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Periodic table - Trends in the chemistry of the elements of Period 3 - revision notes for Advanced A/AS level inorganic chemistry:

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

Part 5 Period 3 Na to Ar

5.1 Survey of the individual elements of Period 3, physical properties and chemical reactions with oxygen, chlorine, water, strong acids and bases

A brief description of the elements that make up Period 3 of the Periodic Table, namely, Na sodium, Mg magnesium, Al aluminium, Si silicon, P phosphorus, S sulfur, Cl chlorine and Ar argon. There physical and chemical properties are summarised below with links to other more detailed notes, including more notes on Period 3.

INORGANIC Part 5 Period 3 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

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

5.1 Survey of the individual elements Na, Mg, Al, Si, P, S, Cl, Ar

Z = 11 Sodium Na in Group 1 Alkali Metals

  • The structure of the element:

    • Giant lattice metallic structure of immobile positive metal ions surrounded by a 'sea' of freely moving mobile electrons (so–called delocalised electrons).

  • Physical properties: 

    • A Soft silvery solid, less dense than water; mpt 98oC; bpt 883oC; good conductor of heat/electricity.

  • Group, electron configuration (and oxidation states): 

    • Gp1 Alkali Metal; e.c.  2,8,1 or 1s22s22p63s1; (only +1 oxidation state) e.g. in NaCl, Na2O etc.

  • Reaction of element with oxygen: 

    • Burns when heated in air to form the ionic white solid oxides, sodium oxide, Na2O or (Na+)2O2–.

      • 4Na(s) + O2(g) ==> 2Na2O(s)  sodium oxide (Na+)2O2–.

        • and the lesser reaction  Na(s) + O2(g) ==> Na2O2(s)  sodium peroxide, (Na+)2[O2]2–

  • Reaction of oxide with water:

    • Na2O is a very basic oxide, readily dissolving/reacting to form an alkaline solution of sodium hydroxide of pH 14.

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

  • Reaction of oxide with acids:

    • Behaves as a basic oxide dissolving to form the chloride, sulfate and nitrate salt in the relevant dilute acid.

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

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

    • Na2O(s) + 2HNO3(aq) ==> 2NaNO3(aq) + H2O(l)

    • In all cases the ionic equation is: Na2O(s) + 2H+(aq) ==> 2Na+(aq) + H2O(l)

  • Reaction of oxide with strong bases/alkalis:

    • None, sodium oxides are ONLY basic.

  • Reaction of element with chlorine: 

    • Burns when heated in chlorine to form white powder/colourless crystals of ionic sodium chloride, NaCl or Na+Cl–.

      • 2Na(s) + Cl2(g) ==> 2NaCl(s) 

  • Reaction of chloride with water:

    • The salt readily dissolves forming a neutral solution of sodium and chloride ions (pH 7).

      • NaCl(s) + aq ==> Na+(aq) + Cl–(aq) 

      • The chloride ion is such a weak base it shows no chemical interaction with water and the sodium ion shows virtually no acidic character, so sodium chloride solution is neutral ~pH 7.

  • Reaction of element with water:

    • The metal reacts quite quickly forming hydrogen gas and alkaline sodium hydroxide, pH 14.

      • 2Na(s) + 2H2O(l) ==> 2NaOH(aq) + H2(g) 

  • Other comments:

    • –

  • Links to other pages on site: 

 

Z = 12 Magnesium Mg in Group 2 Alkaline Earth Metals

  • The structure of the element:

    • Giant lattice metallic structure of immobile positive metal ions surrounded by a 'sea' of freely moving mobile electrons (so–called delocalised electrons).

  • Physical properties: 

    • A moderately hard silvery–white solid; mpt 649oC; bpt 1090oC; good conductor of heat/electricity.

  • Group, electron configuration (and oxidation states): 

    • Gp2 Alkaline Earth Metal; e.c.  2,8,2 or 1s22s22p63s2; (only+2 oxidation state) e.g. MgCl2, MgO.

  • Reaction of element with oxygen: 

    • Burns brightly when heated in air to form a white powder of ionic magnesium oxide Mg2+O2– when heated strongly in air.

      • 2Mg(s) + O2(g) ==> 2MgO(s) 

  • Reaction of oxide with water:

    • It is slightly soluble in water, and is a basic oxide forming an alkaline solution of magnesium hydroxide

    • Mg2+(OH–)2, of about pH12.

