The shape of the lesson
By the end of the lesson, learners can
- Hydrogen and a lone pair: Load water and then hydrogen fluoride in the intermolecular forces simulator and switch on the hydrogen bond view.
- Anomalies it explains: Fig. 1 shows the boiling points of the hydrides rising down each group as the number of electrons rises.
- Bond length and bond energy: Inside a molecule, the covalent bonds have measurable lengths and strengths.
- Donor: H covalently bonded to N, O or F (large δ+).
- Acceptor: a lone pair on N, O or F in another molecule.
- Draw: partial charges, lone pair, dashed line in a straight line.
- Shorter bond, stronger bond: C≡C > C=C > C–C.
One picture every idea returns to
A hydrogen bond forms between a hydrogen atom bonded to a very electronegative atom (N, O or F) in one molecule and a lone pair on N, O or F in a neighbouring molecule. The H carries a large δ+ charge and is tiny, so the lone pair can approach it closely. Hydrogen bonds are the strongest intermolecular forces, which explains the unusually high boiling points of H2O, HF and NH3 and the open, less dense structure of ice. Inside molecules, bond energy is the energy needed to break one mole of a covalent bond in the gas phase; shorter bonds are generally stronger, so C≡C > C=C > C–C.
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, The strongest attraction.. Water should boil at about −80 °C if it followed the trend of its group. 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: Hydrogen and a lone pair | Load water and then hydrogen fluoride in the intermolecular forces simulator and switch on the hydrogen bond view. Project tab 7 of the lab (Chemical bonding lab, intermolecular forces) and let a student drive it while the class predicts. | Whole class, then pairs |
| 13 to 21 min | Teach and check: Anomalies it explains | Fig. 1 shows the boiling points of the hydrides rising down each group as the number of electrons rises. | Pairs |
| 21 to 30 min | Teach and check: Bond length and bond energy | Inside a molecule, the covalent bonds have measurable lengths and strengths. | Anticipation / Reaction Guide |
| 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, 14 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 Bond length and bond energy, 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.
Anticipation / Reaction Guide
Before the lesson, students mark five statements true or false. After the teaching they mark them again and justify any change of mind in one sentence each. The statements come from the examiner traps below.
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 |
|---|---|
| A hydrogen bond is an intermolecular force, not a covalent bond; the covalent O–H bond inside water is about 20 times stronger. | When drawing hydrogen bonding you must show the lone pair and the partial charges, and the dashed line must go from the lone pair to the hydrogen, not to the other atom. |
Hydrogen bonding and bond properties: worksheet
6 questions, 14 marks. Written in the style of the papers; no past paper question is reproduced.
State the two conditions needed for a hydrogen bond to form between two molecules.
Draw a diagram to show a hydrogen bond between two ammonia molecules. Include partial charges and lone pairs.
Using Fig. 1, explain why the boiling points of CH4, SiH4, GeH4 and SnH4 increase steadily while H2O has a much higher boiling point than H2S.
Explain why ice floats on water.
Explain why ethanol, C2H5OH, is completely miscible with water, but ethane is not.
The bond lengths of C–C, C=C and C≡C are 0.154, 0.134 and 0.120 nm. Describe the relationship between bond length and bond energy, and explain it.
Answer key and marking notes
- One molecule has a hydrogen atom covalently bonded to N, O or F.
- The other molecule has a lone pair of electrons on an N, O or F atom.
- One NH3 with Nδ− and Hδ+ labelled; the second NH3 with its lone pair drawn on N.
- A dashed line from the lone pair on N of the second molecule to a δ+ H of the first.
- The N–H···N arrangement drawn roughly linear.
- Down Group 14 the molecules have more electrons, so id-id forces get steadily stronger; CH4 cannot hydrogen bond.
- H2O forms hydrogen bonds because O is very electronegative and has lone pairs; H2S cannot.
- Hydrogen bonds are much stronger than the other intermolecular forces, so much more energy is needed to boil water.
- In ice each water molecule forms four hydrogen bonds, holding the molecules in an open lattice.
- The molecules are further apart than in liquid water, so ice is less dense.
- Ethanol has an O–H group, so it can form hydrogen bonds with water molecules.
- Ethane is non-polar with only id-id forces, which cannot compete with the hydrogen bonds between water molecules.
- The shorter the bond, the higher the bond energy: C≡C > C=C > C–C.
- More shared pairs increase the electron density between the nuclei, pulling them closer and holding them more strongly.
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 conditions are needed for a hydrogen bond to form between two molecules?
- Aa hydrogen atom covalently bonded to N, O or F in one molecule, and a lone pair on an N, O or F atom in the other
- Bany hydrogen atom in one molecule, and any electronegative atom in the other
- Ca hydrogen atom bonded to carbon, and a double bond in the other molecule
- Dtwo hydrogen atoms, one on each molecule, attracted to each other
The graph shows the boiling points of the hydrides of Group 14 and Group 16. Why is the boiling point of water so much higher than the trend in Group 16 suggests?
- Awater molecules form hydrogen bonds with each other, which need more energy to overcome than the forces between the other hydrides
- Bwater has more electrons than the other Group 16 hydrides, so its id-id forces are stronger
- Cthe O–H covalent bonds are very strong and must be broken when water boils
- Dwater is the only Group 16 hydride with permanent dipoles
Which compound forms hydrogen bonds between its own molecules?
- Amethanol, CH3OH
- Bmethoxymethane, CH3OCH3
- Cethanal, CH3CHO
- Dfluoromethane, CH3F
Why is ice less dense than liquid water?
