Ionic and Molecular Lewis Dot Structures — Study Notes

Context and purpose
- Learning intentions: how to draw ionic Lewis dot structures and molecular Lewis dot structures.
- Lewis dot structures help visualize valence electrons, bonding, and the formal charges in compounds.

I. Ionic Lewis Structures

Core idea
- Represent ionic compounds as a transfer of electrons from a cation (metal) to an anion (non-metal), forming cations and anions that are held together by electrostatic attraction.
- Use bracket notation with the ionic charge around each ion. Include coefficients in front to show the number of each ion; a coefficient of 1 is not written.
- Ordering convention: metal followed by non-metal.
Step-by-step procedure (as given)
1. Draw the Lewis dot structure of each atom (valence electrons for each element).
2. Transfer the electron(s) from the cation to the anion to achieve octets where appropriate (or full shells for ions).
3. Complete the Ionic Lewis dot structure: place brackets around each ion and include its charge; add coefficients to indicate the number of ions if needed.
- Example convention: [Metal environment]^{n+} and [Non-metal environment]^{m-}
Key concepts to remember
- Ionic compounds are composed of ions in a lattice; discrete molecules are not formed.
- The overall charge of the formula unit must be zero:
  $ext{Total charge} = ext{(sum of cation charges)} + ext{(sum of anion charges)} = 0.$
- The cation is typically a metal and the anion a non-metal (the common convention).
Example (conceptual): Ca3N2
- Calcium forms Ca^{2+} ions; Nitrogen forms N^{3-} ions.
- Ratio is 3 Ca^{2+} to 2 N^{3-} to balance charges: 3( +2 ) + 2( −3 ) = 0.
- In a Lewis-dot ionic representation, you would show Ca^{2+} and N^{3-} ions with appropriate brackets and charges, indicating the lattice composition and the overall neutrality.

II. Molecular Lewis Structures

Core idea
- Central notion: place the least electronegative atom (often the one that can form the most bonds) in the center. Hydrogen and halogens cannot be central atoms.
- Bond all other atoms to the center atom with single bonds first (2 electrons per bond).
- Determine the total number of valence electrons for the molecule.
- Distribute electrons to outer atoms to satisfy octets (or full shells) first; any remaining electrons go to the center atom.
- If the center atom does not have a full octet after distributing electrons, shift electrons from outer atoms to form double or triple bonds as needed.
Step-by-step procedure
1. Put the least electronegative atom in the center (except H and halogens which cannot be central).
2. Connect all outer atoms to the center with single bonds (each bond uses 2 electrons).
3. Count total valence electrons in the molecule: sum of valence electrons for all atoms.
4. Place electrons to satisfy outer atoms first (give them a full octet), then place any remaining electrons on the center atom.
5. If the center atom lacks a full octet, convert one or more outer single bonds into double or triple bonds by shifting electrons accordingly.
Bonding concepts
- Single bond: 2 electrons total (1 pair) shared between two atoms.
- Double bond: 4 electrons total (two pairs) shared between two atoms.
- Triple bond: 6 electrons total (three pairs) shared between two atoms.
- For each bond, electrons are counted toward both participating atoms’ octets.
Examples (typical outcomes)
- HCN (hydrogen cyanide): central atom is C, bonded to H (single bond) and N (triple bond), giving a linear molecule with the carbon using four bonds total (no lone pairs on C) and N bearing one lone pair.
- CHCl3 (chloroform): central C bonded to one H and three Cl atoms; Cl atoms each have lone pairs to complete octets; C has four single bonds and no lone pairs.
- H2O (water): O is central, bonded to two H atoms via single bonds; O has two lone pairs after bonding.
- CO2 (carbon dioxide): O=C=O with carbon in the center, two double bonds to oxygens; O and C each achieve octets.

III. Bonding: Types and Electron Counts

Summary of bond types with electron counts
- Single bond: 2 electrons total; 1 electron contribution from each atom.
- Double bond: 4 electrons total; 2 electrons from each atom.
- Triple bond: 6 electrons total; 3 electrons from each atom.
- Note: The total electrons involved in a bond are shared between the two atoms and count toward each atom’s octet.