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

  • Reaction of oxide with acids:

    • Behaves as a basic oxide dissolving to form the chloride, sulfate and nitrate salt in the relevant dilute acid.

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

    • MgO(s) + H2SO4(aq) ==> MgSO4(aq) + H2O(l)

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

    • In all cases the ionic equation is: MgO(s) + 2H+(aq) ==> Mg2+(aq) + H2O(l)

  • Reaction of oxide with strong bases/alkalis:

    • None, sodium oxide is ONLY basic.

  • Reaction of element with chlorine: 

    • Forms colourless solid ionic magnesium chloride, Mg2+(Cl–)2, when heated in chlorine.

      • Mg(s) + Cl2(g) ==> MgCl2(s) 

  • Reaction of chloride with water:

    • The salt dissolves in water forming a nearly neutral solution of about pH6.

      • MgCl2(s) + aq ==> Mg2+(aq) + 2Cl–(aq) 

    • or more correctly ...

      • MgCl2(s) + 6H2O(l) ==> [Mg(H2O)6]2+(aq) + 2Cl–(aq) 

    • The solution is slightly acidic, because the hexa–aqa magnesium ion can donate a proton to a water molecule forming the oxonium ion, more simply H+(aq).

      • [Mg(H2O)6]2+(aq) + H2O(l) [Mg(H2O)5OH]+(aq) + H3O+(aq) 

  • Reaction of element with water:

    • Very slow reaction with cold water to form hydrogen (bubbles form slowly on the surface) and alkaline magnesium hydroxide. 

      • Mg(s) + 2H2O(l) ==> Mg(OH)2(aq) + H2(g) 

    • Ignited magnesium will continue to burn in steam to form the white powder of magnesium oxide and hydrogen gas and you get the same reaction if steam is passed over heated magnesium ribbon.

      • Mg(s) + H2O(g) ==> MgO(s) + H2(g) 

  • Other comments:

    • –

  • Links to other pages on site: 

 

Z = 13 Aluminium Al in Group 3/13

  • The structure of the element:

    • Giant lattice metallic structure of immobile positive metal ions surrounded by a 'sea' of freely moving mobile electrons (so–called delocalised electrons).

  • Physical properties: 

    • Moderately hard silvery–white high melting solid; mpt 661oC; bpt 2467oC;  good conductor heat/electricity.

  • Group, electron configuration (and oxidation states): 

    • Gp3; e.c. 2,8,3 or 1s22s22p63s23p1;  (only +3 oxidation state) e.g. Al2O3 and AlCl3.

  • Reaction of element with oxygen: 

    • Aluminium reacts when heated strongly in air to form the white powder of aluminium oxide which has a giant ionic structure, (Al3+)2(O2–)3.

      • 4Al(s) + 3O2(g) ==> 2Al2O3(s)  

      • The above reaction occurs very rapidly on a freshly cut aluminium surface, but the microscopic oxide layer inhibits any further reaction, giving aluminium a 'lower reactivity' than expected, and its excellent anti–corrosion properties.

  • Reaction of oxide with water:

    • Insoluble, no reaction but it is an amphoteric oxide and forms salts with both strong acids and bases (alkalis). See below for more details.

  • Reaction of oxide with acids:

    • It behaves as a basic oxide dissolving to form the chloride, sulfate and nitrate salt in the relevant dilute acid.

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

    • Al2O3(s) + 3H2SO4(aq) ==> Al2(SO4)3(aq) + 3H2O(l) 

    • Al2O3(s) + 6HNO3(aq) ==> 2Al(NO3)3(aq) + 3H2O(l) 

    • ionic equation: Al2O3(s) + 6H+(aq) ==> 2Al3+(aq) + 3H2O(l) 

  • Reaction of oxide with strong bases/alkalis:

    • The oxide also behaves as an acidic oxide by dissolving in strong soluble bases to form aluminate(III) salts.

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

    • forming sodium aluminate(III) with sodium hydroxide.

    • ionic equation: Al2O3(s) + 2OH–(aq) + 3H2O(l) ==> 2[Al(OH)4]–(aq) 

    • Therefore aluminium oxide is an amphoteric oxide, because of this dual acid–base behaviour.

  • Reaction of element with chlorine: 

    • Burns when heated strongly in chlorine gas to form the white* solid aluminium chloride on heating in chlorine gas.

