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The periodic table of the chemical elements is a tabular method of displaying the chemical elements, first devised by English analytical chemist John Newlands in 1863. The periodic table was revised a few years later by Russian Chemist Dmitri Mendeleev, after which it gained scientific acceptance. Newlands 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 new elements have been discovered, and new theoretical models have been developed to explain chemical behavior.
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 117 confirmed elements as of October 16, 2006 (while element 118 has been synthesized, element 117 has not).
- 1 Methods for displaying the periodic table
- 2 Arrangement
- 3 Periodicity of chemical properties
- 4 Structure of the periodic table
- 5 See also
- 6 Further resources
- 7 References
- 8 External links
Methods for displaying the periodic table[edit | edit source]
Standard periodic table[edit | edit source]
- Lanthanides are also 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.
Arrangement[edit | edit source]
The layout of the periodic table illustrates recurring ("periodic") chemical properties. Elements are listed in order of increasing atomic number (i.e. the number of protons in the atomic nucleus). Rows are arranged so that elements with similar properties fall into the same vertical columns ("groups"). According to quantum mechanical theories of electron configuration within atoms, each horizontal row ("period") in the table corresponded to the filling of a quantum shell of electrons. There are progressively longer periods further down the table, grouping the elements into s-, p-, d- and f-blocks to reflect 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.
Periodicity of chemical properties[edit | edit source]
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[edit | edit source]
- A group, is a vertical column in the periodic table of the elements.
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 (unsystematic) names, e.g. the alkali metals, alkaline earth 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), and these have no trivial names and are referred to simply by their group numbers.
- 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 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 lanthanides and actinides form two substantial horizontal series of elements.
Periodic trends of groups[edit | edit source]
Modern quantum mechanical theories of atomic structure explain group trends by proposing that elements within the same group have the same electron configurations in their valence shell, which is the most important factor in accounting for their similar properties. Elements in the same group also show patterns in their atomic radius, ionization energy, and electronegativity. From top to bottom in a group, the atomic radii of the elements increase. Since there are more filled energy levels, electrons are found farther from the nucleus. From the top, each successive element has a lower ionization energy because it is easier to remove an electron since the atoms are less tightly bound. Similarly, a group will also see a top to bottom decrease in electronegativity due to an increasing distance between valence electrons and the nucleus.
Periodic trends of periods[edit | edit source]
Elements in the same period show trends in atomic radius, ionization energy, electron affinity, and electronegativity. Moving left to right across a period, atomic radius usually decreases. This occurs because each successive element has an added proton and electron which causes the electron to be drawn closer to the nucleus. This decrease in atomic radius also causes the ionization energy to increase when moving from left to right across a period. The more tightly bound an element is, the more energy is required to remove an electron. Similarly, electronegativity will increase in the same manner as ionization energy because of the amount of pull that is exerted on the electrons by the nucleus. Electron affinity also shows a slight trend across a period. Metals (left side of a period) generally have a lower electron affinity than nonmetals (right side of a period) with the exception of the noble gases.
Examples[edit | edit source]
Noble gases[edit | edit source]
All the elements of Group 18, 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 much less reactive than other groups. 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 electron shells. However, their reactivity remains low in absolute terms.
Halogens[edit | edit source]
In Group 17, known as the halogens, elements are missing just one electron each to fill their shells. 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, for example, with regard to iodine and fluorine, since a large I- ion is more stable in solution than a small F-, there is less volume in which to disperse the charge.
Transition metals[edit | edit source]
For the transition metals (Groups 3 to 12), horizontal trends across periods are often important as well as vertical trends down groups; the differences between groups adjacent are usually not dramatic. Transition metal reactions often involve coordinated species.
Lanthanides and actinides[edit | edit source]
The chemical properties of the lanthanides (elements 57-71) and the actinides (elements 89-103) are even more similar to each other than the transition metals, and separating a mixture of these can be very difficult. This is important in the chemical purification of uranium concerning nuclear power.
Structure of the periodic table[edit | edit source]
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 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):
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.
See also[edit | edit source]
Further resources[edit | edit source]
- 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).
- Eric R. Scerri, The Periodic Table: Its Story and Its Significance, Oxford University Press, 2006.
References[edit | edit source]
- Theodore L. Brown, H. Eugene LeMay, and Bruce E. Bursten (2005). Chemistry:The Central Science, 10th edition, Prentice Hall. ISBN 0-13-109686-9.
- Helmenstine, Marie Trends in the Periodic Table. About, Inc.. URL accessed on 2007-01-27.
[edit | edit source]
- WebElements periodic table
- The IUPAC periodic table
- Visual Elements. ChemSoc.org.
- Stephen Hawking's Universe - 03: PBS documentary on the history of the periodic table and the cosmic evolution of the elements.
- Heilman, Chris, Alternate Layouts.
- Periodic table. Los Alamos National Laboratory's chemistry division.
- The Wooden Periodic Table Table: actual table containing samples of each naturally occurring element.
- The Periodic Library: downloadable periodic table with extra information, such as electron configuration and crystal structure
- Flash Periodic Table: interactive Flash based periodic table.
- The Periodic Table: informational and interactive periodic table
- The Extended Periodic Table - Analysis of the extended periodic table.
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