The shape of the lesson
By the end of the lesson, learners can
- Radius, charge and shielding: Step across Period 3 in the simulator and watch the atom shrink.
- Structure decides melting point: Switch to the structure view.
- First ionisation energy: The general rise across the period has two dips.
- Radius down; IE and electronegativity up.
- IE dips at Al (3p) and S (paired 3p).
- m.p.: metallic rising, Si peak, simple molecular low.
- Conductivity: Na < Mg < Al; Si semiconductor.
One picture every idea returns to
Across Period 3 the nuclear charge rises while electrons enter the same shell, so shielding is almost constant: the atomic radius decreases, and first ionisation energy and electronegativity generally increase. Melting point depends on structure: it rises from Na to Al as metallic bonding strengthens, peaks at silicon (giant covalent), then drops sharply for the simple molecular P4, S8, Cl2 and monatomic Ar. Conductivity rises from Na to Al (more delocalised electrons); Si is a semiconductor and the rest are insulators.
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, Eight elements, one row.. Walk from sodium to argon and the elements change from soft reactive metals, through a giant covalent semiconductor, to small molecules and a noble gas. 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: Radius, charge and shielding | Step across Period 3 in the simulator and watch the atom shrink. Project tab 1 of the lab (Periodicity lab, trends across period 3) and let a student drive it while the class predicts. | Whole class, then pairs |
| 13 to 21 min | Teach and check: Structure decides melting point | Switch to the structure view. Project tab 2 of the lab (Periodicity lab, structure and bonding) and let a student drive it while the class predicts. | Pairs |
| 21 to 30 min | Teach and check: First ionisation energy | The general rise across the period has two dips. Project tab 5 of the lab (Periodicity lab, ionisation energies) and let a student drive it while the class predicts. | Connect-Extend-Challenge |
| 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, 15 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 First ionisation 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.
Connect-Extend-Challenge
After the new idea, each student writes three lines: how it connects to what they knew, how it extends their thinking, and what still challenges them. Pairs compare, then the challenges become the plenary.
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 |
|---|---|
| Silicon's high melting point is due to its giant covalent structure, not to being a metal. | And the fall from sulfur to chlorine to argon in melting point is about intermolecular forces between small molecules or atoms; no covalent bonds are broken when P, S or Cl 2 melt. |
Period 3: physical properties: worksheet
6 questions, 15 marks. Written in the style of the papers; no past paper question is reproduced.
Explain why atomic radius decreases across Period 3.
Explain the variation in melting point from sodium to argon.
Explain why the first ionisation energy (a) generally increases across Period 3, (b) falls from Mg to Al.
Explain why Na+, Mg2+ and Al3+ decrease in radius.
Explain why silicon is a semiconductor while aluminium is a good conductor.
State and explain the trend in electronegativity from sodium to chlorine.
Answer key and marking notes
- Nuclear charge increases from Na to Cl.
- The electrons are added to the same (third) shell, so shielding hardly changes and the outer electrons are pulled closer.
- Na, Mg, Al: giant metallic; melting point rises as the number of delocalised electrons and the ionic charge increase and the ions get smaller.
- Si: giant covalent, so many strong covalent bonds must be broken; highest melting point.
- P4, S8, Cl2: simple molecular, only weak id-id forces are overcome; S8 highest as it has the most electrons.
- Ar: single atoms with the weakest id-id forces; lowest.
- (a) Nuclear charge increases while shielding stays similar, so the outer electron is held more strongly.
- (b) Al's outer electron is in 3p, which is higher in energy than 3s and partly shielded by the 3s electrons, so less energy is needed to remove it.
- All three have the configuration 1s2 2s2 2p6.
- The nuclear charge increases from 11 to 13, pulling the same number of electrons closer.
- Aluminium has three delocalised electrons per atom free to move.
- In silicon the electrons are held in covalent bonds; only a few gain enough energy to move, so it conducts slightly (more when heated).
- Electronegativity increases.
