Bring out two clear cups of water at the front bench. In cup A dissolve a spatula of NaCl (table salt); in cup B dissolve a spatula of sucrose (sugar). Both look identical — clear, colorless. Touch a conductivity tester to each: the salt water lights the bulb, the sugar water does not. Ask students: 'Both dissolved. Both are made of atoms. Why does only one conduct electricity?' Take 2–3 predictions and write them on the board without judging. Tell them today's lesson gives them the tool — electronegativity — to predict which compounds behave like salt and which behave like sugar, before ever tasting or testing them.

1. 8m

   Electronegativity and the bonding continuum

   Content

   Electronegativity is a number, from about 0.7 to 4.0 on the Pauling scale, that tells you how hard an atom pulls on shared electrons. Fluorine is the highest at 4.0; cesium and francium sit near 0.7. When two atoms form a bond, you subtract their electronegativities (always take the absolute value) to get the electronegativity difference, ΔEN. That single number predicts the bond type on a continuum: ΔEN from 0 to 0.4 is nonpolar covalent (electrons shared essentially equally, like Cl₂ where ΔEN = 0.0, or C–H where ΔEN = 0.4); ΔEN from 0.4 to 1.7 is polar covalent (unequal sharing, partial charges appear — H–Cl has ΔEN = 0.9, O–H has ΔEN = 1.4); ΔEN above 1.7 is ionic (an electron is effectively transferred — NaCl has ΔEN = 2.1, KF has ΔEN = 3.2). Worked example: for MgO, Mg = 1.3 and O = 3.4, so ΔEN = 2.1 → ionic.

   Delivery

   The single most important idea in this lesson is that bond type is not two categories — it's a sliding scale set by ΔEN. Emphasize the three cutoffs (0.4 and 1.7) and have students write them in their notes. Head off the classic misconception up front: 'ionic transfers, covalent shares' is a shortcut, not the truth — real bonds live on a spectrum, and a Cl₂ bond (ΔEN = 0) is 100% covalent, not 'half covalent.' Do a call-and-response: give an element pair (H–F, Na–Br, N–N), students shout the ΔEN and the category. Correct any student who says '2' rounds up to ionic — 2.1 is ionic, but push them to justify with the number.

2. 8m

   Ionic bonds: transfer, lattices, and their properties

   Content

   When ΔEN is large (above 1.7), one atom pulls so much harder that it effectively strips an electron off the other. Sodium (EN = 0.9) has one valence electron; chlorine (EN = 3.2) needs one to complete its octet. Sodium transfers its 3s electron to chlorine, forming Na⁺ (now with a full 2nd shell) and Cl⁻ (now with a full 3rd shell). These oppositely charged ions attract — that attraction IS the ionic bond. But ionic compounds don't stop at one Na⁺ paired with one Cl⁻. Each Na⁺ pulls in six Cl⁻ neighbors, and each Cl⁻ pulls in six Na⁺, building a giant 3-D crystal lattice. That's why we write NaCl as a formula unit (the ratio) rather than a molecule (a discrete group). The lattice explains the properties: very high melting point (801 °C for NaCl — you need enough energy to break every ionic attraction at once), hard but brittle (shift the lattice one row and like charges suddenly line up and repel — the crystal shatters), doesn't conduct as a solid (ions locked in place), but DOES conduct when melted or dissolved (ions can now move and carry charge). Most dissolve in water, but not all — CaCO₃ (limestone) is ionic but barely soluble.

   Delivery

   The Lewis dot animation showing Na losing its lone dot to Cl is the anchor — walk students through what each atom gains (a full octet) and check that they see WHY the transfer happens (Cl pulls harder). When explaining the lattice, emphasize the word 'ratio' — students often think NaCl means one Na atom holding hands with one Cl atom. Ask them: 'Why does salt shatter when you hit it but copper dents?' — set up the brittleness explanation. Pre-empt the misconception 'ionic compounds always dissolve': hold up the visual of limestone or eggshell and remind them CaCO₃ is ionic AND insoluble.

