What this chapter covers9.1, 9.2, 9.3
The Periodic Table is not a filing cabinet. Arranged by proton number, the elements fall into a pattern that repeats: whatever happens to atomic radius, melting point or the acidity of the oxides as you cross one period happens again, in the same shape, as you cross the next. That repetition is what the word periodicity means, and topic 9 is the topic where you are asked to describe it, explain it, and then use it to predict chemistry you have never met.
Period 3 — sodium to argon — is the worked example the syllabus chooses, for a good reason. Eight elements in a row take you from a soft reactive metal through a semiconductor to a gas that reacts with nothing; from a strongly basic oxide to a strongly acidic one; from a chloride that dissolves quietly to one that fumes and hisses. Every idea you have met so far — shells and shielding, ionic and covalent bonding, electronegativity, structure — is needed to explain that, and the reward is that eight elements’ worth of chemistry collapses into about four arguments.
What topic 9 asks you to do
9.1 Periodicity of physical properties of the elements in period 3
Describe qualitatively, and indicate the periodicity in, the variations in atomic radius,
ionic radius, melting point and electrical conductivity of the elements;
and explain the variation in melting point and electrical conductivity in terms of the structure
and bonding of the elements.
9.2 Periodicity of chemical properties of the elements in period 3
Describe and write equations for the reactions of the elements with oxygen, chlorine and water;
state and explain the variation in oxidation number of the oxides and chlorides in terms of their
outer-shell electrons; describe and write equations for the reactions of the oxides with water,
including the likely pHs; describe, explain and write equations for the acid/base behaviour of the
oxides and of the hydroxides NaOH, Mg(OH)2 and Al(OH)3, including amphoteric
behaviour; describe, explain and write equations for the reactions of the chlorides with water,
including the likely pHs; explain all of those variations in terms of bonding and
electronegativity; and suggest the type of bonding in a chloride or an oxide from observations of
its properties.
9.3 Chemical periodicity of other elements
Predict the characteristic properties of an element in a given group using chemical periodicity;
and deduce the nature, possible position in the Periodic Table and identity of unknown elements
from given information about their physical and chemical properties.
Notice the verbs. Describe and explain carry most of the marks, and the explanations are expected in a particular vocabulary — nuclear charge, shielding, structure, bonding, electronegativity. There is almost no calculation in this topic. What there is instead is a requirement to say why, in a sentence that could not have been written about a different element.
What is deliberately not here
Group 2 and group 17 are separate units in this syllabus (topics 10 and 11) and are not on this page, although the habits of thought built here are exactly the ones those topics use. Transition elements, period 4, and the chemistry of phosphorus(III) oxide, the oxides of chlorine, PCl3 and the sulfur chlorides are all outside topic 9 and are left out — the syllabus names the compounds it wants, and this page keeps to that list. Lattice energy, Born–Haber cycles and the quantitative treatment of bond polarity belong to A Level units later in the course.
The eight elements9.1.1
Before any trend makes sense you need to know what each of these elements actually is — what its particles are and what holds them together. Almost every explanation in 9.1 is a consequence of the row below.
| Element | Outer shell | Structure | The particles | Forces holding them |
|---|---|---|---|---|
| Na | 3s1 | giant metallic | Na+ ions in a sea of electrons | metallic bonding, 1 delocalised electron per ion |
| Mg | 3s2 | giant metallic | Mg2+ ions in a sea of electrons | metallic bonding, 2 delocalised electrons per ion |
| Al | 3s23p1 | giant metallic | Al3+ ions in a sea of electrons | metallic bonding, 3 delocalised electrons per ion |
| Si | 3s23p2 | giant covalent | atoms, each bonded to four others | covalent bonds throughout the crystal |
| P | 3s23p3 | simple molecular | P4 molecules | van der Waals forces between molecules |
| S | 3s23p4 | simple molecular | S8 molecules | van der Waals forces between molecules |
| Cl | 3s23p5 | simple molecular | Cl2 molecules | van der Waals forces between molecules |
| Ar | 3s23p6 | monatomic | single Ar atoms | van der Waals forces between atoms |
Three kinds of structure, in one row
Period 3 gives you a giant metallic structure, a giant covalent structure and a simple molecular structure within seven steps. That is why it is the period the syllabus chooses: almost every question about melting points, conductivity, hardness or solubility that can be asked at AS can be asked about these eight elements, and the answer is always found by naming the structure first and then naming the force that has to be overcome.
Everything on this page that is plotted comes from the same small table of measurements, printed in full in the last section. The model below draws any of it, and generates the explanation from the structure of each element rather than from a stored sentence.
Say period 3, not the period
Questions in this topic almost always name the period, and the explanation for a trend across a period is not the explanation for a trend down a group. Across: same shell, rising nuclear charge, shielding unchanged. Down: extra shell each time, so distance and extra shielding beat the rising charge. Mixing the two is the single commonest way to lose marks here.
Atomic radius9.1.1
Sodium’s atoms are 186 pm across the radius; chlorine’s are 99 pm. The atom has almost halved while gaining six protons and six electrons, and explaining that is the model for every other explanation in 9.1.
The explanation, in the words the mark scheme wants
Across period 3 the nuclear charge increases from +11 to +17. The extra electrons are added to the same outer shell, so the shielding by inner-shell electrons stays essentially the same — ten electrons throughout. The effective nuclear charge felt by the outer electrons therefore increases, and they are pulled in more strongly, so the atomic radius decreases.
Three clauses, and all three are needed. "More protons, so smaller" is not enough on its own, because more electrons are being added too; what makes the argument work is that the added electrons are in the shell already being filled and so are no good at shielding one another.
Shielding, counted
Shielding is worth pinning down numerically, because "shielding increases across the period" is a common and costly mistake. The table below counts the electrons that do the shielding — those in shells inside the outer one — for every element in the period.
Trap: "shielding increases across period 3"
Wrong. It is tempting because the number of electrons increases. But shielding is done by
electrons inside the outer shell, and every element from sodium to argon has exactly the
same ten: 1s22s22p6. The electrons being added go into the third
shell, alongside the ones they would have to shield.
Right: shielding is effectively constant across the period, which is precisely why the
rising nuclear charge is felt so strongly.
Argon is not on the line
Argon’s 192 pm is a van der Waals radius — half the distance between two atoms that are merely touching. The other seven are metallic or covalent radii, measured between atoms that are bonded, which is always shorter. Plotting them together produces a dramatic jump at argon that means nothing at all. If a question gives you argon’s radius, say which kind it is before you comment on it.
