Doc Brown's Chemistry  Revising Advanced Level Inorganic Chemistry Periodic Table Revision Notes

Part 5 Period 3 Na to Ar: 5.3 trends in bonding, formulae, oxidation states and reactions

The trends in the physical and chemical character of the elements is discussed first. The trends in bond type and formula of the compounds of the elements of Period 3 of the Periodic Table are tabulated and explained. The trends in valency and oxidation states of the elements in their compounds is also discussed and explanations provided on the basis of their electron configurations. The oxides, chlorides and hydrides of all of the Period 3 elements are tabulated with their formulae and type of bonding. The trends in the Period 3 elements chemical reaction with oxygen, water and chlorine are described and explanations provided with symbol equations, as are the reactions of the Period 3 oxides and chlorides with water and acids or alkalis. Finally, two series of isoelectronic species are tabulated in terms of their nuclear charge, ion charge/neutral atom and atomic or ionic radii.

INORGANIC Part 5 Period 5 survey, group trends page sub-index: 5.1 Period 3 survey of individual elements : 11. sodium : 12. Magnesium : 13. Aluminium : 14. Silicon : 15. Phosphorus : 16. Sulfur : 17. Chlorine : 18. Argon * 5.2 Period 3 element trends & explanations of physical properties * 5.3 Period 3 element trends in bonding, structure, oxidation state, formulae & reactions

Advanced Level Inorganic Chemistry Periodic Table Index * Part 1 Periodic Table history * Part 2 Electron configurations, spectroscopy, hydrogen spectrum, ionisation energies * Part 3 Period 1 survey H to He * Part 4 Period 2 survey Li to Ne * Part 5 Period 3 survey Na to Ar * Part 6 Period 4 survey K to Kr and important trends down a group * Part 7 s-block Groups 1/2 Alkali Metals/Alkaline Earth Metals * Part 8  p-block Groups 3/13 to 0/18 * Part 9 Group 7/17 The Halogens * Part 10 3d block elements & Transition Metal Series * Part 11 Group & Series data & periodicity plots * All 11 Parts have their own sub-indexes near the top of the pages

5. Survey of Period 3: Na across to Ar (8 elements, Z = 11 to 18)

5.3 Period 3 trends in bonding, structure, oxidation state, formulae & reactions

M+ X- ionic bond, Mδ+-Xδ+ polar bond and M-X a relatively non-polar bond (no partial charges shown)

Element Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
old/latest Group 1 2 3/13 4/14 5/15 6/16 7/17 0/18
ZSymbol 11Na 12Mg 13Al 14Si 15P 16S 17Cl 18Ar
Structure of element solid metallic lattice of Na+ and free e- solid metallic lattice of Mg2+ and free e-s solid metallic lattice of Al3+ and free e-s solid giant covalent lattice Sin solid small covalent molecules P4 solid small covalent molecules S8 gaseous small covalent molecules Cl2 gaseous single atoms Ar
electron configuration [Ne]3s1 [Ne]3s2 [Ne]3s23p1 [Ne]3s23p2 [Ne]3s23p3 [Ne]3s23p4 [Ne]3s23p5 [Ne]3s23p6
common oxidation states e.g. in oxides, chlorides, hydrides +1

max +1


max +2


max +3


max +4

+3, +5

max +5

-2, -2, +4, +6

max +6

-1, +1, +3, +5, +7

max +7

at Xe can get max of +8 in some compounds, but not for Ar!

electronegativity of element 0.93 1.31 1.61 1.90 2.19 2.58 3.16 3.20
formula of oxides

(oxidation states of the period 3 element)

Na2O, Na2O2








P4O6 and P4O10

(+3, +6)

SO2, SO3

(+4, +6)

