Chemistry

Periodic table of the elements

Periodic table of the elements



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Atomic radii

The size of the atoms or ions plays a role in physical and chemical properties such as their incorporation in crystal lattice gaps or possible coordination numbers (number of neighboring atoms).

The sizes or radii of atoms are experimentally accessible as

  • covalent atomic radii
  • metallic atomic radii
  • Van der Waals radii.

Covalent and metallic atomic radii are obtained from the X-ray crystal structures of the elements by halving the shortest distances between the nuclei. The larger van der Waals radii result from the core distances of identical atomic neighbors that are not chemically bound.

The atomic radii increase within a group, since with increasing atomic number the heavier homologues also occupy more and therefore further outward electron shells.

The atomic radius decreases from left to right within a period. This observation, which is astonishing at first glance, can be understood if one considers that the atomic number increases steadily in a period, but the electrons that are added do not restart the shells. As a result, the attraction to the electron shell increases and the atomic radius shrinks. The atomic radius is the result of the attraction of the electrons by the nucleus on the one hand and the mutual repulsion of the electrons on the other.


Periodic Table of the Elements - Chemistry and Physics


In school, the elements of main groups I., II., VII. And VIII. Are discussed in more detail, as these form so-called element families, i.e. within these main groups a clear similarity of the individual elements can be ascertained. But first of all there are two exceptions, namely

a) hydrogen : as a gas and non-metal, it does not go well with the other alkali metals
b) Astatine : as a semi-metal it has very little in common with the other halogens of the element family

In chemistry, the letters k, l, m, n, o, p and q are assigned to the respective electron shells.

  • The majority of all elements are metals, since subgroup elements that are not taken into account in the simplified PSE above are exclusively metals.
  • The main group indicates the number of electrons on the most lucrative shell (valence electrons). Only these electrons are involved in chemical reactions.
  • The period, on the other hand, indicates how many occupied electron shells a particular atom has.


We now take a look at the elements of the first 3 periods.
Here we consider period by period, i.e. line by line.

Did you notice something?
Then you should be able to answer the following questions off the cuff:

a) How many protons does the 28th element of the PSE have? How many electrons does it have?
b) What is the name of the element with atomic mass 12.0, which has 6 electrons?
c) Which element has 20 protons?
d) In which main group is the 15th element of the PSE?

Lithium has 3 protons according to the atomic number and because of the difference
[Atomic mass - atomic number (7-3 = 4)] 4 neutrons in the nucleus.
In the atom, the number of protons is always the same as the number of electrons.

  • The atomic number indicates the number of protons and electrons. Electrons have & # 8222 no & # 8220 mass
  • Electrons are located on electron shells at certain distances
    the atomic nucleus. The protons, on the other hand, are found in the atomic nucleus.
  • The neutron number results from the difference between atomic mass and atomic number.
  • Hydrogen is the only element that does not have a neutron
  • Within a period, the number of protons and electrons increases by 1 from left to right
  • The number of neutrons increases from the 1st to the 7th period disproportionately to the number of protons

The 1st and 2nd periods
(see also menu item atomic structure & amp; chemical bond, & acute Bohr's atomic model & acute)

Chemistry worksheets for the classroom:

The elements of the periodic table - a hardware store exploration (RAABE Fachverlag)

Chemistry worksheets to download.
Chemistry teaching material 8.-9. Great

You get into the unit with a hardware store excursion. Worksheet M 1 contains the work order and tips for exploring a hardware store.

The exploration can take place as a class excursion or as a homework assignment, depending on the possibility. It is advantageous if the station groups explore the hardware store together, but this is not always possible depending on the catchment area of ​​the school. In the latter case, however, sufficient lead time must be allowed for.


The periodic table of the elements, which most chemistry books depict, is a special case. Because this tabular overview of the chemical elements, which goes back to Dmitri Mendeleev and Lothar Meyer, and the approaches of other chemists to organize the elements, are different forms of representation of a hidden structure of the chemical elements. This is the conclusion reached by researchers from the Leipzig Max Planck Institute for Mathematics in the Natural Sciences and the University of Leipzig in a current paper. The mathematical approach of the Leipzig scientists is very general and, depending on the order and classification principle, can provide many different periodic systems - not only for chemistry, but also for many other fields of knowledge.

The periodic table of the elements is celebrating its 150th birthday this year. The tabular overview is closely associated with the names of Dmitri Mendeleev and Lothar Meyer - two researchers who in the 1860s created an arrangement of elements based on their atomic masses and similarities. Today they are sorted according to the atomic number, which indicates the number of protons in the atomic nucleus - from light hydrogen (one proton) to the exotic oganesson (118 protons). In addition, the elements are classified into groups: Atoms in the same column usually have the same number of electrons in the outer shell of their electron shell.

At first glance, the periodic table seems to have brought a clear and definitive order into the currently known 118 elements. But appearances are deceptive, because some things are still controversial today: Scientists do not agree on exactly which elements belong to the third group below scandium and yttrium. For example, the correct position of lanthanum and actinium is discussed. If you take a closer look, you will discover slightly different variants of the periodic table in classrooms, lecture halls and textbooks.

Guillermo Restrepo and Wilmer Leal, who work at the Max Planck Institute for Mathematics in the Natural Sciences and at the University of Leipzig, are not surprised. For them there is no clearly correct arrangement of the elements, because depending on the criterion used for the classification, there is a different periodic table. For example, atoms can be subdivided according to their electron configuration, i.e. the number and arrangement of their electrons, their chemical behavior, their solubility or their occurrence in geological deposits. Today it has become widely accepted to line up the chemical elements according to their atomic number and to divide them into groups according to their electronic configuration. But even from this periodic table there are numerous different forms of representation, for example as a spiral with various more or less large bulges, pyramid-shaped or as a three-dimensional flower.

Guillermo Restrepo and Wilmer Leal have now systematically investigated the ambiguity of the periodic table. In doing so, findings have been made that are also of importance beyond chemistry. Accordingly, all forms of representation of the chemical elements are based on a common structure, which mathematicians call an ordered hypergraph. The venerable periodic table of Mendeleev and Meyer thus only offers a representation of the general structure that Guillermo Restrepo and Wilmer Leal now postulate. New arrangements can be derived from this at any time. Guillermo Restrepo therefore compares the order of the chemical elements with a sculpture on which light falls from different directions. “The various shadows that the figure casts are the period tables. That is why there are so many ways to set up these tables. In a sense, the periodic tables are projections. Projections of the internal structure of the periodic table. "

The Leipzig scientists are now trying to determine the hidden mathematical structure on which the well-known periodic tables of chemistry are based. In the meantime, you have defined three conditions that must be met in order to set up a periodic table. First of all, you need objects that should be sorted: For Mendeleev, Meier and the creators of the other known periodic tables of chemistry, these are the chemical elements. It must be possible to organize these objects on the basis of a property, such as atomic mass or atomic number. Finally, a criterion is needed to group the objects in classes. Mendeleev and Meier used chemical similarity for this.

"If these three conditions are met, periodic tables can also be created for other chemical objects and even for objects outside of chemistry," says Guillermo Restrepo. He and Wilmer Leal show this by looking at the chemical bonds between atoms of 94 elements and different partners as objects. They arrange these according to the electronegativity of the element under consideration and its atomic radius in this bond. Fluorine, chlorine or oxygen, for example, are very electronegative and take on relatively small atomic radii in compounds. They then classify the bonds based on whether they are similar.

"We examined almost 5000 substances that consist of two elements in different proportions," explains Guillermo Restrepo. “Then we looked for similarities in these data. For example, sodium and lithium are similar because they combine with the same elements in the same proportions - for example with oxygen or chlorine, bromine and iodine. So we found patterns with which the elements can be classified. "

The 44 classes of chemical compounds have some similarities with the main groups of Mendeleev's and Meier's periodic table. For example, the alkali metals sodium and lithium are in one group because they form the same simple salts with halogens such as chlorine or fluorine. The bonds of the four halogens fluorine, chlorine, bromine and iodine are also found in a group like the elements themselves. But there are also classifications that differ significantly from that in the conventional periodic table. For example, carbon and silicon are no longer in one class because they form very different compounds.

The representation of the periodic table of chemical bonds no longer has anything to do with the familiar matrix-like arrangement of the classical periodic tables of the elements. Instead, the 94 covalent bonds are represented in a network of differently colored circles, with each circle standing for a chemical bond and the color symbolizing membership of one of the 44 groups. Since two criteria are now used for sorting, there is no longer a clear sequence of atoms, as was the case with Mendeleev and Meyer - mathematicians speak of a partial order. The circles are therefore connected to other circles by one or more arrows, creating an ordered hypergraph.

The chemical elements and their compounds can also be represented in completely different periodic tables - depending on the underlying order and classification principle. And what's more: the objects of numerous other sciences and their applications can also be arranged in periodic tables. Ordered hypergraphs are used, for example, in information systems and in web mining. Possible periodic systems also arise when looking at states that can be ordered according to social or economic indicators and classified according to geographical proximity or cultural similarity. Other examples can be found in engineering, environmental sciences, sociology, and many other disciplines. The scientists study periodic systems not only out of an interest in chemistry, but above all because of their applications in many other disciplines.


Table of contents

Basic principle edit

A periodic table is a systematic tabular compilation of the chemical elements in which the elements are arranged according to two principles: On the one hand, they are arranged according to increasing atomic number (i.e. the number of protons in the atomic nucleus that is unique and characteristic for each element). On the other hand, the representation is chosen so that elements with similar chemical behavior are close together. With increasing atomic number, the properties of the elements resemble each other in regular, albeit differently long, periodic intervals. [3] The term "periodic table" indicates that these periodicities are represented by the selected arrangement of the elements.