Where you have met this before
The period 3 radius trend is also outcome 1.1.7 in topic 1, where it is used to introduce effective nuclear charge. It appears twice in the syllabus deliberately: once as a consequence of atomic structure, and here as an example of periodicity. The explanation is identical; what is new is the requirement to notice that period 2 does the same thing, and that period 4 will too.
Ionic radius9.1.1
Ionic radius across period 3 is not one trend but two, with a cliff in the middle. That shape is the whole answer, and the reason for it is that the elements on the left lose their outer shell while those on the right add to theirs.
| Ion | Na+ | Mg2+ | Al3+ | P3− | S2− | Cl− |
|---|---|---|---|---|---|---|
| radius / pm | 102 | 72 | 54 | 212 | 184 | 181 |
| electrons | 10 | 10 | 10 | 18 | 18 | 18 |
| occupied shells | 2 | 2 | 2 | 3 | 3 | 3 |
| protons | 11 | 12 | 13 | 15 | 16 | 17 |
The two halves, and the step between them
Na+ → Al3+: each has lost the entire third shell, so all three have
the ten electrons of neon in two shells. Across the three, the nuclear charge rises from +11 to +13
while the number of electrons stays at ten, so the pull on each electron grows and the ion gets
smaller.
P3− → Cl−: each has gained electrons into the third shell, so all three
have the eighteen electrons of argon in three shells. Across the three, the nuclear charge rises
from +15 to +17 over the same eighteen electrons, so again the ion gets smaller.
The step: two occupied shells becomes three. That, not the charge, is why 54 pm jumps to
212 pm.
Isoelectronic species: same electrons, different size
N3−, O2−, F−, Na+, Mg2+ and Al3+ all have ten electrons, and they get steadily smaller in that order: 171, 140, 133, 102, 72, 54 pm. Nothing is different about them except the number of protons doing the pulling. When a question asks you to put isoelectronic ions in order of size, the only thing to compare is the proton number — most protons, smallest ion.
Trap: "the anion is bigger because it has an extra shell"
Wrong. Cl− has the same three occupied shells as a Cl atom. The extra electron
goes into the 3p sub-shell that was already being filled.
Right: the same nuclear charge now has to hold one more electron, so the attraction per
electron is weaker, the repulsion between the electrons is greater, and the cloud expands. An
extra shell is the right answer for going down a group, not for forming an anion.
First ionisation energy1.4.3, 9.1.1
First ionisation energy is not named in the 9.1 outcomes, but it belongs on any honest account of periodicity: it is the clearest measurable proof that the pattern repeats, and examiners use it freely in questions about the Periodic Table. It is set out formally in topic 1, and what follows is the period 3 half of it, told as a periodic trend.
Definition
The first ionisation energy of an element is the energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions.
X(g) → X+(g) + e−
Every word earns its place: gaseous, one mole, one electron from each atom, and the state symbols on both sides.
Across period 3 the values run 496, 738, 578, 787, 1012, 1000, 1251, 1521 kJ mol−1. The general direction is upward, for exactly the reason the radius falls, with two dips that are the most-asked detail in the whole of topic 1.
The general rise
Nuclear charge increases, shielding is essentially unchanged, and the atomic radius decreases. The outer electron is therefore held more tightly and more energy is needed to remove it.
The dip at aluminium
Magnesium is 1s22s22p63s2; aluminium is …3s23p1. The electron removed from aluminium is the 3p one.
Why Al is lower than Mg
The electron removed from aluminium is in a 3p sub-shell, which is higher in energy and slightly further from the nucleus than the 3s sub-shell, and it is additionally shielded by the 3s2 pair. Less energy is therefore needed to remove it, despite aluminium’s greater nuclear charge.
The dip at sulfur
Phosphorus is …3s23p3: three p orbitals, one electron in each. Sulfur is …3s23p4, so one of those orbitals now holds two.
Why S is lower than P
In sulfur, two electrons occupy the same 3p orbital and repel one another. That mutual repulsion makes one of them easier to remove than the single, unpaired 3p electron of phosphorus, even though sulfur has the greater nuclear charge.
The same two dips, one period up
Period 2 runs Li 520, Be 900, B 801, C 1086, N 1402, O 1314, F 1681, Ne 2081. Boron dips below beryllium — the 2p-below-2s effect — and oxygen dips below nitrogen — the pairing effect. Identical shape, different elements. Switch on "show period 2 as well" in the model in section 2 and the two curves sit one above the other with the same two kinks. That is periodicity: not that the numbers are the same, but that the shape is.
Answering "explain the trend in first ionisation energy across period 3"
Three things are wanted, and a full-mark answer gives all three: the general rise (charge up, shielding constant, radius down), the dip at Al (3p above 3s), and the dip at S (paired 3p electrons repel). Students who describe only the rise routinely lose half the marks available.
Melting point: the metals and silicon9.1.2
The melting points across period 3 do not form a trend at all in the usual sense — they climb to a huge peak at silicon and then collapse. That is not a defect in the pattern; it is the pattern, because what has to be overcome on melting changes three times along the row.
The rule that answers every melting-point question here
Melting a solid means giving the particles enough energy to break free of the forces holding
them in place. So identify the structure first, then name the force:
giant metallic → the attraction between metal ions and delocalised electrons;
giant covalent → covalent bonds, throughout the crystal;
simple molecular → van der Waals forces between molecules;
monatomic → van der Waals forces between single atoms.
Na → Mg → Al: 371 K, 923 K, 933 K
All three are metals, so the same force is being broken and the question is only how strong it is. Three things change together across them, and all three point the same way.
| Na | Mg | Al | |
|---|---|---|---|
| charge on the ion | 1+ | 2+ | 3+ |
| delocalised electrons per ion | 1 | 2 | 3 |
| ionic radius / pm | 102 | 72 | 54 |
| melting point / K | 371 | 923 | 933 |
| enthalpy change of atomisation / kJ mol−1 | +109 | +150 | +314 |
Why the metallic bond strengthens Na → Mg → Al
Each ion carries a higher positive charge, each contributes more delocalised electrons to the sea, and each is smaller. The electrostatic attraction between the lattice of cations and the delocalised electrons is therefore stronger at every step, so more energy is needed to break it.
"Charge density" is the phrase that does all three at once
A higher charge on a smaller ion is a higher charge density, and it is worth getting into the habit of saying so — the same phrase explains the metallic bond here, the hydration of cations in section 18, and the acidity of Al3+(aq). One idea, three appearances.
Silicon: 1687 K, and the reason it is a different kind of number
Silicon is the peak of the period, at nearly twice aluminium's melting point, and the reason is a change of category rather than a change of degree.