      • 2Al(s) + 3Cl2(g) ==> 2AlCl3(s)

      • * It is often a faint yellow in colour, due to traces of iron forming iron(III) chloride.

      • Aluminium chloride is a curious substance in its behaviour. The solid, AlCl3, consists of an ionic lattice of Al3+ ions, each surrounded by six Cl– ions, BUT on heating, at about 180oC, the thermal kinetic energy of vibration of the ions in the lattice is sufficient to cause it break down and sublimation takes place (s ==> g). In doing so the co–ordination number of the aluminium changes from six to four to form the readily vapourised covalent dimer molecule, Al2Cl6, shown above.

  • Reaction of chloride with water:

    • With a little water it rapidly, and exothermically hydrolyses to form aluminium hydroxide and nasty fumes of hydrogen chloride gas.

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

    • However, if a large excess of water is rapidly added, a weakly acidic solution of aluminium chloride is formed, with the minimum of nasty fumes!

      • AlCl3(s) + aq ==> Al3+(aq) + 3Cl–(aq) 

      • or more correctly: AlCl3(s) + 6H2O(l) ==> [Al(H2O)6]3+(aq) + 3Cl–(aq) 

    • The solution is slightly acidic, because the hexa–aqa aluminium ion can donate a proton to a water molecule forming the oxonium ion.

      • [Al(H2O)6]3+(aq) + H2O(l) [Al(H2O)5OH]2+(aq) + H3O+(aq) 

  • Reaction of element with water:

    • None due to protective aluminium oxide layer.

  • Reactions of the hexa–aqua aluminium ion:

    • It gives a gelatinous white precipitate with sodium hydroxide or ammonia solution which displays amphoteric behaviour by dissolving in excess strong alkali (NaOH(aq), NOT NH3(aq)) and acids.

      • Al3+(aq) + 3OH–(aq) ==> Al(OH)3(s) 

      • or [Al(H2O)6]3+(aq)  + 3OH–(aq) ==> [Al(OH)3(H2O)3] + 3H2O(l) 

        • The hydroxide readily dissolves in acids to form salts:

        • Al(OH)3(s) + 3H+(aq) ==> Al3+(aq) + 3H2O(l) 

          • or more elaborately: [Al(OH)3(H2O)3] + 3H3O+(aq) [Al(H2O)6]3+(aq)  + 3H2O(l) 

          • Thus showing amphoteric behaviour, since the hydroxide ppt. also dissolves in excess strong alkali (below).

      • [Al(H2O)6]3+(aq) + 6OH–(aq) ==> [Al(OH)6]3–(aq) + 6H2O(l)  (from original aqueous ion)

        • or [Al(OH)3(H2O)3](s) + 3OH–(aq) ==> [Al(OH)6]3–(aq) + 3H2O(l) (from hydroxide ppt.)

        • or more simply Al(OH)3(s) + 3OH–(aq) ==> [Al(OH)6]3–(aq)  (from hydroxide ppt.)

    • With aqueous sodium carbonate solution, the hydroxide ppt. is formed, and, because of its acidic nature, bubbles of carbon dioxide gas are evolved.

      • 2[Al(H2O)6]3+(aq) + CO32–(aq) 2[Al(H2O)5(OH)]2+(aq) + H2O(l) + CO2(g)   

      • this process of proton donation continues until the gelatinous ppt. [Al(OH)3(H2O)3](s) is formed, but will not dissolve in excess of the weak base/alkali.

      • See Appendix 1 in Transition Metal Chemistry pages for more examples and equations etc.

      • No Cr2(CO3)3 is formed because of the acid–base reaction above, due to the acidity of the chromium(III) ion. Note the similarly highly charged and small ions Cr3+ and Fe3+ behave in the same way.

  • Other comments:

    • –

  • Links to other pages on site: 

 

Z = 14 Silicon Si in Group 4/14

  • The structure of the element:

    • Non–metal existing as a giant covalent lattice, Sin, where n is an extremely large number, held together by tetrahedrally arranged Si–Si bonds.

  • Physical properties: 

    • Hard high melting solid; mpt 1410oC; bpt 2355oC;  poor conductor of heat/electricity, but with other elements added, conducts better, hence use in microchips.

