Convert The Structure Below To A Lewis Structure

Converting Skeletal Structures to Lewis Structures: A Comprehensive Guide

Skeletal structures, also known as condensed structures or line-angle formulas, provide a simplified representation of organic molecules. They show the carbon skeleton and the arrangement of atoms, omitting explicit depiction of carbon atoms and hydrogen atoms directly bonded to carbons. However, to understand the bonding, reactivity, and properties of a molecule, we need the complete Lewis structure, which shows all atoms and valence electrons. This guide will walk you through the step-by-step process of converting skeletal structures into complete Lewis structures.

Understanding Skeletal Structures and Lewis Structures

Before diving into the conversion process, let's review the basics of each structure type:

Skeletal Structures:

• Carbon atoms are implied: Carbon atoms are at the corners and intersections of lines.
• Hydrogen atoms are implied: Hydrogen atoms attached to carbon atoms are not explicitly shown. You need to infer their presence based on the number of bonds each carbon atom can form (four).
• Other atoms are explicitly shown: All atoms other than carbon and hydrogen are explicitly shown in the structure.
• Lines represent bonds: A single line represents a single covalent bond, a double line represents a double bond, and a triple line represents a triple bond.

Lewis Structures:

• All atoms are explicitly shown: Every atom in the molecule is clearly represented.
• All valence electrons are shown: Valence electrons are depicted as dots (·) or lines representing shared electron pairs (bonds). Lone pairs (unshared electron pairs) are crucial for understanding molecular geometry and reactivity.
• Bonds are explicitly shown: Each bond is represented by a line (or a pair of dots).
• Formal charges are indicated: Formal charges are assigned to atoms to reflect the distribution of electrons, ensuring the overall charge of the molecule is neutral or reflects the actual charge.

Step-by-Step Conversion Process

Converting a skeletal structure to a Lewis structure involves a series of logical steps:

1. Count the Total Number of Valence Electrons:

This is the crucial first step. Determine the number of valence electrons contributed by each atom in the molecule. Remember:

• Carbon (C) has 4 valence electrons.
• Hydrogen (H) has 1 valence electron.
• Oxygen (O) has 6 valence electrons.
• Nitrogen (N) has 5 valence electrons.
• Halogens (F, Cl, Br, I) have 7 valence electrons.

Add the valence electrons from all atoms to get the total. If the molecule has a charge, adjust the total accordingly: add one electron for each negative charge and subtract one electron for each positive charge.

2. Identify the Central Atom(s):

The central atom(s) is typically the least electronegative atom (except hydrogen, which is always terminal). Often, this is the atom with the highest bonding capacity.

3. Connect Atoms with Single Bonds:

Connect the atoms using single bonds (one line per bond) to create a basic framework. Start with the central atom(s) and extend bonds to the surrounding atoms as indicated by the skeletal structure.

4. Distribute Remaining Valence Electrons as Lone Pairs:

After connecting all atoms with single bonds, distribute the remaining valence electrons as lone pairs around the atoms, starting with the terminal atoms (atoms bonded to only one other atom) to satisfy the octet rule (or duet rule for hydrogen).

5. Check for Octet Rule Satisfaction:

The octet rule states that atoms (except hydrogen and some exceptions like boron) tend to have eight electrons in their valence shell. Ensure all atoms (other than hydrogen) have eight electrons around them. If an atom lacks an octet, consider forming multiple bonds (double or triple bonds).

6. Form Multiple Bonds (If Necessary):

If an atom does not have a complete octet, form multiple bonds by converting lone pairs into shared pairs (bonds) between adjacent atoms. This is particularly common for carbon, nitrogen, and oxygen.

7. Assign Formal Charges (If Necessary):

Calculate the formal charge on each atom using the formula: Formal Charge = (Valence Electrons) - (Non-bonding Electrons) - (1/2 Bonding Electrons). If the molecule has a formal charge, this should be included in the Lewis structure. Remember that the sum of formal charges should equal the overall charge of the molecule.

8. Minimize Formal Charges:

While some formal charges are unavoidable, strive to minimize their magnitude and distribute them across the molecule logically.

9. Resonance Structures:

If multiple Lewis structures can be drawn that differ only in the placement of electrons (not atoms), these are resonance structures. Resonance structures contribute to the overall picture of the molecule's bonding, and are typically depicted with double-headed arrows.

Example: Converting the skeletal structure of ethanol (CH3CH2OH) to a Lewis structure.

Step 1: Count valence electrons: 2(4) + 6(1) + 6 = 20 electrons

Step 2: Central atoms are the carbons.

Step 3: Connect atoms with single bonds: Create a chain of two carbon atoms, with three hydrogens bonded to the first carbon, two hydrogens bonded to the second carbon and one oxygen bonded to the second carbon.

Step 4: Distribute remaining electrons as lone pairs: This leaves 4 electrons. Place them as two lone pairs on the oxygen atom.

Step 5: Check for octet rule satisfaction: All atoms satisfy the octet rule (or duet for hydrogen).

Step 6: No multiple bonds needed.

Step 7: Assign formal charges: All formal charges are zero.

Step 8: Minimal formal charges already achieved.

Step 9: No resonance structures.

The resulting Lewis structure shows all atoms, bonds, and lone pairs, providing a complete picture of ethanol’s electronic structure.

Advanced Cases and Considerations

• Expanded Octet: Elements in the third period and beyond (like phosphorus and sulfur) can sometimes accommodate more than eight electrons in their valence shell.

• Odd Electron Species: Some molecules have an odd number of valence electrons, resulting in a radical species with an unpaired electron.

• Electron-Deficient Molecules: Some molecules, such as boron trifluoride (BF3), are electron-deficient, meaning they do not satisfy the octet rule.

• Formal Charge and Stability: Lewis structures with lower formal charges are generally more stable.

By following these steps, you can effectively convert any skeletal structure into a Lewis structure, paving the way for a more complete understanding of the molecule's properties and reactivity. Remember, practice is key to mastering this crucial skill in chemistry. Working through numerous examples will help you solidify your understanding and develop a confident approach to drawing Lewis structures from skeletal formulas. The more complex the skeletal structure, the more challenging the conversion may be, but with careful attention to detail and systematic application of the rules, you will be able to accurately represent the molecule's electronic arrangement.