4.10         Periodic Trends

 

The variation in the character of the elements across the third period , sodium to chlorine, as indicated

 

Comparison of the Elements of the Third Period

 

In Period 3, sodium, the left-hand element, is a very reactive metal, whereas chlorine, next to the extreme right is a very reactive non-metal.  In between, the elements show a gradual transition from metals to non-metals.  These periodic changes in the physical properties of the elements across the table are reflected in a periodic change in structure.  The structures of the elements vary from metallic, through giant molecular in the metalloids to simple molecular structure in the non-metals. (see Table 1 )

 

 

The reaction of the elements with water (where appropriate). Water would include steam where this is required for the elements to react.

 

Reaction of the elements with water

 

Na     vigorous reaction forming an alkaline solution of sodium hydroxide Na(s)  +  2H2O(l)                        NaOH(aq)  +  H2(g)      

 

Mg     slow reaction with cold water forming a weakly alkaline solution

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

vigorous reaction with steam

Mg(s)  +  H2O(g)                                 MgO(s)  +  H2(g)   

 

Al      usually inert due to an impervious oxide layer on the surface if the oxide layer is removed using Hg2+(aq) then Al reacts slowly with warm water Al(s)  +  H2O(l)                              Al(OH)3(s)  +  H2(g)

 

Si      no reaction  

 

P        no reaction (stored under water)

 

S       no reaction  

 

Cl       moderately soluble forming an acid solution

Cl2(g)  +  2H2O(l)                      HCl(aq)  +  HOCl(aq)                                                             

                           chloric(I) acid        

 

Ar      no reaction  

 

 

 

The reactions of the elements with oxygen and chlorine to form the oxides and chlorides listed below.

 

Reactions of the Elements in Period 3

 

ELEMENT

Heat element in  dry chlorine

Heat element in  dry oxygen

Na

very vigorous reaction forming Na+Cl

very vigorous reaction  forming (Na+)2O2- and (Na+)2(O22-)

Mg

vigorous reaction forming Mg2+(Cl-)2

very vigorous reaction  forming Mg2+O2-

Al

vigorous reaction  forming Al2Cl6

vigorous reaction at first forming (Al3+)2(O2-)3,  then the oxide layer  prevents further attack

Si

slow reaction  forming SiCl4

slow reaction  forming SiO2

P

slow reaction forming  PCl3 and PCl5

vigorous reaction  forming P4O6 and P4O10

S

slow reaction  forming SCl2

and S2Cl2

slow reaction  forming SO2

Cl

no reaction   

no reaction

Ar

no reaction   

no reaction

                  

The trends in reactivity vary with different reagents.  The reactivity of the elements with oxygen and chlorine gradually falls across the period.  The trends can be related with considerable accuracy to the action of the elements as oxidising or reducing agents.

The metals (Na/Mg/Al) are strong reducing agents and readily give up electrons to form the corresponding ions.

                             M                Mn+  +  ne-

These elements never exist free in nature but only in compounds with Na+, Mg2+ and Al3+ ions.  Thus these three metals react vigorously with non-metals and oxidising agents.

Cl2  +  2e-             2Cl-

O2  +  4e-              2O2-

 

Silicon in  Group IV is a very weak reducing agent.  It will react slowly with chlorine and oxygen, which are strong oxidising agents, but not with weak oxidising agents.

Phosphorus and sulphur are both weak oxidising and weak reducing agents.  They react slowly with oxygen and chlorine (strong oxidising agents) and molten sulphur will react slowly with hydrogen forming hydrogen sulphide.

H2(g)  +  S(l)                   H2S(g)

The sulphur is now acting as the oxidising agent and hydrogen as the reducing agent.

Chlorine is a strong oxidising agent and thus has no reaction with oxygen.  However it will react violently with hydrogen in sunlight to form hydrogen chloride.

H2(g)  +  Cl2(g)               2HCl(g)

Argon on the extreme right shows no reactivity with these reagents.

 

 

Knowledge of the oxides and chlorides limited to formulae, type of bonding , structures (excluding that of phosphorus pentoxide), classification of the oxides as acidic, basic or amphoteric and the reaction of the compounds with water (if any).

Na2O, MgO, Al2O3, SiO2, P4O10, SO2, SO3, Cl2O7, NaCl, MgCl2, Al2Cl6 and PCl5 (detailed structures of individual compounds not required).  Interpretation of acid-base character of oxides and behaviour of chlorides towards water in terms of structure and bonding.

 

 

The reactions of the elements with oxygen and chlorine to form the oxides and chlorides listed below.

 

Comparison of the Oxides Na to P

(See Table 3)

(seeTable 3)

Na2O, MgO, Al2O3

All have giant structures in which the bonding is ionic.  Therefore they have high melting points and boiling points and conduct electricity when molten.

 

SiO2

Silicon is a metalloid and forms the oxide SiO2 which has a giant covalently bonded structure.  It is therefore a solid with high melting point and boiling point and does not conduct electricity when molten.

 

P4O10, SO2, SO3, Cl2O7

The remaining oxides are formed by non-metals.  These oxides are built up from molecules within which the bonding is covalent, but between the molecules the bonding is the weak Van der Waals type.  These oxides therefore have low melting points and boiling points and do not conduct electricity in the liquid state.

The bonding in these oxides becomes less ionic (more covalent) on moving across the period from left to right, the difference in electronegativity between the element and oxygen decreases and the bonding is more covalent (see table).

Most of the oxides have negative values of DHfq i.e.. they are stable compounds.  The DHfq per mole of oxygen atoms gives a measure of the stability of the bonds in the oxide.  Note that DHfq per mole of oxygen atoms shows a general decrease (i.e. becomes less negative) from left to right across the period- the stability of the oxide decreases. 