Cl2O, ClO2, Cl2O7


Ratio of period 3 element to oxygen (formulae in bold) 1 : 0.5 1 : 1 1 : 1.5 1 : 2 1 : 2.5 1 : 2.5 - -
bonding and structure of oxides ionic lattice ionic lattice ionic lattice solid covalent giant structure solid covalent small molecules covalent small gaseous molecules covalent small gas/liquid molecules -
Melting point of oxide in highest oxidation state 1275oC 2852oC 2072oC 1723oC 580oC 17/62oC -91oC -
electronegativity difference X-O (O is 3.44) nature of bond 2.51

Na+ O2-or  O22-


Mg2+ O2-


Al3+ O2-









formula of chlorides NaCl MgCl2 AlCl3 SiCl4 PCl3, PCl5 S2Cl2, SCl2, SCl4 Cl2 -
bonding in chlorides ionic lattice ionic lattice ionic lattice, readily vaporises to covalent dimer molecules Al2Cl6 covalent small liquid molecules liquid covalent small molecules covalent small liquid molecules small diatomic gaseous molecule -
electronegativity difference X-Cl (Cl is 3.16) nature of bond 2.23

Na+ H-


Mg2+ Cl-


Al3+ Cl-









Formula of hydride NaH MgH2 AlH3 SiH4 PH3 H2S HCl -
bonding and structure of hydride ionic lattice 'polymer-like' structure of intermediate ionic/covalent nature 'polymer-like' structure of intermediate ionic/covalent nature small covalent gaseous molecule small covalent gaseous molecule small covalent gaseous molecule small covalent gaseous molecule -
electronegativity difference X-H (H is 2.20) nature of bond 1.27

Na+ H-













  • The structure and physical properties of the elements (see also section 5.1)

    • The trend is metal lattice ==> giant covalent structure ==> small covalent molecules

    • Sodium Na, magnesium Mg and aluminium Al are silvery solids, with a metal lattice structure, high boiling points and are good conductors of heat/electricity due to the delocalised free electrons moving between the immobile metal ions.

      • The melting/boiling points increase from Na  ==> Mg ==> Al due to 1 ==> 2 ==> 3 potential number of delocalised electrons that may contribute to bonding.

    • Si has a non-metallic giant covalent structure based on a tetrahedral arrangement of S-Si bonds and is a poor conductor of heat/electricity.

      • The strong 3D bonding gives silicon a high melting/boiling point and great hardness.

    • Phosphorus P4, sulfur S8 and chlorine Cl2 are simple-small covalent molecules and Ar consists of single atoms. The molecules are only held together by the weakest of the intermolecular forces, namely the instantaneous dipole - induced dipole forces, and consequently have very low melting/boiling points.

      • From left to right the elements become less metallic and more non-metallic.

    • -

  • Electron configuration and oxidation states

    • Electron configurations of  2,8,1 or 1s22s22p63s1 to 2,8,8 or 1s22s22p63s23p6  

      • Filling the s orbital (max 2 e-'s) gives the metallic s-block elements of Groups 1-2,

      • filling the p orbitals gives the predominantly non-metallic p block elements of Group 3-7, 0 (Gps 13-18) bar aluminium for Period 3.

    • Oxidation states in compounds (numerically = valency) are: sodium Na (+1 only), magnesium Mg (+2 only), aluminium Al (+3 only), Si (+4, -4 with electropositive metals), P (usually -3, +3 or +5), S (-2, +4 and +4), Cl (-1, +1, +3, +5 and +7), Ar has no stable compounds due to the full outer quantum level (shell) being full, conferring extra electronic stability on the atom.

      • From Na to Cl the maximum oxidation state is equal to the 'old' group number and the 'highest' oxide formulae can be predicted up to chlorine and the chloride formula up to P (there is no stable SCl6 but there is a ClF7.

        • So in the 'highest' oxides you can go from +1 to +7 for groups 1 to 7/17

          • (at Xe on Period 5 you can reach +8, but not for Ar)

        • Na2O, MgO, Al2O3, SiO2, P4O10 (= P2O5), SO3, Cl2O7 (at Xe you can have XeO4)

        • i.e. using all available 1-7 outer 3s and 3p electrons (valence electrons) are all used in the bonding of the highest possible oxide.

