The Periodic Table: chemical periodicity

Periodicity: the trends in atomic and ionic radius across Period 3

Periodicity is the repeating pattern of physical and chemical properties that appears when the elements are placed in order of increasing proton number. Properties such as atomic radius, melting point and the acid/base character of the oxide recur at regular intervals, because the outer-shell electron configuration recurs at regular intervals. Across a period the outer shell is being filled, so properties change steadily in one direction; down a group a new shell is started with the same outer configuration, so the properties of the first element appear again, modified by the extra shells.

Atomic radius falls steadily from sodium to chlorine. Every element across Period 3 has the same number of inner shells, so the shielding of the outer electrons is much the same, but the nuclear charge increases by one proton at each step. The extra proton pulls the whole electron cloud inwards, and the electron being added enters the same outer shell, so it does not shield its neighbours from the extra charge. The outer shell is therefore drawn closer to the nucleus. Argon is a special case: it is monatomic and its radius is measured in a different way, so it is not compared directly with the metallic or covalent radii of the elements before it.

Ionic radii follow the same logic once you think about the electron configuration. A positive ion is always smaller than its parent atom, because the outer shell has been removed and the remaining electrons are held by a nucleus whose charge is shared between fewer electrons; Na+ is far smaller than Na. A negative ion is always larger than its parent atom, because the extra electrons increase the electron-electron repulsion and the cloud expands; Cl- is far larger than Cl.

Na+, Mg2+ and Al3+ are all 10-electron (neon-type) ions, and their radii fall in the order Na+ > Mg2+ > Al3+. The three ions have exactly the same electron configuration, so the only difference is the nuclear charge: 11 protons in Na+, 12 in Mg2+ and 13 in Al3+. The same ten electrons are pulled in more strongly by each extra proton, and the ion contracts. Comparing these three ions is the cleanest illustration of periodicity in ionic radius, because the shielding is identical.

• Across Period 3: nuclear charge rises, shielding is roughly constant, so atomic radius falls.
• Down a group: a new shell is added, so atomic radius rises.
• Cation smaller than its atom: Na+ < Na. Anion larger than its atom: Cl- > Cl.
• Na+ > Mg2+ > Al3+, all 10-electron ions, because the nuclear charge is 11, 12 and 13.

Melting point and electrical conductivity explained by structure and bonding

Melting point is decided by the strongest force holding the particles together, so it is decided by the structure and bonding of the element, not by its position in the row. Sodium, magnesium and aluminium are metals with giant metallic lattices; silicon is a giant covalent (macromolecular) structure; phosphorus, sulfur and chlorine are simple molecular and argon is monatomic. To melt the solid you must break the strongest attraction present, and that attraction is very different in each of these three structures.

From sodium to aluminium the melting point rises, because the metallic bond becomes stronger. Each metal atom contributes all of its outer-shell electrons to the lattice as delocalised electrons (one from sodium, two from magnesium, three from aluminium) and the positive ions carry a higher charge, so the electrostatic attraction between the ions and the sea of delocalised electrons increases. Sodium melts at about 98 C, magnesium at about 650 C and aluminium at about 660 C.

Silicon has by far the highest melting point in the period, about 1414 C. Silicon is a giant covalent structure in which every atom is joined to four others by strong covalent bonds, and melting it means breaking covalent bonds rather than separating whole molecules. The same idea explains its conductivity: silicon is a semiconductor, meaning it conducts only weakly, but its conductivity rises as the temperature rises because the energy of the lattice vibrations frees more electrons from their bonds. It is the only Period 3 element after the metals that conducts at all.

The remaining elements melt at very low temperatures, and their order follows the number of electrons per particle. Phosphorus (P4, about 44 C), sulfur (S8, about 115 C), chlorine (Cl2, about -101 C) and argon (about -189 C) are held together only by weak London (dispersion) forces between separate molecules or atoms, and these forces are far weaker than a covalent or metallic bond. The more electrons a particle contains, the more polarisable it is and the stronger the London forces, which is why S8 melts above P4 and why Cl2 and Ar, with few electrons, melt far below room temperature.