Why silicon is so high
Silicon has a giant covalent (macromolecular) structure: every atom is joined to four others by strong covalent bonds, in a lattice that continues throughout the crystal. Melting it means breaking covalent bonds, not merely separating molecules, and a great deal of energy is needed to break enough of them.
Trap: "silicon has strong van der Waals forces"
Wrong, and it is a wrong answer that gets written often because "silicon is a non-metal" suggests molecules. There are no molecules in silicon. The word to use is giant or macromolecular, and the force is the covalent bond itself.
Melting point: the non-metals9.1.2
After silicon the melting points fall off a cliff: P 317 K, S 388 K, Cl 172 K, Ar 84 K. All four of these are held together by van der Waals forces alone, so within this group the question is only which molecule is the biggest.
Ranking molecular solids
Van der Waals forces arise from instantaneous dipoles in the electron cloud. The more electrons a molecule has and the larger its surface, the stronger those forces are and the more energy is needed to separate the molecules. So the order follows the size of the molecule: S8 (128 electrons) > P4 (60) > Cl2 (34) > Ar (18).
| P | S | Cl | Ar | |
|---|---|---|---|---|
| molecule | P4, tetrahedral | S8, a puckered ring | Cl2, diatomic | Ar, single atoms |
| electrons per molecule | 60 | 128 | 34 | 18 |
| melting point / K | 317 | 388 | 172 | 84 |
| boiling point / K | 553 | 718 | 239 | 87 |
| state at room temperature | solid | solid | gas | gas |
Trap: "chlorine has more electrons than phosphorus, so it should melt higher"
An atom of chlorine does have more electrons than an atom of phosphorus — 17 against 15. But van der Waals forces act between molecules, and the molecules are Cl2 with 34 electrons against P4 with 60. Always count per molecule, never per atom.
The model below does the whole of sections 6 and 7 in one piece. It predicts the melting order from structure alone — giant covalent, then metallic, then molecular, then monatomic, with the metals ranked by charge density and the molecules by electron count — and then puts that prediction next to the measured values.
Where the prediction fails, and why that is worth knowing
The structure argument gets seven of the eight positions right and puts sodium and sulfur the wrong way round: sulfur melts at 388 K and sodium at 371 K. Sodium has the weakest metallic bonding in the period — one small delocalised electron per ion and the largest metal ion of the three — and sulfur has the largest molecule. It is the one place where the two ranges overlap. An exam answer that says "metallic bonding is stronger than van der Waals forces" is right as a generalisation and would be marked right; knowing that the generalisation has a soft edge, and where, is what stops you being ambushed by a data question.
Boiling points tell the same story, more starkly
Melting only loosens a lattice; boiling separates every particle from every other. So the gap between the giant structures and the molecular ones widens: silicon boils at 3538 K and chlorine at 239 K. If a question gives you boiling points rather than melting points, the explanation is word for word the same.
Electrical conductivity9.1.1, 9.1.2
Conductivity across period 3 spans about twenty-five powers of ten, which is why any plot of it needs a logarithmic axis and why the description is usually given in words: good, good, good, poor, none, none, none, none.
The one requirement for conduction
A substance conducts electricity if it contains charged particles that are free to move. In a solid element that means delocalised electrons — there are no mobile ions in an element. So the question for each element is simply: are any of its electrons free of a particular atom?
| Element | Structure | Mobile charge? | Conductivity |
|---|---|---|---|
| Na, Mg, Al | giant metallic | yes — 1, 2 and 3 delocalised electrons per ion | good, and rising Na → Al |
| Si | giant covalent | almost none — every outer electron is in a localised bond | poor; a semiconductor |
| P, S, Cl, Ar | simple molecular / monatomic | none — all electrons are in bonds or lone pairs within a molecule | none; insulators |
Why conductivity rises Na → Mg → Al
Each atom contributes more electrons to the delocalised sea — one, two, then three — so there are more mobile charge carriers per unit volume. (The ions are also smaller and more highly charged, which is what raises the melting point; the same change explains both properties.)
Silicon, and the one property that runs backwards
A semiconductor has no free electrons at ordinary temperatures: every outer electron is in a covalent bond. Heat promotes a few of them into a level where they can move, so silicon’s conductivity rises as it is heated. A metal does the opposite — heating makes the ions vibrate more and obstruct the moving electrons, so its conductivity falls. If an exam question gives you an element whose conductivity increases with temperature, it is telling you the element is a semiconductor.
Trap: "silicon conducts when molten because it has ions"
Wrong. Silicon is an element made of atoms; there are no ions in it in any state. The test "conducts when molten" separates ionic compounds from covalent ones, and applying it to an element is a category error. Molten silicon does conduct rather better than solid silicon, but that is metallic behaviour of the liquid, not ions.
Why it is called periodicity9.1.1
Everything so far has been one period. What makes it periodicity rather than a list of facts about sodium to argon is that the next period does the same thing.
| Property | Period 2 (Li → Ne) | Period 3 (Na → Ar) |
|---|---|---|
| atomic radius | falls, 152 → 64 pm | falls, 186 → 99 pm |
| first ionisation energy | rises, with dips at B and O | rises, with dips at Al and S |
| structure | metallic → giant covalent → molecular → monatomic | metallic → giant covalent → molecular → monatomic |
| melting point | peaks at carbon, then collapses | peaks at silicon, then collapses |
| oxides | basic → amphoteric → acidic | basic → amphoteric → acidic |
Periodicity
Periodicity is the regular, repeating variation of properties of the elements with increasing proton number. A property is periodic if, plotted against proton number, the same shape reappears in each successive period.
Turn on the period 2 overlay in the model in section 2 and the two curves have the same shape at different heights. Period 2's values are smaller for radius and larger for ionisation energy — its outer electrons are in the second shell, closer in and less shielded — but every feature is in the same place relative to the start of the period.
Why this matters more than it looks
Periodicity is the reason a table exists at all. Mendeleev left gaps for elements nobody had seen and predicted their densities and the formulae of their oxides, and he was right — because the properties of an element are set by where it sits, and position is set by the structure of the atom. That argument is what outcome 9.3 asks you to run for yourself in sections 21 and 22.
The elements with oxygen9.2.1
Topic 9.2 begins by putting each element in turn with the same three reagents, which is the cleanest way to see chemical periodicity: one variable — the element — and a fixed test. The syllabus names the products it wants, and it is worth learning that list exactly, because several of these elements form more than one oxide and only one of each is asked for.