        • and similarly in the 'highest' chlorides (up to Group 5), and fluorides (for Groups 6 and 7) you also go from maximum oxidation state of +1 to +7 in the halide compounds irrespective of bond character.

        • NaCl, MgCl2, AlCl3, SiCl4, PCl5, SF6, ClF7 (at Xe you can have XeF8)

    • -

  • Reaction of element with oxygen and the structure of the oxide (see also section 5.1)

(Gp 1) 4Na(s) + O2(g) ==> 2Na2O(s) and Na2O2 on heating the metal in air (Gp 2) 2Mg(s) + O2(g) ==> 2MgO(s) on heating metal in air
(Gp 3) 4Al(s) + 3O2(g) ==> 2Al2O3(s) needs high temperature (Gp 4) Si(s) + O2(g) ==> SiO2(g) needs high temperature
(Gp 5) P4(s) + 5O2(g) ==> P4O10(s) on heating in air (Gp 6) S(s) + O2(g) ==> SO2(g) and a little SO3 on heating in air
(Gp 7) Chlorine - no reaction (Gp 0) Argon - no reaction
  • Reaction with oxygen and oxide structure (see also section 5.1)

    • The metals Na, Mg and Al burn to form a giant ionic oxide lattices

      • sodium oxide/peroxide, magnesium oxide and aluminium oxide ...

      • (Na+)2O2- and (Na+)2O22-, Mg2+O2- and (Al3+)2(O2-)3 respectively.

    • Silicon Si forms a giant covalent lattice of (SiO2)n where n is very larger number

    • Phosphorus P forms two simple molecular covalent solid oxides P4O6 and P4O10.

    • Sulfur S can form two simple molecular covalent gas molecules SO2 and SO3

    • Chlorine Cl forms oxide molecules of Cl2O, Cl2O7 (and others).

    • Argon has no reaction.

    • The overall pattern, from left to right is

      • giant ionic lattice => giant covalent lattice ==> small covalent molecules.

      • The first three elements (Na, Mg, Al) all have high melting points, as you would expect from strong stable giant ionic lattices.

      • Silicon also has a very high melting point, but this is due a vey strong 3D giant covalent bond network in crystalline silicon (giant covalent lattice).

      • However, from phosphorus to chlorine the oxides have much lower and decreasing melting points as they relatively small covalent molecules with much weaker intermolecular forces attracting the molecules together in the liquid or solid.

        • Intermolecular bonding is much weaker than covalent or ionic bonding, so the particles need much less kinetic energy to overcome the attractive bonding forces.

    • The change in bonding character from ionic to covalent in the oxide, follows the decreasing difference in electronegativity between that of the period 3 element and oxygen.

    • More details on the oxides including the oxide of the element in its highest oxidation state (see table):

      1. sodium oxide - a giant ionic lattice of sodium ions (Na+) and oxide ions (O2-). The strong electrostatic forces give the structure a high melting point.

      2. magnesium oxide - a giant ionic lattice of magnesium ions (Mg2+) and oxide ions (O2-). Again, the strong electrostatic forces give the structure a high melting point. The melting point is higher for magnesium oxide than sodium oxide because of the greater charge on the magnesium ion.

      3. aluminium oxide - essentially an ionic lattice of aluminium ions (Al3+), but despite the higher charge of the aluminium ion, although the melting point is very high, it is not as high as magnesium oxide. The aluminium ion is highly polarising and distorts the electron clouds of the oxide ions towards it, thereby producing some covalent character and a lowering of the melting point though lower than the preceding magnesium and aluminium oxides.

      4. (c) doc b silicon(IV) oxide (silica) - this is not an ionic lattice, but a giant covalent structure in which the giant lattice is held together by a 3D network of strong S-O covalent bonds. This gives silica great thermal stability and high melting point

      5. phosphorus(V) oxide, Mr(P4O10) = 284 - is relatively simple covalent molecule with weak intermolecular forces (Van der Waals forces) attracting the molecules together, hence it has a much lower melting point than the four preceding giant structures (1. - 4.).