Edit representation

There are different variants of periodic tables. The best-known representation arranges the elements in a two-dimensional tabular grid grid, taking into account the periodicities, in which a grid box corresponds to each element. The horizontal lines of the display are referred to as periods, the vertical columns as groups.

Within each period the atomic number of the elements increases from left to right. The line breaks are chosen so that chemically similar elements are each in the same column (group). The elements of a group therefore show similar chemical behavior. For example, there is the group of chemically inert noble gases or the group of reactive halogens.

The periods have different lengths. The first period comprises only two elements. This is followed by two periods with eight elements each, two further periods with eighteen elements each and finally two periods with thirty-two elements each.

The long form of the periodic table, in which the last two periods are displayed as continuous lines, is often unfavorable because of the required width of the display. In the mostly used medium-length form [4], element groups cut out of these periods are shown below the main system in a space-saving manner. In this form the periodic table has seven periods and eighteen groups. There is also an even more compact but seldom used short form of the periodic table.

Edit information content

Usually the elements are listed with their atomic number and their element symbol. [5] Depending on the application, further information on the element such as full name, mass, melting temperature, density and physical state can be given. Any information on "shells" relates to the shell model of atomic physics. Color codings are often used to represent different properties, for example the belonging to the metals, semi-metals or non-metals.

The peculiarity of the periodic table compared to a mere tabular listing of element properties, however, lies in the information about the relationships between the elements, which results from the placement of the elements in question. The fact that an element belongs to a certain group immediately allows conclusions to be drawn about the essential chemical characteristics of the element, such as its reactivity or preferred binding partners. The positioning within the overall system allows conclusions to be drawn regarding those properties that show a systematic trend in the periodic table, such as the ionization energy.

Edit scope

With the most recent expansion of the periodic table in 2015, the elements 1 (hydrogen) to 118 (oganesson) have now been completely discovered or created and described. [6] [7] In nature, the elements with atomic numbers 1 to 94 occur, with technetium (atomic number 43), promethium (61), astatine (85), neptunium (93) and plutonium (94) in such small amounts of course it can happen that they were first artificially created and described. [8] Of these 94 natural elements, 83 are primordial, meaning that they have existed since the earth was formed. The original stocks of the remaining 11 have long since decayed because of their shorter half-lives, but they are constantly being rebuilt through radioactive decay in the natural decay series of the primordial elements. [9]

The elements of the ordinal numbers 95 to 118 were created exclusively artificially. [9] The last discovered elements 113, 115, 117 and 118 were confirmed by the IUPAC on December 30, 2015, with which the seventh period of the periodic table is now complete. [10]

Images of the respective elements can be found in the table of the chemical elements.

Structure of an atom edit

All substances are made up of atoms. An atom consists of protons and neutrons, which form the atomic nucleus, and of electrons, which surround the atomic nucleus as an "electron shell". The protons each carry a positive and the electrons a negative elementary charge, so the number of electrons in the electron shell must be equal to the number of protons in the atomic nucleus if the atom is to be electrically neutral. The number of protons or electrons of an electrically neutral atom is called its "atomic number".

Chemical compounds are substances that are made up of two or more types of atoms. The atoms combine to form molecules. The binding forces that hold the atoms together in a molecule are mediated by interactions between the electrons. The decisive factor for the properties of the binding forces are mainly the properties of the electrons in the outer area of ​​the shell, the valence electrons.

The chemical behavior of an atom - for example its tendency to preferentially form compounds with certain other types of atoms - is therefore largely determined by the structure of the electron shell and especially the valence electrons.This structure is always the same for a given number of electrons, so that the atomic number determines the chemical behavior of the atom.

Atoms with the same atomic number and therefore the same behavior in chemical reactions are called chemical elements. [11] In the periodic table, all existing elements are arranged in such a way that the laws resulting from the structure of the atoms in the chemical and atomic physical properties of the elements can be recognized.

Structure of the electron shell edit

The electron shell of an atom has structures that are examined and described by quantum mechanics. It can be divided into main shells [12]. Each main shell can in turn be subdivided into subshells [13], which consist of orbitals [14]. The quantum mechanical state in which a given electron is located is described by four quantum numbers: [15] The main quantum number, the secondary quantum number, the magnetic quantum number and the spin quantum number.

The principal quantum number n = 1, 2, 3, ... numbers the main shells. Alternatively, these trays can be used as KBowl (for n = 1), L.Bowl (for n = 2), M.Bowl (for n = 3) and so on. The diameter of the main shells increases as the main quantum number increases.

A main shell with the main quantum number n owns n Subshells that differ in their secondary quantum number. The lower shells are with the letters s, p, d, f and so on [13] (the choice of these letters is historically determined). A given subshell in a particular main shell is identified by its letter preceded by the main quantum number, for example 2p for the p- Lower shell in the L.-Peel (n = 2).

The individual subshells are divided into orbitals, which are differentiated by the magnetic quantum number. Every s- Lower shell contains an orbital, each p- Lower shell contains three orbitals, each d- Lower shell contains five orbitals and each f- Lower shell contains seven orbitals. [13]

The spin quantum number describes the two possible spin orientations of the electron.

The Pauline Exclusion Principle says that no two electrons in an atom can match in all four quantum numbers. [15] Two electrons that are in the same orbital already agree in three quantum numbers (namely those that describe this orbital). The two electrons must therefore differ in the fourth quantum number, their spin alignment. This exhausts the possible variations for the quantum numbers in this orbital, so each individual orbital can be occupied by a maximum of two electrons. [15] The following maximum electron numbers result for the various shells:

  • the K-Peel (n = 1) has only one lower shell (1s) and this only a single orbital. Since this can be filled with a maximum of two electrons, the K-Shell a maximum of two electrons.
  • the L.-Peel (n = 2) has two lower shells 2s and 2p, which consist of one or three orbitals. It can hold a maximum of eight electrons in its four orbitals.
  • the M.-Peel (n = 3) has three lower shells 3s, 3p and 3d, can hold a maximum of 18 electrons in its nine orbitals.
  • the N-Peel (n = 4) can be in its four lower shells 4s until 4f take up a maximum of 32 electrons and so on.

In general, a shell with the principal quantum number n a maximum of 2n Pick up 2 electrons. [16]

In general, electrons that are on a higher main shell are more energetic than electrons on shells further inside. The main shells can, however, overlap energetically, since the energy of the lower shells is in the sense within a main shell spdf increases [17] and more energetic subshells of a given main shell can have a higher energy than the lowest energy subshells of the next main shell. [17] This has consequences for the systematic structure of the periodic table.

The diagram opposite shows a schematic representation, not to scale, of the energy levels in the electron shell. The lines on the left symbolize the main shells, the lines on the right their lower shells. The boxes represent the orbitals in each subshell, each of which can be assigned two electrons (“spin up” and “spin down”). From the main shell n= 3 the sub-shells of successive main shells overlap energetically.

If one imagines the atoms of the various elements to be generated one after the other in such a way that a proton in the nucleus and an electron in the shell (as well as the neutrons required, if applicable) are added to an atom of the previous element, then the added electron always occupies the lowest energy of the remaining elements free orbitals (“construction principle”). Since the occupation pattern of the individual orbitals repeats itself with the beginning of each new shell during the successive filling, the structures of the valence electrons repeat themselves and thereby the chemical properties of the atoms.

For the sake of clarity, in the following text each element name is preceded by its ordinal number as an index. The color of the element boxes indicates the shell that is currently being refilled.

First period: 1Hydrogen up to 2Helium edit

The simplest atom is that 1Hydrogen atom with a proton in the nucleus and an electron in the shell (there are also isotopes with one or two neutrons). The electron is in the s- Lower shell of the K-Peel.

It follows that 2Helium atom with two protons (as well as one or two neutrons) and two electrons. The added electron occupies the free space in the only orbital of the s- lower shell. So that is K- Bowl exhausted and filled the first period of the periodic table. [18]

Second period: 3Lithium up to 10Neon edit

The filling up of the begins with the next electron L.-Peel: 3Lithium has an electron in it 2s-Orbital, 4Beryllium has a second electron in the 2s-Orbital, which is completely filled with it.

Now the filling of the 2p-Orbitals: 5Boron has in addition to the filled 2s-Orbital an electron in 2p-Orbital. Following 6Carbon, 7Nitrogen, 8Oxygen, 9Fluorine and 10Neon. With these eight elements is also the L.- Bowl completely filled and the second period ended. [18]

Third period: 11Sodium up 18Argon edit

The filling up of the M.- Shell starts with the same pattern. [18] When considering the respective configurations of the valence electrons it becomes clear that, for example, the first element of this period (11Sodium, with one valence electron) chemical similarities to the first element of the previous period (3Lithium, also with a valence electron) will have.

Fourth period: 19Potassium up to 36Krypton edit

19
K
20
Approx

21
Sc
22
Ti
23
V
24
Cr
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn

31
Ga
32
Ge
33
As
34
Se
35
Br
36
Kr
4s- 3d- lower shell 4p- lower shell

After the eighth element of the third period, the 18Argon, however, there is an interruption in the regularity. Until then, the 3s- and 3p- Undershells of the M.-Shell filled, there are ten places left in theirs 3d- Lower shell free. However, since that 4s-Orbital of the next higher shell (N, n = 4) has a lower energy than that 3d-Orbitals of the M.Bowl, first this will be 4s-Orbital filled with two electrons (19Potassium, 20Calcium). That 19Potassium has a valence electron and is therefore chemically similar to 11Sodium and 3Lithium. Since the periodic table is intended to highlight these and other similarities, the 19Potassium started a new period.