Reactions of the Period 3 elements with oxygen and with chlorine

With oxygen. Sodium burns in a limited supply of air with an orange flame to give the white oxide: 4Na + O2 -> 2Na2O. (In excess oxygen the peroxide Na2O2 can also form, but the syllabus oxide for sodium is Na2O.) Magnesium burns with a brilliant white light: 2Mg + O2 -> 2MgO. Aluminium burns to give the amphoteric oxide: 4Al + 3O2 -> 2Al2O3, and the same oxide forms as a tough, self-sealing layer whenever aluminium is left in air. Silicon reacts only at high temperature: Si + O2 -> SiO2. Phosphorus catches fire spontaneously in air, producing white clouds of phosphorus(V) oxide: P4 + 5O2 -> P4O10. Sulfur burns with a blue flame to give the dioxide: S + O2 -> SO2, and in excess air over a vanadium(V) oxide catalyst this is oxidised further to SO3. Argon does not react at all.

With chlorine. Sodium burns in chlorine with a bright yellow-orange flame: 2Na + Cl2 -> 2NaCl. Magnesium reacts on heating: Mg + Cl2 -> MgCl2. Aluminium gives white fumes of the chloride: 2Al + 3Cl2 -> 2AlCl3. Silicon reacts on strong heating: Si + 2Cl2 -> SiCl4. Phosphorus reacts vigorously, and with an excess of chlorine the pentachloride is formed: P4 + 10Cl2 -> 4PCl5 (in a limited supply of chlorine the trichloride PCl3 is obtained instead). Argon is unreactive, and the elements to its right in the period are not part of this pattern.

The trend behind both sets of reactions is electronegativity. Sodium has an electronegativity of about 0.9 and oxygen about 3.5, so the bonding in Na2O is strongly ionic; chlorine at about 3.0 gives an equally ionic chloride. As you move to the right, the element becomes more electronegative, the difference from oxygen or chlorine becomes smaller, and the bonding becomes increasingly covalent: SiO2 is giant covalent and SiCl4 is a simple molecular liquid. Moving right also makes the elements less willing to give up electrons, so the oxide becomes less basic, and magnesium and aluminium need heating before they will burn.

• Oxygen: 4Na + O2 -> 2Na2O, 2Mg + O2 -> 2MgO, 4Al + 3O2 -> 2Al2O3, Si + O2 -> SiO2, P4 + 5O2 -> P4O10, S + O2 -> SO2.
• Chlorine: 2Na + Cl2 -> 2NaCl, Mg + Cl2 -> MgCl2, 2Al + 3Cl2 -> 2AlCl3, Si + 2Cl2 -> SiCl4, P4 + 10Cl2 -> 4PCl5.
• The oxidation numbers in these products rise across the period: +1, +2, +3, +4, +5.
• The bonding changes from ionic to covalent as the electronegativity difference shrinks.

Sodium and magnesium with water

Sodium reacts vigorously with cold water: 2Na + 2H2O -> 2NaOH + H2. The metal skims across the surface as a molten ball, fizzing and often igniting with an orange flame, and the hydrogen is evolved fast enough to be lit with a splint. The heat that melts the sodium is the heat of the reaction itself. The solution left behind is strongly alkaline, pH about 13-14, because sodium hydroxide is a strong base and is fully ionised.

Magnesium reacts very slowly with cold water: Mg + 2H2O -> Mg(OH)2 + H2. Only a few bubbles of hydrogen are seen, because magnesium hydroxide is only sparingly soluble and the thin layer that forms coats the metal and blocks the water from reaching fresh magnesium underneath.

With steam, magnesium reacts much faster, and the product is the oxide rather than the hydroxide: Mg + H2O -> MgO + H2. The magnesium glows brightly and the hydrogen is collected over water. The change of product is not an accident: at steam temperature the magnesium oxide is formed directly instead of being hydrolysed by liquid water.

Why only sodium and magnesium? Reactivity with water depends on how easily the element loses its outer-shell electrons, and that ability falls sharply as the nuclear charge increases and the atom becomes smaller across the period. Aluminium is protected by its impervious oxide layer, and the elements from silicon onwards do not react with water under these conditions at all. Chlorine does react with water, but by disproportionation, not by forming hydrogen, so it is not part of this pattern.

Oxidation numbers of the oxides and chlorides from outer-shell electrons

The oxidation number of the Period 3 element in its oxide or chloride is simply the number of its outer-shell (valence) electrons that it has used in bonding, which is why it climbs smoothly across the period. Sodium has one outer electron, so sodium is +1 in Na2O and in NaCl. Magnesium has two, so it is +2 in MgO and MgCl2. Aluminium has three, so it is +3 in Al2O3 and AlCl3. Silicon has four, so it is +4 in SiO2 and SiCl4. Phosphorus has five, so it is +5 in P4O10 and PCl5. Each element gives away or shares every one of its outer-shell electrons, and no more.