IV. Worked Examples from the transcript

Example 1: Ca3N2 (ionic compound)
- Valence/ion viewpoint: Ca forms Ca^{2+}, N forms N^{3-}.
- Ionic ratio ensures charge balance: 3(+2) from Ca and 2(−3) from N sum to zero.
- Final depiction (conceptual): a lattice of Ca^{2+} and N^{3-} ions in the proper stoichiometric ratio; brackets around ions with the appropriate charges, e.g., [Ca^{2+}] and [N^{3-}].
Example 2: HCN (molecular compound)
- Valence electron count: H (1) + C (4) + N (5) = 10 electrons.
- Structure: H–C≡N with a triple bond between C and N; N carries one lone pair; carbon has no lone pairs.
- Final Lewis structure (conceptual): H–C≡N with N bearing one lone pair.
Example 3: CHCl3 and H2O
- CHCl3 (chloroform): central C bonded to H and three Cl atoms; each Cl has three lone pairs; C has no lone pairs.
- H2O (water): O central, bonded to two H; O has two lone pairs.
Example 4: CaO and MgBr2 (ionic compounds)
- CaO: [Ca^{2+}] [O^{2-}] — ionic lattice; overall neutral.
- MgBr2: [Mg^{2+}] [Br^-]2 — one Mg^{2+} cation with two Br^- anions.
Example 5: CHFO and CO2
- CHFO (formyl fluoride, H–C(=O)–F): central C bonded to H, O (double bond), and F (single bond); O has two lone pairs; F has three lone pairs; total valence electrons = 18.
- CO2 (carbon dioxide): O=C=O with two double bonds; O and C octets satisfied; total valence electrons = 16.
- Final structures reflect octet completion and formal charges where applicable (no formal charges in these neutral, stable species).
Example 6: NH4+ (ammonium)
- Central N bonded to four H atoms with single bonds; no lone pairs on N; overall charge +1 on the ion.
- Lewis structure (with brackets): [NH4]+ showing N bonded to four H atoms; no lone pairs on N; +1 formal charge on the species.

V. Success Criteria and Practical Applications

Skills to demonstrate
- I can draw the Lewis dot structure of an ionic substance (e.g., Ca3N2, Al2O3).
- I can draw the Lewis dot structure of a molecular substance (e.g., SF2, HCN, CO2, H2O, CHFO).
Common targets mentioned in the transcript
- Ionic: Al2O3 (goal to represent the lattice with Al^{3+} and O^{2-} ions).
- Molecular: SF2 (draw the correct Lewis structure with appropriate lone pairs and bonding).

VI. Quick Reference Rules (summary)

For ionic Lewis structures:
- Identify metal (cation) and non-metal (anion).
- Transfer electrons to form ions with full shells; apply brackets and charges.
- Balance the overall charge with appropriate coefficients.
For molecular Lewis structures:
- Place the least electronegative atom in the center (except H and halogens cannot be central).
- Use single bonds to connect outer atoms to the center first; count valence electrons.
- Fulfill octets for outer atoms first, then the center; if needed, form multiple bonds by shifting electrons.
Valence electron bookkeeping
- Total valence electrons = sum of valence electrons from all atoms in the formula.
- Ensure that your final structure honors the octet rule for second-row elements, and account for charge where applicable.
Notation and formatting
- Use bracket notation with charges for ions: [X^{m+}] or [X^{n-}].
- In a lattice, show the repeating unit with appropriate stoichiometric coefficients.

VII. Reference to the Transcript Sections

Page 2: Learning intentions – focus on ionic and molecular Lewis structures.
Page 3: Ionic Lewis Structures – steps and bracket/charge convention.
Page 4: Molecular Lewis Structures – center atom rules, stepwise electron distribution, octet completion.
Page 5: Bond types and electron counts for single/double/triple bonds.
Pages 6–11: Worked examples (Ca3N2; HCN; CHCl3 and H2O; CaO and MgBr2; CHFO and CO2; NH4+).
Page 12: Success criteria – what students should be able to do (ionic and molecular structures for provided substances).

Notes: Some symbols in Page 1 of the transcript were garbled, showing various fragments of Lewis structures. The essential concepts and standard practice for ionic and molecular Lewis structures are covered in the sections above, with the clear, commonly accepted rules and example outcomes.