Definitions

periodic law

periodic law

[peer-ee-od-ik, peer-]
periodic law, statement of a periodic recurrence of chemical and physical properties of the elements when the elements are arranged in order of increasing atomic number. Such an arrangement in the form of a table in which the groupings of elements having similar properties are easily identified is called the periodic system or the periodic table.

Pioneering Periodic Arrangements of the Elements

Laws of Triads and of Octaves

Early in the 19th cent., a number of chemists had noticed certain relationships between the properties of elements and their atomic weight. In 1829 J. W. Döbereiner stated that there existed some three-element groups, or triads, in which the atomic weight of the middle element was the average of the other two and the properties of this element lay between those of the other two. For example, calcium, strontium, and barium form a triad; lithium, sodium, and potassium, another. The English chemist J. A. Newlands found (1863-65) that if the elements are listed according to atomic weight starting with the second, the 8th element following any given element has similar chemical properties, and so does the 16th. This became known as the law of octaves. About the same time, A. E. de Chancourtois arranged the elements according to increasing atomic weight in the form of a vertical helix with eight elements in a turn, so that elements having similar properties fell along vertical lines.

The Periodic Table

D. I. Mendeleev was the first to state the periodic law close to its present form. He proposed in 1869 that the properties of elements are periodic functions of the atomic weight and grouped the elements accordingly in a periodic system. Working independently and not aware of Mendeleev's work, Lothar Meyer arrived at a similar system, publishing his results about a year after Mendeleev's. When Mendeleev devised his periodic table a number of positions could not be fitted by any of the then known elements. Mendeleev suggested that these empty spaces represented undiscovered elements and by means of his system accurately predicted their general properties and atomic weights.

Introduction of Atomic Numbers

The work (1913-14) of H. G. Moseley on the X-ray spectra of elements (see X ray) led to the present form of the periodic law. He found that the wavelength of the X-radiation of elements decreased with increasing atomic weight. However, the relationship was not a strict one. He assigned a new set of numbers, called atomic numbers, to the elements he had studied, so that there was a relation between the wavelength and the atomic number. The atomic number is the number of positive charges, or protons, contained in the atomic nucleus (see atom) or, equivalently, the number of negative charges, or electrons, outside the nucleus in a neutral atom. The periodic law can be explained on the basis of the electronic structure of the atom, which is believed to be the main factor underlying the chemical properties and many of the physical properties of the elements. In turn, the electronic structures of atoms have been successfully accounted for by the quantum theory.

In spite of its great success, the periodic system that had been introduced by Mendeleev had some discrepancies. Arranged strictly according to atomic weight, not all elements fell into their proper groups. Better arrangement could be made if the positions of certain neighboring couples were interchanged. For example, to suit the chemical order of the table, the inert gas argon (at. wt. 39.948) should come before the chemically active metal potassium (at. wt. 39.0983). Through Moseley's work, it was found that although the atomic number of an element is roughly half its atomic weight, the atomic weight does not always increase with increasing atomic number. The discrepancies occur just for those elements where Mendeleev's law failed. Based on atomic number, the periodic law now has no exceptions. Although all the missing elements in the periodic table have been found (with the aid of the periodic table itself), the table retains its usefulness to the chemist as a reliable check for disputed or uncertain data concerning some of the known elements.

The periodic table is a tabular method of displaying the chemical elements. Although earlier precursors exist, its invention is generally credited to the Russian chemist Dmitri Mendeleev in 1869. The table is a visual representation of the periodic law which states that certain properties of elements repeat periodically when arranged by atomic number. The table arranges elements into vertical columns (Groups) and horizontal rows (Periods) to display these commonalities.

In the beginning

People have known about basic chemical elements such as gold, silver, and copper from antiquity, as these can all be discovered in nature in native form and are relatively simple to mine with primitive tools. Aristotle, a philosopher, theorised that everything is made up of a mixture of one or more of four elements. They were earth, water, air, and fire. This was more like the four states of matter (in the same order): solid, liquid, gas, and plasma, though he also theorised that they change into new substances to form what we see.

Hennig Brand was the first person recorded to have discovered a new element. Brand was a bankrupt German merchant who was trying to discover the Philosopher's Stone — a mythical object that was supposed to turn inexpensive base metals into gold. He experimented with distilling human urine until in 1649 he finally obtained a glowing white substance which he named phosphorus. He kept his discovery secret, until 1680 when Robert Boyle rediscovered it and it became public.

By 1869, a total of 63 elements had been discovered. As the number of known elements grew, scientists began to recognize patterns in the way chemicals reacted and began to devise ways to classify the elements.

Alexandre-Emile Béguyer de Chancourtois

Alexandre-Emile Béguyer de Chancourtois, a French geologist, was the first person to notice the periodicity of the elements — similar elements seem to occur at regular intervals when they are ordered by their atomic weights. He devised an early form of periodic table, which he called the telluric helix. With the elements arranged in a spiral on a cylinder by order of increasing atomic weight, de Chancourtois saw that elements with similar properties lined up vertically. His chart included some ions and compounds in addition to elements. His paper was published in 1862, but used geological rather than chemical terms and did not include a diagram; as a result, it received little attention until the work of Dmitri Mendeleev.