- Increasing nuclear charge and decreasing atomic radius mean the bonding pair is attracted 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 row correctly describes how atomic radius and electronegativity change across Period 3 from sodium to chlorine?
- Aatomic radius decreases; electronegativity increases
- Batomic radius increases; electronegativity increases
- Catomic radius decreases; electronegativity decreases
- Datomic radius increases; electronegativity decreases
Which list shows the ions in order of increasing ionic radius?
- AAl3+ < Mg2+ < Na+ < Cl−
- BNa+ < Mg2+ < Al3+ < Cl−
- CCl− < Al3+ < Mg2+ < Na+
- DAl3+ < Cl− < Mg2+ < Na+
The chart shows the melting points of the Period 3 elements. Why does silicon have the highest melting point?
- Asilicon has a giant covalent structure, so a very large number of strong covalent bonds must be broken to melt it
- Bsilicon has the strongest metallic bonding, because each atom gives four electrons to the sea of delocalised electrons
- Csilicon is a non-metal made of small molecules with very strong intermolecular forces between them
- Dsilicon has the largest atoms in Period 3, so it has the strongest van der Waals forces
The graph shows the first ionisation energies of the Period 3 elements. Why is the first ionisation energy of aluminium lower than that of magnesium?
- Aaluminium loses a 3p electron, which is higher in energy than 3s and partly shielded by the 3s electrons
- Baluminium has fewer protons than magnesium, so its outer electron is held less strongly
- Caluminium's electron is removed from a 3p orbital that already contains two paired electrons, which repel
- Dthe graph shows no fall at aluminium; the fall is at magnesium, because magnesium has a full 3s subshell
Which statement about the electrical conductivity of the Period 3 elements is correct?
- Aaluminium conducts better than magnesium, which conducts better than sodium; silicon is a semiconductor and phosphorus to argon do not conduct
- Bsodium conducts best, because it is the most reactive metal and loses its outer electron most easily
- Csodium, magnesium and aluminium conduct equally well, because they are all metals
- Dsilicon conducts best, because it has the most outer-shell electrons of the solid elements
The first ionisation energy of sulfur (1000 kJ mol−1) is slightly lower than that of phosphorus (1012 kJ mol−1). What is the explanation?
- Ain sulfur (3p4) one 3p orbital holds a pair of electrons, and the repulsion between them makes one easier to remove
- Bsulfur's outer electron is in a new, higher-energy subshell, so it is further from the nucleus
- Csulfur has a smaller nuclear charge than phosphorus
- Dsulfur atoms are larger than phosphorus atoms, so the outer electron is further from the nucleus
Phosphorus, sulfur, chlorine and argon all have simple molecular or monatomic structures. Which gives their melting points in decreasing order, with the correct reason?
- AS8 > P4 > Cl2 > Ar, because larger molecules with more electrons have stronger instantaneous dipole-induced dipole forces
- BP4 > S8 > Cl2 > Ar, because phosphorus is nearer the metals and has some metallic bonding
- CCl2 > S8 > P4 > Ar, because chlorine is the most electronegative and so has permanent dipole-dipole forces
- DAr > Cl2 > S8 > P4, because argon atoms are the smallest and pack most closely
The melting point of magnesium (650 °C) is much higher than that of sodium (98 °C). Which explanation is correct?
- AMg2+ ions are smaller and more highly charged than Na+, and each Mg atom provides two delocalised electrons, so the metallic bonding is stronger
- Bmagnesium has a giant covalent structure, whereas sodium has metallic bonding
- Cmagnesium atoms are larger than sodium atoms, so there are stronger van der Waals forces between them
- Dmagnesium has fewer delocalised electrons than sodium, so the positive ions are packed more closely
Quiz key
Across the period the nuclear charge rises by one each step, while the extra electrons enter the same outer shell, so shielding stays roughly the same. The outer electrons are pulled closer, so atomic radius decreases, and the nucleus attracts a bonding pair more strongly, so electronegativity increases (chlorine is the most electronegative Period 3 element). Atomic radius does not increase: more electrons does not mean a bigger atom when no new shell is added. Electronegativity does not decrease; that is the trend down a group, where the radius and shielding increase. The last row reverses both trends, which is the pattern going down a group, not across a period.
