The Photographic Periodic Table of the Elements

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Periodic Table of Elements

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Periodic table

Postby Yocage В» 22.01.2019


The periodic table , also known as the periodic table of elements , is a tabular display of the chemical elements , which are arranged by atomic number , electron configuration , and recurring chemical properties.

The structure of the table shows periodic trends. The seven rows of the table, called periods , generally have metals on the left and nonmetals on the right.

The columns, called groups , contain elements with similar chemical behaviours. Six groups have accepted names as well as assigned numbers: for example, group 17 elements are the halogens ; and group 18 are the noble gases. Also displayed are four simple rectangular areas or blocks associated with the filling of different atomic orbitals. The elements from atomic numbers 1 hydrogen through oganesson have been discovered or synthesized, completing seven full rows of the periodic table.

Numerous synthetic radionuclides of naturally occurring elements have also been produced in laboratories. The organization of the periodic table can be used to derive relationships between the various element properties, and also to predict chemical properties and behaviours of undiscovered or newly synthesized elements. Russian chemist Dmitri Mendeleev published the first recognizable periodic table in , developed mainly to illustrate periodic trends of the then-known elements.

He also predicted some properties of unidentified elements that were expected to fill gaps within the table. Most of his forecasts proved to be correct. Mendeleev's idea has been slowly expanded and refined with the discovery or synthesis of further new elements and the development of new theoretical models to explain chemical behaviour.

The modern periodic table now provides a useful framework for analyzing chemical reactions , and continues to be widely used in chemistry , nuclear physics and other sciences.

Each chemical element has a unique atomic number Z representing the number of protons in its nucleus. For example, carbon has three naturally occurring isotopes: all of its atoms have six protons and most have six neutrons as well, but about one per cent have seven neutrons, and a very small fraction have eight neutrons. Isotopes are never separated in the periodic table; they are always grouped together under a single element.

Elements with no stable isotopes have the atomic masses of their most stable isotopes, where such masses are shown, listed in parentheses. In the standard periodic table, the elements are listed in order of increasing atomic number Z the number of protons in the nucleus of an atom. A new row period is started when a new electron shell has its first electron. Columns groups are determined by the electron configuration of the atom; elements with the same number of electrons in a particular subshell fall into the same columns e.

Elements with similar chemical properties generally fall into the same group in the periodic table, although in the f-block, and to some respect in the d-block, the elements in the same period tend to have similar properties, as well. Thus, it is relatively easy to predict the chemical properties of an element if one knows the properties of the elements around it.

Since , the periodic table has confirmed elements, from element 1 hydrogen to oganesson. The first 94 elements occur naturally; the remaining 24, americium to oganesson 95— , occur only when synthesized in laboratories. Of the 94 naturally occurring elements, 83 are primordial and 11 occur only in decay chains of primordial elements. A group or family is a vertical column in the periodic table.

Groups usually have more significant periodic trends than periods and blocks, explained below. Modern quantum mechanical theories of atomic structure explain group trends by proposing that elements within the same group generally have the same electron configurations in their valence shell.

Under an international naming convention, the groups are numbered numerically from 1 to 18 from the leftmost column the alkali metals to the rightmost column the noble gases. In America, the roman numerals were followed by either an "A" if the group was in the s- or p-block , or a "B" if the group was in the d-block. The roman numerals used correspond to the last digit of today's naming convention e. In Europe, the lettering was similar, except that "A" was used if the group was before group 10 , and "B" was used for groups including and after group In addition, groups 8, 9 and 10 used to be treated as one triple-sized group, known collectively in both notations as group VIII.

Some of these groups have been given trivial unsystematic names , as seen in the table below, although some are rarely used.

Elements in the same group tend to show patterns in 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, valence 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 has a top-to-bottom decrease in electronegativity due to an increasing distance between valence electrons and the nucleus. A period is a horizontal row in the periodic table. Although groups generally have more significant periodic trends, there are regions where horizontal trends are more significant than vertical group trends, such as the f-block, where the lanthanides and actinides form two substantial horizontal series of elements.

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.

The more tightly bound an element is, the more energy is required to remove an electron. Electronegativity increases in the same manner as ionization energy because of the pull exerted on the electrons by the nucleus. 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. Specific regions of the periodic table can be referred to as blocks in recognition of the sequence in which the electron shells of the elements are filled.

Each block is named according to the subshell in which the "last" electron notionally resides. The f-block , often offset below the rest of the periodic table, has no group numbers and comprises lanthanides and actinides. According to their shared physical and chemical properties, the elements can be classified into the major categories of metals , metalloids and nonmetals.

Metals are generally shiny, highly conducting solids that form alloys with one another and salt-like ionic compounds with nonmetals other than noble gases.

A majority of nonmetals are coloured or colourless insulating gases; nonmetals that form compounds with other nonmetals feature covalent bonding.

In between metals and nonmetals are metalloids, which have intermediate or mixed properties. Metal and nonmetals can be further classified into subcategories that show a gradation from metallic to non-metallic properties, when going left to right in the rows. The metals may be subdivided into the highly reactive alkali metals, through the less reactive alkaline earth metals, lanthanides and actinides, via the archetypal transition metals, and ending in the physically and chemically weak post-transition metals.