Only after 19Potassium and 20Calcium becomes the 3d- Lower shell of the M.- Bowl filled, this is done from 21Scandium by 30Zinc. [19] These elements "inserted" in the periodic table all have a filled one 4s- Lower shell and only differ in the degree of filling of the one below M.-Peel. They therefore show only relatively minor chemical differences; they belong to the "transition metals". With the 30Zinc is the one M.- The bowl is now completely filled, the rest of the ones can then be refilled NBowl with the elements 31Gallium up 36Krypton. [18]

Fifth period: 37Rubidium to 54Xenon edit

37
Rb
38
Sr

39
Y
40
Zr
41
Nb
42
Mon
43
Tc
44
Ru
45
Rh
46
Pd
47
Ag
48
CD

49
In
50
Sn
51
Sb
52
Te
53
I.
54
Xe
5s- 4d- lower shell 5p- lower shell

The filling up of the N-Shell will however after the 36Krypton interrupted again. With the 36Krypton is that 4p- Lower shell completed, and there are still the lower shells 4d and 4f to fill. Again, however, the s- Lower shell of the next higher shell (O, n = 5) has a lower energy and is preferably filled up (37Rubidium, 38Strontium), with which you can start a new period again. Then come the ten transition metals 39Yttrium to 48Cadmium, with which the remaining 4d- Lower shell is filled [19] and then the six elements 49Indium to 54Xenon with which the 5p- Lower shell is filled. [19]

Sixth period: 55Cesium to 86Radon edit

55
Cs
56
Ba

57
La

72
Hf
73
Ta
74
W.
75
re
76
Os
77
Ir
78
Pt
79
Au
80
Ed

81
Tl
82
Pb
83
Bi
84
Po
85
At
86
Marg
6s- 5d- 5d- lower shell 6p- lower shell
58
Ce
59
Pr
60
Nd
61
Pm
62
Sm
63
Eu
64
Gd
65
Tb
66
Dy
67
Ho
68
He
69
Tm
70
Yb
71
Lu
4f- lower shell

This scheme, which is determined by the energetic position of the respective lower shells, is repeated in the following periods. In the sixth period, the following sub-shells are filled one after the other: 6s (55Cesium and 56Barium), 5d (57Lanthanum), 4f (58Cerium up 71Lutetium), 5d (72Hafnium up to 80Mercury) and 6p (81Thallium to 86Radon). [19]

In the diagram above, the padding is the 4f- Lower shell shown as an inset to limit the width of the diagram.

Seventh period: 87Francium to 118Oganesson edit

87
Fr.
88
Ra

89
Ac

104
Rf
105
Db
106
Sg
107
Bra
108
Hs
109
Mt
110
Ds
111
Rg
112
Cn

113
Nh
114
Fl
115
Mc
116
Lv
117
Ts
118
Above
7s- 6d- 6d- lower shell 7p- lower shell
90
Th
91
Pa
92
U
93
Np
94
Pooh
95
At the
96
Cm
97
Bk
98
Cf
99
It
100
Fm
101
Md
102
No
103
Lr
5f- lower shell

In the seventh period the following are filled: 7s (87Francium and 88Radium), 6d (89Actinium), 5f (90Thorium up 103Lawrencium), 6d (104Rutherfordium to 112Copernicium) and 7p (113Nihonium up 118Oganesson). [19]

For the sake of simplicity, some irregularities when filling the individual lower shells are not shown here. While the 4d-Shell moves for example one of the elements in some elements s-Electrons in the d- lower shell. So has that 47Silver does not have two electrons in the as expected 5sLower shell and nine electrons in the 4d- Lower shell, but only one 5s-Electron and ten for it 4d-Electrons. A list of these exceptions can be found in the article on the structure principle.

In summary, the following filling pattern results (shown in the long form of the periodic table):

Filling the shell
1 H Hey K (n = 1)
2 Li Be B. C. N O F. No L. (n = 2)
3 N / A Mg Al Si P. S. Cl Ar M. (n = 3)
4 K Approx Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr N (n = 4)
5 Rb Sr Y Zr Nb Mon Tc Ru Rh Pd Ag CD In Sn Sb Te I. Xe O (n = 5)
6 Cs Ba La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho He Tm Yb Lu Hf Ta W. re Os Ir Pt Au Ed Tl Pb Bi Po At Marg P. (n = 6)
7 Fr. Ra Ac Th Pa U Np Pooh At the Cm Bk Cf It Fm Md No Lr Rf Db Sg Bra Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Above Q (n = 7)

In representations of the periodic table, the periods are usually numbered consecutively with Arabic numerals from one to seven. The period number is also the main quantum number of the outermost main shell covered with electrons. [18]

A main shell can while it is the outermost, only containing up to eight electrons (the K-Shell: up to two). The next electron added creates a new main shell, which now becomes the new outermost one. The main shell considered is only the second outermost, third outermost, and so on during its further filling. Regardless of the capacity of its outermost shell, every element only has between one and eight valence electrons.

The systematic structure of the periodic table described above was done in such a way that the elements were arranged in the order of ascending ordinal numbers and a new line ("period") was started with certain elements. The criterion for the beginning of a new period was not the physical criterion of the degree of filling of the respective main shell, but the chemical similarity to the elements above it from the previous period, i.e. the same number of valence electrons. From this follows the structure of the periodic table, which is designed to make these relationships visible. The following division of the periodic table results into different blocks:

Main groups edit

In the first two columns ("groups") of the periodic table, the two orbitals are the s- Lower shell of the current main shell filled up (s-Block). In the last six groups there will be six p- Lower shells of the current main shell filled up (p-Block). These eight groups are the main groups of the periodic table. The number of valence electrons increases by one from one main group to the next. For the 50 main group elements, the number of their valence electrons and thus their chemical behavior in its essential features is immediately apparent from their group membership. If the material properties of the elements are determined by the valence electrons, there are many similarities in the elements of the same group. The group number, usually written in Roman numerals, is also the number of electrons in the respective outermost main shell [18] (with the exception of the 2Helium, which is a noble gas in main group VIII, but has only two electrons).

The elements of the first main group each have a valence electron. With the exception of the 1Hydrogen) around soft, silvery-white and very reactive metals, [20] the alkali metals. An example of the chemical similarity of the alkali metals is the fact that they all have 17Chlorine reigns to form colorless salts that crystallize in cube form. The formulas of these compounds also correspond to one another: LiCl, NaCl, KCl, RbCl, CsCl and FrCl. [21]

The alkaline earth metals follow as the second main group. The boron group is the third main group, the carbon group the fourth and the nitrogen group the fifth. The chalcogens represent the sixth main group and the halogens the seventh.

As can be justified quantum mechanically, not only closed main shells, but also closed lower shells are particularly stable. The elements in the eighth main group all have a closed main or lower shell: At 2Helium is the first main shell and therefore also its only subshell 1s completed. With the other elements 10Neon, 18Argon, 36Krypton, 54Xenon and 86Radon is always - if the main shell is not yet complete - the p- Lower shell completed, these elements have eight valence electrons (one octet). Because of the stability of their valence electron structures, these elements form almost no chemical bonds. They are all gaseous and are called noble gases.

Other elements can also reach noble gas shells and thus particularly stable states by releasing or absorbing electrons. The alkali metals give up their single valence electron easily [22] and then appear as monovalent cations, for example 3Li +, 11Na + etc. [22] The alkaline earth metals reach the noble gas configuration by releasing their two valence electrons and then form divalent cations, for example 4Be ++, 12Mg ++ etc. [22] The halogens, on the other hand, lack an electron to complete an octet. They therefore preferentially accept one electron, the result is the monovalent anions 9F -, 17Cl - etc. [22]

A block with subgroups is inserted between the main groups:

Subgroups: Edit outer transition metals

In the last four periods, the filling of the respective outermost main shell was interrupted by the d-Fill the lower shell of the second outermost main shell. the d-Subshells each hold 10 electrons, so there is an additional block with 10 groups in these four periods. All 40 elements in this one d-Block are metals, [23] they are called "outer transition metals" [23].They all have two valence electrons in the outermost shell (for exceptions see → construction principle) and therefore show fewer differences in their chemical behavior than the main group elements. The existing differences are due to the different electronic structures of the next lower main shell. Corresponding to the repetitive filling pattern, the elements in this block also show clear similarities in their chemical properties. [23]

Subgroups: Edit Inner Transition Metals

In the last two periods, the d-Subshells of the second outermost main shell interrupted by the filling of the f- Lower shells of the third outermost main shell. the f-Subshells each hold fourteen electrons, so there is an additional block with 14 groups in these two periods. The 28 elements in this fBlock are saved as inner transition elements designated. They have two valence electrons in the outermost main shell, one electron in the d- Lower shell of the penultimate main shell and only differ in the degree of filling of the third from last main shell (for exceptions see → construction principle). Their chemical differences are correspondingly small. [24]

The on that 57Lanthanum following 14 inner transition metals from 58Cerium up 71Lutetium in the sixth period are also called lanthanoids. The on that 89Actinium following 14 inner transition metals of 90Thorium up 103Lawrencium in the seventh period are also called actinides.

Some properties of the elements vary in a systematic way with their position in the periodic table. If you move from one main group to the next within a period (“from left to right”), the physical and chemical properties change in a systematic, characteristic way, because the number of valence electrons increases by one at a time. [25] In the next period, if they are determined by the number of valence electrons, the properties repeat in a similar way, because the number of valence electrons increases again in the same way. [25]

If you go from one period to the next within a main group (“from top to bottom”), the properties in question are usually similar (same number of valence electrons), but gradually different (different main shells as the outermost shell). [25]

Atomic radius edit

The atomic radius generally decreases within a period from left to right, [26] because the electrons are drawn closer and closer to the nucleus due to the increasing atomic number. During the transition to the next period, the atomic radius increases sharply again, because the occupation of the next outer main shell begins.