Sulfur has two oxidation numbers in its oxides because it has six outer electrons and can use either four or all six of them. Sulfur is +4 in SO2 and +6 in SO3, and the same +4 to +6 change is what the Contact process achieves when SO2 is oxidised over a vanadium(V) oxide catalyst.

Working out an oxidation number is a two-step calculation, and the marks are for the working. Put the oxidation number of the oxygen as -2 (or the chlorine as -1), remember that a neutral molecule adds up to zero, and solve for the unknown. In P4O10: 4x + 10(-2) = 0, so 4x = 20 and x = +5. In SO3: x + 3(-2) = 0, so x = +6. In Al2O3: 2x + 3(-2) = 0, so 2x = 6 and x = +3.

The pattern is worth memorising as a check on your own answers. If you write an oxide or chloride of a Period 3 element with an oxidation number that is not the number of its outer-shell electrons, either the formula or the oxidation number is wrong. That is why P4O6 and AlO, both of which students sometimes write, cannot be the oxides asked for in this topic.

The oxides and hydroxides with water, and acid/base behaviour

Reactions with water. Na2O + H2O -> 2NaOH, a vigorous reaction giving a strongly alkaline solution of pH about 13-14. MgO + H2O -> Mg(OH)2, a slow reaction giving a sparingly soluble hydroxide and a pH of about 9-10. Al2O3 does not react with water at all, because its lattice is far too strong, so a mixture of aluminium oxide and water stays at pH 7. SiO2 is also insoluble in water and also leaves the pH at about 7. P4O10 reacts violently and exothermically: P4O10 + 6H2O -> 4H3PO4, giving a solution of pH about 1-2. Sulfur dioxide dissolves and reacts reversibly: SO2 + H2O <-> H2SO3, giving the weak acid sulfurous acid, pH about 2-3. Sulfur trioxide reacts so violently that it produces a corrosive mist: SO3 + H2O -> H2SO4, pH about 0-1.

Acid/base character across the period. A basic oxide reacts with an acid to form a salt and water, and its solution contains hydroxide ions. A acidic oxide reacts with a base to form a salt and water, and in water it produces hydrogen ions. Sodium oxide and magnesium oxide are basic; aluminium oxide is in between; silicon dioxide, phosphorus(V) oxide, sulfur dioxide and sulfur trioxide are acidic, and the acidity increases from SiO2 to SO3 as the element becomes more electronegative.

Amphoteric behaviour. Aluminium oxide and aluminium hydroxide are amphoteric, which means they react with both acids and bases. With hydrochloric acid: Al2O3 + 6HCl -> 2AlCl3 + 3H2O and Al(OH)3 + 3HCl -> AlCl3 + 3H2O. With sodium hydroxide: Al2O3 + 2NaOH + 3H2O -> 2NaAl(OH)4 and Al(OH)3 + NaOH -> NaAl(OH)4 (older books write these products as NaAlO2, sodium aluminate). Al(OH)3 is obtained as a white gelatinous precipitate when sodium hydroxide solution is added to a solution of an aluminium salt, and the precipitate redissolves in excess sodium hydroxide but not in excess ammonia solution, which is the standard test for an aluminium ion.

The hydroxides follow the same order. NaOH is a strong base, completely ionised in water, giving pH 13-14. Mg(OH)2 is a weak base and only sparingly soluble, so even a saturated solution is only mildly alkaline. Al(OH)3 is amphoteric, and its precipitate dissolves in excess sodium hydroxide solution but not in excess ammonia solution, which is the standard test for an aluminium ion. Since a base is a proton acceptor, the change from strong base through weak base to amphoteric mirrors the change from basic oxide through amphoteric oxide to acidic oxide. Silicon dioxide completes the pattern by showing that a weakly acidic oxide reacts with a base even when it does not react with water: hot concentrated sodium hydroxide solution dissolves it, SiO2 + 2NaOH -> Na2SiO3 + H2O, and at high temperature it displaces carbon dioxide from calcium carbonate, SiO2 + CaCO3 -> CaSiO3 + CO2, which is how sand is removed as slag in a blast furnace.

The chlorides with water: hydrolysis, pH, bonding and electronegativity

Sodium chloride dissolves to give a neutral solution, pH about 7, because the ions are simply hydrated by water molecules and neither Na+ nor Cl- reacts with water, so no hydrogen ions are produced. Magnesium chloride is only slightly acidic, pH about 6-7, because the small, doubly charged Mg2+ ion pulls electron density from the oxygen of a water molecule, which weakens an O-H bond and releases a few hydrogen ions.