Nonmetals may be simply subdivided into the polyatomic nonmetals , being nearer to the metalloids and show some incipient metallic character; the essentially nonmetallic diatomic nonmetals , nonmetallic and the almost completely inert, monatomic noble gases. Specialized groupings such as refractory metals and noble metals , are examples of subsets of transition metals, also known [30] and occasionally denoted. Placing elements into categories and subcategories based just on shared properties is imperfect.

There is a large disparity of properties within each category with notable overlaps at the boundaries, as is the case with most classification schemes. Radon is classified as a nonmetallic noble gas yet has some cationic chemistry that is characteristic of metals.

Other classification schemes are possible such as the division of the elements into mineralogical occurrence categories , or crystalline structures. Categorizing the elements in this fashion dates back to at least when Hinrichs [33] wrote that simple boundary lines could be placed on the periodic table to show elements having shared properties, such as metals, nonmetals, or gaseous elements.

The electron configuration or organisation of electrons orbiting neutral atoms shows a recurring pattern or periodicity. The electrons occupy a series of electron shells numbered 1, 2, and so on. Each shell consists of one or more subshells named s, p, d, f and g. As atomic number increases, electrons progressively fill these shells and subshells more or less according to the Madelung rule or energy ordering rule, as shown in the diagram.

The electron configuration for neon , for example, is 1s 2 2s 2 2p 6. With an atomic number of ten, neon has two electrons in the first shell, and eight electrons in the second shell; there are two electrons in the s subshell and six in the p subshell.

In periodic table terms, the first time an electron occupies a new shell corresponds to the start of each new period, these positions being occupied by hydrogen and the alkali metals. Since the properties of an element are mostly determined by its electron configuration, the properties of the elements likewise show recurring patterns or periodic behaviour, some examples of which are shown in the diagrams below for atomic radii, ionization energy and electron affinity. It is this periodicity of properties, manifestations of which were noticed well before the underlying theory was developed , that led to the establishment of the periodic law the properties of the elements recur at varying intervals and the formulation of the first periodic tables.

Atomic radii vary in a predictable and explainable manner across the periodic table. For instance, the radii generally decrease along each period of the table, from the alkali metals to the noble gases ; and increase down each group. The radius increases sharply between the noble gas at the end of each period and the alkali metal at the beginning of the next period.

These trends of the atomic radii and of various other chemical and physical properties of the elements can be explained by the electron shell theory of the atom; they provided important evidence for the development and confirmation of quantum theory. The electrons in the 4f-subshell, which is progressively filled across the lanthanide series, are not particularly effective at shielding the increasing nuclear charge from the sub-shells further out. The elements immediately following the lanthanides have atomic radii that are smaller than would be expected and that are almost identical to the atomic radii of the elements immediately above them.

This is known as the lanthanide contraction. The effect of the lanthanide contraction is noticeable up to platinum element 78 , after which it is masked by a relativistic effect known as the inert pair effect. The first ionization energy is the energy it takes to remove one electron from an atom, the second ionization energy is the energy it takes to remove a second electron from the atom, and so on.

For a given atom, successive ionization energies increase with the degree of ionization. Electrons in the closer orbitals experience greater forces of electrostatic attraction; thus, their removal requires increasingly more energy. Ionization energy becomes greater up and to the right of the periodic table. Large jumps in the successive molar ionization energies occur when removing an electron from a noble gas complete electron shell configuration. Similar jumps occur in the ionization energies of other third-row atoms.

Electronegativity is the tendency of an atom to attract a shared pair of electrons. The higher its electronegativity, the more an element attracts electrons. It was first proposed by Linus Pauling in Hence, fluorine is the most electronegative of the elements, [n 5] while caesium is the least, at least of those elements for which substantial data is available. There are some exceptions to this general rule. Gallium and germanium have higher electronegativities than aluminium and silicon respectively because of the d-block contraction.

Elements of the fourth period immediately after the first row of the transition metals have unusually small atomic radii because the 3d-electrons are not effective at shielding the increased nuclear charge, and smaller atomic size correlates with higher electronegativity.

The electron affinity of an atom is the amount of energy released when an electron is added to a neutral atom to form a negative ion.

Although electron affinity varies greatly, some patterns emerge. Generally, nonmetals have more positive electron affinity values than metals. Chlorine most strongly attracts an extra electron. The electron affinities of the noble gases have not been measured conclusively, so they may or may not have slightly negative values.

The Periodic Table Song (2018 UPDATE!), time: 3:05
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Re: periodic table

Postby Shalar В» 22.01.2019

Atomic weight worked well enough to allow Mendeleev to accurately predict the properties of elements. Adams perioduc the rare earths and the "radioactive elements" i. The Atom in the History of Human Thought. New York: Nova Science.

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Re: periodic table

Postby Nikonris В» 22.01.2019

Http:// York: W. Seaborg first published his table in a classified report dated Nonmetals may be simply subdivided into the polyatomic nonmetalsbeing nearer to the metalloids and show some incipient metallic character; the essentially nonmetallic diatomic nonmetalsnonmetallic and the almost completely inert, monatomic noble gases.

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