Within a group, the radius usually increases from top to bottom, [26] because a main shell is added in each case.

First ionization energy edit

The "first ionization energy" is the energy that has to be expended to remove an electron from the electron shell so that the neutral atom becomes a simply positively charged ion. The individual valence electron of the alkali metals is particularly loosely bound and can be easily detached. As the atomic number progresses within a period, an ever greater ionization energy has to be expended until it reaches the maximum value of the period in the case of the noble gas with its particularly stable octet configuration.

Electron affinity edit

The electron affinity is the binding energy that is released when an atom binds an additional electron to itself, so that the neutral atom becomes a simply negatively charged ion. [22] The halogens have a particularly high electron affinity, [22] because they can complete their electron octet by accepting an electron.

Electronegativity edit

If two atoms of different elements are chemically bonded to each other, one of the two usually attracts the electrons of the common electron shell more strongly, so that the center of gravity of the electron shell shifts towards this atom. The ability of an atom to attract electrons in a bond is measured by its electronegativity.

The electronegativity of the main group elements increases from left to right within a period because the nuclear charge increases. [27] Within a group it usually grows from bottom to top, because in this direction the number of occupied main shells decreases [27] and with it the shielding of the nuclear charge by the internal electrons. The element with the smallest electronegativity (0.7 according to Pauling) [28] is the one at the bottom left of the periodic table 55Cesium. The element with the greatest electronegativity (4.0 according to Pauling) [28] is the one on the top right 9Fluor, followed by its left neighbor, the 8Oxygen (3.5). [28] 1Hydrogen and the semi-metals occupy a middle position with values ​​around 2. [28] Most metals have values ​​around 1.7 or less. [28]

The shift in the center of gravity of the charge in the molecule depends on the difference in the electronegativities of the two atoms. [28] The more the center of gravity of the charge is shifted, the greater the ionic part of the bond, [28] because the electrostatic attraction of the two unlike partial charges contributes to the bond. The ionic bond character is particularly pronounced because of the described tendency towards electronegativities in bonds in which one bond partner is on the left and the other on the right in the periodic table. [29] An example of this is sodium chloride N a C l < displaystyle mathrm > .

Bonds in which both partners come from the left half of the periodic table and therefore both belong to the metals (see below) are metallic bonds. [30] Bonds in which both partners originate from the right side are mainly covalent bonds. [29]

Value edit

One of the most characteristic features of an element is its valence, [31] that is, its property to combine with certain preferred numbers of atoms of the various partner elements when a chemical compound is formed.

An atom that still lacks an electron to complete a valence electron octet can bond with a single 1Hydrogen atom enter in order to use the single valence electron of hydrogen in the common electron shell to complete its own octet. An atom two electrons missing will tend to bond with two 1Hydrogen atoms enter into. As these examples show, a connection between the preferred number of binding partners and the structure of the valence electron shell - i.e. the group membership in the periodic table - is to be expected in general. However, the relationships are often much more complex than in the examples shown here.

A simple measure of the value of an element is the number of 1Hydrogen atoms that binds the element to itself in a binary hydride. [31] Another possible measure is twice the number of 8Oxygen atoms that the element binds in its oxide. [31]

The elements of the first and penultimate main group (the alkali metals or halogens) have the valence 1, so their hydrides have the formulas [32]

The elements of the second and third from last main group (the alkaline earth metals and the oxygen group) generally have the valence 2, so their hydrides are [32]

In the other main groups the bonding possibilities become more diverse (so there are countless hydrocarbon compounds), but one also encounters A s H 3 < displaystyle textstyle mathrm in the nitrogen group, for example >> or S b H 3 < displaystyle textstyle mathrm >> and in the carbon group on P b H 4 < displaystyle textstyle mathrm > > . [32]

Of the 8Oxygen is bivalent, so typical oxides of the monovalent alkali metals are [33]

and typical oxides of the divalent alkaline earth metals are [34]

but there are also other oxidation states. The last three oxides mentioned were the starting point for Döbereiner's triad system (see below).

Basicity edit

The basicity of the oxides and hydroxides of the elements increases from top to bottom and decreases from left to right. Oxides and hydroxides of metals dissolved in water (see below) form alkalis, while oxides and hydroxides of non-metals dissolved in water form acids. [35]

Calcium oxide dissolved in water, for example, forms lime water. [36] The same result is obtained when calcium hydroxide is dissolved in water. [36] Both sodium oxide and sodium hydroxide give sodium hydroxide solution when dissolved in water. [36] Both potassium oxide and potassium hydroxide, when dissolved in water, produce potassium hydroxide solution. [37]

The metals from the first main group even dissolve as elements in water and result in basic ("alkaline") solutions. [37] [38] They are therefore called alkali metals. Dissolved in water 11Sodium, for example, results in sodium hydroxide solution dissolved in water 19Potassium makes potassium hydroxide solution. [39]

Edit examples of other regularities

The most reactive elements are found in main groups I and VII (alkali metals or halogens), because these elements have a particularly strong tendency to give up (with alkali metals) or take up (with halogens) an electron a complete octet of electrons to get. [25]

The enthalpy of atomization, i.e. the energy required to break down a molecule E formed from an element E.x is required, shows a clear periodicity for the main group elements depending on the also periodic cohesion of the elements, because of this the number x of bound atoms. The enthalpy of atomization shows minima for the 0-valent noble gases and maxima for the tetravalent elements of main group IV. [25]

The density of the main group elements shows the same course because it is closely related to the bond of the respective element: The alkali metals have particularly small bonds and densities, the highest values ​​are for the elements of the middle groups.

A similar pattern can be seen in the dissociation enthalpies of E2-Molecules: The minima are again with the noble gases, the maxima now with the elements of the V main group (N.2, P2 etc.), according to the bonds possible in diatomic molecules. [25]

The melting and boiling temperatures, the melting and evaporation heats are further examples of physical properties of the elements that show a periodic behavior. [42] This even applies to the relevant properties of simple binary compounds, for example the melting temperatures or heats of fusion of hydrides, fluorides, chlorides, bromides, iodides, oxides, sulfides and so on. [42]

Metals, semi-metals and non-metals machining

Categorizations of non-metals
Reactive non-metals Noble gases
H, C, N, P, O, S, Se, F, Cl, Br, I He, Ne, Ar, Kr, Xe, Rn
Non-metals Halogens Noble gases
H, C, N, P, O, S, (Se) F, Cl, Br, I, At He, Ne, Ar, Kr, Xe, Rn
Fixed Fluid Gaseous
C, P, S, Se, I, At Br H, N, O, F, Cl, He, Ne, Ar, Kr, Xe, Rn
Electronegatives
Non-metals
Strongly electronegative
Non-metals
Noble gases
H, C, P, S, Se, I N, O, F, Cl, Br He, Ne, Ar, Kr, Xe, Rn
Polyatomic
elements
Diatomic
elements
Monatomic
Elements (noble gases)
C, P, S, Se H, N, O, F, Cl, Br, I He, Ne, Ar, Kr, Xe, Rn

The vast majority of elements are metals. They are usually silvery, shiny, malleable, not very volatile, and conduct electricity and heat. [25] The metal character is most clearly pronounced with the elements in the lower left of the periodic table and decreases towards the upper right. The semimetals are connected in this direction (matt gray, shiny, brittle, slightly volatile, only moderately conductive and heat-conducting [25]). At the top right of the periodic table are the non-metals (colored, not shiny, brittle, mostly volatile, not conductive and only poorly thermally conductive [25]).

The first two main groups (the alkali and alkaline earth metals) therefore contain exclusively metals, the last two main groups (the halogens and noble gases) only non-metals. [25] The boundary between metals and non-metals, marked by the semi-metals, runs diagonally through the middle main groups, so that these generally contain non-metals in the upper part, semi-metals below and metals in the lower part. Typical semi-metals are around 5Boron, 14Silicon or 32Germanium. Elements located on the border can even change their affiliation depending on the present modification: That located on the border between metals and semimetals 50As white β – tin, tin is a metal; as gray α – tin, it is a semimetal. [25] The one lying on the border between semi-metals and non-metals 6As graphite, carbon is a semi-metal, and as diamond, it is a non-metal. [25]

In the V. and VI. Main group, the transition that takes place within a group can be easily observed: The elements in the groups above 7Nitrogen, 8Oxygen and 16Sulfur are distinct non-metals. The elements below 15Phosphorus, 33Arsenic and 34Selenium occurs in non-metallic modifications (white, red and purple phosphorus, yellow arsenic, red selenium [25]) as well as in semiconducting modifications (black phosphorus, gray arsenic, gray selenium [25]). The items in the groups below 51Antimony, 52Tellurium, 83Bismuth and 84Polonium occurs preferentially in semi-metallic or metallic form. [25]

The typical representatives of the metals on the left-hand side of the periodic table always have only a small number of valence electrons and willingly give them up (low ionization energy, see above) in order to achieve an octet of valence electrons. When metal atoms combine to form a metal lattice by means of chemical bonds, the released valence electrons form an "electron gas" that embeds the positively charged metal atoms and holds them together. This is what is known as the metallic bond. From the properties of this type of bond, the characteristic properties of the metals, such as their shine or their ease of deformability, follow. In particular, the large number of freely moving electrons leads to high electrical conductivity.

