Periodic table

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The periodic table of the chemical elements is a tabular method of displaying the chemical elements, first devised in 1869 by the Russian chemist Dmitri Mendeleev. Mendeleev intended the table to illustrate recurring ("periodic") trends in the properties of the elements. The layout of the table has been refined and extended over time, as many new elements have been discovered, and new theoretical models have been developed to explain chemical behaviour. Various different layouts are possible, to emphasize different aspects of behaviour; the most common forms, however, are still quite similar to Mendeleev's original. The periodic table is now ubiquitous within the academic discipline of chemistry, providing an extremely useful framework to classify, systematize and compare all the many different forms of chemical behaviour. The table has also found wide application in physics, biology, engineering, and industry.

Arrangement

Earlier attempts to list the elements had usually simply put them in order of atomic mass. Mendeleev's key insight in devising the periodic table was to lay out the elements to illustrate recurring ("periodic") chemical properties (even if this meant some of them were not in mass order), and to leave gaps for "missing" elements. Mendeleev used his table to predict the properties of these "missing elements", and many of them were indeed discovered and fitted the predictions well.

With the development of theories of atomic structure (for instance by Henry Moseley) it became apparent that Mendeleev had listed the elements in order of increasing atomic number (i.e. the number of protons in the atomic nucleus).

In order to illustrate recurring properties, Mendeleev began new rows in his table so that elements with similar properties fell into the same vertical columns ("groups").

With the development of modern quantum mechanical theories of electron configuration within atoms, it became apparent that each horizontal row ("period") in the table corresponded to the filling of a quantum shell of electrons. In Mendeleev's original table, each period was the same length. Modern tables have progressively longer periods further down the table, and group the elements into s-, p-, d- and f-blocks to reflect our understanding of their electron configuration.

In printed tables, each element is usually listed with its element symbol and atomic number; many versions of the table also include the element's atomic mass and other information, such as its abbreviated electron configuration, electronegativity and most common valence numbers. As of 2005, the table contains 116 chemical elements whose discoveries have been confirmed. 94 are found naturally on Earth, and the rest are synthetic elements that have been produced artificially in particle accelerators.

Periodicity of chemical properties

The main value of the periodic table is the ability to predict the chemical properties of an element based on its position in the table. It should be noted that the properties vary differently when moving vertically along the columns of the table, than when moving horizontally along the rows.

Groups and periods

Groups

• A group, also known as a family, is a vertical column in the periodic table of the elements.

Groups are considered the most important way of classifying the elements. In some groups, the elements have very similar properties and exhibit a clear trend in properties down the group — these groups tend to be given trivial (non-scientific) names, e.g. the alkali metals, halogens and noble gases. Some other groups in the periodic table display fewer similarities and/or vertical trends (for example Groups 14 and 15). Modern quantum mechanical theories of atomic structure explain that elements within the same group have the same electron configurations in their valence shell, which is the largest factor in accounting for their similar chemical properties.

Periods

• A period is a horizontal row in the periodic table of the elements.

Although groups are the most common way of classifying elements, there are some regions of the period table where the horizontal trends and similarities in properties are more significant than vertical group trends. This can be true in the d-block (or "transition metals"), and especially for the f-block, where the lanthanoids and actinoids form two substantial horizontal series of elements. The period number also shows how many electron shells there are in an element.

Examples

Noble gases

All the elements of group 18 (8 or 0 if discounting transition elements), the noble gases, have full valence shells. This means they do not need to react with other elements to attain a full shell, and are therefore unreactive, monoatomic gases. Helium is the most inert element among noble gases, since reactivity, in this group, increases with the periods: it is possible to make heavy noble gases react since they have much larger electronic shells. However, their reactivity remains low in absolute terms.

Halogens

In group 17, known as the halogens, elements are missing just one electron to fill their shell. Therefore, in chemical reactions they tend to acquire electrons (the tendency to acquire electrons is called electronegativity). This property is most evident for fluorine (the most electronegative element of the whole table), and it diminishes with increasing period.

As a result, all halogens form acids with hydrogen, such as hydrofluoric acid, hydrochloric acid, hydrobromic acid and hydroiodic acid, all in the form HX. Their acidity increases with higher period, since a large I^- ion is more stable in solution than a small F^-, that has less volume in which to disperse the charge.

Transition metals

In transition metals (groups 3 to 12, see transition metal), the differences between groups are usually not dramatic, and the reactions involve coordinated species. However, it is still possible to make useful predictions.

Lanthanides and actinides

The chemical properties of the lanthanides (elements 57-71) and the actinides (elements 89-103) are even more similar to each other than in transition metals, and separating a mixture of these can be very difficult. This is important in the chemical purification of uranium, important for nuclear power.

Methods for displaying the periodic table

Standard periodic table

Notes

• Actinides and lanthanides are collectively known as "rare earth elements", a deprecated term. Regarding group membership of these elements, see here.
• Alkali metals, alkaline earth metals, transition metals, actinides, lanthanides, and poor metals are all collectively known as "metals".
• Halogens and noble gases are also non-metals.

