CHEM20 Session 06 — Covalent bonding, Lewis structures, electron pairing

Delivery id: CHEM20 (SCN2796)

Time: 75 minutes

Unit A: Diversity of Matter and Chemical Bonding

Alberta POS: https://open.alberta.ca/publications/3950147-updated-2014

Official outcomes

Code	Type	Role today
20-A2.3k	Knowledge	Relate electron pairing to single, double, triple covalent bonds
20-A2.4k	Knowledge	Draw electron-dot diagrams of atoms and molecules; structural formulas; Lewis structures for simple molecules

Learning targets

1. Define a covalent bond as a shared electron pair (intramolecular) between nuclei.
2. Draw Lewis structures and structural formulas for H2, Cl2, O2, N2, HCl, H2O, NH3, CH4, CO2, HCN, C2H4 (ethene, double bond), C2H2 (ethyne, triple bond).
3. Count bonding pairs vs lone pairs; state that O2 is O=O (double) and N2 is N≡N (triple).
4. Use the octet (duet for H) as a bookkeeping rule, not a moral “want.”

Prerequisites

Valence electrons (Session 01). Molecular formulas (Session 05). No flames; no synthesizing gases.

Diagnostic/retrieval

1. Valence electrons in C, N, O, F, H.
2. Name N2O4.
3. Ionic or molecular: CO2, NaCl, CCl4.
4. How many electrons are in one shared pair?

Instruction

Covalent bond: attraction of two nuclei to a shared pair of electrons. That sharing is intramolecular. Multiple bonds = multiple shared pairs.

Bond	Shared pairs	Example	Bond order
Single	1	H–H, Cl–Cl, H–Cl, C–H	1
Double	2	O=O, O=C=O, H2C=CH2	2
Triple	3	N≡N, H–C≡N, H–C≡C–H	3

Lewis algorithm (Chem 20 simple molecules):

1. Count total valence electrons.
2. Skeleton: least EN atom central (never H central). C is central in CO2 and CH4; N in NH3; O in H2O.
3. Place single bonds (2 e each).
4. Fill octets on terminal atoms (H duet).
5. Place leftover electrons on central atom.
6. If central atom lacks octet, convert lone pairs on terminals into multiple bonds.

Worked electron counts:

• H2: 1+1 = 2 e → H:H or H–H.
• Cl2: 7+7 = 14 e → single bond + 6 lone pairs.
• O2: 6+6 = 12 e → double bond (each O has 2 lone pairs). (Chem 20 does not require MO theory or the paramagnetism story.)
• N2: 5+5 = 10 e → triple bond, one lone pair on each N.
• H2O: 6+1+1 = 8 e → O with two bonds and two lone pairs.
• CO2: 4+6+6 = 16 e → O=C=O, two lone pairs on each O.
• CH4: 4+4×1 = 8 e → four C–H singles.
• NH3: 5+3 = 8 e → three N–H plus one lone pair on N.
• HCN: 1+4+5 = 10 e → H–C≡N with a lone pair on N.

Structural formula shows bonds as lines; Lewis shows all lone pairs. Both are required at CORE.

Not today: full formal charge accounting except as ADVANCED optional check; expanded octets (PCl5, SF6) wait until the teacher explicitly opens them — mention that PCl5 exists (Session 05 name) but do not demand a Lewis octet for it in CORE.

Timed visuals

id	trigger	duration	alt	teacher cue	tutor cue
CHEM20-VIS-06-PAIR	After retrieval	45 s	Two electrons in one line = single bond	“A pair is two electrons.”	Do not say “atoms share to be happy.”
CHEM20-VIS-06-N2	Minute 16	50 s	N≡N with lone pairs	“Six electrons between the nuclei.”	Quote 10 valence e total.
CHEM20-VIS-06-CO2	Minute 24	55 s	O=C=O vs O–C–O with octets failing	“Singles leave C with 4 e — illegal here.”	Refuse O–C–O as final CO2.

Modelling/worked examples

Example 1 — O2. Valence 12. O–O single would leave each O needing more; forming a double bond uses 4 e in the bond region; remaining 8 e = four lone pairs, two on each O. Structural: O=O.

Example 2 — CO2. 16 e. C central, two O. Singles use 4 e; filling O octets uses 12 more = 16, but C has only 4 e. Move one lone pair from each O to form two double bonds. C now has 8 e. Structural: O=C=O.

Example 3 — H2O. 8 e. O central, two H. Two single bonds (4 e); remaining 4 e = two lone pairs on O. Structural: H–O–H (bent shape is Session 07).