      6. sulfur(VI) oxide -  Mr{(SO3)3} = 240 - again another relatively simple covalent molecule with weak intermolecular forces (Van der Waals forces) attracting the molecules together, hence it has a much lower melting point than the giant structures (1. - 4.). With a smaller relative molecular mass, its melting point is lower than P4O10.

        • You don't need to know this but, out of idle curiosity reading ...!

        • In the gaseous phase sulfur(VI) oxide is a trigonal planar 'monomer' (SO3) molecule, but if very pure sulfur(VI) oxide is condensed below 27oC. it forms crystals of the trimer molecule that melt at 17oC (unstable gamma form). Above 27oC, condensation of SO3 vapour produces fibrous stable crystals of the alpha form, which has a polymeric and more stable structure.

      7. chlorine(VII) oxide - Mr(Cl2O7) = 183 - again another relatively simple covalent molecule with weak intermolecular forces (Van der Waals forces) attracting the molecules together, hence it has a much lower melting point than the giant structures (1. - 4.). With an even smaller relative molecular mass, the intermolecular forces are further reduced and so its melting point is lower than P4O10 or (SO3)3.

      8. Argon cannot form a stable oxide

  • Reaction of the oxides with water, acids and alkalis (see also section 5.1)

(Gp 1) Na2O(s) + H2O(l) ==> 2NaOH(aq)

 pH 13-14 strong base  from ionic oxide

ionic equation:

 (Na+)2O(s) + H2O(l) ==> 2Na+(aq) + 2OH-(aq)

or Na2O2(s) + 2H2O(l) ==> 2NaOH(aq) + H2O2(aq)

(Gp 2) MgO(s) + H2O(l) ==> Mg(OH)2(aq/s)

 ~pH 11-12 weak base  from ionic oxide

ionic equation:

Mg2+(s) + H2O(l) ==> Mg2+(aq) + 2OH-(aq)

(Gp 3) Al2O3, insoluble, no reaction with water (pH remains at 7), but amphoteric with respect to strong acids and strong bases (alkalis) (Gp 4) SiO2, insoluble, no reaction with water (pH remains at 7), but weakly acidic and will dissolve a little in strong bases (alkalis) e.g. conc. NaOH(aq)
(Gp 5) P4O6(s) + 6H2O(l) ==> 4H3PO3(aq)

  ~pH 2 weak acid  from covalent oxide

P4O10(s) + 6H2O(l) ==> 4H3PO4(aq)

 pH 0-1 strong acid  from covalent oxide

Because phosphoric(V) acid is tribasic, there are three possible 'phosphate' anions

H2PO4-HPO42-  and  PO43-

(Gp 6) SO2(aq) + H2O(l) ===> H2SO3(aq)

sulfurous acid, theoretically dibasic, but the main ionic reaction is

SO2(aq) + H2O(l) H+(aq) + HSO3-(aq)

  pH 2-3 weak acid  from covalent oxide

There are two possible anions

HSO3- hydrogensulfite or hydrogensulfate(IV) ion

and SO32- sulfite or sulfate(IV) ion

SO3(g) + H2O(l) ==> H2SO4(aq)

  pH 0-1 strong acid  from covalent oxide

Sulfuric(VI) is dibasic so there are two possible anions

HSO4- hydrogensulfate or hydrogensulfate(VI) ion

SO42- sulfate or sulfate(VI) ion

(Gp 7) Cl2O(g) + H2O(l) ==> 2HClO(aq)

  ~pH 3? weak acid  from covalent oxide

Cl2O7(l) + H2O(l) ==> 2HClO4(aq)

  pH 1 strong acid  from covalent oxide

(Gp 0) argon has no oxide
  • The chemical character of the oxides - reaction of the Period 3 oxides with water, acids or alkalis.

    • Relating bonding character to acid - base character of the oxide.

      • The ionically bonded oxides of sodium and magnesium form bases - alkaline solutions in water.