Edit more complex relationships

Special position of the head elements Edit

The periodic table arranges the elements in such a way that the elements belonging to a group are chemically and physically similar to one another. The degree of similarity varies from case to case, but it is noticeable that the first members of each main group (the "head elements" [44] Li, Be, B, C, N, O, F) have less resemblance to the rest As such, members of their group each have among themselves. [44] Reasons for this are, among other things, that due to the small atomic radii, the valence electrons of these atoms are particularly strongly bound to the nuclei, and that the head elements, in contrast to the other group members in the outer shell, cannot exceed an electron octet. [44]

An example of this special position is the gaseous nature of 7Nitrogen and 8Oxygen in contrast to other representatives of the respective group. [45] Another example is the fact that the head elements cannot assume any higher oxidation numbers than their valence electron structure corresponds to. So can he 8Oxygen assume at most the oxidation number +2, while the other members of the oxygen group often have the oxidation numbers +4 and +6, which they due to the participation of the oxygen missing d-Acquire orbitals at the respective bond. [45]

The special position of the head element is particularly pronounced in the s-Block of the periodic table (especially if you have the 1Hydrogen instead of the 3Lithium counts as a head element), less pronounced in the p-Block, although present but only slightly pronounced in the d-Block and even less in the f-Block. [45]

Edit helical relationships

The named head elements are more similar to the main group elements on the right below them in the periodic table than their own group members and are therefore examples of oblique relationships. This particularly applies to similarities between 3Lithium and 12Magnesium, 4Beryllium and 13Aluminum, 5Boron and 14Silicon. [46] [47] The reason for this is that some important trends in elemental properties such as electronegativity, ionization energy or atomic radii run from bottom left to top right and thus “obliquely” in the periodic table. If you move down the periodic table, for example, the electronegativity decreases. If you move to the right, it increases. When moving down to the right, the two trends approximately cancel each other out and the electronegativity is only slightly changed. [46]

Knight relationship edit

An unusual relationship between elements is the knight relationship according to Michael Laing, which is characterized in analogy to the chess piece of the knight in that some metallic elements from the fourth period in some characteristics (e.g. melting points and boiling points) have similar properties to a metallic one Have element that is one period below and two groups to the right. [45] are examples 30Zinc and 50Tin, which has the same properties in an alloy with copper, in the coating of steel and in its biological importance as a trace element. [45] Other examples are 48Cadmium and 82Lead, 47Silver and 81Thallium, as well 31Gallium and 51Antimony. [45]

Edit relationships between main and subgroups

There are many similarities between a given group n and the group ten columns further to the right n + 10. [49] Are a striking example 12Magnesium from the second and 30Zinc from the twelfth group, whose sulfates, hydroxides, carbonates and chlorides behave very similarly. [50] Other distinct examples are 21Scandium from the third group and 13Aluminum from the thirteenth group as well 22Titan from the fourth group and 50Tin from the fourteenth group. [50] Only between the alkali metals in the first group and the noble metals (29Copper, 47Silver, 79Gold) in the eleventh group there is no resemblance. [50]

In the medium-length form of the periodic table in use today, these relationships are not very obvious. However, they were well known to the early pioneers of the periodic table, who could only orientate themselves on chemical similarities. The relationships lead to the fact that the "long" periods four to seven (without the separately presented lanthanoids and actinides) have a double periodicity: Both their left half (up to the precious metals) and their right half (up to the Noble gases have properties that tend to run parallel to the main groups in the short periods two and three.

The so-called short period system takes these similarities into account by presenting the two short periods two and three as a closed block (not divided into two as is otherwise the case), while it divides the four long periods and lists their left and right halves as separate lines below one another. [51] For this purpose, the elements of the iron, cobalt and nickel groups are added to each of the long periods containing 18 elements one Group together. These periods can then be divided into two halves of eight groups each (one of which is a group of three), which are arranged one below the other in the short period system. The short period system therefore only has 8 columns. Because of the existence of a group of three, however, corresponds despite the eight Columns the transition to the element one row below the transition to the one in the long form ten Groups further to the right element. The original main and subgroups can be distinguished by indenting differently [51].

The short form of the periodic table shows in particular the parallel course of valencies (more precisely: the maximum oxidation numbers) between secondary and main groups, which has been lost in the long forms and only survives there in the form of group numbering (see next section). On the other hand, the short form is less clear than the long forms, and it also emphasizes the similarities between main and subgroups more than they actually are. [51]

Excursus: Editing the numbering of groups

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
Yes IIa IIIb IVb Vb VIb VIIb VIIIb Ib IIb IIIa IVa Va Via VIIa VIIIa
Yes IIa IIIa IVa Va Via VIIa VIIIa Ib IIb IIIb IVb Vb VIb VIIb VIIIb
H ⏞ < displaystyle overbrace < qquad >> Hey
Li Be B. C. N O F. No
N / A Mg Al Si P. S. Cl Ar
K Approx Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Rb Sr Y Zr Nb Mon Tc Ru Rh Pd Ag CD In Sn Sb Te I. Xe
Cs Ba * Hf Ta W. re Os Ir Pt Au Ed Tl Pb Bi Po At Marg
Fr. Ra * Rf Db Sg Bra Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Above

Two of the three common numbering systems for the groups go back to the group arrangement in the short period system just described.

The eight groups of the short period system are numbered with Roman numerals from I to VIII. If one pulls the short period system apart again to the long form, the main and sub-group elements combined in a group of the short form must be divided again into two separate groups in the long form. If you want to keep the group numbering in the short form, each group number is doubled. To distinguish, add an a or b to the group number.

In the convention, which is mainly used in the USA, the main group elements receive an a, the subgroup elements a b. [53] The result is the numbering sequence (main groups shown in bold):

Ia IIa IIIb IVb Vb VIb VIb VIIIb Ib IIb IIIa IVa Va VIa VIIa VIIIa

In the convention, which is mainly used in Europe, the first series from I to VIII is consistently given an a, the second a b. [53] The result is the numbering sequence

Ia IIa IIIa IVa Va VIa VIIa VIIIa Ib IIb IIIb IVb Vb VIb VIIb VIIIb

The advantage of the two numbering systems derived from the short form is that the group number for the main groups is identical to the number of valence electrons. [38] It is immediately apparent that, for example, the elements of main group IV have four valence electrons.

IUPAC recommends numbering the groups sequentially with Arabic numerals from 1 to 18. [38] [54] While this numbering is transparent and unambiguous, the connection between the group number and the number of valence electrons is lost. For example, the elements with four valence electrons are in group 14.

Additional influences edit

The properties of unknown elements can be approximately predicted if the properties of the surrounding elements in the periodic table are known. The regular variation of the properties within the groups and periods is, however, interrupted by numerous exceptions, which add to the complexity of the field of chemistry. [53] The higher the atomic number, the less suitable the systematics of the periodic table is for predicting material properties, since the higher charge of the atomic nucleus increases the speed of electrons close to the nucleus and thus increases relativistic effects. [55] In the case of elements from the fourth period onwards, the electrons of the innermost shells (especially the s-Orbitals) due to the increasing number of positive charges in the atomic nucleus closer to the atomic nucleus, whereby the speed of these electrons almost reaches the speed of light. As a result, contrary to the general tendency, the ion radius decreases and the ionization energy for these electrons increases (effect of the inert electron pair).

The radioactive elements are marked as additional information that has nothing to do with the electron configuration and therefore with the position in the PSE:

Element 82 (lead) is the last element of which stable, i.e. non-radioactive, isotopes exist. [56] All of the following (atomic number 83 and higher) show exclusively unstable and thus radioactive isotopes. 83 (bismuth) is a borderline case. It only has unstable isotopes, including one with an extremely long half-life (209 Bi with t 1/2 = 1, 9 ⋅ 10 19 < displaystyle t_ <1/2> = 1 <,> 9 cdot 10 ^ <19 >> a). Also below element 82 there are two elements with exclusively unstable isotopes: 43 (technetium) and 61 (promethium). [8th]

This actually leaves only 80 stable elements that occur in nature - all others are radioactive elements. Of the radioactive elements, only bismuth, thorium and uranium are naturally present in larger quantities, [57] since these elements have half-lives in the order of magnitude of the age of the earth or longer. With the exception of one isotope of plutonium, all other radioactive elements are either, like radium, intermediate decay products of one of the three natural radioactive decay series or arise from rare natural nuclear reactions or from spontaneous splitting of uranium and thorium. Elements with atomic numbers above 94 can only be produced artificially. Although they are also created during the element synthesis in a supernova, no traces of them have been found in nature due to their short half-lives. The last element detected so far is Oganesson with the ordinal number 118, but this only has a half-life of 0.89 ms. Presumably there is an island of stability with higher ordinal numbers. [58]

Since the number of protons in the atomic nucleus is identical to the atomic number, the atomic mass increases with the atomic number. While the atomic number increases by one unit from one element to the next, the increase in the atomic mass is much more irregular.

