The periodic law of the chemical elements: ‘The new system of atomic weights which renders evident the analogies which exist between bodies' [^]

Abstract

The historical roots, the discovery and the modern relevance of Dmitri Mendeleev's remarkable advance have been the subject of numerous scholarly works. Here, with a brief overview, we hope to provide a link into the contents of this special issue honouring the great scientist. Mendeleev's advance, announced in March 1869, as he put it in 1889, to the ‘…then youthful Russian Chemical Society…’, first set out the very basis of the periodic law of the chemical elements, the natural relation between the properties of the elements and their atomic weights. This was, and still is, the centrepiece of a historical journey for chemistry to today's position as a pre-eminent science.

This article is part of the theme issue ‘Mendeleev and the periodic table'.

1. Introduction: The New Light

Every so often in science there is a discovery or advance whose brilliance is such that one's entire knowledge, experience, conceptions and also intuition even the day after such an event simply bear no relation whatsoever to those that went the day before.

So it was on the 6th of March 1869 at a meeting of the Russian Chemical Society in St Petersburg when a paper by Dmitri Mendeleev with the title ‘Relation of the Properties to the Atomic Weights of the Elements' was read to the audience by Nikolai Menshutkin, an associate of Mendeleev's (figure 1). Apparently the author himself was away on a trip to inspect the cheese-making procedures employed in the Russian countryside!

Figure 1.

A portrait of the 27-year-old Dmitri Mendeleev (Credit: Wikimedia Commons) and his 1869 handwritten version of the periodic law of the chemical elements (Credit: Science Photo Library).

Chemistry is not merely an immense collection of facts, but more an exact science that teaches us to classify and arrange these facts, and that classification must begin with the chemical elements themselves.

Thus, in 1869, Mendeleev's advance, unlike many attempts of his many predecessors, used two sets of data for a complete classification of the chemical elements, namely, elements' atomic weights and their inherent similarities in chemical properties [2]. With this epoch-making advance, the resulting periodic law of the chemical elements was born. Not only did Mendeleev show that a remarkable, natural periodicity existed in the chemical properties of the elements then known, but he also had the courage and the vision to state that this method of classification constituted a fundamental law of nature and identified gaps in the classification as then-undiscovered elements (figure 2).

Figure 2.

Mendeleev's 1870 original version of the periodic system. (Credit: Wikimedia Commons).

The resulting, modern periodic table of the elements is surely one of the most powerful—and easily recognizable—icons that pervade all of science [3] (figure 3).

Figure 3.

A representation of the modern periodic table of the chemical elements based here through a representation of the covalent radii of the elements. Remarkably, even without designating the chemical elements, this is instantly recognized as a representation of the periodic table. The metallic and non-metallic status of the chemical elements (as reflected in the shading of the spheres) is taken from G. T. Seaborg, Dalton Trans., 1996, 3899 [4]. Credit for this representation to Karl Harrison, University of Oxford [5]. (Online version in colour.)

This single representation, over one and a half centuries after its first appearance, still consolidates and represents so much of our modern knowledge of the world around us. Aside from providing a natural order to the chemical elements known at that time, Mendeleev's underpinning periodic law allowed for the prediction of the existence and remarkably the atomic order of chemical elements not then known, but discovered soon after.

Even today, nothing quite like the periodic table exists in any other disciplines of science. Reznick in this volume highlights the parallels—and also a significant difference—between the discovery and subsequent development of the periodic table and Darwin's discovery of evolution and the subsequent development of evolutionary biology [6].

To highlight the cascading impact of Mendeleev's 1869 advance, we reproduce here simply a single commentary, taken from a standard scientific text of that time; this is ‘Atomic Theory' by Ad Wurtz [1], published just over a decade after the first publication of this advance:

The work of Mendeleev has lately thrown a new light upon the relations existing between the atomic weights of the elements and their properties. The latter are a function of the atomic weight, which function is periodic. It is not limited to such and such a group of elements, but embraces all of the elementary bodies of chemistry …thus dealing with the most varied and the most profound questions of science…in a word, regard the facts of chemistry from a lofty and comprehensive point of view.

Mendeleev first challenged the world and then led us to confront how prepared were our minds to recognize an advance of sheer brilliance—a genuine seminal advance—which, quite simply, changed our world the day after its appearance in 1869.