        • The large difference in electronegativity (see table) allows the ionic bond and hence the oxide ion to exist.

        • The oxide ion is a strong Bronsted-Lowry base i.e. a strong attractor of protons and the oxide ion abstracts a proton from a water molecule to give the hydroxide ion, hence an alkaline solution is formed.

        • O2-(s) + H2O(l) ==> 2OH-(aq)

      • Aluminium oxide, ionic with some covalent character, is amphoteric.

      • The covalent oxides of the non-metals, Si, P, S and Cl, all form acids (weak or strong).

      • So the trend in oxide character across period 3 is from strongly basic oxides of metals to strongly acidic oxides of non-metals.

    • Sodium oxide/peroxide Na2O/Na2O2 and magnesium oxide MgO are basic and form an alkali in water (see above) and salts with acids (examples below).

      • Na2O(s) + 2HCl(aq) ==> 2NaCl(aq) + H2O(l)

      • Na2O(s) + H2SO4(aq) ==> Na2SO4(aq) + H2O(l)

      • MgO(s) + 2HCl(aq) ==> MgCl2(aq) + H2O(l)

      • MgO(s) + 2HNO3(aq) ==> Mg(NO3)2(aq) + H2O(l)

      • See reactions of acids for lots more examples with sulfuric acid and nitric acid too with oxides and hydroxides.

      • The equations with phosphoric(V) acid are the trickiest to balance and the examples below assume total neutralisation of the tribasic phosphoric(V) acid!

      • 3Na2O(s) + 2H3PO4(aq) ==> 2Na3PO4(aq) + 3H2O(l)

      • 3MgO(s)  + 2H3PO4(aq) ==>  Mg3(PO4)2(aq) + 3H2O(l)

    • Aluminium oxide Al2O3 has no reaction with water, insoluble, but is amphoteric and forms salts with (i) acids and (ii) alkalis.

      • (i) Al2O3(s) + 6HCl(aq) ==> 2AlCl3(aq) + 3H2O(l)

      • (ii) Al2O3(s) + 2NaOH(aq) + 3H2O(l) ==> 2Na[Al(OH)4](aq) 

    • Silicon(IV) oxide (silicon dioxide) SiO2 has no reaction but is weakly acidic forming salts with alkalis.

      • Silicon dioxide will slowly dissolve in hot concentrated sodium hydroxide (strong base)

      • SiO2(s) + 2NaOH(aq) ==> Na2SiO3(aq) + H2O(l)

    • Phosphorus(III) oxide P4O6 and phosphorus(V) oxide P4O10 are moderately-strong acidic oxides forming phosphoric(III) acid H3PO3 and phosphoric(V) acid H3PO4 on reaction with water.

      • Therefore the oxides are acidic and form salts with bases.

      • You can formally write an equation for dissolving phosphorus(V) oxide in sodium hydroxide to form trisodium phosphate(V), but its not as simple as this!

        • P4O10(s) + 12NaOH(aq) ==> 4Na3PO4(aq) + 6H2O(l)

      • Generally speaking, in a series of oxides for the same element, the higher the oxidation state of X in a 'XxOy' series, the more acidic is the oxide, so H3PO4 is a stronger acid than H3PO3.

      • The oxides or acids are readily neutralised to give phosphate salts e.g.

      • H3PO4 (aq) + NaOH(aq) ==> NaH2PO4(aq) + H2O(l)

      • Two further reactions are possible with the sodium hydroxide to give Na2HPO4 and Na3PO4.

    • Sulphur dioxide will dissolve in alkalis to form sulfate(IV) salts (sulfites) e.g.

      • SO2(g) + 2NaOH(aq) ==> Na2SO3(aq) + H2O(l)

      • Although weakly acidic, it will displace weaker acidic oxides e.g.

      • SO2(g) + CaCO3(s) ==> CaSO3(s) + CO2(g)

      • This reaction forms part of the chemistry of desulfurization of flue gases from fossil fuel power station furnaces.