The five oxides named in 9.2.1
Na2O, MgO, Al2O3, P4O10 and SO2. Sodium peroxide, phosphorus(III) oxide and sulfur trioxide are real compounds but are not what this outcome asks for — SO3 does appear later, in 9.2.2 and 9.2.3, because it is one of the oxides whose reactions with water and alkali you need.
| Element | Equation | Conditions and observations |
|---|---|---|
| Na | 4Na + O2 → 2Na2O | burns in air or oxygen with a strong orange flame; a white solid |
| Mg | 2Mg + O2 → 2MgO | burns with a brilliant white flame; a white ash |
| Al | 4Al + 3O2 → 2Al2O3 | the surface oxide layer protects the bulk metal; powder or shaken foil burns with white sparks |
| Si | Si + O2 → SiO2 | needs strong heating; a white solid |
| P | 4P + 5O2 → P4O10 | white phosphorus ignites spontaneously in air; dense white smoke |
| S | S + O2 → SO2 | burns on gentle heating with a pale blue flame; a colourless, choking gas |
Sulfur only reaches +4, and that is allowed
Burning sulfur in air gives SO2, not SO3: sulfur uses four of its six outer electrons. An element does not have to use all of them. SO3 is made industrially from SO2 and more oxygen over a vanadium(V) oxide catalyst — that is the Contact process, in topic 7 — and not by burning sulfur harder.
The elements with chlorine9.2.1
Chlorine is the second fixed test, and it is a cleaner one than oxygen because chlorine is always −1 in these compounds. That makes the formula of each chloride a direct readout of how many electrons the other element is using.
| Element | Equation | Conditions and observations |
|---|---|---|
| Na | 2Na + Cl2 → 2NaCl | heated sodium lowered into chlorine burns with an orange flame; white smoke |
| Mg | Mg + Cl2 → MgCl2 | heated magnesium burns with a brilliant white flame; a white solid |
| Al | 2Al + 3Cl2 → 2AlCl3 | heated in a stream of dry chlorine; a pale yellow solid that sublimes |
| Si | Si + 2Cl2 → SiCl4 | heated in a stream of dry chlorine; a colourless liquid that fumes in moist air |
| P | P4 + 10Cl2 → 4PCl5 | white phosphorus burns in chlorine; in excess chlorine the product is the (V) chloride, an off-white solid |
Sulfur is not on the list
Sulfur does react with chlorine, but the product is S2Cl2, and 9.2.1 names five chlorides only: NaCl, MgCl2, AlCl3, SiCl4 and PCl5. Phosphorus(III) chloride, formed in a limited supply of chlorine, is also outside the outcome.
Balancing these without thinking about it
Write the product first, from the oxidation numbers — aluminium is +3, chlorine −1, so AlCl3 — and only then balance. Chlorine and oxygen arrive as diatomic molecules, so if the product needs an odd number of them, double everything: AlCl3 needs 1½ Cl2, so the equation is written with 2AlCl3 and 3Cl2. The model below builds each equation that way and then counts the atoms on both sides to confirm it.
The elements with water9.2.1
The third test is the shortest, because the syllabus asks for sodium and magnesium only.
Sodium with cold water
2Na + 2H2O → 2NaOH + H2
Vigorous and immediate. The metal melts into a ball from the heat released, skates across the surface on the hydrogen it is producing, and leaves a strongly alkaline solution — pH 13–14, and universal indicator goes purple.
Magnesium with cold water, and with steam
Mg + 2H2O → Mg(OH)2 + H2 (cold water, very slow)
Mg + H2O → MgO + H2 (steam, vigorous)
With cold water a few bubbles appear over several days and the solution becomes faintly alkaline — Mg(OH)2 is only sparingly soluble, so it cannot get the pH much above 10. With steam the heated ribbon glows brightly and the product is the oxide, not the hydroxide.
Two different products, and the condition decides which
Cold water gives the hydroxide; steam gives the oxide. This is asked often enough that it is worth over-learning, and the giveaway in the question stem is the word steam or a drawing of a heated tube with damp mineral wool at one end.
Why the list stops at magnesium
Aluminium's oxide layer is tough, continuous and reforms instantly, so bulk aluminium appears unreactive with water even though the metal beneath is not. Silicon, phosphorus and sulfur are not attacked by water at all. So the reactivity with water falls away sharply after magnesium, which is itself a periodic trend: the ease of losing electrons decreases from left to right across the period, exactly as the ionisation energies said it would.
Oxidation numbers in the oxides and chlorides9.2.2
Line up the oxides and chlorides of period 3 and the oxidation number of the period 3 element climbs by one at every step. That is not a coincidence to be memorised; it is a direct count of the outer-shell electrons.
| Na | Mg | Al | Si | P | S | |
|---|---|---|---|---|---|---|
| outer-shell electrons | 1 | 2 | 3 | 4 | 5 | 6 |
| oxide | Na2O | MgO | Al2O3 | SiO2 | P4O10 | SO3 |
| oxidation number | +1 | +2 | +3 | +4 | +5 | +6 |
| chloride | NaCl | MgCl2 | AlCl3 | SiCl4 | PCl5 | — |
| oxidation number | +1 | +2 | +3 | +4 | +5 | — |
The statement 9.2.2 wants
The maximum oxidation number an element in period 3 reaches in its oxide or chloride is equal to the number of electrons in its outer shell, because that is the number of electrons it has available to lose or to share. Across the period the number of outer-shell electrons increases by one at each element, so the maximum oxidation number increases by one as well.
Ionic on the left, shared on the right — same count either way
In Na2O and MgO the electrons really are transferred, and the oxidation number is the charge on the ion. In SiO2 and SO3 nothing is transferred at all; the electrons are shared with a more electronegative atom, and the oxidation number is a bookkeeping device that counts how many of them have been pulled away. The count is the same in both cases, which is exactly why oxidation number is a useful idea for a whole period at once when "charge on the ion" only works for part of it.
SO2 and SO3
Sulfur has six outer electrons, so its maximum is +6, reached in SO3. In SO2 it is +4 — only four of the six are used. An element may use fewer than all of its outer electrons; it can never use more. 9.2.2 names both, and the pair is the syllabus's chosen illustration of exactly this point.
Trap: P4O10 looks like it breaks the pattern
The subscripts are large because the molecule is a real P4O10 cage rather than an empirical formula. Divide through: P2O5, so each phosphorus is balanced by 2½ oxygens at −2, giving +5. The oxidation number is +5, and it fits the sequence perfectly.
Bonding and electronegativity9.2.6
Outcome 9.2.6 asks you to explain every trend in 9.2.2 to 9.2.5 in terms of bonding and electronegativity. It is the hinge of the whole topic, and it rests on one number per element.