John Newlands

John Newlands was an English chemist who in 1865 classified the 56 elements that had been discovered at the time into 11 groups which were based on similar physical properties.

Newlands noted that many pairs of similar elements existed which differed by some multiple of eight in atomic weight. However, his law of octaves, likening this periodicity of eights to the musical scale, was ridiculed by his contemporaries. It was not until the following century, with Gilbert N. Lewis' valence bond theory (1916) and Irving Langmuir's octet theory of chemical bonding (1919) that the importance of the periodicity of eight would be accepted.

The first periodic table

Dmitri Mendeleev, also spelt Dmitry Mendeleyev, middle name (patronymic) Ivanovich, a Siberian-born Russian chemist, was the first scientist to make a periodic table much like the one we use today. Mendeleev arranged the elements in a table ordered by atomic mass. It is sometimes said that he played "chemical solitaire" on long train rides using cards with various facts of known elements. On March 6, 1869, a formal presentation was made to the Russian Chemical Society, entitled The Dependence Between the Properties of the Atomic Weights of the Elements. His table was published in an obscure Russian journal but quickly republished in a German journal, Zeitschrift für Chemie (Eng., "Chemistry Magazine"), in 1869. It stated:

  1. The elements, if arranged according to their atomic weights, exhibit an apparent periodicity of properties.
  2. Elements which are similar as regards to their chemical properties have atomic weights which are either of nearly the same value (e.g., Pt, Ir, Os) or which increase regularly (e.g., K, Rb, Cs).
  3. The arrangement of the elements, or of groups of elements in the order of their atomic weights, corresponds to their so-called valencies, as well as, to some extent, to their distinctive chemical properties; as is apparent among other series in that of Li, Be, Ba, C, N, O, and Sn.
  4. The elements which are the most widely diffused have small atomic weights.
  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–for example, elements analogous to aluminium and silicon–whose atomic weight would be between 65 and 75.
  7. The atomic weight of an element may sometimes be amended by a knowledge of those of its contiguous elements. Thus the atomic weight of tellurium must lie between 123 and 126, and cannot be 128.
  8. Certain characteristic properties of elements can be foretold from their atomic weights.

Advantages

  • Mendeleev predicted the discovery of other elements and left space for these new elements, namely eka-silicon (germanium), eka-aluminium (gallium), and eka-boron (scandium). Thus, there was no disturbance in the periodic table.
  • He pointed out that some of the then current atomic weights were incorrect.
  • He provided for variance from atomic weight order.

Disadvantages

  • There was no place for the isotopes of the various elements.
  • His table did not include any of the noble gases, which hadn't been discovered. But these were added by Sir William Ramsay as Group 0, without any disturbance to the basic concept of the periodic table.

Unknown to Mendeleev, Lothar Meyer was also working on a periodic table. In his work published in 1864, Meyer presented only 28 elements, classified not by atomic weight but by valence alone. Also, Meyer never came to the idea of predicting new elements and correcting atomic weights. Only a few months after Mendeleev published his periodic table of all known elements (and predicted several new elements to complete the table, plus some corrected atomic weights), Meyer published a virtually identical table. Some people consider Meyer and Mendeleev the cocreators of the periodic table, although most agree that Mendeleev's accurate prediction of the qualities of the undiscovered elements lands him the larger share of credit. In any case, at the time Mendeleev's predictions greatly impressed his contemporaries and were eventually found to be correct. An English chemist, William Odling, also drew up a table that is remarkably similar to that of Mendeleev in 1864.

Henry Moseley

In 1914 Henry Moseley found a relationship between an element's X-ray wavelength and its atomic number and therefore resequenced the table by electronic charge rather than atomic weight. Before this discovery, atomic numbers were just sequential numbers based on an element's atomic weight. Moseley's discovery showed that atomic numbers had an experimentally measurable basis.

Moseley's research also showed that there were gaps in his table at atomic numbers 43 and 61 which are now known to be Technitium and Promethium, respectively, both radioactive and not naturally occurring. Following in the footsteps of Dmitri Mendeleyev, Henry Moseley also predicted new elements.

Glenn T. Seaborg

During his Manhattan Project research in 1943 Glenn T. Seaborg experienced unexpected difficulty isolating Americium (95) and Curium (96). He began wondering if these elements more properly belonged to a different series which would explain why the expected chemical properties of the new elements were different. In 1945, he went against the advice of colleagues and proposed a significant change to Mendeleev's table: the actinoid series (previously called the actinide series).

Seaborg's actinide concept of heavy element electronic structure, predicting that the actinides form a transition series analogous to the rare earth series of lanthanide elements, is now well accepted in the scientific community and included in all standard configurations of the periodic table. The actinide series are the second row of the f-block (5f series) and comprise the elements from Actinium to Lawrencium. Seaborg's subsequent elaborations of the actinide concept theorized a series of superheavy elements in a transactinide series comprising elements 104 through 121 and a superactinide series inclusive of elements 122 through 153.

Curiosities

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