Periodic table of elements

The Periodic Table of the Elements is a tabular method of displaying the chemical elements. A chemical element is a fundamental classification of atomic matter where differentiation of particles is based on the number of protons found in their nuclei, the so-called atomic number. So far, 118 elements are known to exist. They are either produced artificially or found naturally in the environment.

The component parts of the table are known as Rows and Groups (the later referring to columns). In the table, elements are sorted in ascending order of atomic number. The physical state of matter that they exist in at a temperature of 273.15 K (0 °C) and at a pressure of 100,000 Pa (≈ 1 atm) is usually also provided.

Contents 1 Typical version 2 Elements by periodic table group (vertical column) 3 Elements by periodic table row 4 Elements classified alphabetically 5 History 6 References

[edit] Typical version

The table shown below is only one of many versions. A number of others — often structured in a different manner — are available on-line, for example ptable.com.

[edit] Elements by periodic table group (vertical column)

See Atomic electron configuration for the orbital occupancies of ground state atoms.

Elements in any one group behave in a similar way and show the same overall general properties. The number of electrons in the outer shells of the electron orbitals is the same within a group, only the principal quantum number n describing the orbitals increases. For instance, the right column is occupied by the noble gases. Because they have an outer shell of electrons that is completely filled, they are inert in behavior. (See the article atomic orbital for a more detailed explanation of the building up of the electronic shells.) 

On the extreme left there is a group of metals, called the alkali metals. They are characterized by having a single electron in their outer valence orbital. Going down from Li, Na, K to Fr the reactivity decreases, but these metals react easily – for example, with water, in a very exothermic way. This lowered reactivity is related to the atomic weight. As the mass of the nucleus increases with each neutron and proton, so also does the number of electrons (to balance the electric charge). Because these electrons are all at a greater distance from the nucleus, the energy gained from removing one electron diminishes according to the proportionality of n⁻². The one free electron in the outer valence shell is in a so-called ns-orbital. In all these metals there is only one free electron and these ball-shaped orbitals are denoted by 1s, 2s, 3s, etc., all with an increasing average radius from the nucleus.

For all the subsequent groups, characterizations can be formulated based upon similarities in reactivity—the way they react, how many electrons can be shared, etc. As a result, the likelihood of reactants to be able to react together can be determined from the Periodic Table of Elements.

The way the element at the top of a column reacts and the way it produces new chemicals give a clue to how other elements in that same group (column) will react. For instance carbon and hydrogen (H₂) react to form methane. The bonds of the carbon (C) take a tetrahedral or pyramidal shape. Silicon (Si) responds in the same way and silicon chemistry is to a significant extent similar to carbon chemistry.

[edit] Elements by periodic table row

Elements in a row show different periodicities – such as electronegativities, increasing atomic mass, increasing number of protons and neutrons. The stability of these elements diminishes however with an increasing number of nucleons (protons and neutrons). These instabilities lead to radioactive decay as is evident from the actinides and lanthanides, the groups of the rare earths. They contain elements such as plutonium, radon and uranium, and, notorious since the poisoning scandal in the UK, polonium.

The ratio of protons to neutrons is important in determining the stability of elements. If the ratio is greater than 1, the probability increases that an element (or an isotope of an element) will be radioactive and unstable. The greater the ratio, the greater the probability that the element will be radioactive.

Other properties vary differently from the column-wise view of the periodic table of elements. Going downward, the reactivity decreases. Going along a row from left to right, the reactivity increases, and properties such as electronegativity increase as well. For the latter reason, the tendency to lose an electron to form a covalent bond decreases and turns into the tendency of needing one or more electrons to form a stable covalent bond, producing a more stable chemical substance or molecule.

The alkali metals are the most eager to lose an electron (they have the lowest electronegativity in the row) making them very reactive with water. The halogens are reactive for the opposite reason – they need to attract that electron to obtain a more stable electron-configuration. hydrofluoric acid is among the strongest acids known to mankind, where sodium hydroxide is one of the strongest bases — as an example. Electronegativity also is the source for the occurrence of electrical polarity in substances. In every molecule of water (H₂O), the hydrogen is less inclined to keep its electron tightly bound. The oxygen, on the other hand, being more electronegative, likes to share electrons to create a more stable electron configuration. This results in the protons'(hydrogen) side of the molecule being slightly more positively charged ∆+, whereas the oxygen side is slightly more negatively charged 2∆−. Therefore water is a highly polar fluid with the capability of facilitating hydrogen bonding in solutions, which is a somewhat non-standard behaviour.

[edit] Elements classified alphabetically

See Chemical elements

[edit] History

In the early part of the 19th century it was known that matter was composed of a limited number of types of atoms or elementary substances ("elements")^[1] [2] and it was observed that many of those elements resembled one another. So it was only natural for chemists to search for a method of classification of the elements. Mass being a property common to all matter, and the relative masses of the elements being known, it was natural to seek some relationship between mass and properties. Early attempts at classification on this basis were hampered by the fact that a number of elements had not been discovered, and by a number of gross inaccuracies in the established atomic weights. In spite of these handicaps, the earlier attempts at classification were so obviously based on some fundamental truth that the modern chemist can only wonder at the negligible attention they attracted, or the ridicule with which they were received.