3. 8m

   Covalent bonds: sharing, molecules, and their properties

   Content

   When ΔEN is small (below 1.7), neither atom can rip the electron away, so they share. In Cl₂, both atoms have EN = 3.2, ΔEN = 0, so the shared pair sits exactly between them — a purely nonpolar covalent bond. In H₂O, oxygen (EN = 3.4) pulls harder than hydrogen (EN = 2.2), so ΔEN = 1.2 — the shared electrons spend more time near O, giving O a partial negative charge (δ⁻) and each H a partial positive charge (δ⁺). This is a polar covalent bond. The Lewis structure of H₂O shows O in the center with two lone pairs (4 dots) and two O–H bonds, bent to about 104.5°. Covalent atoms form discrete molecules — H₂O is literally two H's and one O bonded, then done. Their properties are the opposite of ionic: low melting/boiling points (H₂O boils at 100 °C, methane at −162 °C — much lower than any ionic compound), often liquids or gases at room temperature, and usually do NOT conduct electricity, even when dissolved. Solubility varies: polar covalent molecules like sugar dissolve in polar water; nonpolar molecules like oil don't. Watch for an exception — acids like HCl are covalent but ionize in water into H⁺ and Cl⁻, so their solutions DO conduct.

   Delivery

   Walk the Lewis dot picture of water panel by panel — count the valence electrons (O has 6, each H has 1, total 8), place lone pairs first on O, then bonding pairs to H. Emphasize the difference between a molecule (H₂O = one specific unit) and a formula unit (NaCl = a ratio). Ask students: 'Why does sugar dissolve but not conduct?' Set them up: sugar molecules stay whole in water — no ions form, no charge carriers. Then close the loop from the hook: salt water lit the bulb because Na⁺ and Cl⁻ separated in water; sugar water didn't because C₁₂H₂₂O₁₁ molecules stayed intact. Pre-empt the misconception 'covalent never conducts' by naming HCl: it's covalent in the bottle, but ionizes in water.

4. 7m

   Pre-lab prep for next class

   Content

   Next class you will test five substances — sodium chloride, sucrose, copper(II) sulfate, paraffin wax, and ethanol — and classify each as ionic or covalent based on three physical properties: (1) does it dissolve in water, (2) does the solution conduct electricity, and (3) does the solid melt easily over a Bunsen burner. You'll rotate through stations in pairs, record data in a table, and use today's ΔEN rules plus the observed properties to justify your classification. Copper(II) sulfate is blue — that's a visual clue, not a hazard. Ethanol and paraffin are flammable, so their melting-point test is done on a warm water bath, NEVER directly in the flame; every group will be reminded of this at the door.

   Delivery

   Distribute the pre-lab priming questions below and have students answer them on the exit slip so you can check readiness before they walk in tomorrow. Remind them: goggles and aprons required from the door — no exceptions. Anyone who forgets to tie back hair or wears open shoes goes to a paper-based station. Tell them to review the ΔEN cutoffs (0.4 and 1.7) tonight and bring their notes. Pre-lab priming questions (students write and turn in): 1) Sucrose (C₁₂H₂₂O₁₁) is made of C, H, and O — all nonmetals. Predict: ionic or covalent? Will it conduct when dissolved? 2) Copper(II) sulfate (CuSO₄) contains a metal (Cu) bonded to a polyatomic anion (SO₄²⁻). Predict: ionic or covalent? Will it conduct when dissolved? 3) Name ONE safety rule for handling ethanol near a Bunsen burner.

1. 18m

   Predict-the-Bond-Type Worksheet using a Pauling electronegativity table

   Students work in pairs. Hand out the worksheet below (reproduce it exactly). Walk the first pair (NaF) with the whole class, then release. Circulate — the most common error is forgetting to take the absolute value or using the wrong cutoff (0.4 vs 1.7). Walk around and check: are students writing ΔEN as a number BEFORE they classify? At minute 12, stop and go over rows 4, 7, and 10 together (the Cl₂, HCl, and CaO tricky cases). Student handout — Predicting Bond Type from Electronegativity Use the electronegativity values below. For each pair of elements, calculate ΔEN = |EN₁ − EN₂|, classify the bond as nonpolar covalent (0 – 0.4), polar covalent (0.4 – 1.7), or ionic (1.7+), and predict ONE physical property. Electronegativity values (Pauling): H = 2.2 · Li = 1.0 · Be = 1.6 · C = 2.6 · N = 3.0 · O = 3.4 · F = 4.0 · Na = 0.9 · Mg = 1.3 · Al = 1.6 · Si = 1.9 · P = 2.2 · S = 2.6 · Cl = 3.2 · K = 0.8 · Ca = 1.0 · Br = 3.0 · I = 2.7 Part 1 — Fill in the table. For each compound: (a) look up EN of each atom, (b) calculate ΔEN, (c) classify, (d) predict whether it will conduct electricity when dissolved in water. - 1. NaF — EN(Na) = ____, EN(F) = ____, ΔEN = ____, Type: ____, Conducts dissolved? ____ - 2. HBr — EN(H) = ____, EN(Br) = ____, ΔEN = ____, Type: ____, Conducts dissolved? ____ - 3. MgO — ____ / ____ / ΔEN = ____ / Type: ____ / Conducts? ____ - 4. Cl₂ — ____ / ____ / ΔEN = ____ / Type: ____ / Conducts? ____ - 5. H₂O (per O–H bond) — ____ / ____ / ΔEN = ____ / Type: ____ - 6. KI — ____ / ____ / ΔEN = ____ / Type: ____ / Conducts? ____ - 7. HCl — ____ / ____ / ΔEN = ____ / Type: ____ / Conducts dissolved? Careful — think! - 8. CH₄ (per C–H bond) — ____ / ____ / ΔEN = ____ / Type: ____ - 9. AlF₃ — ____ / ____ / ΔEN = ____ / Type: ____ / Conducts? ____ - 10. CaO — ____ / ____ / ΔEN = ____ / Type: ____ / Melting point: high or low? ____ Part 2 — Justify. Answer in complete sentences. - A) Row 4 (Cl₂) has ΔEN = 0. Explain why this bond is purely covalent, not 'half covalent.' - B) Row 7 (HCl) is covalent, yet HCl solution lights a conductivity bulb. Explain why. - C) Compare rows 3 (MgO) and 8 (CH₄). Which has the higher melting point, and why does bond type predict this?