Besides the discovery of the periodic law of the chemical elements, Mendeleev also made other critical contributions to chemical problems of broad scope. For example, it is perhaps not well recognized that Mendeleev studied the origins of petroleum and launched the idea of the so-called ‘abiotic origin’ hypothesis that hydrocarbons originated from iron carbides by the action of water vapour in the deep interior of the earth. He wrote: ‘The capital fact to note is that petroleum was born in the depths of the earth, and it is only there that we must seek its origin' [7]. Remarkably, Mendeleev's insights relating to abundant, highly-active iron based catalysts now come to the fore not only in modern studies of the utilization of CO₂ but also in their effective ‘hydrogen stripping' from fossil fuels, with the energy carrier hydrogen now advanced as a solution to the world's looming climate emergency, as highlighted in the article by Yao et al. [8].

2. The simple or natural elementary bodies (the situation pre-1869)

Dalton, still widely regarded as ‘the immortal author', recognized chemical combinations as being formed by the union or addition of elementary atoms, the relative weights of which he endeavoured to determine, referring those weights to one of the elements—hydrogen—as unity. Dalton revived and advanced the hypothesis of atoms to explain the fact that in chemical combinations, elements unite in fixed proportions, and in certain cases in multiple proportions [9].

Thus, prior to Mendeleev's 1869 advance, all substances were universally placed as belonging to only two classes—simple or elementary bodies or substances and compound substances. A listing taken from a textbook [10] just one year before Mendeleev's discovery vividly highlights the situation at that time (figure 4). This remarkable compilation reveals the totality of the world's then-known Elementary Bodies, together with their appropriate characteristics. At that stage, upon these 64 elementary bodies, the entire fabric and classification of the science—and indeed the application—of chemistry was based.

Figure 4.

The 1868 List of the Natural Elementary Bodies with their Symbols, Combining weights and Specific Gravities. The black letters indicate the non-metallic substances; the italics, the commonly occurring metals and the ordinary type, the rarer metals. (Taken from reference [10]).

3. The ‘lofty point of view' of Mendeleev (1869 and beyond)

Mendeleev's advance was both simple in its principle and stunningly productive in its results. In his Faraday Lecture of the Chemical Society of London presented at the Royal Institution on June 4th, 1889, Mendeleev announced that he could, even at that stage—note, only the twentieth anniversary of his discovery—conclude that the periodic law was now generally considered proven. At that lecture, he presented his conclusions relating to the periodic law of the chemical elements, noting, ‘It was in March, 1869, that I ventured to lay before the then youthful Russian Chemical Society the ideas upon the same subject' [11].

We reproduce here just four of his eight remarkable conclusions that one must surely still find startling in the sheer power and scale of that advance ….‘in the very words I (he) used' (at the Russian Chemical Society) …

‘1. The elements, if arranged according to their atomic weights, exhibit an evident periodicity of properties'.

‘5. The magnitude of the atomic weight determines the character of the element just as the magnitude of the molecule determines the character of a compound body'.

‘6. We must expect the discovery of many yet unknown elements…'

and

‘8. Certain characteristic properties of the elements can be foretold from their atomic weights’.

Mendeleev concluded, ‘Today, 20 years after the above conclusions were formulated, they may still be considered as expressing the essence of the now well-known periodic law'.

4. The contribution of Lothar Meyer

Among the characteristic properties dependent upon, and foretold from atomic weight, one finds, remarkably, even the characteristic density of an element. The fact that even a chemical element's density was subject to periodic variations with the increasing value of the atomic weights was first brilliantly highlighted by Lothar Meyer. For such a graphic and enduring construction, we must also remain indebted to Meyer who independently contributed a highly important development to the periodic law of the chemical elements [12].

Thus, if the chemical elements are arranged along the axis of the abscissae, at distances from zero and proportional to their atomic weights, each element thereby occupies a fixed point along the axis. If an ordinate is then drawn and placed on that the atomic volume of the given element, this graphic description reveals at once that the variations of the atomic volumes (and consequently of the density of an element) are periodic [13]. With this seminal advance, Meyer proved that the position occupied by the element on this curve is in relation to its physical property of density (figure 5).

Figure 5.

The relationships between the atomic weights of the elements and their atomic volumes (after Meyer). Taken from reference [1].