Also the change in structure and bonding of the oxides leads to a marked difference in their reaction with water, acid and alkali i.e.. their acid/base character.

 

Reaction with Water

Na2O

Ionic oxide  with Na+ and O2- ions in the crystal lattice.  Reacts vigorously with water forming an alkaline solution.

O2- (aq) +  H2O (l)                   2OH- (aq)

 

MgO

Reacts slowly with water but readily with dilute acids. i.e.. less basic than Na2O.

MgO(s)  +  H2O(l)                     Mg2+(aq)  +  2OH-(aq)

MgO(s)  +  2H+(aq)                    Mg2+(aq)  +  H2O(l)

 

Al2O3

This is an amphoteric oxide.  It has no reaction with water but reacts with both H+ and OH-.

Al2O3(s)  +  6H+(aq)                  2Al3+(aq)  +  3H2O(l)

Al2O3  +  2OH-(aq)  +  3H2O(l)                     2Al(OH)4-(aq)

                                                                   aluminate ion

 

SiO2

No reaction with water but reacts with concentrated alkali forming the silicate ion (SiO32-).  Weakly acidic.

SiO2(s)  +  2OH-(aq)                  SiO32-(aq)  +  H2O(l)

 

P4O6, P4O10, SO2, SO3, Cl2O7

These are all strongly acidic oxides and react readily with water forming strong acid solutions.

                    P4O6(s)  +  6H2O(l)                  4H3PO3(aq)

                    P4O10(s)  +  6H2O(l)                 4H3PO4(aq)

                    SO2(g)  +  H2O(l)                     H2SO3(aq)

                    sulphur (IV) oxide

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

                    sulphur (VI) oxide

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

                    chlorine (VII) oxide                  chloric(VII) acid

 

Note the correlation on moving left to right across Period 3 between structure and bonding and acid/base character of the oxides.( Table 2)

 

                                                         

 

L

H

S

The ionic, metallic oxides  are basic (Na2O, MgO) or  amphoteric (Al2O3).

The oxide of the  metalloid Si, SiO2 has  a giant covalent  lattice structure and  is weakly acidic.

The remaining  oxides, formed by  non-metals, have  simple molecular structures and are  acidic.

R

H

S         

 

                                     

Comparison of the Chlorides of the Elements From Na to P

(See Table 4)

 

Structure and Bonding

 

NaCl, MgCl2

These chlorides have giant structures.  NaCl has an ionic lattice built up from Na+ and Cl- ions and MgCl2 has a layer lattice with some ionic bonding.  Therefore these chlorides are solids with high melting points and boiling points and conduct electricity when molten or in solution.

 

AlCl3

This has a layer lattice with covalent/ionic bonding. 

Note: In the liquid and vapour phases it exists as the dimer Al2Cl6 with covalent and dative bonding

Solid AlCl3 has a mixture of covalent and ionic bonding.  With the increased covalent bonding its melting point, boiling point and electrical conductivity when molten are all much lower than for NaCl and MgCl2.

 

SiCl4/PCl5/S2Cl2/Cl2

The remaining chlorides have structures built up from molecules within which there is covalent bonding but the bonding between the molecules is the weak Van der Waals type.  Therefore these chlorides have low melting points and as liquids do not conduct electricity.

 

The bond nature of these chlorides becomes less ionic (i.e.. more covalent) on moving across the period from left to right.  For any chloride, the smaller the difference in electronegativity between the element and chlorine, the more covalent the bonding.  Since the electronegativity increases from left to right, the electronegativity difference between the element and chlorine decreases.  Therefore the bonding is more covalent. (see Table 4)

All the chlorides have a negative DHfq and are therefore stable.

Note the decrease in the heat evolved when the elements react with one mole of Cl atoms.  This is the gradation expected from the reaction between the electronegative Cl and the strongly electropositive sodium on the left to the reaction between the electronegative non-metals Cl and P on the right.

 

 

Reaction with Water

 

NaCl

Dissolves to form the hydrated Na+(aq) and Cl-(aq).  The solution is neutral.

MgCl2

Similar effect to sodium.  Mg2+(aq) and Cl-(aq) formed.

AlCl3

This gives Al3+(aq) and Cl-(aq) but in this case the hydrated cation dissociates as shown.

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

Brønsted       Brønsted

acid              base

 

The solution is acidic (pH~3)

The remaining chlorides hydrolyse in water forming acid solutions.

 

Commercial aspects

 

Laboratory aspects

 

Reactions of the acids and their salts based on redox properties, stability of oxidation states, acid strength. The variation in strength of acid with the change in oxidation number of the central atom. Detailed knowledge required only for named acids.

 

 

Nitrous acid and nitrites: reaction with iron (II), iodide, manganate (VII), bromine water and aluminium in alkali.

 

 

 

 

Nitric acid and nitrates: reaction of concentrated acid with iron (II) , iodide, copper metal and aluminium (in alkali).

 

 

 

Phosphoric acid:

 

 

 

Sulphurous acid, sulphite and sulphur dioxide:

 

 

Sulphuric acid and sulphates: reaction of concentrated acid with copper metal, halide ions (see Module 1), sulphates with barium ions.

 

Conc. sulphuric acid and copper metal

Reacts forming a blue solution of copper sulphate and releasing sulphur dioxide gas

                    Cu  +  2H2SO4                 CuSO4  +  SO2  +  2H2O

 

Halide ions

Reacts forming the hydrogen halide

                    H2SO4  +  2NaX                       Na2SO4  +  2HX

                   

Thiosulphates:

 

With halogens

 

With silver ions Ag+