   Materials

   • Printed worksheet (content below)
   • Periodic table with electronegativity values
   • Pencil

   Example outputs

   • Row 1 (NaF): EN(Na)=0.9, EN(F)=4.0, ΔEN = 3.1 → ionic → conducts when dissolved (Na⁺ and F⁻ separate in water).
   • Row 4 (Cl₂): EN(Cl)=3.2, EN(Cl)=3.2, ΔEN = 0.0 → nonpolar covalent → does not conduct. Purely covalent because identical atoms pull equally — 'half covalent' is not a real category.
   • Row 7 (HCl): EN(H)=2.2, EN(Cl)=3.2, ΔEN = 1.0 → polar covalent, BUT HCl ionizes in water into H⁺ + Cl⁻, so the solution conducts. This is the exception to 'covalent doesn't conduct.'
   • Row 10 (CaO): ΔEN = 2.4 → ionic → high melting point (~2600 °C) because breaking the 3-D lattice requires enormous energy.
   • Part 2C: MgO (ionic, ΔEN = 2.1) melts at ~2850 °C; CH₄ (nonpolar covalent, ΔEN = 0.4) melts at −182 °C. Ionic lattices require breaking every ion–ion attraction, while covalent CH₄ only needs to overcome weak intermolecular forces.
   • example_outputs

1. Given: EN(K) = 0.8 and EN(Br) = 3.0. Calculate ΔEN and predict the bond type in KBr. Justify with the cutoff you used.

   calculationΔEN = |0.8 − 3.0| = 2.2. Because 2.2 > 1.7, KBr is ionic.

2. A white solid melts at 850 °C, does not conduct as a solid, but conducts strongly when dissolved in water. Is it most likely ionic or covalent? Justify.

   short answerIonic. High melting point indicates a strong 3-D lattice; no conduction as a solid because ions are locked in place; conduction when dissolved because ions separate and can carry charge. All three properties match ionic bonding.

3. Which statement is TRUE? A) A bond between two identical atoms (like Cl₂) is half ionic. B) All covalent compounds fail to conduct electricity in water. C) A bond with ΔEN = 0 is purely nonpolar covalent. D) All ionic compounds dissolve in water.

   multiple choiceC. Two identical atoms have ΔEN = 0, so electrons are shared perfectly equally — this is purely covalent. A is wrong (identical atoms are 100% covalent, not half ionic). B is wrong (HCl and other acids ionize and conduct). D is wrong (CaCO₃ is ionic but barely soluble).

4. Draw the Lewis structure of H₂O and identify whether the O–H bond is nonpolar covalent, polar covalent, or ionic. Use EN(H) = 2.2 and EN(O) = 3.4.

   short answerLewis structure: central O with 2 lone pairs (4 dots) and two single bonds to H atoms, bent shape. ΔEN = |2.2 − 3.4| = 1.2, which is between 0.4 and 1.7, so the O–H bond is polar covalent. O carries δ⁻ and each H carries δ⁺.

• ☐Printed 'Predict the Bond Type' worksheet (one per student)
• ☐Periodic table handout with Pauling electronegativity values
• ☐Front-bench demo: 2 clear cups, distilled water, spatula of NaCl, spatula of sucrose, conductivity tester (battery + bulb probe)
• ☐Pencils