[edit] Methods for displaying the periodic table
[edit] Standard periodic table
Notes
- 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.
| |
| | Natural occurrence | Undiscovered | Synthetic | From decay | Primordial | |
[edit] Alternative versions (Layout/view of the table)
[edit] Arrangement
The layout of the periodic table demonstrates 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.
As of 2006, the table contains 117 chemical elements whose discoveries have been confirmed. Ninety-two are found naturally on Earth, and the rest are
synthetic elements that have been produced artificially in
particle accelerators. Elements 43 (technetium) and 61 (promethium), although of lower atomic number than the naturally occurring element 92, uranium, are synthetic; elements 93 (neptunium) and 94 (plutonium) are listed with the synthetic elements, but have been found in trace amounts on earth.
[edit] 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.
[edit] Groups and periods
- 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.
[edit] Periodic trends of groups
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.
[edit] Periodic trends of periods
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.
[edit] Examples
[edit] Noble gases
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.
[edit] Halogens
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.
[edit] Transition metals
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.
[edit] 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 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.
[edit] Structure of the periodic table
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):
| 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 |
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.
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.
[edit] History
-
In Ancient Greece, the influential Greek philosopher
Aristotle proposed that there were four main elements: air, fire, earth and water. All of these elements could be reacted to create another one;
e.g., earth and fire combined to form lava. However, this theory was dismissed when the real chemical elements started being discovered. Scientists needed an easily accessible, well organized database with which information about 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) |
| chlorine | 35.453 | 0.0032 |
| bromine | 79.904 | 3.1028 |
| iodine | 126.90447 | 4.933 |
| |
| calcium | 40.078 | 1.55 |
| strontium | 87.62 | 2.54 |
| barium | 137.327 | 3.594 |
In 1829 Döbereiner proposed the Law of Triads: The middle element in the triad had atomic weight that was the average of the other two members. The densities of some triads followed a similar pattern. Soon other scientists found chemical relationships extended beyond triads. Fluorine was added to Cl/Br/I group; sulfur, oxygen, selenium and tellurium were grouped into a family; nitrogen, phosphorus, arsenic, antimony, and bismuth were classified as another group.
This was followed by the English chemist
John Newlands, who noticed in 1865 that when placed in order of increasing atomic weight, elements of similar physical and chemical properties recurred at intervals of eight,
[citation needed] which he likened to the
octaves of music, though his
law of octaves was ridiculed by his contemporaries. However, while successful for some elements, Newlands' law of octaves failed for two reasons:
- It was not valid for elements that had atomic masses higher than Ca.
- When further elements were discovered, such as the noble gases (He, Ne, Ar), they could not be accommodated in his table.
Finally, in 1869 the Russian chemistry professor
Dmitri Ivanovich Mendeleev and four months 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.
Earlier attempts to list the elements to show the relationships between them (for example by
Newlands) 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 fit 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 net amount of positive charge on 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 the 1940s
Glenn T. Seaborg identified the
transuranic lanthanides and the actinides, which may be placed within the table, or below (as shown above). Element 106, seaborgium, is the only element that was named after a then living person.
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