The mass of a proton is 1.0073 atomic mass units (1 u = 1.66 · 10 −27 kg), that of a neutron is 1.0087 u. [59] The mass of an electron of 0.0005 u is mostly negligible. The mass of a hydrogen atom consisting of one proton and one electron is 1.0078 u. Since all atoms have an integer number of protons and neutrons (each with about 1 u mass) in the nucleus, they basically also have an integer atomic mass to a good approximation , which corresponds rounded to the number of protons and neutrons contained in the nucleus [60] (the atomic masses are usually somewhat smaller than an integer, the "mass defect" corresponds to the binding energy released during the formation of the nucleus [61]). In the apparent contradiction to this, however, some of the mass specifications in the periodic table differ significantly from the integer. For chlorine, for example, there is an indication of 35.45 and [62]

Ele-
ment
Mass
number
(Isotope)
Natural
Frequency [63]
Atom-
mass [63]
(u)
Medium
Atomic mass
(u)
15P. 31 100 % ,000 30,97 30,97
16S. 32 0 95,02 % 0 31,97 32,06
33 00 0,75 % 0 32,97
34 00 4,21 % 0 33,97
36 00 0,02 % 0 35,97
17Cl 35 0 75,77 % 0 34,97 35,45
37 0 24,23 % 0 36,97
18Ar 36 00 0,337 % 35,97 39,95
38 00 0,063 % 37,96
40 0 99,600 % 39,96

The reason for this is that two atoms with the same number of protons can have different numbers of neutrons. Such atoms have the same atomic number and thus the same chemical behavior, so by definition belong to the same chemical element and are therefore in the same place in the periodic table. But because they have different numbers of neutrons, they are different "isotopes" of this element (from ancient Greek ἴσος ísos “Equal” and τόπος tópos "Place").

consist of only one naturally occurring isotope, they are pure elements. [60] The other elements are mixed elements; their natural occurrence consists of a mixture of different isotopes. [60] For these mixed elements in the periodic table is the mean atomic mass the naturally occurring isotope mixture entered. The naturally occurring chlorine, for example, consists of 75.77% of the chlorine isotope with the mass number 35 (with 17 protons and 18 neutrons in the nucleus) and 24.23% of the chlorine isotope 37 (17 protons and 20 neutrons). [64] Its mean atomic mass is the mean weighted with the frequency of the (almost integer) atomic masses 34.97 u and 36.97 u, [63] therefore the above-mentioned 35.45 u.

If the isotopes of two successive elements in the periodic table have very different frequency distributions, it can happen that the mean atomic mass decreases from one element to the next. So that has to do with that 18Argon following 19Potassium has a higher atomic number, but a smaller mean atomic mass. The same applies to 27Cobalt and 28Nickel, 52Tellurium and 53Iodine, as well 90Thorium and 91Protactinium.

Since the atomic mass (apart from the exceptions mentioned) grow fairly regularly with the ordinal number, in the 19th century they could successfully be used as a basis for the search for regularities instead of the actual principle of order, the as yet unknown ordinal number.

Edit elements

In ancient Greece and in ancient China it was speculated more than 2000 years ago that the multitude of phenomena in nature must be traced back to a small number of "elements". In Greece Empedocles represented the four-element doctrine with the elements fire, water, earth and air. [65] In China there were the elements wood, fire, earth, metal and water in the five-element teaching.

The current concept of an element as a substance that cannot be further broken down goes back to Joachim Jungius and Robert Boyle. [65] In 1789, Antoine Laurent de Lavoisier presented a first systematic, 33-entry table with presumed "simple substances", [66] of which 21 were in fact already elements in today's sense. [65] However, there was still complete uncertainty about the internal structure of the elements and thus all matter in general. According to John Dalton's atomic hypothesis (1808), all substances are made up of the smallest, indivisible "atoms", whereby the atoms of one chemical element are identical to one another, but differ in shape and weight from the atoms of another element. [67] According to the hypothesis, chemical reactions were to be regarded as regrouping of indestructible atoms, and the laws of constant proportions [68] and multiple proportions [69] could thus be easily explained. Although atoms were accepted as a working hypothesis by many chemists, there was no evidence that they existed.

Atomic mass edit

While the densities of the various elements had been known for a long time, the lack of knowledge about the number and size of the atoms made it impossible to determine their absolute masses. Dalton had already drawn up a fairly imprecise list of the ratios of the atomic masses based on constant proportions, comprising 14 elements. [69]

William Prout noticed that many atomic masses were roughly integer multiples of the atomic mass of hydrogen and in 1815 he hypothesized that all elements were composed of corresponding amounts of hydrogen as the "original substance". [70] The atomic masses previously listed as non-whole numbers would turn out to be whole numbers if more precise measurements were taken. [70] Prout's hypothesis caused more precise mass determinations, mainly by Jöns Jakob Berzelius and Jean Servais Stas, who confirmed the non-integer nature of many atomic masses and thus refuted Prout's hypothesis, but also served as the basis for more reliable investigations because of their significantly improved accuracy. [70] The reason for the strikingly large number of elements with approximately whole-numbered atomic masses remained unclear. [71]

In the 1850s, Stanislao Cannizzaro revisited the hypothesis that Amedeo Avogadro had put forward in 1811, but had hitherto been ignored, that the same volume of different gases contain the same number of particles at the same temperature and pressure. This hypothesis made it possible to systematically compare the masses of the same (albeit unknown) numbers of atoms in gaseous compounds and to determine the relative atomic masses of the elements with reference to a reference element. [72] With their help, numerous previously incorrectly assumed proportions in chemical compounds could be corrected. On this basis, Cannizzaro published between 1858 and 1860 in preparation for the Karlsruhe Congress (which Meyer and Mendelejew also attended) more reliable and consistent atomic masses, which allowed rapid progress in the development of periodic systems in the 1860s. [74]

Forerunners of the periodic table edit

At the beginning of the 19th century, regularities in the relationships between the elements were searched for. Obstacles were, among other things, the uncertainties in the atomic mass and the fact that numerous elements were not even known. [75] Döbereiner was the first to establish a connection between atomic mass and the chemical properties of individual elements. In 1824 Falckner published a system of natural element families. [76] Gmelin created a tabular sorting of the elements in 1843. [77] Other pioneers, who also knew Mendeleev, were Pettenkofer (1850), Odling (1857), Dumas (1858) and Lenßen (1857). [78] In 1862, Chancourtois developed a three-dimensional representation, in which he arranged the elements helically on a cylinder according to increasing atomic masses. [78] Attempts were also made by Hinrichs (1864), Baumhauer (1867) and Quaglio (1871) to depict the system in a spiral. [78] In 1863/64 Newlands put up a table of the elements in groups of eight (law of octaves) arranged according to atomic mass. [78]

Johann Wolfgang Döbereiner (Triad System) edit

From this he initially drew the suspicion that strontium consisted of barium and calcium, which he did not find confirmed in corresponding experiments. [81]

From a modern point of view, calcium, strontium and barium are three elements from the group of alkaline earth metals that are one below the other in the periodic table, which gives their identical valences and therefore their similarity in the formation of oxides x + O < displaystyle mathrm > justified. Since the period length in this area of ​​the periodic table is 18 elements (a period here comprises eight main groups and ten subgroups), they show the same difference in ordinal numbers among each other (18, Döbereiner still unknown):

and therefore roughly the same difference in atomic masses (just under 50 u).

Leopold Gmelin noted in his 1827 "Handbook of Theoretical Chemistry" With regard to the atomic mass, "some strange relationships, which are no doubt related to the innermost nature of substances." [82] Among other things, he pointed to another group of three, namely lithium, sodium and potassium. If one forms the arithmetic mean of the atomic masses of lithium and potassium, "one obtains fairly precisely [the atomic mass] of the sodium, which metal also comes into its chemical relationships between the two mentioned." [82]

H Hey
Li Be B. C. N O F. No
N / A Mg Al Si P. S. Cl Ar
K Approx . Ga Ge As Se Br Kr
Rb Sr . In Sn Sb Te I. Xe
Cs Ba . Tl Pb Bi Po At Marg
Fr. Ra . Nh Fl Mc Lv Ts Above
Position of the four triads in the modern periodic table

Döbereiner published a more detailed one in 1829 "Attempt to group the elementary substances according to their analogy". [83] A newly discovered triad contained "three salt formers" with chlorine and iodine as well as bromine, which was only isolated in the previous year [84]. The comparison using the atomic masses determined by Berzelius gave [85]

Another newly found triad comprised sulfur, selenium and tellurium, all of which "combine with hydrogen to form peculiar hydrogen acids":

In his attempts at order, Döbereiner insisted that the elements combined in a triad actually showed chemical similarity: “The fact that the arithmetic mean of the atomic weights of oxygen = 16.026 and carbon = 12.256 expresses the atomic weight of nitrogen = 14.138 can be used here are out of the question because there is no analogy between these three substances. ”He also insisted on the special meaning of the number three. The elements iron, manganese, nickel, cobalt, zinc and copper, which are very similar to one another, presented a problem for him, because "how should they be arranged if the triad is accepted as the principle of grouping?"

O N H
F. Cl Br J L. N / A K
S. Se Te Mg Approx Sr Ba
P. As Sb G Y Ce La
C. B. Si Zr Th Al
Ti Ta W. Sn CD Zn
Mon V Cr U Mn Co Ni Fe
Bi Pb Ag Hg Cu
Os Ir R Pt Pd Au

In 1827 Gmelin had shown the then known 51 elements individually in a V-shaped arrangement in order to clearly show their "relationship and difference", [86] he went over in 1843 to the 55 known elements "depending on their physical and chemical conditions" to be summarized in groups of mostly three elements, which in turn were arranged "according to their similarities" in a V-shaped scheme in order of increasing electropositivity. [87] Today's main groups can be recognized in some of Gmelin's groups (R. = Rhodium, today Rh L. = Lithium, today Li G = Glycium, today Beryllium Be).

