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Periodic Table

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See the standard periodic table below.
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 since Mendeleev's time, and new theoretical models have been developed to explain chemical behavior. Various different layouts are possible to emphasize different aspects of behavior; the most common forms, however, are still quite similar to Mendeleev's original design.

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 behavior. The table has also found wide application in physics, biology, engineering, and industry. The current standard table contains 116 confirmed elements, and two further speculated elements.

Contents

Arrangement

Earlier attempts to list the elements to show the relationships between them had usually involved putting 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). This sequence is nearly identical to that resulting from ascending atomic mass.

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 list 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 location on 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 and Periods are the two main ways to view the Periodic Table of the Elements. Take Group 2 for example, all of the elements in that group have similar characteristics as the others.

Groups

Groups are considered the most important method 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, alkaline earth metals, transition 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

Although groups are the most common way of classifying elements, there are some regions of the periodic 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. 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, (7 if discounting transition metals) 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

Other depictions

Other alternative periodic tables exist as well.

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 occupying 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 the family, or group, to which 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):

Subshell: S G F D P
Period
1 1s
2 2s 2p
3 3s 3p
4 4s 3d 4p
5 5s 4d 5p
6 6s 4f 5d 6p
7 7s 5f 6d 7p
8 8s 5g 6f 7d 8p
9 9s 6g 7f 8d 9p

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

In Ancient Greece, it was believed that there were four elements. These comprised Air, Fire, Earth and Water. All of these elements could be reacted to create another one...eg. earth and fire made air. However, this theory was dismissed when the real chemical elements started being discovered. Scientists needed an easily accessible, well organized database through which the elements could be recorded and accessed. This was to be known as 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:

 Some triads
Element Molar mass
(g/mol) 		Density
(g/cm³) 		Quotient
(cm³/mol)
chlorine 35.4527 0.003214 11030
bromine 79.904 3.122 25.6
iodine 126.90447 4.93 25.7
 
calcium 40.078 1.54 26.0
strontium 87.62 2.64 33.2
barium 137.327 3.594 38.2

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 Russian chemistry professor Dmitri Ivanovich Mendeleev and four month later the German Julius Lothar Meyer independently 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 neighbors 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

See also

References

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External links

Periodic tables
Standard table | Vertical table | Table with names | Names and atomic masses (large) | Names and atomic masses (small) | Names and atomic masses (text only) | Inline F-block | Elements to 218 | Electron configurations | Metals and nonmetals | Table by blocks
Lists of elements
Name | Atomic symbol | Atomic number | Boiling point | Melting point | Density | Atomic mass
Groups:   1 -  2 -  3 -  4 -  5 -  6 -  7 -  8 -  9 - 10 - 11 - 12 - 13 - 14 - 15 - 16 - 17 - 18
Periods:  1  -  2  -  3  -  4  -  5  -  6  -  7  -  8  -  9
Series:   Alkalis  -  Alkaline earths  -  Lanthanides  -  Actinides  -  Transition metals  -  Poor metals  -  Metalloids  -  Nonmetals  -  Halogens  -  Noble gases
Blocks:  s-block  -  p-block  -  d-block  -  f-block  -  g-block

 

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