The shape of the lesson
By the end of the lesson, learners can
- A tug of war: Pick pairs of atoms in the simulator and watch where the bonding pair sits.
- Trends in the Periodic Table: Across a period the nuclear charge rises while shielding stays almost the same, so the atoms shrink and pull harder.
- From covalent to ionic: Subtract the two electronegativities.
- Definition: attraction for the bonding pair in a covalent bond.
- Across a period: increases; down a group: decreases.
- Difference 0: non-polar; small: polar; large: ionic.
- δ− goes on the more electronegative atom.
One picture every idea returns to
Electronegativity is the power of an atom to attract the pair of electrons in a covalent bond towards itself. It increases with nuclear charge and decreases as the atomic radius and shielding increase, so it rises across a period and falls down a group; fluorine (4.0 on the Pauling scale) is the most electronegative element. When two bonded atoms have equal electronegativities the bond is pure covalent. A small difference gives a polar covalent bond with δ+ and δ− ends; a large difference (roughly 2 or more) means the electron transfer is essentially complete and the bonding is ionic.
Forty-five minutes, phase by phase
| Time | Phase | What happens in the room | Grouping |
|---|---|---|---|
| 0 to 5 min | Hook: predict | Open with the question in the title, Who pulls the pair?. Every bond is a tug of war over a shared pair of electrons. Put the figure on the board: it is the one picture every idea in the lesson returns to. | Whole class, then pairs |
| 5 to 13 min | Teach and check: A tug of war | Pick pairs of atoms in the simulator and watch where the bonding pair sits. Project tab 5 of the lab (Chemical bonding lab, polarity) and let a student drive it while the class predicts. | Whole class, then pairs |
| 13 to 21 min | Teach and check: Trends in the Periodic Table | Across a period the nuclear charge rises while shielding stays almost the same, so the atoms shrink and pull harder. | Pairs |
| 21 to 30 min | Teach and check: From covalent to ionic | Subtract the two electronegativities. | Cooperative Concept Mapping |
| 30 to 36 min | Check | The four Stage 1 questions from the student lesson, answered on mini whiteboards or by show of fingers. Re-teach the one most get wrong. | Individual, then whole class |
| 36 to 43 min | Plenary | The exam-style questions from Stage 2, one at a time: thirty seconds alone, then a random call. The feedback lines in the key give the examiner's wording. | Individual, random call |
| 43 to 45 min | Exit and homework | Set the worksheet below (6 questions, 13 marks). Learners self-assess against the objectives before they leave. | Individual |
Protect the parts that carry the learning
The sequence carries about fifty minutes of material, so the lab is driven by one student while the class predicts rather than explored freely. Times are a guide.
Protected: the check at 30 minutes. It is the only individual read on whether the key idea landed, and it is not cut.
Cut first if the build overruns: the last check in From covalent to ionic, never the check or the plenary.
In a 60 minute block: run worksheet questions 1 to 3 in class as a Rally Coach set and model both the hardest worked example and its check live.
Cooperative Concept Mapping
Teams receive the key terms on cards and arrange them on a large sheet, drawing and labelling the links. Compare maps between teams before showing the model map.
The full structure, with its variations and the other twelve, is on the cooperative-learning page.
What to head off, and how
| Trap learners fall into | Teaching move that pre-empts it |
|---|---|
| Electronegativity is about electrons in a bond. | Do not confuse it with electron affinity or ionisation energy, which describe isolated gaseous atoms. And a molecule can contain polar bonds yet be non-polar overall if the dipoles cancel by symmetry, as in CO 2; that idea returns in the intermolecular forces lesson. |
Electronegativity and bonding: worksheet
6 questions, 13 marks. Written in the style of the papers; no past paper question is reproduced.
Define electronegativity.
Explain why the electronegativity of chlorine (3.0) is greater than that of bromine (2.8).
Explain why electronegativity increases from lithium to fluorine across Period 2.