Electronegativity across period 3
Electronegativity is the ability of an atom to attract the bonding electrons in a covalent bond. Across period 3 it increases — Na 0.93, Mg 1.31, Al 1.61, Si 1.90, P 2.19, S 2.58, Cl 3.16 — for the same reason the atoms get smaller: rising nuclear charge, unchanged shielding, a shorter distance from the nucleus to the bonding pair.
The bonding in a compound is decided by the difference in electronegativity between the two elements. Oxygen is 3.44 and chlorine 3.16, both large. So as the period 3 partner becomes more electronegative, the difference narrows, and the bonding slides from ionic to covalent.
| Compound | Δχ with O (3.44) | bonding | Compound | Δχ with Cl (3.16) | bonding |
|---|---|---|---|---|---|
| Na2O | 2.51 | ionic | NaCl | 2.23 | ionic |
| MgO | 2.13 | ionic | MgCl2 | 1.85 | ionic, with covalent character |
| Al2O3 | 1.83 | ionic, with covalent character | AlCl3 | 1.55 | covalent, borderline |
| SiO2 | 1.54 | covalent (giant) | SiCl4 | 1.26 | covalent |
| P4O10 | 1.25 | covalent | PCl5 | 0.97 | covalent |
| SO2, SO3 | 0.86 | covalent | — | — | — |
One number, four outcomes
That column of differences is doing the work in the rest of this chapter. A large difference means ionic, which means a giant lattice, a high melting point, conduction when molten, a basic oxide, and a chloride that merely dissolves. A small difference means covalent, which means discrete molecules, a low melting point, no conduction, an acidic oxide, and a chloride that is hydrolysed. Everything in sections 15 to 19 is that sentence, spelled out.
The numbers are a guide, not a law
The usual A Level rule of thumb is that a difference above about 1.7–2.0 means ionic and below it covalent, and it is a rule of thumb: there is no sharp line in nature, which is exactly why aluminium's compounds are awkward. What matters in an answer is the direction — the difference falls across the period, so the bonding becomes more covalent — not the precise threshold. Quote a value only if the question gives you one.
The oxides with water9.2.3
Seven oxides, one reagent, and a pH that falls from 14 to nearly 0 across the period. This is the single most striking piece of chemical periodicity at AS, and it follows entirely from the bonding table in section 14.
| Oxide | Bonding | With water | Equation | pH |
|---|---|---|---|---|
| Na2O | ionic | reacts readily | Na2O + H2O → 2NaOH | 13–14 |
| MgO | ionic | reacts slightly | MgO + H2O → Mg(OH)2 | 9–10 |
| Al2O3 | ionic with covalent character | no reaction; insoluble | — | 7 |
| SiO2 | giant covalent | no reaction; insoluble | — | 7 |
| P4O10 | covalent, molecular | reacts violently | P4O10 + 6H2O → 4H3PO4 | 0–1 |
| SO2 | covalent, molecular | dissolves and reacts | SO2 + H2O → H2SO3 | 2–3 |
| SO3 | covalent, molecular | reacts violently | SO3 + H2O → H2SO4 | 0–1 |
Why the ionic oxides give alkaline solutions
An ionic oxide contains the oxide ion, O2−, and the oxide ion is a powerful base: it takes a proton from water.
O2− + H2O → 2OH−
Every oxide ion that gets into solution produces two hydroxide ions, which is why sodium oxide produces so strongly alkaline a solution from so little solid.
Why the covalent oxides give acidic solutions
A covalent oxide has no ions to give away. It reacts with water instead, forming an oxo-acid — a molecule with —OH groups attached to the central atom — and that acid then releases H+ into the solution. The more electronegative the central atom and the higher its oxidation number, the more it pulls electrons away from the O—H bonds and the more readily the proton leaves, so the acid gets stronger across the period.
Why MgO gives pH 9–10 and Na2O gives 13–14
Two reasons, and the better answer gives both. Solubility: Mg(OH)2 is only sparingly soluble, so very few hydroxide ions reach the solution — what does dissolve is fully ionised, but there is hardly any of it. Lattice strength: MgO is built of 2+ and 2− ions rather than 1+ and 2−, so the attractions in the lattice are much stronger and water prises the ions out only slowly. Both come back to the charge on the magnesium ion.
Trap: "Al2O3 and SiO2 are neutral oxides"
Wrong. Neither reacts with water, so neither changes the pH — but a neutral oxide is one that is neither acidic nor basic to anything, such as CO or NO. Aluminium oxide is amphoteric and silicon dioxide is acidic; both prove it by reacting with sodium hydroxide. "No reaction with water" and "neutral" are different statements, and only the first is true here.
Acid–base behaviour of the oxides9.2.4
Water is a weak test — two of the seven oxides do not respond to it at all. Acid and alkali are the sharper test, and under it every oxide declares itself.
| Oxide | Nature | With dilute HCl | With NaOH(aq) |
|---|---|---|---|
| Na2O | basic | Na2O + 2HCl → 2NaCl + H2O | no reaction |
| MgO | basic | MgO + 2HCl → MgCl2 + H2O | no reaction |
| Al2O3 | amphoteric | Al2O3 + 6HCl → 2AlCl3 + 3H2O | Al2O3 + 2NaOH + 3H2O → 2NaAl(OH)4 |
| SiO2 | acidic | no reaction | SiO2 + 2NaOH → Na2SiO3 + H2O |
| P4O10 | acidic | no reaction | P4O10 + 12NaOH → 4Na3PO4 + 6H2O |
| SO2 | acidic | no reaction | SO2 + 2NaOH → Na2SO3 + H2O |
| SO3 | acidic | no reaction | SO3 + 2NaOH → Na2SO4 + H2O |
The trend to state
Across period 3 the oxides change from basic (Na2O, MgO) through amphoteric (Al2O3) to acidic (SiO2, P4O10, SO2, SO3). The explanation is the bonding: the ionic oxides of the metals contain O2− and are basic; the covalent oxides of the non-metals have no O2− and instead form acids; aluminium oxide sits between the two and does both.
SiO2 needs hot concentrated alkali
It is acidic, but it is also a giant covalent lattice, so the reaction has to break covalent bonds throughout the crystal. In practice it needs hot, concentrated NaOH — or molten alkali. State the condition if you are asked for the reaction; it is the difference between silicon dioxide and a gas like SO2 that reacts with cold dilute alkali at once.