As early as 1829^[3] Johann Wolfgang Döbereiner (1780–1849) pointed out that there were "triads" of similar elements in which the atomic weight of one was the arithmetic mean of the other two, e.g., calcium (Ca), strontium (Sr), and barium (Ba); lithium (Li),sodium (Na), and potassium (K); chlorine (Cl), bromine (Br), and iodine (I). In 1850 Max Josef von Pettenkofer (1818–1901) developed this idea further by drawing attention to series of similar elements in which the atomic weight of any element differed from that of the next lighter or next heavier in the series by 8 or a multiple of 8, e.g., Li, Na, K; or magnesium (Mg), Sr, Ba; oxygen,sulfur,selenium, tellurium. Until then little notice was taken of such regularities.

In 1860 a chemical conference (the first ever) was held in Karlsruhe (Germany) where Stanislao Cannizzaro had presented a more accurate list of atomic weights than had previously been available. Not only had some values been slightly wrong through inaccurate measurements but some were half or a third of the correct value through false reasoning. The availability of these corrected weights inspired the introduction of the first "periodic table" of the now familiar type. It was drawn up by Alexandre-Emile Béguyer de Chancourtois in 1863, who arranged the then known elements in sixteen columns in order of atomic weight, and found that by doing so many similar elements were in the same column (group). Independently of Chancourtois, John A.R. Newlands was thinking along similar lines. He also arranged the known elements in order of atomic weight, and then noticed that in many cases there was a repetition of chemical properties at each 8th element; he called this the "Law of Octaves".

When Newlands described his ideas in a lecture at a Chemical Society meeting (see picture on the right) the lack of spaces for undiscovered elements and the placing of two elements in one box were criticized, and Professor G. F. Foster humorously inquired if he had examined the elements according to their initial letter. Apparently the idea of a periodic table was in the air, for not only Newlands drew up a table, but simultaneously William Odling, also in the UK, designed one. Both tables show already some of the now familiar characteristics of the periodic table.

While Newlands and Odling were developing their table in England, Lothar Meyer and Dmitri Mendeleev in Germany and Russia, respectively, were thinking on similar lines, and independently in 1868–1870 they published tables now known as the Mendeleev table. Mendeleev published his table in Russian, but an abstract appeared in German^[4]. In this work Mendeleev placed groups of similar elements in the same horizontal line and not, as later became the custom, in the same column. Meyer's table is known to have been in existence in manuscript form at least as early as 1868, but it was not published until 1870.

In publishing his table Mendeleev recognized that there must be a number of undiscovered elements, and he predicted the properties of some of these. The spectacular confirmation of these predictions by the discovery of gallium (1875) and scandium (1879) had excellent publicity value, and in 1882 the Royal Society of London awarded the Davy Medal jointly to Mendeleev and Lothar Meyer in recognition of their work, ignoring the claims of Newlands. At a later date (1887) they made some reparation to Newlands by awarding him the Davy Medal.

The Mendeleev classification proved to be extremely useful, but it contained some anomalies, that were resolved much later (1913) by Henry Moseley (1887–1915). By studying the X-ray spectra of a variety of elements, Moseley discovered^[5] that the frequencies of the K-lines differ from element to element in a predictable and consistent fashion. He found that the frequency is approximately proportional to (Z–b)², where Z is the atomic number of the element and b is an electric screening constant. Moseley was able to identify the atomic number Z with the total number of electrons in electrically neutral atoms. Remember that in a neutral atom, the sum of charges e of the electrons is by definition equal to minus the positive charge of the nucleus and hence Z is the nuclear charge expressed in units of |e|. In the X-ray spectrum the nuclear charge is partly screened by the electrons; this explains the presence of b. From Moseley's experiments, it became clear that the atomic number, not the atomic weight (averaged over isotopes), must govern the position of the atom in the periodic system. In fact, Moseley confirmed experimentally the earlier theoretical hypothesis of A. J. van den Broek that the classification of elements should be based on atomic number Z and not on atomic weight. Moseley's work revealed that some heavier isotopes were intruding in the table at the wrong places. Because atoms of different isotopic composition—but same nuclear charge Z—show the same chemical behavior, it is common today to arrange the elements in the periodic table by Z.

The electronic structure of the different atoms in the periodic table was first discussed in 1922 by Niels Bohr,^[6] who had incorporated Moseley's findings. The novel idea, that a chemical element is defined by its atomic number rather than by being non-separable by chemical means, was hard to swallow for the more conservative members of the chemistry community. However, an important argument, in favor of the atomic number definition, was Bohr's 1922 prediction of the properties of element number 72. From its electronic structure, Bohr predicted that the element must be in the same group (IV) as titanium and zirconium, i.e., be directly below zirconium in the periodic table. This meant that element 72 could not be a lanthanoid (aka a rare earth metal). Most chemists at the time held that element 72, whose properties were not yet established, would be a lanthanoid. In the same year 1922 coworkers of Bohr found the element 72 in zirconium minerals from unmistakable X-ray signatures; they could isolate enough of it to study some of its chemical properties, and they announced in January 1923 that element 72 had been discovered. The element was called hafnium, after the Latin name (Hafnia) of Copenhagen, and placed in group IV below zirconium.

The discovery caused some tension^[7] between physicists, who accepted Bohr's reasoning, and traditional chemists who believed that element 72 was a lanthanoid, tentatively called celtium. Only in 1930 did the the International Committee on Chemical Elements (a branch of IUPAC, the International Union of Pure and Applied Chemistry) recognize hafnium as the one and only element of atomic number 72. This was a great triumph for the "atomic" definition of chemical element and basically settled the dispute.

[edit] References

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