- Ain ice each molecule is hydrogen bonded to four others in an open lattice, which holds the molecules further apart than in the liquid
- Bice contains trapped air bubbles that lower its density
- Cthe covalent O–H bonds become longer when water freezes
- Dhydrogen bonds break when water freezes, so the molecules move apart
The bond lengths of C–C, C=C and C≡C are 0.154 nm, 0.134 nm and 0.120 nm. Which statement about their bond energies is correct?
- AC≡C has the largest bond energy, because more shared electrons pull the nuclei closer and hold them more strongly
- BC–C has the largest bond energy, because it is the longest bond
- Call three have the same bond energy, because the same two elements are bonded
- DC=C has exactly twice the bond energy of C–C
Bond energies such as C–H (413 kJ mol−1) are quoted as mean values in data tables. Why?
- Athe energy of a particular bond varies slightly with the molecule it is in, so an average over many compounds is used
- Bthe value is the mean of the bond making and bond breaking energies
- Cthe value is averaged over the three states of matter
- Dbond energies cannot be measured, so they are estimated from atomic radii
Fluorine is more electronegative than oxygen, and the hydrogen bond F–H···F is stronger than O–H···O. Yet hydrogen fluoride boils at 20 °C and water at 100 °C. What is the best explanation?
- Aeach water molecule has two H atoms and two lone pairs, so it forms on average twice as many hydrogen bonds as an HF molecule, which has only one H atom
- BHF molecules have more electrons than water, so their id-id forces are weaker
- CHF does not form hydrogen bonds because fluorine has no lone pairs
- Dthe H–F covalent bond is weaker than the O–H bond, so HF boils more easily
In the graph, the Group 14 hydrides show a steady rise in boiling point with no anomaly at methane, CH4. What is the explanation?
- Acarbon is not electronegative enough to form hydrogen bonds, and the rise is due to id-id forces increasing with the number of electrons
- Bmethane forms hydrogen bonds, but they are cancelled by its tetrahedral shape
- Cthe Group 14 hydrides are ionic, so their boiling points depend on ion size
- Dmethane has the strongest permanent dipoles, which are balanced by its small size
Quiz key
Only N, O and F are electronegative enough to leave a bonded hydrogen strongly δ+ with almost no shielding; this hydrogen is then attracted to a lone pair on an N, O or F atom of a neighbouring molecule. Any hydrogen is not enough: in C–H the hydrogen is barely δ+ and does not hydrogen bond, and chlorine, though electronegative, is too large to form hydrogen bonds in the AS model. A C–H hydrogen and a double bond do not meet either condition. Two δ+ hydrogen atoms would repel, not attract.
Oxygen is small and highly electronegative, so the δ+ hydrogen of one water molecule is strongly attracted to a lone pair on the oxygen of another: this is hydrogen bonding, much stronger than the other intermolecular forces, so water boils at 100 °C instead of the roughly −80 °C the trend predicts. Water has the fewest electrons of the four (10), so its id-id forces are the weakest. Boiling separates molecules; it does not break O–H covalent bonds. H2S, H2Se and H2Te are bent molecules with permanent dipoles too.
Methanol has an O–H group, so its δ+ hydrogen can bond to a lone pair on the oxygen of another methanol molecule. In methoxymethane and ethanal the oxygen has lone pairs, but all the hydrogen atoms are bonded to carbon, so there is no suitable δ+ hydrogen; these compounds can accept hydrogen bonds from water but cannot hydrogen bond with themselves. In fluoromethane the fluorine is bonded to carbon, and again every hydrogen is on carbon.
On freezing, each water molecule forms the maximum of four hydrogen bonds (two through its H atoms and two through its lone pairs), arranged tetrahedrally in an open hexagonal lattice with a lot of empty space. When ice melts, some hydrogen bonds break and the molecules can pack more closely, so the liquid is denser. Pure ice without any bubbles is still less dense than water. The O–H covalent bond length hardly changes on freezing. Freezing forms more hydrogen bonds, not fewer.
More shared pairs give more electron density between the nuclei, so the nuclei are pulled closer (a shorter bond) and held more strongly: mean bond energies are about 347, 612 and 838 kJ mol−1 for C–C, C=C and C≡C. A longer bond means the nuclei are further from the shared electrons, so it is weaker, not stronger. Bond energy depends on the number of shared pairs, not only on which elements are bonded. C=C (612) is less than twice C–C (694) because the π bond is weaker than the σ bond.
The exact energy needed to break a C–H bond depends on its surroundings: the four successive C–H bond dissociations in methane all differ, and C–H bonds in other compounds differ again. Tables therefore give a mean value over a range of compounds, which is why calculations using them give only approximate enthalpy changes. Making and breaking the same bond involve the same amount of energy, with opposite signs, so there is nothing to average. Bond energies refer to gaseous molecules only, not to all three states. Bond energies are found from experimental enthalpy data, not from atomic radii.
An HF molecule has three lone pairs but only one hydrogen, so on average it can form only one hydrogen bond per molecule (chains). A water molecule has two hydrogens and two lone pairs, so it can form up to four hydrogen bonds, an average of two per molecule, making the total intermolecular attraction greater. HF and H2O both have 10 electrons, and more electrons would strengthen, not weaken, id-id forces. Fluorine in HF has three lone pairs, and HF does form hydrogen bonds, which is why it boils much higher than HCl. Boiling does not break the covalent H–F bond, which in fact is stronger than O–H.
C–H bonds are almost non-polar and carbon has no lone pair in CH4, so methane cannot hydrogen bond and there is no anomaly. All four hydrides are tetrahedral and non-polar, so only id-id forces act; these increase from CH4 (10 electrons) to SnH4 (54 electrons), giving the steady rise. Shape cancels dipoles, not hydrogen bonds, and methane has no hydrogen bonds to cancel. The hydrides are covalent molecules, not ionic. Being symmetrical, the Group 14 hydrides have no permanent dipoles at 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.