Amphoteric, and why it is aluminium
Aluminium sits at the boundary between the metals and the non-metals, and its electronegativity difference from oxygen — 1.83 — sits in the band between ionic and covalent. The oxide has enough ionic character to donate oxide ions to an acid and enough covalent character to accept hydroxide ions from an alkali. That is not a coincidence to be learned; it is what "boundary" means, and it is also why the elements on the metal/non-metal diagonal of the table (Be, Al, Zn, Pb) are the ones with amphoteric oxides.
The hydroxides9.2.4
Outcome 9.2.4 names three hydroxides alongside the oxides, and they behave exactly as their oxides do — which is the point of including them.
| Hydroxide | Solubility and appearance | Nature | Reactions |
|---|---|---|---|
| NaOH | very soluble; a strong alkali | basic | NaOH + HCl → NaCl + H2O |
| Mg(OH)2 | sparingly soluble; a white suspension, pH about 10 | basic | Mg(OH)2 + 2HCl → MgCl2 + 2H2O |
| Al(OH)3 | insoluble; a gelatinous white precipitate | amphoteric | Al(OH)3 + 3HCl → AlCl3 + 3H2O Al(OH)3 + NaOH → NaAl(OH)4 |
The practical test that identifies aluminium
Add sodium hydroxide solution drop by drop to a solution of an aluminium salt and a white gelatinous precipitate of Al(OH)3 appears. Keep adding, and it dissolves again to a colourless solution of the aluminate. Do the same with a magnesium salt and the precipitate of Mg(OH)2 stays put however much alkali you add. That "precipitate, then redissolves in excess" behaviour is the laboratory signature of an amphoteric hydroxide.
The aluminate formula to use
Write it as NaAl(OH)4, sodium tetrahydroxoaluminate. Some books write NaAlO2 with water on the other side, which is the same substance dehydrated; both appear in mark schemes, but NaAl(OH)4 is the one this syllabus uses. Whichever you pick, balance it.
The chlorides with water9.2.5
Five chlorides, one reagent, and three quite different things happen. Sorting out which is which is the most demanding part of topic 9, and it is entirely decided by the bonding and by the charge density of the cation.
| Chloride | Bonding | What happens in water | pH |
|---|---|---|---|
| NaCl | ionic | dissolves; the ions separate and nothing else happens | 7 |
| MgCl2 | ionic | dissolves; the hydrated Mg2+ ion is very weakly acidic | 6–6.5 |
| AlCl3 | covalent (borderline) | reacts vigorously; the hydrated Al3+ ion is appreciably acidic | 2–3 |
| SiCl4 | covalent | hydrolysed violently; fumes of HCl and a white solid | 1–2 |
| PCl5 | covalent | hydrolysed violently; fumes of HCl | 1–2 |
Route one: it just dissolves
NaCl(s) + aq → Na+(aq) + Cl−(aq)
Sodium chloride's ions separate and each is surrounded by water molecules. Na+ carries one unit of charge on a fairly large ion, so its pull on the water molecules attached to it is far too weak to disturb them. The solution is neutral.
Route two: the hydrated cation is an acid
AlCl3 + 6H2O → [Al(H2O)6]3+ + 3Cl−
[Al(H2O)6]3+ ⇌ [Al(H2O)5(OH)]2+ + H+
Why a metal ion can make a solution acidic
A small, highly charged cation pulls strongly on the lone pairs of the water molecules bonded to it. That pull is passed on: electrons are drawn out of the O—H bonds of those water molecules, weakening them, until one of them releases a proton to a free water molecule. The higher the charge density of the cation — charge divided by size — the further this goes and the lower the pH.
That is why the three metal chlorides run 7, 6–6.5, 2–3 and not all 7. Na+ is 1+ and large; Mg2+ is 2+ and smaller; Al3+ is 3+ and smaller still. The model plots the charge densities so you can see how steeply that rises.
Route three: the water attacks the central atom
SiCl4 + 2H2O → SiO2 + 4HCl
PCl5 + H2O → POCl3 + 2HCl (cold water)
PCl5 + 4H2O → H3PO4 + 5HCl (excess or hot water)
Hydrolysis
In a covalent chloride there is no ion to hydrate. Instead a water molecule uses a lone pair to attack the slightly positive central atom, and a chlorine leaves as HCl. The HCl dissolves in the water present, which is what makes the solution strongly acidic — the acid comes from the hydrogen chloride, not from a metal ion.
Why SiCl4 is hydrolysed and CCl4 is not
Both are tetrahedral covalent chlorides of a group 14 element, and one is destroyed by water while the other floats in it unchanged. The attacking water molecule needs somewhere to bond before a chlorine leaves. Silicon is a bigger atom, less crowded by its four chlorines, and has empty 3d orbitals of the right energy available to accept the lone pair. Carbon, in the second shell, has neither the room nor the available orbitals. This comparison is a favourite because it tests whether you know why hydrolysis happens rather than that it does.
Trap: "the chloride dissolves to give chloride ions, so the solution is acidic"
Wrong. Chloride ions are the conjugate base of a strong acid and are effectively neutral in water — sodium chloride proves it. The acidity in the other four cases comes from one of two places: a highly charged cation losing a proton from its water ligands, or a covalent chloride being hydrolysed to HCl. Name whichever applies.
Physical properties of the chlorides9.2.7
The same bonding trend that decides what happens in water also decides the melting points, and the two sets of evidence agree — which is what makes the deduction in the next section possible.
| Chloride | Melting point / K | Structure | Conducts when molten? |
|---|---|---|---|
| NaCl | 1074 | giant ionic lattice | yes |
| MgCl2 | 987 | giant ionic lattice, some covalent character | yes |
| AlCl3 | 451 (sublimes) | ionic solid that becomes molecular Al2Cl6 on heating | poorly |
| SiCl4 | 203 | simple molecular — a liquid at room temperature | no |
| PCl5 | 435 (sublimes) | solid of PCl4+ and PCl6− ions; molecular on heating | — |
Aluminium chloride is the awkward one, and it is awkward for a reason
It sublimes at 451 K, which is far too low for an ionic lattice and rather high for a small molecule; the solid has substantial ionic character and the vapour is a covalent Al2Cl6 dimer. If asked to classify it, say covalent with some ionic character, cite the low sublimation temperature and the violent hydrolysis, and note that Δχ = 1.55 puts it right on the boundary. Being unable to answer cleanly is the correct response to a compound that is genuinely on the line — but say why it is on the line.
PCl5: ionic where you would least expect it
Solid phosphorus(V) chloride is built of PCl4+ and PCl6− ions, which is why it is a solid at all when SiCl4 is a liquid — and it becomes ordinary covalent PCl5 molecules in the vapour. It is a reminder that "ionic" and "covalent" are descriptions of a structure in a particular state, not permanent labels attached to a formula.