  • Group, electron configuration (and oxidation states): 

    • Gp4; e.c. 2,8,4 or 1s22s22p63s23p2;  (only +4 oxidation state) e.g. SiO2 and SiCl4 etc.

  • Reaction of element with oxygen: 

    • Reacts when strongly heated in air to form silicon dioxide (silica, silicon(IV) oxide).

      • Si(s) + O2(g) ==> SiO2(g) 

      • Like aluminium, silicon is protected by a thin oxide layer.

  • Reaction of oxide with water:

    • None and insoluble.

  • Reaction of oxide with acids:

    • None, only acidic in nature.

  • Reaction of oxide with bases/alkalis:

    • It is a weakly acidic oxide dissolving very slowly in hot concentrated sodium hydroxide solution to form sodium silicate.

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

    • or simplified ionic equation: SiO2(s) + 2OH–(aq) ==> SiO32–(aq) + H2O(l)

  • Reaction of element with chlorine: 

    • On heating in chlorine forms the covalent liquid silicon tetrachloride.

      • Si(s) + 2Cl2(g) ==> SiCl4(l)

  • Reaction of chloride with water:

    • Hydrolyses to form gelatinous hydrated silicon oxide and hydrochloric acid.

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

  • Reaction of element with water:

    • None - due to protective silicon oxide layer.

  • Other comments:

    • –

  • Links to other pages on site: 

 

Z = 15 Phosphorus P in Group 5/15

  • The structure of the element:

    • Two solid allotropes (red and white) consisting of P4 molecules, also a polymer form.

  • Physical properties: 

    • Colourless gas; mpt 44oC; bpt 280oC; poor conductor of heat/electricity.

  • Group, electron configuration (and oxidation states): 

    • Gp5; e.c. 2,8,5  or 1s22s22p63s23p3; Variety of oxidation states from –3 to +5 e.g.

    • PH3 (–3), P4O6 (+3), P4O10, PCl5 and H3PO4 (+5).

    • The maximum oxidation state is +5.

  • Reaction of element with oxygen: 

    • White phosphorus ignites spontaneously in air, which is why it is kept under water!)

    • The reaction is vigorous, very exothermic and the phosphorus burns with a bright flame.

    • With limited air/oxygen, on heating the phosphorus, the covalent white solid phosphorus(III) oxide is formed.

      • P4(s) + 3O2(g) ==> P4O6(s) 

    • With excess air/oxygen, on heating the phosphorus, the covalent white solid phosphorus(V) oxide is formed.

      • P4(s) + 5O2(g) ==> P4O10(s) 

  • Reaction of the oxides with water: Both oxides are acidic, typical non–metallic element behaviour, and both phosphorus oxides dissolve in water to form acidic solutions.

    • Phosphorus(III) oxide forms (i) phosphonic acid and (II) a little phosphoric(III) acid, which is an isomer.

      • (i) major reaction: P4O6(s) + 6H2O(l) ==> 4HPO(OH)2(aq)

        • [a dibasic/diprotic acid, often written as H3PO3, but behaves as O=PH(OH)2]

        • Note that phosphorus still has an oxidation state equivalent to +3 in phosphonic acid.

      • (ii) minor reaction: P4O6(s) + 6H2O(l) ==> 4H3PO3(aq)   (which could be written as P(OH)3)

    • Phosphorus(V) oxide forms phosphoric(V) acid, no complications here!

      • P4O10(s) + 6H2O(l) ==> 4H3PO4(aq)   [a tribasic/triprotic acid, O=P(OH)3]

  • Reaction of element with chlorine: 

    • With limited chlorine, on heating the phosphorus, the covalent liquid phosphorus(III) chloride is formed.

      • P4(s) + 3Cl2(g) ==> 4PCl3(l) 

    • With excess chlorine, on heating the phosphorus, the ionic* solid phosphorus(III) chloride is formed.

      • P4(s) + 5Cl2(g) ==> 4PCl5(s) 

      • * PCl5 is a bit unusual for an 'expected covalent' liquid chloride.

        • It is an ionic solid with the structure [PCl4]+[PCl6]– 

          • Hence its melting point is much greater than the liquid phosphorus(III) chloride, where the molecules are only held together by the inter–molecular forces.

          • However, gaseous phosphorus(V) chloride consists of PCl5 covalent molecules.

  • Reaction of oxide with acids:

    • None, only acidic in nature.

