At the beginning of the nineteenth century, chemists began to realize that some further classification of the chemical elements, other than the centuries-old grouping into metals and non-metals was possible. During that century, tables of atomic weights (some of which were in reality equivalent weights) were built up, and a rudimentary theory of valency came into being. In 1869, almost simultaneously, Lothar Meyer and Mendeleef independently put forward the theory that: the properties of the elements are periodic functions of their atomic weights. Thus, elements such as iron, cobalt and nickel, having similar atomic weights, possessed similar properties; an increase in atomic weight in passing from one element to another contributed a definite change in physical and chemical properties.

From this periodic law, both scientists drew up tables of classification. That of Mendeleef was based on a consideration of the valencies of the elements as exhibit in their oxides, chlorides and hydrides; a consideration of these enabled him to arrange the elements in eight groups, some of which were sub-divided into sub-groups.

A modified form of Mendeleef’s Periodic Classification is shown with all the elements now known included. (Group 0 is new; the noble gases had not been discovered in Mendeleef’s time.) The atomic weight is put under the symbol; the number on the left is the atomic number. The vertical columns are called groups and the horizontal rows series or periods. The letter after a group number refers to the original sub-grouping of Mendeleef.

It will be seen that certain groups possess elements in Periods 2 and 3 – these are called head elements – whereas other groups in the centre of the table have no head elements. The groups with head elements – the main groups are often referred to simply as Group I (i.e. Group IA), Group II, etc.

Period 1 of the table contains only hydrogen and helium; there is no wholly satisfactory way of grouping hydrogen and its symbol is placed in Group I, with the alkali metals, and Group VII with the halogens, since it has properties in common with both of these groups of elements, e.g. it is monovalent in its oxide, H2O, (cf. sodium monoxide, Na2O) but is a diatomic gas (cf. fluorine, F₂, and chlorine, Cl₂).

Periods 2 and 3 each contains eight elements (including the noble gases) and are called the ‘short periods’.

Periods 4 and 5 contain eighteen each; the series from scandium, Sc (at. No. 21), to zinc, Zn (at. No. 30) is called a transition series, as is that immediately below it, from yttrium, Y, to cadmium, Cd.

Periods 6 and 7 each contain thirty two elements, although they appear in the table to contain only eighteen. This is because in Group III there is a series of fifteen elements beginning at lanthanum in Period 6, and another similar series beginning at actinium in Period 7, these series should go in horizontally, each to form a series within a transition series’. To avoid inconvenient sideways expansion of the table, these two series are printed below it, as shown. The first series the lanthanoids, were originally called the ‘rare earth metals’; all these resemble one another and closely cover a small range of atomic weights. The second series is called the actinides, and it includes the recently-discovered heavy elements, e.g. plutonium and neptunium.

The periodicity of properties in the table may now be considered. Along a period, e.g. period 3, there is a gradual change in the nature of the typical oxides (i.e. those in which the elements exhibit their group valency);



The same gradation can be found in the lower periods, e.g. from potassium to bromine, if we disregard the transition series in between. The elements at the extreme left and extreme right are the most reactive (Group 0 accepted) those in Group IA being the most electropositive and those in Group VIIB the most electronegative.

Passing down any main group possessing head elements, i.e. in Groups 1A, IIA, IIIB, IVB, VB, VIB and VIIB there is a marked resemblance between in electropositive character down the group; with a lead are much more electropositive than carbon and silicon. Similarly, in Group VIIB, iodine is more electropositive than fluorine or chlorine. The top element of each group i.e. in Period 2, is often not strictly representative of that group; than Lithium (Group IA) resembles magnesium in Group IIA in many respects and boron has some similar diagonal resemblances to silicon. To the right of the table, the abnormality of the top element can be generalized by saying that its maximum valency is less than that of other members of the group; thus carbon can never exceed a valency of 4 (and this is also true of the other Period 2 elements), whereas silicon can have a covalency of 6 (e.g. in the complex fluorosilicate ion [SiF6]2- as can other elements in group IVB.

Among the elements in the middle of the Table (often called the transition elements) regularities of behaviour are not so readily observed. At the extreme left and right – i.e. in Group IIIA and Group IIB – there is still a strong resemblance between the elements in a group but in the middle, particularly in Group VIII, there is little or no’ group resemblance’. Nor is there a marked increase in electropositive character as we go down a group. Thus in group IB, containing the metals copper, silver and gold (the ‘coinage metals’) there is some resemblance between the metals themselves, but they differ considerably in chemical properties, and gold is actually less electropositive than copper. It is now known that there are gradation of properties as we cross a transition series horizontally e.g. from scandium to zinc.

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