Thus, as far as the densities are concerned, it is evident from the very principle upon which the curve is constructed that the light chemical elements (having considerable atomic volumes) occupy the maxima and the heavy metals (possessing low atomic volumes) the minima. Thus, the alkali elements (Li, Na, K, Rb and Cs) make up the peaks of the curve while the alkaline earth metals (Be, Mg, Ca, Sr and Ba) are found on the descending slope and the halogens (F, Cl, Br and I) on the ascending part of the curve. The importance of an element's molar volume (directly related to its density) in dictating whether it is a metal or non-metal was first highlighted by D. A. Goldhammer in 1913 and independently by K. F. Herzfeld in 1927 [14,15]. Thus, Herzfeld noted, for example, if the element Ag had the large atomic volume of K, it would not be a metal. On the other hand, if the noble gas Xe had, in the solid state, the low molar volume of Cu, it would be a metal. The entire issue of the occurrence of metals and non-metals in the periodic table forms the basis of the contribution in this volume of Yao et al. [16]. It is salutary to note that Mendeleev wrote ‘The preparation of metallic sodium by Humphrey Davy was one of the greatest discoveries in chemistry', reflecting the fact that this was the first example of a ‘light' (i.e. low density, large molar volume) metallic element. Prior to this monumental discovery, the established elemental metals were heavy, high density elements, such as Au, Hg and Pb. Thus, as Mendeleev noted, ‘…through it (Davy's discovery) the conception of elements became broader and more correct'. Truly, one of the greatest discoveries in science.

Meyer further noted that many other physical properties of the elements, such as ductility, fusibility and volatility, are similarly related to their atomic weights and are subject to periodic variations with the increase of their atomic weights.

Here again the obvious periodicity in the variations of basic physical properties was a striking manifestation of Mendeleev's periodic law of the chemical elements. Thus, there can be no doubt that Meyer independently discovered the key, central principles of the periodic law of the chemical elements. This graphical representation of the atomic volume as a function of the atomic weight of the solid elements by Meyer also constituted a seminal advance.

In relation to the extent and impact of the contributions of both scientists in the formulation of the periodic law, we note that in 1882 the Royal Society of London awarded their gold Davy Medal, one of their highest distinctions, to Mendeleev and Meyer jointly, ‘For their discovery of the periodic relations of the atomic weights’ [17].

5. Life after the 1869 periodic law

After the publication of Mendeleev's seminal papers, the periodic table increasingly became part of the teaching of chemistry in universities throughout the world toward the end of the nineteenth century, and periodic table wallcharts became available commercially. The discovery of a periodic table chart dating from 1885 at the University of St Andrews, Scotland, thought to be the oldest in the world, is reviewed in the paper by Aitken and colleagues [18]. The relationship between pictorial representation of Mendeleev's ideas in his Table and the visual appeal of Charles Darwin's Tree is explored in the paper by Reznick [6].

In terms of theoretical advances, Mendeleev, of course, could not know about the existence of sub-atomic particles and therefore he arranged the periodic table strictly by atomic weight, which implied that something about atomic weight dictated the physical and chemical properties of the elements. Many decades passed before it was realized that it was the number and nature of electrons (and nucleons) that really determined the properties of elements.

The journey toward rationalizing the structure of the periodic table in terms of the electronic configurations of atoms began of course with the discovery of the electron by J. J. Thomson in 1897 [19]. In 1901, Jean Perrin first suggested a planetary model of the atom [20], with orbiting electrons, but Thomson himself favoured the so-called plum pudding model [21]. He was keenly interested in the link between atomic structure and chemical periodicity and in 1904 began to speculate as to how the arrangement of the plums within the pudding might be related to the chemistry of the elements [22]. These ideas never really developed into anything useful. Chemists Gilbert N. Lewis and Irving Langmuir also proposed static models with characteristic arrangements of electrons within the nucleus [23].

The most significant steps leading to our current understanding of chemical periodicity were all linked to the Manchester group of Ernest Rutherford, whose analysis of the α-particle scattering experiments performed by Geiger and Marsden led to his nuclear model of the atom [24], giving a more secure foundation to some of the ideas proposed by Perrin. Working as a demonstrator and research fellow within Rutherford's group, Henry Moseley went on to show that the charge on the atomic nucleus could be equated with the order number of an element in the periodic table [25,26]. Moseley's conclusions were based on the measurement of the frequencies of X-rays emitted from different elements under bombardment by cathode rays, as discussed in the paper by Egdell and Bruton in this volume [27]. In parallel with Moseley's experiments, the Danish physicist Niels Bohr, who was a regular visitor to Manchester, was working on his famous trilogy of papers published in 1913 [28]. In the first of these Bohr introduced the ideas of concentric circular rings that could be occupied by the single electron in a hydrogen atom, each ring having a characteristic quantized energy and angular momentum determined by a single quantum number n [29]. In his second paper, Bohr began to grapple with the occupancy of the different rings in atoms with more than one electron [30]. He had hoped that the screening constants in Moseley's formulae for the frequencies of K- and L-type X-rays would provide some guidance as to ring occupancy [31], but in the end, he proposed configurations for elements up to Z = 24 on a fairly ad hoc basis. Bohr continued to explore the relationship between periodicity and electronic structure, later incorporating a second or subsidiary quantum number k (this would now be labelled l) introduced by Sommerfeld based on consideration of elliptical orbits with differing eccentricities. In his 1922 Nobel Lecture Bohr summarized his work to date on periodicity and specified electron configurations for most elements up to Z = 90 (thorium) [23]. The next major contribution to electron configurations based on the ‘old' quantum theory of the Bohr–Sommerfeld model came from Edmund Stoner, who introduced the idea of a third quantum number j (this would now be called m_l) which could take on 2k+1 integer values ranging from –k to +k [32].