In 1857 Ernst Lenßen was able to divide practically all elements known at the time into 20 triads [88] [89] (but was less strict than Döbereiner with regard to chemical similarity). He even put groups of three of triads together to form "enneads" (groups of nine), in which the atomic masses of the respective middle triads were in turn related by the mean value rule. [88] Using his system, he predicted, among other things, the atomic masses of the elements erbium and terbium, which had already been discovered but not yet isolated, but none of his predictions were successful. [88] He also tried to establish connections with other physical and chemical properties. [89] [90]

John A. R. Newlands (Law of Octaves) Edit

The previous attempts at order were largely limited to finding isolated groups with similar elements. [91] John Alexander Reina Newlands published a table with 24 elements (and a space for a supposedly undiscovered element) in 1864, in which the elements were arranged as usual in the order of increasing atomic masses, but in which he did not rely on patterns in the Atomic mass differences, but pointed to repetitive differences in the place numbers of similar elements. [92] This was the first periodic system, a compilation of elements that shows that the properties of the elements repeat themselves after certain regular intervals. [92] Newlands was also the first to interchange the sequence of the elements iodine and tellurium based on the atomic mass and prefer the arrangement based on the chemical properties. [92]

In 1865, Newlands developed another system that now comprised 65 elements. [92] It should show that the chemical properties are repeated in every eighth position, which he compared to the octaves from music. [92] (Since the noble gases had not yet been discovered, the period length in the first periods of Newland's table was actually seven elements both Similar elements were counted, just as one also counts the octave in music e.g. B. counting from one C to the next C inclusive results in a period length of 8.) [92]

Newlands called this relationship between the elements the "law of octaves", which was the first time that the repetition pattern in the element properties was regarded as a law of nature. [92] The law of octaves can be perfectly applied to the first two periods, but because then (as we know today) the periods become longer, the law was less successful in the following periods. [92]

The first accurate prediction of an as yet undiscovered element goes back to Newlands: Due to a gap in one of his tables, he predicted the existence of an element with atomic mass 73 between silicon and tin in 1864. [93] This corresponds to germanium, discovered in 1886, in the announced position and with an atomic mass of 72.61. [93] However, his predictions of still unknown elements between rhodium and iridium and between palladium and platinum did not come true. [93]

The discovery of periodicity is occasionally attributed to Alexandre-Emile Béguyer de Chancourtois, who in 1862 arranged the elements according to increasing atomic mass along a three-dimensional screw so that a screw turn corresponded to 16 units, i.e. elements at a distance of 16 units came to stand vertically on top of each other. [94] However, his system was largely ignored and he did not develop it further. [94]

Dmitri Mendeleev and Lothar Meyer (Periodic Table) Edit

The modern periodic table was developed by Lothar Meyer and Dmitri Iwanowitsch Mendeleev. [95] Both published their results in 1869 and jointly received the Davy Medal of the British Royal Society in 1882 for their work.

Mendeleev is mentioned more often than Meyer as the founder of today's periodic table. On the one hand, because Meyer's periodic table was published a few months later, on the other hand, because Mendeleev made predictions about the properties of the as yet undiscovered elements. [96] In Russia the periodic table is still used today as a Tablitsa Mendeleeva ("Mendeleev's Table"). Neither Mendeleev nor Meyer knew each other's work on the periodic table. [97] [78] The works of Béguyer de Chancourtois from 1862, Newlands from 1863/64 or Hinrichs from 1866/67 were also unknown to Mendeleev. [78]

Lothar Meyer edit

In his textbook published in 1864 "The modern theories of chemistry" Meyer presented a table containing 28 elements and sorted according to increasing atomic mass. [98] The division into rows was made so that each column (corresponding to today's main groups) contained elements of the same value and the value changed by one unit from one column to the next. Meyer pointed out that the atomic mass difference between the first and second elements of each column was about 16, the next two differences fluctuated by about 46, and the last difference was always about 87 to 90. [99] He suggested that this could - similar to homologous series of molecules - indicate the systematic structure of atoms from smaller components. [100] [101]

Meyer had swapped the elements tellurium and iodine, according to their chemical properties, in relation to the sequence based on the atomic masses. [102] Meyer had to leave some gaps in the table, including one between silicon and tin, in which, according to his difference scheme, an element of atomic mass 73 was to be expected. [102] The missing element was germanium, discovered in 1886, with an atomic mass of 72.61. Another table, not arranged according to atomic mass, contained 22 elements that Meyer had not included in his scheme - they correspond to today's transition metals. [102]

In 1870 (submitted in December 1869, less than a year after Mendeleev's first publication of a periodic table), Meyer published an expanded version of his table in which, using updated atomic masses, he had succeeded in “arranging all the elements known to date in the same scheme”. [103] The periods in this system ran vertically, the groups horizontally. The (not yet so called) transition metals were now part of the table. Similar to a short period system, they were arranged in periods that alternated with the (not yet so-called) main groups.

To illustrate the variation of the properties along the periods, Meyer added a diagram that shows the periodically varying atomic volumes as a function of the atomic mass (similar to the diagram in the section Atomic Radii). This illustration contributed significantly to the acceptance of the periodic table. [104] Meyer discussed various physical properties of the atoms that run parallel to the atomic volumes and are therefore also periodic, such as density, volatility, ductility, brittleness or specific heat. [103]

Dmitri Mendeleev edit

Mendeleev's name is mainly associated with the periodic table in its present form. [106] His periodic table was more complete than other systems of that time, [107] he promoted and defended his system with commitment, [106] worked it out over decades [106] and used it for far more extensive and detailed predictions than other authors periodic systems. [108]

In search of a structure for his chemistry textbook, Mendeleev created Jul. / March 1, 1869 greg. a first draft of his version of the periodic table. [109] In March [110] he published his system with a detailed explanation in the journal of the Russian Chemical Society. [111]

He expressly pointed out that most of the properties of the elements are not suitable as a clear ordering principle. For example, most elements can have different values. Most of the properties of the free elements depend on the modification present (graphite and diamond, for example, are modifications of carbon with significantly different properties), and so on. The only unambiguous and numerically measurable property of an element, which is preserved in all modifications of the free element as well as in all its compounds, is its atomic mass [111] (the atomic number as another such property was still unknown to Mendeleev).

He arranged the elements of the "natural groups" already known as belonging together (such as the halogens, the alkaline earth metals, the nitrogen group, etc.) according to their atomic mass and found that this arrangement corresponded to "the natural resemblance prevailing among the elements" without further assistance. [111] He stated: "The elements arranged according to the size of their atomic weight show a clear periodicity of their properties," [111] and on this basis tried to fit the other elements into the scheme according to their chemical behavior.

In this article, Mendeleev already predicted the existence of two new elements with atomic masses between 65 and 75, which were supposed to resemble aluminum and silicon, based on gaps that had remained in his system. [112] Like some of his predecessors, Mendeleev, too, had swapped tellurium and iodine in relation to the sequence derived from the atomic masses. His prediction that the atomic mass of tellurium had to be corrected because according to his system it could not be 128 and must rather lie between 123 and 126, [112] did not come true - here there is actually an irregularity of the atomic mass. In the same year, two short German-language descriptions of the new system were published. [112] [113]

In 1871 an extensive article appeared in which Mendeleev presented two further developed versions of his periodic table. [114] One of these variants was the first short-period system. In this article he demonstrated, among other things, how the atomic mass of an element could be determined or corrected using the periodic table if its chemical behavior was known. The article also contains the three most famous predictions about the properties of as yet unknown elements, the existence of which Mendeleev discovered from remaining gaps in his periodic table. By skillful interpolation between the physical and chemical properties of the neighboring elements, he was able to accurately predict numerous properties of the as yet unknown elements. [116]

Mendeleev named the unknown elements after the element above the respective gap in his short period table with the addition of the prefix Eka (sanscr. "one"). Ekaaluminium was discovered by Paul Émile Lecoq de Boisbaudran in 1875 and named gallium after France, the land of discovery. Ekabor was discovered in 1879 by Lars Fredrik Nilson and - after Scandinavia - given the name Scandium. Ecasilicon was discovered by Clemens Winkler in 1886 and was named germanium after the country of discovery, Germany.

Comparison of the predictions for e-silicon "Es" (1871) and the findings on germanium "Ge" discovered in 1886 [117]
(Selection)
element oxide chloride Ethyl compound
Atom-
Dimensions
density
(g / cm³)
Heat cap.
J / (kg K)
colour formula density
(g / cm³)
formula Boiling
Point
density
(g / cm³)
formula Boiling
Point
forecast 72 5,5 306 dark gray EsO2 4,7 EsCl4 100 ° C 1,9 It (C2H5)4 160 ° C
found 72,32 5,47 318 greyish white GeO2 4,703 GeCl4 0 86 ° C 1,887 Ge (C2H5)4 160 ° C

Not all of Mendeleev's predictions were that successful. Overall, only about half of his predictions of new elements were true. [118]

The noble gas argon, discovered in 1894, appeared to pose a significant threat to the general validity of Mendeleev's periodic table, as it could not be integrated into the existing system. [119] However, when further noble gases were discovered in quick succession (1895 helium, 1898 neon, krypton and xenon, 1900 radon) [120] it became clear that the periodic table only had to be expanded to include a new group of elements between the halogens and the alkali metals, to be able to absorb them all. [121] Mendeleev spoke of a “critical test” which his periodic table “survived brilliantly”. [121]

Mendeleev published about thirty versions of the periodic table over the years, and another thirty are available in manuscript. [122] The oldest surviving display board in the periodic table dates from the period between 1879 and 1886 and is located in the University of St. Andrews. [123]

