The standard table provides the necessary
basics. There are also other methods for displaying the chemical
elements for more details or different perspectives.Standard
Color coding for atomic numbers
Series of the Periodic Table
- Elements numbered in blue are
liquids at standard temperature and pressure (STP);
- those in green are gases at STP;
- those in black are solid at STP;
- those in red are synthetic (all
are solid at STP).
- those in gray have not yet
been discovered (they also have muted fill colors indicating
the likely chemical series they would fall under).
A group is a vertical column in the periodic table of the
elements. There are 18 groups in the standard periodic table.
Elements in a group have similar configurations of their valence
shell electrons, which gives them similar properties.
There are three systems of group numbers; one using Arabic
numerals and the other two using Roman numerals. The Roman
numeral names are the original traditional names of the groups;
the Arabic numeral names are a newer naming scheme recommended
by International Union of Pure and Applied Chemistry (IUPAC).
The IUPAC scheme was developed to replace both older Roman
numeral systems as they confusingly used the same names to
mean different things.
Explanation of the structure of the periodic table
The number of electron shells
an atom has determines what period it belongs to. Each shell
is divided into different subshells, which as atomic number
increases are filled in roughly this order:
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p 8s 5g 6f 7d 8p ...
Hence the structure of the table. Since the outermost electrons
determine chemical properties, those tend to be similar within
groups. Elements adjacent to one another within a group have
similar physical properties, despite their significant differences
in mass. Elements adjacent
to one another within a period have similar mass but different
For example, very near to nitrogen (N) in the second period of the chart are carbon (C) and oxygen (O). Despite their similarities in mass (they differ
by only a few atomic mass units), they have extremely different
properties, as can be seen by looking at their allotropes:
diatomic oxygen is a gas that supports burning, diatomic nitrogen
is a gas that does not support burning, and carbon is a solid
which can be burnt (yes, diamonds
can be burnt!).
In contrast, very near to chlorine (Cl) in the next-to-last group in the chart (the
halogens) are fluorine (F) and bromine (Br). Despite their dramatic differences
in mass within the group, their allotropes have very similar
properties: They are all highly corrosive (meaning they combine
readily with metals to form metal halide salts); chlorine
and fluorine are gases, while bromine is a very low-boiling
liquid; chlorine and bromine at least are highly colored.
History of the periodic table
The original table was created without a knowledge of the
inner structure of atoms:
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 Alexander Reina
Newlands, who in 1865 noticed 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 Lothar
Meyer and the Russian chemist Dmitry 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.
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