Deducing the bonding from the evidence9.2.7
Outcome 9.2.7 turns the whole chapter round. Instead of being told a compound and asked what it does, you are given what it does and asked what it is — and the compound need not be one of the twelve on this page.
Three measurements, and what each one rules out
Melting point. High (roughly above 1000 K) means a giant structure; low means
simple molecular, because only weak forces between separate molecules are being
broken.
Conduction when molten. Conducting means ions that become mobile once the lattice
breaks up. Not conducting when molten rules ionic out, however high the melting point.
The solution in water. Dissolving to a neutral solution suggests an ionic compound of a
low-charge-density cation; reacting violently to an acidic solution suggests a covalent compound
being hydrolysed.
The first two together separate the three cases completely, which is why they are the two an examiner will give you.
| giant ionic | giant covalent | simple molecular | |
|---|---|---|---|
| melting point | high | very high | low |
| conducts as a solid | no | no | no |
| conducts when molten | yes | no | no |
| conducts in solution | yes | insoluble | only if it hydrolyses to an acid |
| in water | dissolves, usually near neutral | insoluble, no reaction | often hydrolysed; acidic |
| period 3 examples | NaCl, MgCl2, Na2O, MgO | SiO2 | SiCl4, PCl5, P4O10, SO2 |
Worked example: the classic three-line question
A white solid X melts at 1900 K. It does not conduct electricity as a solid or when molten, and it is insoluble in water. It dissolves slowly in hot concentrated sodium hydroxide. Deduce the type of bonding in X and suggest its identity.
Step 1. A melting point of 1900 K means a giant structure — far
too high for molecules held together by van der Waals forces.
Step 2. It does not conduct when molten, so there are no ions. That rules out giant ionic,
and the only giant structure left is giant covalent.
Step 3. Reacting with alkali but not with acid makes it an acidic oxide. A giant covalent
acidic oxide in period 3 is SiO2.
Trap: "high melting point, so it must be ionic"
The commonest wrong answer in this outcome, and the reason the "conducts when molten" line is always in the data. Giant covalent substances have the highest melting points of all — silicon dioxide beats sodium chloride comfortably — and contain no ions whatsoever. High melting point tells you giant. Only the conduction test tells you ionic.
Two more pieces of evidence worth recognising
Fumes in moist air. A covalent chloride being hydrolysed by the moisture, giving HCl. No
ionic chloride does this.
A white precipitate with silver nitrate after hydrolysis. Confirms that chloride ions are
now in the solution — useful for identifying the compound, but it says nothing about the bonding it
had before it met the water, because a hydrolysed covalent chloride produces chloride ions
too.
Predicting from a position9.3.1
Sub-topic 9.3 is short in the syllabus and long in the exam. It asks you to use what period 3 has taught you about why elements behave as they do, and apply it to elements you have not studied.
The chain of reasoning
Position → outer-shell electrons → the ion it forms and the oxidation numbers it reaches → the type of bonding in its compounds → their structures → their physical properties and their behaviour in water, acid and alkali.
Every link in that chain has been established somewhere on this page. 9.3.1 is the chain run forwards; 9.3.2 is the same chain run backwards.
Worked example: an element in group 2, period 5
Predict the properties of strontium and of its oxide and chloride.
Two outer electrons, so it is a metal forming Sr2+ and reaching +2. Metallic structure, so it conducts and has a moderately high melting point — though lower than magnesium's, because the ion is much larger and the metallic bonding correspondingly weaker. Further down the group than magnesium, so the outer electrons are further out and better shielded and therefore more easily lost: it reacts with cold water readily, giving hydrogen and an alkaline solution. Its oxide, SrO, is ionic and basic, reacting with water to give Sr(OH)2 and with acid to give a salt and water. Its chloride, SrCl2, is ionic, a high-melting solid that conducts when molten and dissolves to a roughly neutral solution — Sr2+ is large, so its charge density is too low for it to be appreciably acidic.
Not one of those statements required knowing anything about strontium. Every one came from the group number and a period 3 argument.
Two directions, two sets of reasons — keep them apart
Across a period: same shell, nuclear charge rising, shielding unchanged → atoms smaller,
ionisation energy and electronegativity rising, metals giving way to non-metals, oxides going from
basic to acidic.
Down a group: an extra shell each time, so although the nuclear charge rises steeply the
extra shielding nearly cancels it → atoms larger, ionisation energy and electronegativity falling,
metallic character increasing, the metals more reactive and the non-metals less.
Identifying an unknown element9.3.2
Outcome 9.3.2 gives you properties and asks for the element, or at least for its part of the table. These questions look like puzzles, and they are — but they are puzzles with a fixed set of clues, because only a handful of observations can be reported at AS.
| Observation | What it tells you |
|---|---|
| conducts as a solid, malleable | a metal, so left-hand side or the middle of the table |
| fizzes with cold water, alkaline solution left behind | group 1, or group 2 low down; the more vigorous, the further down |
| the oxide dissolves to give pH 13–14 | an ionic oxide of a reactive metal, so group 1 |
| the oxide dissolves in both acid and alkali | amphoteric — Al, or another element on the metal/non-metal diagonal |
| the oxide is insoluble and reacts only with hot concentrated alkali | an acidic giant covalent oxide, so group 14 at the top |
| the chloride fumes in air and hydrolyses to pH 1–2 | a covalent chloride, so a non-metal or a small highly charged metal ion |
| the chloride dissolves to a neutral solution | an ionic chloride of a large, low-charge cation — group 1, or group 2 low down |
| melts very high, does not conduct when molten | giant covalent — carbon, silicon, or an oxide of one of them |
| the highest oxide has the formula X2O5 or XO3 | count back: 5 or 6 outer electrons, so group 15 or group 16 |
The formula of the highest oxide is a group number in disguise
If an element's highest oxide is X2O5, then X is +5, so X has five outer-shell electrons and is in group 15. This is the quickest route into most 9.3.2 questions, and it works equally well from the chloride: XCl4 means four outer electrons, group 14.
Answer at the right level of precision
Most of these questions ask for the group or the kind of element, not for a name, because the evidence rarely pins down one element. "A group 2 metal — the vigour with cold water suggests calcium or below" is a better answer than a bare "calcium", and a far better answer than "a metal". Say what the evidence supports and no more, and say which observation each part of your answer rests on.