Using the values H 2.1, C 2.5, N 3.0, O 3.5 and Cl 3.0, place these bonds in order of increasing polarity: C–H, O–H, C–Cl, N–H.
Draw the C=O bond in propanone and the C–Br bond in bromoethane, labelling the partial charges. (C 2.5, O 3.5, Br 2.8)
Magnesium oxide is ionic but carbon monoxide is covalent. Use electronegativity values (Mg 1.2, C 2.5, O 3.5) to explain this difference.
Answer key and marking notes
The power of an atom to attract the pair of electrons in a covalent bond towards itself.
- A chlorine atom has fewer electron shells, so its atomic radius is smaller.
- There is less shielding of the nuclear charge by inner electrons.
- The bonding pair is closer to the nucleus and more strongly attracted, which outweighs bromine's larger nuclear charge.
- The nuclear charge increases while the extra electrons go into the same shell, so shielding is almost constant.
- The atomic radius decreases, so the bonding pair is attracted more strongly.
- Differences: C–H 0.4, C–Cl 0.5, N–H 0.9, O–H 1.4.
- Order: C–H < C–Cl < N–H < O–H.
- Cδ+=Oδ−: oxygen is more electronegative.
- Cδ+–Brδ−: bromine is more electronegative, although the difference is only 0.3.
- Mg and O differ by 2.3; this is large enough for oxygen to take the electrons completely, forming Mg2+ and O2− ions.
- C and O differ by only 1.0, so the electrons are shared, although unequally.
- The bond in CO is therefore polar covalent, not ionic.
8 multiple-choice questions on this lesson
From the topic question bank. Assign them in the app for a marked score in your gradebook, or run the topic as a class game.
Which statement is the definition of electronegativity?
- Athe power of an atom to attract the pair of electrons in a covalent bond towards itself
- Bthe energy needed to remove one electron from each atom in one mole of gaseous atoms
- Cthe energy released when one mole of gaseous atoms each gain one electron
- Dthe tendency of an atom to lose electrons and form a positive ion
Electronegativity increases across Period 3 from sodium to chlorine. What is the main reason?
- Anuclear charge increases while shielding stays about the same, so atomic radius decreases and the bonding pair is attracted more strongly
- Beach element has one more electron shell than the one before, so the nucleus is better shielded
- Cshielding increases across the period, so the outer electrons are held more tightly
- Dthe atoms have more outer electrons, which repel the bonding pair towards the nucleus
Why does electronegativity decrease down Group 17 from fluorine to iodine?
- Aatomic radius and shielding increase, which outweighs the increase in nuclear charge
- Bnuclear charge decreases down the group
- Cthe atoms have fewer outer-shell electrons further down the group
- Diodine is a solid, so its electrons are less free to move towards another atom
Which bond is the most polar? (Electronegativities: H 2.1, C 2.5, N 3.0, O 3.5, F 4.0)
- AH–F
- BO–H
- CN–H
- DC–H
Which correctly shows the partial charges in the C–Cl bond of chloromethane? (Electronegativities: C 2.5, Cl 3.0)
- AC is δ+ and Cl is δ−
- BC is δ− and Cl is δ+
- Cthere are no partial charges because both atoms are non-metals
- DC is +1 and Cl is −1 because an electron is transferred
Which list shows the bonds in order of increasing polarity? (Electronegativities: H 2.1, C 2.5, Cl 3.0, O 3.5, F 4.0)
- AC–H < C–Cl < C–O < C–F
- BC–F < C–O < C–Cl < C–H
- CC–H < C–O < C–F < C–Cl
- DC–Cl < C–H < C–O < C–F
In which compound does the hydrogen carry a negative charge? (Electronegativities: Na 0.9, H 2.1, C 2.5, Cl 3.0, O 3.5)
- ANaH
- BHCl
- CH2O
- DCH4
The chart shows Pauling electronegativities across Period 2. No value is given for neon. What is the best explanation?