      • Sulfur in fossil fuels burns to form sulfur dioxide - a major air pollutant.

      • The flue gases are 'scrubbed' mixing them with air and passing the gas mixture through a slurry of wet limestone powder to form harmless calcium sulfate (which can be used as the commercial product gypsum).

      • SO2(g) + CaCO3(s) + 0.5O2(g) ==> CaSO4(aq/s) + CO2(g)

      • The gypsum is formed as calcium sulfate dihydrate CaSO4.2H2O(s)

    • Chlorine(I) oxide Cl2O and chlorine(VII) oxide Cl2O7 are moderate to strong acidic in water.

    • The overall patterns, from left to right across Period 3 is ...

    • giant ionic lattice ==> giant covalent lattice ==> small covalent molecules

      • The change in bonding character from ionic to covalent in the oxide follows the decreasing difference in electronegativity between that of the element and oxygen.

    • In terms of overall chemical character ...

      • metal basic oxides ==> amphoteric oxides ==> non-metal oxides

    • This is chemically characteristic of metallic ==> non-metallic element character.

    • -

  • Reaction of element with chlorine and the structure of the chloride (see also section 5.1)

(Gp 1) 2Na(s) + Cl2(g) ==> 2NaCl(s)  (Gp 2) Mg(s) + Cl2(g) ==> MgCl2(s) 
(Gp 3) 2Al(s) + 3Cl2(g) ==> 2AlCl3(s) (Gp 4) Si(s) + 2Cl2(g) ==> SiCl4(l)
(Gp 5) P4(s) + 3Cl2(g) ==> 4PCl3(l)

P4(s) + 5Cl2(g) ==> 4PCl5(s) 

(Gp 6) 2S(s) + Cl2(g) ==> S2Cl2(l)  also unstable SiCl2, SiCl4
(Gp 7) chlorine itself (Gp 0) no reaction with argon
  • Reaction with chlorine and chloride structure

    • All of Na to S will combine directly on heating in chlorine to give the chloride.

    • Sodium, magnesium and aluminium give giant ionic lattices

      • sodium chloride Na+Cl-, magnesium chloride Mg2+(Cl-)2 and aluminium chloride Al3+(Cl-)3 respectively.

      • Note that aluminium chloride on heating sublimes above 180oC to form small Al2Cl6 covalent dimer molecules.

    • The non-metal elements give covalent chlorides.

    • Silicon forms the molecular covalent liquid silicon(IV) chloride SiCl4 (silicon tetrachloride)

    • Phosphorus forms phosphorus(III) chloride PCl3 (phosphorus trichloride) with limited chlorine

      • and phosphorus(V) chloride PCl5 (phosphorus pentachloride) with excess chlorine.

    • Sulfur gives disulfur dichloride S2Cl2. by direct combination (and unstable SCl2 and SCl4 can also be formed).

    • There is no stable argon chloride.

    • The overall pattern, from left to right across period 3 is ...

      • giant ionic lattice => polymeric covalent lattice ==> small covalent molecules.

      • This is chemically characteristic of metallic ==> non-metallic element character.

      • The change in bonding character from ionic to covalent in the chloride, follows the decreasing difference in electronegativity between that of the element and oxygen, as in the case of oxides.

      • The formulae largely follow a pattern of rising formulae based on the use of all outer electrons in bonding (1-5) and then a decline in valency (oxidation state of the Period 3 element).