State at standard temperature and pressure (0°C and 1 atm)

• those numbered in Template:Element color/Gas are gaseous
• those numbered in Template:Element color/Liquid are liquid
• those numbered in Template:Element color/Solid are solid

Natural occurrence

• those with solid borders are primordial elements, which have stable isotopes older than the Earth
• those with dashed borders arise naturally from decay of other elements, and have no isotopes with a half-life comparable to the age of the Earth; however, some are found in trace amounts in radioactive ores
• those with dotted borders are synthetic elements, which do not occur naturally
  • Note: Although californium (Cf, 98) does not occur naturally on Earth, Cf and its decay products occur in the Universe; their electromagnetic emissions are regularly observed in supernova spectra
• those without borders have not been discovered or synthesised yet

Other depictions

• The standard table (same as above) provides the basics.
• A vertical table for improved readability in web browsers.
• The big table provides the basics and full element names.
• The huge table provides the basics plus full element names and atomic masses.
• A table with an inline F-block inserts the lanthanides and actinides into their correct place in the table.
• Electron configurations
• Metals and non-metals
• Periodic table filled by blocks
• Table in Chinese
• List of elements by name
• List of elements by symbol
• List of elements by atomic number
• List of elements by boiling point
• List of elements by melting point
• List of elements by density
• List of elements by atomic mass
• List of elements by electronegativity

Other alternative periodic tables exist.

Periodic table structure reflects electron configuration

The primary determinant of an element's chemical properties is its electron configuration, particularly the valence shell electrons. For instance, any atoms with four valence electrons occuping p orbitals will exhibit with some similarity. The type of orbital in which the atom's outermost electrons reside determines the "block" to which it belongs. The number of valence shell electrons determines which family, or group, the element belongs.

The total number of electron shells an atom has determines the period to which it belongs. Each shell is divided into different subshells, which as atomic number increases are filled in roughly this order (the Aufbau principle):

Hence the structure of the table. Since the outermost electrons determine chemical properties, those with the same number of valence electrons are grouped together.

Progressing through a group from lightest element to heaviest element, the outer-shell electrons (those most readily accessible for participation in chemical reactions) are all in the same type of orbital, with a similar shape, but with increasingly higher energy and average distance from the nucleus. For instance, the outer-shell (or "valence") electrons of the first group, headed by hydrogen all have one electron in an s orbital. In hydrogen, that s orbital is in the lowest possible energy state of any atom, the first-shell orbital (and represented by hydrogen's position in the first period of the table). In francium, the heaviest element of the group, the outer-shell electron is in the seventh-shell orbital, significantly further out on average from the nucleus than those electrons filling all the shells below it in energy. As another example, both carbon and lead have four electrons in their outer shell orbitals.

Note that as atomic number (i.e. charge on the atomic nucleus) increases, this leads to greater spin-orbit coupling between the nucleus and the electrons, reducing the validity of the quantum mechanical orbital approximation model, which considers each atomic orbital as a separate entity.

Because of the importance of the outermost shell, the different regions of the periodic table are sometimes referred to as periodic table blocks, named according to the sub-shell in which the "last" electron resides, e.g. the s-block, the p-block, the d-block, etc.

History

Main article: History of the periodic table

The original table was created before the discovery of subatomic particles or the formulation of current quantum mechanical theories of atomic structure. If one orders the elements by atomic mass, and then plots certain other properties against atomic mass, one sees an undulation or periodicity to these properties as a function of atomic mass. The first to recognize these regularities was the German chemist Johann Wolfgang Döbereiner who, in 1829, noticed a number of triads of similar elements:

This was followed by the English chemist John Newlands, who noticed in 1865 that the elements of similar type recurred at intervals of eight, which he likened to the octaves of music, though his law of octaves was ridiculed by his contemporaries. Finally, in 1869, the German Julius Lothar Meyer and the Russian chemistry professor Dmitri Ivanovich Mendeleev almost simultaneously developed the first periodic table, arranging the elements by mass. However, Mendeleev plotted a few elements out of strict mass sequence in order to make a better match to the properties of their neighbours in the table, corrected mistakes in the values of several atomic masses, and predicted the existence and properties of a few new elements in the empty cells of his table. Mendeleev was later vindicated by the discovery of the electronic structure of the elements in the late 19th and early 20th century.

In the 1940s Glenn T. Seaborg identified the transuranic lanthanides and the actinides, which may be placed within the table, or below (as shown above).

Further resources

• [1] Scerri, E.R., references to several scholarly articles by this author.
• Mazurs, E.G., "Graphical Representations of the Periodic System During One Hundred Years". University of Alabama Press, Alabama. 1974.
• Bouma, J., "An Application-Oriented Periodic Table of the Elements". J. Chem. Ed., 66 741 (1989).

See also

• Atomic electron configuration table
• Isotope table (complete)
• Isotope table (divided)
• Discoveries of the chemical elements
• Abundance of the chemical elements
• Tom Lehrer's song The Elements
• IUPAC's systematic element names
• Cosmochemical Periodic Table of the Elements in the Solar System
• Table of chemical elements

References

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External links

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• "WebElements periodic table
• The IUPAC periodic table
• "Visual Elements". ChemSoc.org.
• Another table with pictures and phase diagrams.
• Counterman, Craig, "Periodic Table of the Elements : For each of many properties a separate periodic table and a graph showing the relation with the atomic number". MIT Course 3.091.
• Heilman, Chris, "Alternate Layouts".
• "Periodic table". Los Alamos National Laboratory's Chemistry Division.
• Dayah, Michael, "Periodic Table". Large full-color scalable text-only rendering.
• Magnetic resonance table
• Colorful table with pictures of usages - Printable PDF with complete data and understandable pictures of usages
• List of periodic tables.

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