- Aneon does not form covalent bonds, so it has no bonding pair to attract and a Pauling value cannot be assigned
- Bneon's nuclear charge is completely shielded, so its electronegativity is exactly zero
- Cneon's value would be above 4.0, which is too large to fit on the Pauling scale
- Dneon atoms are too small for their electronegativity to be measured
Quiz key
Electronegativity describes how strongly an atom pulls the shared (bonding) pair of electrons in a covalent bond towards itself; it is a relative number with no units. Removing an electron from gaseous atoms is the first ionisation energy, measured in kJ mol−1. Energy released on gaining an electron is the electron affinity, which also involves complete transfer rather than sharing. Losing electrons to form positive ions is electropositive (metallic) character, roughly the opposite of electronegativity.
Across Period 3 the proton number rises from 11 to 17 but the extra electrons enter the same (third) shell, so shielding by the inner electrons stays roughly constant. The larger nuclear charge pulls the shells in, the radius falls, and the bonding pair sits closer to a more strongly charged nucleus. The number of shells does not change across a period; it changes down a group. Increased shielding would weaken, not strengthen, the attraction. Outer electrons do not push the bonding pair towards the nucleus; the pull comes from the nuclear charge.
Each step down the group adds a shell, so the bonding pair is further from the nucleus and more inner electrons shield it. These effects outweigh the larger number of protons, so the attraction for the bonding pair falls (F 4.0, Cl 3.0, Br 2.8, I 2.5). Nuclear charge actually increases down the group. All the halogens have seven outer electrons. The physical state of the element is caused by intermolecular forces and has nothing to do with the attraction of one atom for a bonding pair.
The electronegativity differences are H–F 1.9, O–H 1.4, N–H 0.9 and C–H 0.4, so H–F is the most polar. O–H is strongly polar but less so than H–F because oxygen is less electronegative than fluorine. N–H is only moderately polar. C–H has such a small difference that it is usually treated as almost non-polar.
Chlorine (3.0) is more electronegative than carbon (2.5), so the bonding pair is drawn towards Cl, giving Cl a small negative charge (δ−) and leaving C slightly positive (δ+). Reversing the charges ignores which atom has the greater pull. Two different non-metals still usually differ in electronegativity, so their bond is polar; only identical atoms share exactly equally. Full charges of +1 and −1 describe ionic bonding; with a difference of only 0.5 the electron pair is shared unequally, not transferred.
The differences are C–H 0.4, C–Cl 0.5, C–O 1.0 and C–F 1.5, so polarity increases in that order. The second option is the same list reversed (decreasing polarity). Placing C–Cl last confuses atom size or number of electrons with electronegativity; chlorine is larger than fluorine but much less electronegative. Putting C–Cl below C–H misreads the differences; in C–H carbon is the δ− atom, but the difference of 0.4 is still smaller than 0.5.
Sodium (0.9) is much less electronegative than hydrogen (2.1), so electron density moves towards H; sodium hydride is ionic, Na+H−, containing the hydride ion. In HCl and H2O the other atom is more electronegative (3.0 and 3.5), so H is δ+. In CH4 carbon (2.5) is slightly more electronegative than hydrogen, so any partial charge on H is positive, not negative.
Electronegativity is defined for an atom in a covalent bond, and Pauling values were derived from bond energies. Neon has a full outer shell and forms no stable covalent compounds, so there is no bonding pair and no value is assigned. Its nuclear charge is not fully shielded; neon has a very high first ionisation energy, which shows its electrons are held strongly, so zero would be wrong. The scale has no upper limit that prevents a larger number; fluorine is simply the most electronegative bonded element. Small size does not stop a value being found: fluorine is also a very small atom and has the highest value of all.
What to project
The student lesson embeds these tabs. Open them full screen on the projector; every slider has a drag handle on the stage.
Slides for this lesson
A projectable deck for every pack is on the way. Until then, present the student lesson full screen: its sections, figure and lab tabs follow the same order as the plan above.