      • NaCl (+1), MgCl2 (+2), AlCl3 (+3), SiCl4 (+4), PCl5 (+5), S2Cl2 (+1), Cl2, Ar no chloride

      • So the number of atoms of chlorine combined with the Period 3 element (the valency) follows the pattern

        • 1  2  3  4  5  1  1  0

    • -

  • Reaction of the chlorides with water (see also section 5.1)

(Gp 1) NaCl(s) + aq ==> Na+(aq) + Cl-(aq)

 just dissolves, ~pH 7

(Gp 2) MgCl2(s) + aq ==> Mg2+(aq) + 2Cl-(aq)

 just dissolves, ~pH 7

(Gp 3) AlCl3(s) + 3H2O(l) ==> Al(OH)3(s) + 3HCl(g)

 with limited water you get hydrolysis to give acid fumes 

AlCl3(s) + aq ==> Al3+(aq) + 3Cl-(aq)

 excess water, weakly acidic solution due to the acidity of [Al(H2O)6]3+

(Gp 4) SiCl4(l) + 2H2O(l) ==> SiO2(s) + 4HCl(aq)

 hydrolysis to give strongly acid solution

(Gp 5) PCl3(l) + 3H2O(l) ==> H3PO3(aq) + 3HCl(aq)

 hydrolysis to give weakly acid solution

PCl5(s) + 4H2O(l) ==> H3PO4(aq) + 5HCl(aq) 

 hydrolysis to give strongly acid solution

(Gp 6) S2Cl2(g) + H2O(l) ==> HCl(aq), S(s), SO2(aq), H2SO3(aq),  H2SO4(aq), H2S(aq)

 complex redox - hydrolysis reaction but final solution is quite acidic

(Gp 7) chlorine itself Gp 0 argon has no chloride
  • Reaction of the chloride with water

    • The ionic sodium chloride NaCl and magnesium chloride MgCl2 dissolve in water to form a nearly neutral solution of hydrated ions.

    • The ionic AlCl3 and the covalent all hydrolyse to form acid solutions.

      • aluminium chloride Al2Cl6 ==> hydrochloric acid or weakly acidic aluminium ion

      • silicon(IV) chloride SiCl4, (silicon tetrachloride) ==> hydrated silicon dioxide + hydrochloric acid

      • phosphorus(III) chloride PCl3 (phosphorus trichloride) ==> phosphoric(III) acid + hydrochloric acid

      • with phosphorus(V) PCl5 (phosphorus pentachloride) ==> phosphoric(V) acid + hydrochloric acid

        • Phosphorus(III) chloride hydrolyses rapidly and exothermically to form phosphoric(III) acid.

          • PCl3(l) + 3H2O(l) ==> H3PO3(aq) + 3HCl(aq)

        • Phosphorus(V) chloride initially hydrolyses to form phosphorus oxychloride and hydrochloric acid.

          • (i) PCl5(s) + H2O(l) ==> POCl3(aq) + 2HCl(aq) 

          • If the aqueous solution is boiled, phosphoric(V) acid is formed and more hydrochloric acid.

          • (ii) POCl3(aq) + 3H2O(l) ==> H3PO4(aq) + 3HCl(aq) 

          • overall (i) + (ii): PCl5(s) + 4H2O(l) ==> H3PO4(aq) + 5HCl(aq) 

      • and disulfur dichloride S2Cl2 ==> a variety products including acidic sulfur dioxide and hydrochloric acid.

    • The general trend is for ionic metal chloride salts to give nearly neutral solutions => metal/non-metal covalent chlorides that hydrolyse to give acidic solutions.

    • -

  • Reaction of element with water

(Gp 1) 2Na(s) + 2H2O(l) ==> 2NaOH(aq) + H2(g)  (Gp 2) Mg(s) + 2H2O(l) ==> Mg(OH)2(aq) + H2(g) 
(Gp 3) aluminium has no reaction with water (Gp 4) silicon has no reaction with water
(Gp 5) phosphorus has no reaction with water (Gp 6) sulfur has no reaction with water
(Gp 7) Cl2(g) + H2O(l) HClO(aq) + HCl(aq)  (Gp 0) argon has no reaction with water
  • Reaction of element with water

    • The reactive metal sodium Na rapidly gives the alkaline sodium hydroxide and hydrogen,

    • as does magnesium Mg BUT much more slowly.

    • Aluminium Al, silicon Si, phosphorus P and sulfur S have no reaction with water.

    • Chlorine, Cl2 forms a weakly acidic solution in water.

    • Argon has no reaction.

