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

ON THIS PAGE: 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.

Sub-index for this page

See also on other pages

5.1 Period 3 survey of individual elements

5.2 Period 3 element trends & explanations of physical properties

6.4 Important element trends down a Group

Revision notes for Advanced A level inorganic chemistry suitable for AQA Advanced A level chemistry Edexcel Advanced A level chemistry OCR Advanced A level chemistry A OCR Advanced A level chemistry A Salters chemistry IB chemistry

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, M^δ+-X^δ+ polar bond and M-X a relatively non-polar bond (no partial charges shown)

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• 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)

• 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⁺)₂O^2- and (Na⁺)₂O₂^2-, Mg²⁺O^2- and (Al³⁺)₂(O^2-)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 (O^2-). 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 (O^2-). 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⁺) and oxide ions (O^2-), 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 and lower than the preceding magnesium oxide, but still higher than the giant covalent lattice. You would expect aluminium oxide, with the greater cationic charge and smaller cation radius, to have a higher melting point than magnesium oxide considering a purely ionic bonding model.

    4. 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, M_r(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 -  M_r{(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 27^oC. it forms crystals of the trimer molecule that melt at 17^oC (unstable gamma form). Above 27^oC, condensation of SO₃ vapour produces fibrous stable crystals of the alpha form, which has a polymeric and more stable structure.

    7. chlorine(VII) oxide - M_r(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

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

• 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.

      • O^2-(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) Al₂O_3(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 'X_xO_y' 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.

  • -

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

• 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 180^oC 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)

• 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.

        • PCl_3(l) + 3H₂O_(l) ==> H₃PO_3(aq) + 3HCl_(aq)

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

        • (i) PCl_5(s) + H₂O_(l) ==> POCl_3(aq) + 2HCl_(aq) 

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

        • (ii) POCl_3(aq) + 3H₂O_(l) ==> H₃PO_4(aq) + 3HCl_(aq) 

        • overall (i) + (ii): PCl_5(s) + 4H₂O_(l) ==> H₃PO_4(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 MH_x

  • 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) + H_2(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^- ==> X^δ+-H^δ  ==> X-H ==> X^δ-H^δ+

  • -

• 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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