Self-test9.1, 9.2, 9.3
Thirty-three questions across the whole of topic 9. Each explains itself once you have answered, and the explanation is usually the part worth reading — several are built directly on the traps in the sections above.
Definitions and standard answers9.1, 9.2, 9.3
| Term | Definition |
|---|---|
| Periodicity | the regular, repeating variation in the properties of the elements with increasing proton number |
| Atomic radius | half the distance between the nuclei of two bonded atoms of the element (metallic radius for a metal, covalent radius for a non-metal) |
| Ionic radius | the radius of the ion the element forms, measured in a crystal of one of its compounds |
| Shielding | the reduction in the attraction an outer electron feels from the nucleus, caused by repulsion from the electrons in the shells between it and the nucleus |
| Effective nuclear charge | the net positive charge an outer electron experiences: the nuclear charge reduced by the shielding of the inner electrons |
| First ionisation energy | the energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions |
| Electronegativity | the ability of an atom to attract the bonding electrons in a covalent bond |
| Oxidation number | the charge an atom would carry if all the bonds in the species were treated as fully ionic |
| Basic oxide | an oxide that reacts with an acid to give a salt and water; the ionic oxides of metals |
| Acidic oxide | an oxide that reacts with a base to give a salt and water; the covalent oxides of non-metals |
| Amphoteric | able to react with both acids and bases, giving a salt and water in each case |
| Hydrolysis | the reaction of a compound with water in which the compound is broken up — here, a covalent chloride attacked by water to give HCl |
| Charge density | the charge on an ion in relation to its size; a small, highly charged cation has a high charge density |
| Giant covalent (macromolecular) | a structure in which covalent bonds continue throughout the crystal, so there are no separate molecules |
The five answers that come up again and again
Why does the atomic radius decrease across period 3? Nuclear charge increases, the extra electrons enter the same shell so shielding is essentially unchanged, effective nuclear charge increases, the outer shell is pulled in.
Why does the melting point rise from Na to Al? The ions have a higher charge and are smaller, and each contributes more delocalised electrons, so the attraction between the ion lattice and the electron sea is stronger.
Why is silicon's melting point so high? Giant covalent structure; melting requires strong covalent bonds to be broken throughout the lattice.
Why do the oxides change from basic to acidic? The bonding changes from ionic to covalent as the electronegativity difference with oxygen falls; ionic oxides contain O2− and are basic, covalent oxides form oxo-acids with water.
Why is the solution of AlCl3 acidic and that of NaCl neutral? Al3+ is small and highly charged, so the hydrated ion polarises its water ligands enough to release H+; Na+ has too low a charge density to do so.
Data used on this page9.1, 9.2
Read this before quoting a number from here in an exam
These are reference values assembled for this page, not the official Cambridge data booklet. Sources differ by a few units — melting points of sulfur are variously quoted as 386, 388 and 392 K, and electronegativities depend on which scale and which revision you use. Where an animation on this page shows a value, the table below uses the same one, so that nothing on the page contradicts anything else. In an examination, use the values printed on the paper.
The elements of period 3
| Na | Mg | Al | Si | P | S | Cl | Ar | |
|---|---|---|---|---|---|---|---|---|
| proton number | 11 | 12 | 13 | 14 | 15 | 16 | 17 | 18 |
| atomic radius / pm | 186 | 160 | 143 | 117 | 110 | 104 | 99 | 192* |
| ion formed | Na+ | Mg2+ | Al3+ | — | P3− | S2− | Cl− | — |
| ionic radius / pm | 102 | 72 | 54 | — | 212 | 184 | 181 | — |
| 1st ionisation energy / kJ mol−1 | 496 | 738 | 578 | 787 | 1012 | 1000 | 1251 | 1521 |
| electronegativity (Pauling) | 0.93 | 1.31 | 1.61 | 1.90 | 2.19 | 2.58 | 3.16 | — |
| melting point / K | 371 | 923 | 933 | 1687 | 317 | 388 | 172 | 84 |
| boiling point / K | 1163 | 1380 | 2740 | 3538 | 553 | 718 | 239 | 87 |
| conductivity / S m−1 | 2×107 | 2×107 | 4×107 | 10−3 | 10−9 | 10−15 | 10−16 | 10−17 |
| structure | metallic | metallic | metallic | giant covalent | P4 | S8 | Cl2 | atoms |
* Argon's 192 pm is a van der Waals radius and is not comparable with the metallic and covalent radii of the other seven. The conductivities of the non-metals are order-of-magnitude figures; they are quoted only to show the size of the collapse, and the qualitative statement — good, poor, none — is what topic 9 asks for.
The oxides
| Oxide | Ox. no. | Δχ with O | Structure | m.p. / K | With water | pH | Nature |
|---|---|---|---|---|---|---|---|
| Na2O | +1 | 2.51 | giant ionic | 1405 | reacts | 13–14 | basic |
| MgO | +2 | 2.13 | giant ionic | 3125 | slight | 9–10 | basic |
| Al2O3 | +3 | 1.83 | giant ionic, covalent character | 2345 | none | 7 | amphoteric |
| SiO2 | +4 | 1.54 | giant covalent | 1883 | none | 7 | acidic |
| P4O10 | +5 | 1.25 | simple molecular | 573 | violent | 0–1 | acidic |
| SO2 | +4 | 0.86 | simple molecular | 200 | dissolves | 2–3 | acidic |
| SO3 | +6 | 0.86 | simple molecular | 290 | violent | 0–1 | acidic |
The chlorides
| Chloride | Ox. no. | Δχ with Cl | Structure | m.p. / K | With water | pH |
|---|---|---|---|---|---|---|
| NaCl | +1 | 2.23 | giant ionic | 1074 | dissolves | 7 |
| MgCl2 | +2 | 1.85 | giant ionic | 987 | dissolves; weakly acidic | 6–6.5 |
| AlCl3 | +3 | 1.55 | ionic solid → molecular Al2Cl6 | 451 subl. | vigorous; acidic | 2–3 |
| SiCl4 | +4 | 1.26 | simple molecular | 203 | hydrolysed violently | 1–2 |
| PCl5 | +5 | 0.97 | PCl4+PCl6− → molecular | 435 subl. | hydrolysed violently | 1–2 |
Enthalpy changes of atomisation, for the metallic bonding argument
| Na | Mg | Al | |
|---|---|---|---|
| ΔHat / kJ mol−1 | +109 | +150 | +314 |
The enthalpy change of atomisation is the energy needed to turn one mole of the solid into one mole of separate gaseous atoms, so it measures the strength of the metallic bonding directly — and it triples from sodium to aluminium, which is the numerical form of the argument in section 6.