  • Reaction of phosphorus oxides with strong bases/alkalis:

    • Both oxides dissolve in alkalis to form a whole series of phosphate(III) and phosphate(V) salts.

    • So, with strong bases like sodium hydroxide, the simplified equations are:

      • initially: P4O6(s) + 4NaOH(aq) + 2H2O(l) ==> 4NaH2PO3(aq)

        • forming the mono sodium salt of phosphonic acid,

        • then, with excess sodium hydroxide, you get the disodium salt of phosphonic acid

          • NaH2PO3(aq) + NaOH(aq) ==> Na2HPO3(aq) + H2O(l)

          • and there is NO trisodium salt because H3PO3 behaves as dibasic/diprotic O=PH(OH)2

          • So, overall with excess sodium hydroxide the reaction is ...

          • P4O6(s) + 8NaOH(aq) ==> 4Na2HPO3(aq) + 2H2O(l)

      • P4O10(s) + 12NaOH(aq) ==> 4Na3PO4(aq) + 6H2O(l) sodium phosphate(V) formed from phosphorus(V) oxide

        • ionic equation: P4O10(s) + 12OH–(aq) ==> 4PO43–(aq) + 6H2O(l)

      • If the empirical formulae P2O3 and P2O5 are used, just halve all the balancing numbers.

      • Other than using excess sodium hydroxide solution, other salts can be formed.

      • e.g. P4O10(s) + 4NaOH(aq) + 2H2O(l) ==> 4NaH2PO4(aq)      sodium dihydrogen phosphate(V)

      • or  P4O10(s) + 8NaOH(aq) ==> 4Na2HPO4(aq) + 2H2O(l)        disodium hydrogen phosphate(V)

  • Reaction of the chlorides with water:

    • Phosphorus(III) chloride hydrolyses rapidly and exothermically to form phosphoric(III) acid or phosphonic acid (see reaction of oxides with water).

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

        • or     PCl3(l) + 3H2O(l) ==> HPO(OH)2(aq) + 3HCl(aq)

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

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

      • Then on boiling the aqueous solution, phosphoric(V) acid is formed and more hydrochloric acid.

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

      • overall the hydrolysis reaction is: PCl5(s) + 4H2O(l) ==> H3PO4(aq) + 5HCl(aq) 

  • Reaction of element with water:

    • None.

  • Other comments:

    • –

  • Links to other pages on site: 

 

Z = 16 Sulfur S in Group 6/16

  • The structure of the element:

    • Three solid allotropes. Two are crystalline lattices based on S8 molecules (rhombic and monoclinic sulfur). A 3rd form is an unstable dark brown–black polymeric form called plastic sulfur, formed when boiling sulfur is poured onto cold water, great fun, but of little use!

  • Physical properties: 

    • Colourless gas; mpt 117oC; bpt 445oC; poor conductor of heat/electricity.

  • Group, electron configuration (and oxidation states): 

    • Gp6; e.c. 2,6  or 1s22s22p63s23p4; oxidation states range from –2 to +6 e.g.

    • Na2S (–2), S2Cl2 (+1), SO2 (+4) and H2SO4, SF6, SO3 (all +6).

  • Reaction of element with oxygen: 

    • Sulfur burns in air with a pale blue flame to form sulfur dioxide (sulfur(IV) oxide), with a little sulfur trioxide.

      • S(s) + O2(g) ==> SO2(g) 

      • Sulfur(IV) oxide is a nasty choking gas, a lung irritant that can cause asthma attacks, so only burn sulfur in a well ventilated fume cupboard.

    • Sulfur trioxide (sulfur(VI) oxide) has to be made by the industrial Contact Process.

      • 2SO2(g) + O2(g) == V2O5 catalyst, 450oC ==> 2SO3(g) 

  • Reaction of the oxides with water: Both dissolve to form acid solutions.

    • Sulfur dioxide forms the weak 'fictitious' sulfurous acid.

      • SO2(g) + H2O(l) ==> H2SO3(aq) 

        • the reaction is better represented ionically as ..

        • SO2(aq) + H2O(l) H+(aq) + HSO3–(aq) 

    • Sulfur trioxide reacts very violently and exothermically to form the oily liquid, strong sulfuric acid.

      • SO3(g) + H2O(l) ==> H2SO4(l) 

      • In water, the sulfuric acid is almost fully ionised.