The final steps in the journey followed the introduction of the new quantum mechanics of Heisenberg and Schrödinger; the (approximate) solution of the Schrödinger equation for many-electron atoms, which provided a rigorous basis for the quantum numbers l and m_l and an understanding of why states with the same n but different l have different energies; and finally exposition of the Pauli exclusion principle, which allowed a maximum occupancy of two for each state specified by the quantum numbers n, l and m_l [33].

At the time of Bohr's Nobel Lecture in 1922, the heaviest known element was uranium with Z = 92. However, in his table of electron configurations, Bohr suggested (without comment as to why) an electron configuration for element 118, realizing that it should have an arrangement with eight electrons in the outer shell, similar to that of the noble gases with Z values of 10 (Ne), 18 (Ar), 36 (Kr) and 54 (Xe). There is something strikingly prescient about Bohr's inclusion of element 118. As discussed in the paper by Chapman in this volume [34], the periodic table now extends to element 118 (oganesson) and no further. This is thanks to several decades of laboratory based research into the nucleosynthesis of super heavy elements, culminating in the acceptance by IUPAC in 2015 of elements 113, 115, 117 and 118. In parallel with the synthesis of elements beyond uranium, there has been sustained effort over the past 50 years in exploring the effects of relativity on the electronic structure of heavy and superheavy elements. The paper by Pyper analyses the influence of relativistic effects on atomic properties and bonding and shows how the periodic trends exhibited by the lighter elements begin to change in the bottom row of the periodic table [35]. This in turn poses new challenges as to how best to present the periodic table when known superheavy elements and beyond are included, a matter discussed in detail in the paper by Scerri. This paper also deals with the thorny issue of how best to incorporate the lanthanides and actinides in the table [36].

Most of the new elements discussed by Chapman were prepared by carefully designed fusion reactions. The chapter by Johnson, Fields and Thompson considers the much broader topic of the origin of the naturally occurring elements [37]. The state of knowledge in this field is reviewed and the level of confidence in each of the proposed mechanisms is given. In particular, the role of neutron star mergers in producing the heaviest elements through the so-called r-process is critically assessed.

6. Concluding remarks

A ‘Modern periodic table’ (admittedly from 1996) is shown in figure 6 [4]; this is from an article from another of the ‘Greats', the chemist Glenn T Seaborg, who pioneered the study of the synthesis of the new chemical elements. Seaborg, the first living person to be honoured by the naming of a chemical element after him, synthesized 10 new elements heavier than uranium. Seaborg noted ‘The chemical elements are the building blocks of nature. All substances are combinations of these elements'. The periodic law represents perhaps the most decisive progress ever made in the application of theory to the subject of chemistry. The key element in Mendeleev's scheme was its basis in experimental (i.e. observed) patterns and trends in chemical properties—and not, as chemists and physicists unconsciously (and frequently) seem to assume, quantum mechanics.

Figure 6.

Modern (1996) periodic table (atomic numbers of undiscovered elements in parentheses). Taken from reference [4]. (Online version in colour.)

The periodic law of the chemical elements of Mendeleev, without question, is the most fundamental natural system of classification ever devised. It represents a triumph of one of the great organizing themes of science. We hope that the articles in this special issue are a noble tribute and accolade to that triumph, and genius, to the glorious name of Dmitri Mendeleev in establishing for the world the periodic law of the chemical elements.

Data accessibility

This article has no additional data.

Authors' contributions

P.P.E. worked with R.G.E., D.F. and B.Y. on conception and design of the article; P.P.E. worked on drafting the article, and all the authors worked on revising the article for important intellectual content. P.P.E. is the final approver of the version to be published.

Competing interests

We declare we have no competing interests.

Funding

We gratefully acknowledge the financial support from EPSRC (EP/N009924/1) and KACST, Saudi Arabia.