Na+, Mg2+ and Al3+ are isoelectronic (2,8): they have lost their whole third shell, and the nuclear charge rises from 11 to 13, so the radius falls: Al3+ is smallest. Cl− keeps its third shell and has gained an electron (2,8,8), so it is much larger than any of the cations. Putting Na+ smallest reverses the cation order, ignoring the rising nuclear charge. Putting Cl− first treats the anion as small because chlorine is on the right of the period, but anions have an extra shell compared with these cations. Cl− cannot fall between Al3+ and Mg2+ because it has one more occupied shell than both.
Silicon forms a giant covalent (macromolecular) lattice like diamond, each atom bonded to four others, so melting means breaking many strong covalent bonds: melting point 1410 °C. Silicon is not a metal and has no sea of delocalised electrons; the metallic bonding trend stops at aluminium (660 °C). Being a non-metal does not make silicon molecular: the simple molecular elements (P4, S8, Cl2, Ar) are the ones with low melting points, 115 °C or less. Atomic radius decreases across the period, so silicon does not have the largest atoms, and it is not held together by van der Waals forces.
Magnesium's outer electron is 3s; aluminium's is 3p1. The 3p subshell is higher in energy and its electron is partly shielded by the 3s2 pair, so less energy is needed to remove it: 578 kJ mol−1 for Al against 738 kJ mol−1 for Mg, even though Al has one more proton. Aluminium has 13 protons, one more than magnesium, so fewer protons is wrong. Paired 3p electrons do not appear until sulfur (3p4); that is the reason for the second fall, at S. The graph clearly rises from Na to Mg and falls at Al; a full 3s subshell makes Mg harder, not easier, to ionise.
Na, Mg and Al provide 1, 2 and 3 delocalised electrons per atom, so conductivity rises Na < Mg < Al. Silicon is a semiconductor: its electrons are held in covalent bonds, and only a few are freed at room temperature. Phosphorus, sulfur, chlorine and argon are molecular or monatomic with no mobile charge carriers, so they do not conduct. Reactivity is not the same as conductivity; sodium has the fewest delocalised electrons per atom, so it conducts least of the three metals. Being metals does not make them equal, because the number of delocalised electrons differs. Silicon's four outer electrons are localised in covalent bonds, so it conducts far less than the metals.
Phosphorus is 3p3, with one electron in each 3p orbital. In sulfur the fourth 3p electron must pair up in an orbital that is already occupied, and the repulsion between the paired electrons makes one of them easier to remove, outweighing the extra proton. The new-subshell reason explains the fall at aluminium (3s to 3p), not at sulfur, whose electron is still 3p. Sulfur has 16 protons, one more than phosphorus, so its nuclear charge is larger. Atomic radius decreases across the period, so sulfur atoms are slightly smaller than phosphorus atoms.
Melting points: S 115 °C, P 44 °C, Cl −101 °C, Ar −189 °C. Only instantaneous dipole-induced dipole (London) forces act between these non-polar particles, and they are stronger for particles with more electrons: S8 (128 electrons) > P4 (60) > Cl2 (34) > Ar (18). White phosphorus is simple molecular and has no metallic bonding. Cl2 has two identical atoms, so it is non-polar and has no permanent dipole; electronegativity only matters between different atoms. Argon has the fewest electrons and so the weakest forces, giving the lowest melting point, whatever its packing.
Metallic bond strength depends on the charge and size of the ions and on the number of delocalised electrons. Mg2+ has twice the charge of Na+ and is smaller, and there are twice as many delocalised electrons per ion, so the electrostatic attraction is much stronger and more energy is needed to melt the metal. Magnesium is a metal with metallic bonding, not a giant covalent structure (the first giant covalent element is silicon). Magnesium atoms are smaller than sodium atoms, and metals are not held together by van der Waals forces. Magnesium has more delocalised electrons per atom (two) than sodium (one), not fewer.
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.