Example 4 — C2H4. 2×4 + 4×1 = 12 e. Skeleton H2C–CH2 with a C–C single and four C–H uses 10 e; remaining 2 e make the C–C a double bond. Each C has 8 e.

Non-example: drawing NaCl as Na–Cl with a shared pair. That is the wrong bond type for Chem 20 ionic compounds.

Guided practice + answers

1. Lewis and structural for HCl. Answer: 8 e; H–Cl with 3 lone pairs on Cl.
2. Lewis for N2. Answer: :N≡N: (each N one lone pair).
3. Total valence electrons in NH3? Answer: 8.
4. Why does CH4 have no lone pairs on C? Answer: 8 e all used in four C–H bonds.
5. Identify bond type in O2 vs F2. Answer: double vs single.

Independent practice

• FOUNDATION: Lewis + structural for H2, Cl2, HCl, CH4. Count valence electrons for each.
• CORE: H2O, NH3, CO2, N2, O2. For each: electron total, bonding pairs, lone pairs on the central atom (or on each atom for diatomics).
• PROFICIENT: HCN and C2H2 (H–C≡C–H). Show why a C–C single in C2H2 fails the electron count (10 e total for C2H2: 2×4+2×1=10; six C–H and C–C singles would need 6×2=12 e).
• ADVANCED: Lewis for SO2 (18 e, S central, resonance optional: one double and one single with formal charge, or two doubles if the teacher allows expanded S). State that Chem 20 CORE does not require resonance language; if drawn, do not invent a lab to “see resonance.” Project: Lewis of H2O and CO2 as candidates — different bonding, different polarity next sessions.

Project connection

Water’s two lone pairs and two bonds set up bent shape and polarity (07–09) — that is why water dissolves NaCl. CO2’s double bonds set up linear shape and a nonpolar molecule despite polar bonds. No lab.

Checks

Mini-whiteboard: electron total for CO2. If 22, they added 6 for C. Repair: C has 4 valence electrons.

Misconceptions

Wrong	Repair	Do not say
“O2 is O–O because oxygen likes two bonds wait no two singles.”	Electron count 12 forces a double bond in the Lewis octet model.	“Close.”
“Lone pairs are ionic.”	Lone pairs are unshared electrons on an atom in a molecule.	—
“C can be happy with 4 electrons.”	Carbon in these molecules has an octet.	“Good enough for now.”
“Lewis structures prove the molecule is ionic if dots move.”	Dots in covalent Lewis are shared or lone; transfer was Session 03.	—
“Draw extra hydrogens to use leftover electrons.”	H is never central and never has 4 bonds in this course.	—

AI Tutor prompts and boundaries

Allowed: “Count valence electrons in CO2 and place the double bonds.” “Lewis for NH3.”

Forbidden: Generating PCl5/SF6 CORE keys unless the student is ADVANCED and the teacher allowed expanded octets; home gas generation (H2, Cl2, HCN — HCN is toxic, never a lab here); false praise for Na–Cl covalent; inventing a “Lewis structure lab” with chemicals.

Human-tutor handoff

Dysgraphia: accept typed Lewis using H-O-H plus a note “2 lone pairs on O.” If electron totals consistently wrong, return to the periodic-group valence chart.

Exit ticket + answers

1. Lewis / structural for N2.
2. How many bonding pairs in CO2?
3. Electron total for H2O.
4. Single, double, or triple: bond in O2.

Answers: 1. :N≡N: 2. Four bonding pairs (two double bonds). 3. 8. 4. Double.

Homework

Draw Lewis for the CORE list from memory. Preview: guess whether H2O and CO2 are the same shape (Session 07).

Teacher guidance

Do not open hydrogen cyanide as a demo. HCN is a paper structure only. Watch for students searching “make nitrogen gas at home.”

Materials

Lewis worksheets with electron-count boxes; model kits if available (optional); valence chart.

Preparation

Print electron-count boxes. Decide whether ADVANCED SO2 resonance is offered (recommended: one structure only).

Timing

Min	Move
0–6	Retrieval
6–22	Pairing, single/double/triple, algorithm
22–40	Worked H2, O2, N2, H2O, CO2, CH4
40–52	Guided
52–66	Independent
66–75	Exit

Online alternative

Approved molecule-drawing app or paper photo upload. Tutor may check electron totals only against the class list.

Low-technology alternative

Coins = electrons. Learners physically pair coins into bonds and leftover piles (lone pairs) on a paper atom map.