
• Step 5: Resonance Structures: Ozone exhibits resonance. The double bond can be moved to the other oxygen atom, creating a second valid Lewis structure. Both structures are important contributors to the overall structure of ozone.
• Conclusion: The proposed Lewis structure is reasonable, but it’s important to recognize that it's just one resonance structure of ozone. The actual ozone molecule is a hybrid of both resonance forms.

Example 3: Nitrogen Dioxide (NO2)

• Proposed Lewis Structure: O=N-O (one double bond, one single bond)
• Step 1: Valence Electrons: N (5) + 2 * O (6) = 17 valence electrons. This is an odd number of electrons.
• Step 2: Connectivity: Oxygen atoms bonded to a central Nitrogen atom is reasonable.
• Step 3: Octet Rule: It is impossible to satisfy the octet rule for all atoms in NO2 because there are an odd number of electrons. Nitrogen will have one unpaired electron.
• Step 4: Formal Charges: Even with the unpaired electron, calculating formal charges helps to assess the "best" arrangement of the electrons.
• Conclusion: The proposed Lewis structure is reasonable, given the constraint of having an odd number of valence electrons. NO2 is a radical species, meaning it has an unpaired electron. This unpaired electron makes it highly reactive.

Example 4: Sulfate Ion (SO4^2-)

• Proposed Lewis Structure (Version 1): Sulfur single-bonded to four Oxygen atoms, each Oxygen with three lone pairs, and a 2- charge on the overall ion.
• Proposed Lewis Structure (Version 2): Sulfur double-bonded to two Oxygen atoms and single-bonded to two Oxygen atoms, with appropriate lone pairs to complete octets, and a 2- charge on the overall ion.
• Step 1: Valence Electrons: S (6) + 4 * O (6) + 2 (charge) = 32 valence electrons.
• Step 2: Connectivity: Sulfur as the central atom bonded to four oxygen atoms is reasonable.
• Step 3: Octet Rule: In Version 1, all atoms satisfy the octet rule. In Version 2, all atoms also satisfy the octet rule.
• Step 4: Formal Charges:
  • Version 1: Sulfur has a +2 formal charge, and each oxygen has a -1/2 formal charge.
  • Version 2: Sulfur has a 0 formal charge, two oxygen atoms have a 0 formal charge, and two oxygen atoms have a -1 formal charge.
• Step 5: Expanded Octet & Resonance: Sulfur, being in the third period, can expand its octet. Version 2 is a better Lewis structure because it minimizes formal charges. Furthermore, multiple resonance structures can be drawn for Version 2, where the double bonds are moved around to different oxygen atoms. In reality, sulfate is a resonance hybrid of all these structures, with the sulfur-oxygen bonds having a bond order between 1 and 2.
• Conclusion: Version 2 is a more reasonable representation of the sulfate ion than Version 1 because it minimizes formal charges and takes into account the possibility of sulfur expanding its octet. The concept of resonance is also crucial for accurately describing the bonding in sulfate.

Common Mistakes to Avoid

• Incorrectly Counting Valence Electrons: This is the most common error. Double-check the periodic table and remember to account for ionic charges.
• Forgetting Lone Pairs: Lone pairs are essential for satisfying the octet rule.
• Ignoring Formal Charges: While minimizing formal charges isn't always possible, ignoring them can lead to incorrect Lewis structures.
• Not Considering Resonance: In molecules where resonance is possible, drawing only one Lewis structure gives an incomplete picture.
• Assuming the Octet Rule Always Applies: Be aware of the exceptions to the octet rule, especially for elements in the third period and beyond.
• Thinking a Lewis Structure is the "Truth": Lewis structures are simplified models. They are useful for understanding bonding, but they don't perfectly represent the actual distribution of electrons in a molecule.

The Importance of Assessing Lewis Structures

Being able to assess the reasonableness of a Lewis structure is a fundamental skill in chemistry. It allows you to:

• Predict Molecular Geometry: The arrangement of atoms and lone pairs around a central atom (as depicted in the Lewis structure) determines the molecule's shape (using VSEPR theory).
• Understand Polarity: Bond polarity and overall molecular polarity depend on the electronegativity differences between atoms and the molecule's geometry.
• Predict Reactivity: The presence of lone pairs, formal charges, and bond strengths can influence a molecule's reactivity.
• Explain Physical Properties: Intermolecular forces (like dipole-dipole interactions and hydrogen bonding) are related to molecular polarity, which is influenced by the Lewis structure.
• Communicate Chemical Information: Lewis structures are a common and efficient way to communicate information about molecular structure.

Advanced Considerations

Beyond the basic steps outlined above, more advanced concepts can further refine your assessment of Lewis structures:

• Hypervalency: Elements in the third period and beyond can sometimes form compounds where they have more than eight electrons in their valence shell (hypervalency). This is due to the availability of d orbitals for bonding.
• Molecular Orbital Theory: While Lewis structures are based on localized bonding, molecular orbital (MO) theory provides a more accurate description of electron distribution by considering the entire molecule.
• Computational Chemistry: Software programs can calculate the electronic structure of molecules, providing information about bond lengths, bond angles, and electron densities that can be compared to predictions based on Lewis structures.
• Experimental Data: Techniques like X-ray crystallography and spectroscopy can provide experimental data that confirms or refines the proposed structure of a molecule.

Conclusion

Deciding whether a proposed Lewis structure is reasonable involves a systematic assessment of valence electrons, connectivity, the octet rule, formal charges, and the potential for resonance. By mastering these principles and avoiding common mistakes, you can confidently evaluate Lewis structures and gain a deeper understanding of chemical bonding and molecular properties. This skill is not just about drawing diagrams; it's about developing a fundamental understanding of how molecules are put together and how their structure influences their behavior. So, practice these steps with different molecules and ions, and you'll become proficient in evaluating the reasonableness of Lewis structures. Remember to always check your work and be mindful of the exceptions to the rules. Happy drawing!