Lewis Structures

Related Examples and Practice Problems

Additional Worked Out Examples/ Practice

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  • Separation techniques: Selected and explaining limitation of appropriate separation

  • Relating Properties to Composition: Predicting classification based on descriptive properties

Topic Summary & Highlights
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Core Concept

Lewis structures are a visual representation on how the valence electrons are arranged in a compound in bonding.

Practice Tips

  • Miscounting valence electrons: Always double-check your total electron count.

  • Forgetting formal charges: Use formal charges to find the most stable structure, especially for ions.

  • Ignoring octet rule exceptions: Remember that hydrogen, boron, and elements in period 3 or higher may not follow the octet rule.

  • Not drawing all resonance structures: If there are multiple ways to arrange double bonds or lone pairs, draw all resonance structures.

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Step-by-Step Guide

Step 1: Count Total Valence Electrons

  • Determine the number of valence electrons for each atom in the molecule. This information is based on the group number for each element in the periodic table.

  • Add up the valence electrons for all atoms in the molecule. For polyatomic ions:

    • Add electrons if the ion is negatively charged.

    • Subtract electrons if the ion is positively charged.

Step 2: Arrange the Atoms

  • Place the least electronegative atom (usually the atom that can form the most bonds) in the center. Hydrogen is always a terminal (outer) atom, as it can only form one bond.

  • Arrange other atoms around the central atom, symmetrically if possible.

Step 3: Form Single Bonds

  • Connect each outer atom to the central atom using a single line, which represents a single bond (2 electrons).

  • Subtract the bonding electrons (2 electrons per bond) from the total valence electron count.

Step 4: Distribute Remaining Electrons as Lone Pairs

  • Place lone pairs of electrons around each outer atom to fulfill the octet rule (8 electrons), starting with the most electronegative atoms.

  • Any remaining electrons should go to the central atom.

Step 5: Check the Octet Rule

  • Ensure each atom (except hydrogen) has a complete octet (8 electrons around it). Hydrogen should have only 2 electrons.

  • If the central atom has fewer than 8 electrons, consider forming double or triple bonds by sharing lone pairs from surrounding atoms.

Step 6: Verify and Adjust

  • Count the total electrons in the structure to ensure it matches the original total valence electrons.

  • For polyatomic ions, add brackets around the structure and indicate the charge outside the brackets (e.g., $[ \text{NO}_3^- ]$).

Example: Drawing the Lewis Structure for CO₂

  1. Count Valence Electrons: Carbon has 4, and each oxygen has 6. Total = 4+6+6=16 electrons.

  2. Arrange the Atoms: Place carbon (the least electronegative) in the center with oxygens on each side.

  3. Form Single Bonds: Connect each oxygen to carbon with a single bond. This uses 2×2=4 electrons, leaving 16 − 4 = 12 electrons.

  4. Distribute Remaining Electrons: Place 6 lone pairs (12 electrons) around the oxygens to fulfill their octets.

  5. Check the Octet Rule for Carbon: Carbon has only 4 electrons, so form double bonds by converting lone pairs from each oxygen into bonding pairs.

  6. Verify: The structure should have two double bonds, with each atom following the octet rule.

The final structure is: O=C=O

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