        • H2SO4(aq) + 2H2O(l) ==> 2H3O+(aq) + SO42–(aq) 

  • Reaction of oxide with acids:

    • None, only acidic in nature.

  • Reaction of oxide with bases/alkalis:

    • Sulfur dioxide dissolves in strong bases to form sulfites/sulfate(IV)s

    • 2NaOH(aq) + SO2(g) ==> Na2SO3(aq) + H2O(l) formation of sodium sulfite/sulfate(IV)

    • ionic equation: 2OH–(aq) + SO2(g) ==> SO32–(aq) + H2O(l)

    • You would NOT attempt to react sulfur trioxide with water, the reaction is very violent and exothermic.

    • but theoretically: 2NaOH(aq) + SO3(g) ==> Na2SO4(aq) + H2O(l)

  • Reaction of element with chlorine: 

    • When chlorine is passed over molten sulfur a variety of chlorides are formed.

    • The main product is disulfur dichloride

    • 2S(s) + Cl2(g) ==> S2Cl2(l)  (SiCl2, SiCl4 also possible)

  • Reaction of chloride with water:

    • Slowly hydrolyses in water, via a complex reaction, to form an acid solution of several products (not meant to be a balanced equation).

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

      • (the equation is not balanced because its a complex redox - hydrolysis reaction!)

  • Reaction of element with water:

    • None.

  • Other comments:

    • –

  • Links to other pages on site: 

Z = 17 Chlorine Cl in Group 7/17 The Halogens

  • The structure of the element:

    • Non–metal existing as covalent diatomic molecule, Cl2, with a single bond.

  • Physical properties: 

    • Pale green gas; mpt –101oC; bpt –34oC;  poor conductor of heat/electricity.

  • Group, electron configuration (and oxidation states): 

    • Gp7 Halogens; e.c. 2,8,7  or 1s22s22p63s23p5;  (oxidation state ranges from –1 to +7) e.g.

    • HCl and NaCl (–1), NaClO and Cl2O (+1), NaClO2 (+3), KClO3 (+5), Cl2O7 and HClO4 (+7).

  • Reaction of element with oxygen: 

    • None, but there are several oxides made indirectly.

  • Reaction of the oxides with water:

    • Chlorine(I) oxide forms weak chloric(I) acid.

      • Cl2O(g) + H2O(l) ==> 2HClO(aq)  

    • Chlorine(V) oxide forms the strong chloric(VII) acid.

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

  • Reaction of oxide with acids:

    • None, only acidic in nature.

  • Reaction of oxide with bases/alkalis:

    • chlorine(I) oxide forms sodium chlorate(I) with sodium hydroxide,

      • Cl2O(g) + 2NaOH(aq) ==> 2NaClO(aq) + H2O(l)

      • ionic equation: Cl2O(g) + 2OH–(aq) ==> 2ClO–(aq) + H2O(l)

    • and chlorine(VII) oxide will dissolve to form sodium chlorate(VII)

      • Cl2O7(l) + 2NaOH(aq) ==> 2NaClO4(aq) + H2O(l)

      • ionic equation: Cl2O7(l) + 2OH–(aq) ==> 2ClO4–(aq) + H2O(l)

  • Reaction of element with water:

    • Slightly reacts to form an acid mixture of chloric(I) acid and hydrochloric acid, but the position of the equilibrium is very much on the left.

      • Cl2(g) + H2O(l) HClO(aq) + HCl(aq) 

      • or more accurately and ionically ...

      • Cl2(g) + 2H2O(l)  HClO(aq) + H3O+(aq) + Cl–(aq) 

  • Other comments:

    • –

  • Links to other pages on site: 

Z = 18 Argon Ar in Group 0/18 The Noble Gases

 

Some trend comments on the reactions of Period 3 elements (from left to right across period 3)

  • In terms of the reactions with oxygen - you start very reactive metals (Na, Mg), then moderately reactive metal (Al), a relatively unreactive non-metal (Si), a very reactive non-metal (P), moderately reactive non-metal (S) and neither chlorine nor argon react directly with oxygen.

    • Its not the greatest of trends, but the individuality of each reaction, or lack of reaction, is important.

  • However, the reaction of the elements with water AND more crucially, the reaction of the oxides with water clearly produce a trend from left to right of basic (alkaline) to acidic character - see pH values in section 5.3.

  • These points are discussed in more detail in section 5.3.

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, 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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