7.5                                           Inorganic Chemistry

 

7.5.1           Periodic Trends

 

 

The organisation of elements in the Periodic Table according to their proton numbers and electronic structures.

 

 

Periodicity and the Elements of the Periodic Table

 

Write notes on the organisation of the elements in the Periodic Table.

These should include ;

 

•    Limitations of classifying elements into metals and non-metals

•    Work of Döbereiner on triads

•    Contribution of John Newlands and the Law of Octaves.

•    Work of Lother Mayer and Dmitri Mendeléev on periodicity of properties.

•    The development of the modern Periodic Table based on atomic number.

 

Because of limitations to the overall classification of elements as metals and non-metals, chemists began to search for trends and similarities in the properties of much smaller groups of elements.

Do the elements have similar properties?

Do they react in the same way with oxygen in the air?

What colour are the oxides, chlorides, sulphates etc.?

How do they react with water?  Are the reactions and products similar?

Are the formulae of their compounds similar?

 

Early in the 19th century, the German chemist Döbereiner pointed out that many of the known elements could be arranged in groups of 3 similar elements.  He called these families of three elements triads.  Two of Döbereiner's triads were lithium, sodium and potassium and chlorine, bromine and iodine.  He showed that when the three elements in each triad were written in order of relative atomic mass, the middle element had properties in between those of the other two elements.  In the Cl, Br, I triad the relative atomic mass of bromine(79.9) is close to the average of the relative atomic masses of Cl and I (35.5+216.9/2=81.2)

 

In 1866, John Newlands an English chemist suggested that when the elements were arranged in order of their relative atomic masses, any one element had properties similar to the elements eight places in front of it and eight places behind it in the list.  Newlands called this a law of octaves.  He claimed that the eighth element was a kind of repetition of the first.  The periodic repetition of similar elements at regular intervals in Newlands octaves led to the name 'periodic table'.  However Newlands ideas were rejected for other reasons.

 

In 1869 the German chemist Julius Lother Mayer, and the Russian chemist Dmitri Mendeléev, separately published results which supported the idea of periodicity. Lother Mayer plotted various physical properties of the known elements against their relative atomic masses.  Most of the credit went to Mendeléev for arranging the elements.  He arranged the 60 or so elements known to him in order of increasing relative atomic mass and showed that elements with similar properties fall in the same vertical columns.  These vertical columns are called groups and the horizontal rows are called periods.  His periodic law stated that the properties of the elements are a periodic function of their relative atomic masses.  This law fulfilled two important functions.

1.                                                                                                                                                                It summarised the properties of elements and classified them into groups  with similar properties.

2                                                                                                                                                                It enabled predictions to be made about the properties of known and unknown elements and led to considerable research activity.

 

The reactive metals are found in Groups I and II (the s-block elements).  These metals are all high in the reactivity series.  They have lower densities, lower melting points and lower boiling points than most other metals and they form stable involatile ionic compounds.

The transition metals occupy two rectangles between Groups II and III. These metals are much less reactive than the metals in Groups I and II.  In this block of elements there is also a marked horizontal similarity in properties as well as the usual vertical likeness (The d-block elements)

The non-metals also form a triangular block in the Periodic Table (The p-block elements).

The noble gases have completely filled s and p subshells of electrons and are very unreactive.

 

 

The variation in the character of the elements across the third period , sodium to chlorine, as indicated by a study of their compounds, limited to formulae, type of bonding and behaviour with water of the following oxides and chlorides:  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.

 

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 )

 

Comparison of the Oxides of the Elements From Na to P

(seeTable 3)

 

Structure and Bonding

 

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)

 

                                                          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

(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+

Brřnsted       Brřnsted

 

acid              base

 

The solution is acidic (pH~3)

The remaining chlorides hydrolyse in water forming acid solutions.

 

The reactions of the elements with oxygen and chlorine to form the oxides and chlorides listed above where appropriate. 

 

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.

 

The reactions of the elements in the third period with water where appropriate.

 

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