Aluminium chloride hydrolyses vigorously: AlCl3 + 3H2O -> Al(OH)3 + 3HCl. It dissolves with a hiss, giving off fumes of hydrogen chloride, a white precipitate of aluminium hydroxide and a solution of pH about 3. Adding water to anhydrous AlCl3 is a violent reaction, and the fumes are the reason it is described as fuming in moist air.

Silicon tetrachloride and phosphorus pentachloride are completely hydrolysed. SiCl4 + 2H2O -> SiO2 + 4HCl (often written SiCl4 + 4H2O -> Si(OH)4 + 4HCl) gives a white gelatinous solid and an acidic solution of pH about 1-2, and PCl5 + 4H2O -> H3PO4 + 5HCl gives a strongly acidic solution of pH about 1 that contains phosphoric acid as well as hydrochloric acid.

The explanation is electronegativity and charge density. In every one of these chlorides the atom bonded to chlorine carries a partial positive charge, because chlorine is the more electronegative partner. A water molecule attacks that partially positive (delta-positive) atom, and the chloride is displaced as HCl; the more polarised the bond and the more positive the atom, the more vigorously the hydrolysis runs. Sodium is the least electronegative element in the period and Na+ has a low charge density, so there is no hydrolysis; Mg2+ is smaller and doubly charged, so hydrolysis is slight; Al3+ has a much higher charge density, and the delta-positive silicon and phosphorus atoms in SiCl4 and PCl5 are attacked so violently that the hydrolysis is effectively complete. The overall order is therefore NaCl (neutral) < MgCl2 (very weakly acidic) < AlCl3 (acidic) < SiCl4 and PCl5 (strongly acidic).

Deducing the bonding from the observations. NaCl and MgCl2 are ionic: they have high melting points, dissolve in water to give conducting solutions, and conduct when molten, because the ions become mobile. AlCl3 is best described as covalent: it sublimes at a low temperature (about 180 C), exists as Al2Cl6 molecules in the vapour, does not conduct in the solid or molten state, and reacts with water instead of simply dissolving in it. SiCl4 and PCl5 are simple molecular covalent compounds with very low melting points that do not conduct in any state. The rule to apply is simple: a high melting point plus conduction when molten or in solution means an ionic lattice; a low melting point plus no conduction in any state means molecules, held together only by weak London forces.

Using periodicity to predict a group, and to identify an unknown element

Periodicity lets you predict, not just recall. Because properties repeat down a group, knowing the pattern tells you what an element you have never met will do. Francium sits at the bottom of Group 1, below caesium, so you can predict that it will react with cold water even more violently than caesium, that it will form an oxide Fr2O, a chloride FrCl and a strongly basic hydroxide FrOH, and that it will be soft and low-melting. The predictions all come from the trend, not from the element itself.

Deducing an unknown element from data. Work through the evidence in a fixed order and write down what each piece tells you. The formula of the oxide or chloride gives the group: an element forming XO2 that is a non-metal belongs in Group 14, while an element forming XCl2 whose oxide gives an alkaline solution belongs in Group 2. The acid/base character of the oxide places the element left or right: basic to the left, amphoteric in the middle, acidic to the right of the period. The melting point and the conductivity tell you the structure and bonding. The reaction with water, if any, gives a further check, and the electron configuration or the position in a trend gives the period.

A worked deduction. An element E is a shiny, brittle solid that conducts electricity weakly, and its conductivity rises when it is heated. Its oxide EO2 has a very high melting point, does not dissolve in water and reacts with hot concentrated sodium hydroxide solution. Its chloride ECl4 is a low-boiling liquid that is hydrolysed violently by water to give an acidic solution and a white solid. Every observation matches silicon: the weak, temperature-dependent conductivity marks it as a semiconductor and therefore a giant covalent element, the formula EO2 places it in Group 14, the acidic oxide that is attacked by alkali places it on the non-metal side of the period, and the hydrolysed ECl4 confirms a small, highly charged atom. E is silicon, in Group 14 (Group IV) and Period 3.

Use the data you are given. In this sort of question the marks are for linking each observation to a deduction, so write the reason next to the result. Simply stating the name of the element, even when it is correct, will not score the marks that are awarded for the reasoning.

• Oxide or chloride formula -> group number.
• Acid/base character of the oxide -> position left, middle or right of the period.
• Melting point and conductivity -> ionic, giant covalent, metallic or simple molecular.
• Reaction with water -> how readily the element gives up its outer-shell electrons.
• Give a reason for every deduction, not just the name of the element.