    • The 'limited' pattern for period 3 (or any other period), is to have reactive metals on the left forming an alkaline solution and a reactive non-metal on the right forming an acid solutions IF they react with water.

    • -

  • The hydrides MHx

    • For hydrides the difference in electronegativity works both ways!

    • From left to right across the period you change from an

      • ionic sodium hydride crystal lattice Na+H-

      • to small non-polar molecule covalent hydrides (silane SiH4 and phosphine PH3)

      • and then a weakly acidic polar covalent hydride molecule (hydrogen sulfide H2S)

      • and finally a strongly acidic polar covalent molecule (hydrogen chloride HCl).

    • The formulae follow a simple period pattern of rising and falling valency for the Period 3 elements.

      • NaH  MgH2  AlH3  SiH4  PH3 H2S  HCl (Ar)

      • the element valency pattern being 1  2  3  4  3  2  1  0

    • On reaction with water, the ionic metal hydrides at the start of the period give an alkaline solution

      • e.g. NaH(s) + H2O(l) ==> NaOH(aq) + H2(g)

    • In the middle are neutral hydrides like phosphine which in contact with water do not change the pH.

    • Then you get weakly acidic ==> strongly acidic hydrides when they dissolve in water e.g.

      • weak acid: H2S(aq) + H2O(l) <=> H3O+(aq) + HS-(aq)

      • strong acid: HCl(aq) + H2O(l) ==> H3O+(aq) + Cl-(aq)

    • So things are a bit complicated with hydrides on period 3 due to the left and right sided differences in electronegativity!

      • You go from  X+H- ==> Xδ+-Hδ  ==> X-H ==> Xδ-Hδ+

    • -

  • Radii of isoelectronic ions

    • Isoelectronic means species having the same total number of electrons.

    • The table below considers the isoelectronic cations and anions associated with Periods 2, 3 and 4.

  • isoelectronic system Group 4/14 Group 5/15 Group 6/16 Group 7/17 (Group 0/18) Group 1 Group 2 Group 3/13
    Period Period 2 Period 3
    [Ne] 10e 1s22s22p6 C4- N3- O2- F- (Ne) Na+ Mg2+ Al3+
    total nuclear charge +6 +7 +8 +9 (+10) +11 +12 +13
    radius in picometre (pm) 260 171 140 136 (38-112*) 95 65 50
    name of ion carbide nitride oxide fluoride (neon) sodium magnesium aluminium
    Period Period 3 Period 4
    [Ar] 18e 1s22s22p63s23p6 Si4- P3- S2- Cl- (Ar) K+ Ca2+ Sc3+
    nuclear charge +14 +15 +16 +17 (+18) +19 +20 +21
    radius  in picometre (pm) 271 212 184 181 (71-154*) 133 99 81
    name of ion silicide phosphide sulfide chloride (argon) potassium calcium scandium
    • Excluding the noble gases themselves where this is frankly, something of a data problem!,

      • there is a clear pattern of decreasing ionic radius with increase in total nuclear charge (+ atomic/proton number) for the two isoelectronic series tabulated above,

        •  based on the electron configurations of neon and argon.

    • From left to right the proton/electron ratio is steadily increasing so that the electrons are experiencing an increasingly greater attractive force of the nucleus, hence the steady decrease in radii for an isoelectronic series.

    • * all sorts of values are quoted for noble gas radii e.g. atomic or covalent and ionic, the higher values fit into the pattern above which is quite clear for all the cations and anions listed.

See also 4.1 Period 2 Survey of the individual elements, 4.2 Period 2 element trends and explanations of physical properties * 4.3 Period 2 element trends in bonding, structure, oxidation state, formulae & reactions, 5.1 Period 3 survey of individual elements, 5.2 Period 3 element trends & explanations of physical properties, 6.1 Survey of Period 4 elements, 6.2 Period 4 element trends in physical properties, 6.3 Period 4 element trends in bonding, formulae and oxidation state and 6.4 Important element trends down a Group

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