Periodic table according to Mendeleev, 1869 [112] Modern periodic table to uranium, arranged according to Mendeleev's scheme
Sc = 45 0 Y = 89 0 La - Lu Ac, Th, Pa, U
Ti = 50.0 Zr = 90.0 ? = 180 ,0 Ti = 48 Zr = 91 0 Hf = 178
I V = 51.0 Nb = 94.0 Ta = 182.0 I V = 51 0 Nb = 930 Ta = 181
Cr = 52.0 Mo = 96.0 I W = 186 Cr = 52 0 Mo = 96 0 I W = 184
Mn = 55.0 Rh = 104.4 Pt = 197.4 Mn = 55.0 I Tc = 97 0 Re = 186
Fe = 56.0 Ru = 104.4 Ir = 198.0 Fe = 56 Ru = 101 Os = 190
Ni = 59.0
Co = 59.0
Pd = 106.6 Os = 199.0 Co = 59 Rh = 103 Ir = 192
Ni = 59 Pd = 106 Pt = 195
H = 1 Cu = 63.4 Ag = 108.0 Hg = 200.0 Cu = 64 Ag = 108 Au = 197
Be = 9.4 Mg = 24.0 Zn = 65.2 Cd = 112.0 Zn = 65 Cd = 112 Hg = 201
B = 11.0 Al = 27.4 ? = 68 ,0 Ur = 116.0 Au = 197? , B = 11 Al = 27 Ga = 70 In = 115 Tl = 204
C = 12.0 Si = 28.0 ? = 70 ,0 Sn = 118.0 C = 12 Si = 28 Ge = 73 Sn = 119 Pb = 207
N = 14.0 P = 31.0 As = 75.0 Sb = 122.0 Bi = 210? , N = 14 P = 31 As = 75 Sb = 122 Bi = 209
O = 16.0 S = 32.0 Se = 79.4 Te = 128? , O = 16 S = 32 Se = 79 Te = 128 Po = 209
F = 19.0 Cl = 35.5 Br = 80.0 J = 127.0 F = 19 Cl = 35 Br = 80 I = 127 At = 210
Li = 7 Na = 23.0 K = 39.0 Rb = 85.4 Cs = 133.0 Tl = 204.0 He = 4 Ne = 20 Ar = 40 Kr = 84 Xe = 131 Rn = 222
Ca = 40.0 Sr = 87.6 Ba = 137.0 Pb = 207.0 H = 1 Li = 7 Na = 23 K = 39 Rb = 85 Cs = 133 Fr = 223
? = 45 ,0 Ce = 92.0 Be = 9 Mg = 24 Ca = 40 Sr = 88 Ba = 137 Ra = 226
? Er = 56.0 La = 94.0
? Yt = 60.0 Di = 95.0
? In = 75.6 Th = 118?

The colors indicate the current assignment of the elements:
Alkali metals, alkaline earth metals, 3rd main group, 4th main group, 5th main group, 6th main group, halogens, noble gases, transition metals, lanthanoids, actinides. The supposed element didymium (Di) later turned out to be a mixture of the rare earths praseodymium and neodymium. [124] In order to convert the modern periodic table shown on the right from Mendeleev's arrangement into the order commonly used today, the last two lines are to be added at the top, shifted one box to the right, and the whole system is to be attached to the diagonal running from top left to bottom right reflect. In the modern periodic table shown, the atomic masses are rounded to whole numbers for the sake of clarity.

Henri Becquerel (radioactivity) edit

Henri Becquerel discovered in 1896 that uranium emitted a previously unknown radiation. [125] The uranium mineral pitchblende emitted significantly more radiation than would have corresponded to the uranium content. [126] Marie and Pierre Curie discovered the new and radioactive elements polonium and radium in the pitchblende in 1898. They also recognized the element thorium as radioactive. [128]

Ernest Rutherford (atomic nucleus) edit

Joseph John Thomson established in 1897 that the cathode rays observed in gas discharge tubes were light material particles and not ether waves. [129] Thomson could do the relationship e / m of the charge and mass of the "electrons" called particles and determined that it was independent of the cathode material, filling gas and other circumstances, that the electrons were apparently universal components of the atoms. [129] Thomson created the plum pudding model in 1904, [130] according to which the electrons were embedded in a uniformly positively charged sphere.

When examining radioactive substances, different types of radiation could be distinguished: deflection in the magnetic field showed that the penetrating beta rays were negatively charged. Becquerel finally identified them as electrons. [131] Ernest Rutherford and Thomas Royds established in 1908 that the less penetrating alpha radiation consisted of doubly positively charged helium ions.

Rutherford's scattering experiments, in which he bombarded metal foils with alpha particles, showed in 1911 that the positive charges of the atoms are concentrated in a small nucleus [132] and the electrons are located outside the nucleus - their arrangement and number were, however, still unknown.

Henry Moseley (ordinal number) edit

The analysis of his scattering experiments led Rutherford to the conclusion in 1911 that the positive charge of the atomic nucleus corresponds to about half the atomic mass: Z ≈ A / 2 < displaystyle textstyle Z approx A / 2>. [133] Antonius van den Broek pointed out that the atomic mass increases by two units from one element to the next, that according to Rutherford's formula, the number of charges in the nucleus increases by one from one element to the next. [134] The number of possible elements is therefore equal to the number of possible nuclear charges and each possible nuclear charge corresponds to a possible element. [134] The atomic number therefore also determines the position of each element in the periodic table. [134] (The increase in atomic masses by two units only applies to a rough approximation. Van den Broek was influenced here by his assumption that all atoms are made up of half alpha particles of mass number 2. [134])

Henry Moseley confirmed that the atomic number (also: atomic number) is a more suitable ordering principle for the elements than the atomic mass. [135] He made use of the fact that materials bombarded with electrons emit not only the braking spectrum (Röntgen 1895 [136]) but also X-rays with a wavelength that is characteristic of the material [137] (Barkla, approx. 1906 [138]) and that the wavelength this radiation can be determined by means of diffraction on crystals (von Laue 1912 [139]). Moseley determined the wavelengths of the characteristic radiation of various elements and found that the frequencies of these radiations were proportional to the square of an integer that described the position of the element in question in the periodic table (Moseley's law). [140] He recognized this number as the number of charges in the atomic nucleus. [140] It was thus possible to easily determine the atomic number of an element experimentally.

Moseley demonstrated that many of the 70 or so allegedly newly discovered elements that competed for the 16 gaps to be filled in Mendeleev's periodic table could not exist because there was no space for them in the grid of ordinal numbers. [141] [142]

The good ten “rare earths” (their exact number was not known at the time) are chemically difficult to separate from one another because they are very similar to one another. Mendeleev had found no place for them in his scheme. [142] Moseley's ordinal number clearly assigned them the places 57 to 71. [142]

The pioneers of the periodic table had to correct the occasional mass inversions (such as between iodine and tellurium) by swapping the relevant elements in the atomic mass scheme, without being able to give a reason for this, except that they fit better into the scheme of chemical similarities. Moseley's atomic number confirmed the correct order of the exchanged elements, [142] the atomic masses had been misleading here.

When the First World War broke out, Moseley volunteered for military service and fell in the Battle of Gallipoli. Moseley's successor completed the systematic measurements and determined that uranium (the heaviest known element up to that point) has the atomic number 92, [143] that there are exactly 92 elements in the element series from hydrogen to uranium. Gaps were recognized in the ordinal numbers 43, 61, 72, 75, 85, 87 and 91, which could be filled in the following decades with the relevant new discoveries. [144]

Frederick Soddy (isotopes) edit

In 1902, Ernest Rutherford and Frederick Soddy determined that the radioactive elements not only gave off radiation, but consisted of unstable atoms that spontaneously converted into new elements by emitting alpha, beta or gamma radiation - in obvious contradiction to what had been previously assumed Indivisibility and immutability of atoms. [145] Starting with actinium (Debierne 1899 [146]), numerous other radioactive substances were quickly discovered after polonium and radon (in 1912 their number had grown to about 30). [145] The new substances were initially viewed as separate elements and it seemed that the periodic table could not accommodate them all. However, it turned out that practically all of them were chemically indistinguishable from elements already known and as such already had a place in the periodic table. For example, a decay product, initially called “radium D”, could not be distinguished from lead. [145] On the other hand, exact determinations of the atomic mass showed that lead samples from different sources could have different atomic masses. [147] Theodore William Richards found an atomic mass of 206.4 for lead made from pitchblende and 208.4 for lead made from thorite. [147]

From several similar cases in 1911 Soddy drew the conclusion that one and the same element could be a mixture of different atomic masses and in 1913 he coined the term "isotopes" for atoms with the same atomic number but different mass numbers. [148] Soddy and Kasimir Fajans set up the displacement theorems according to which an atom of a given element loses two nuclear charges when an alpha particle is emitted and is transferred to the element that is two places further to the left in the periodic table, while when a beta particle is emitted it Particle moves one place further to the right. For the elements between lead and uranium it was obvious why they could exist with different atomic masses. A thorium atom (atomic number 90) can, for example, have arisen from uranium-235 (atomic number 92) through an alpha decay and then has the mass 231. However, it can also have arisen from a beta decay from actinium-230 (atomic number 89) and then has the mass 230. [150] Soddy was able to predict the mass number 206 for lead from uranium ores and the mass number 208 for lead from thorium ores, even before Richards presented the results of his measurements. [147]

In the 1920s, the large number of newly discovered isotopes plunged the periodic table into a crisis, as it seemed that instead of the elements, the much larger number of isotopes with their different atomic masses had to be brought into a systematic order. [96] Fritz Paneth and George de Hevesy showed, however, that the chemical properties of the isotopes of an element were practically identical, so that it was justified to regard them as the same element according to their common atomic number (as a new ordering criterion instead of atomic mass) [96 ] and so keep the periodic table.

Francis William Aston developed the first mass spectrograph in 1919 [151] and found that the other elements that did not come from decay series could also be a mixture of different isotopes. [152] This clarified why the (mean) atomic masses of some elements, such as chlorine, deviated so clearly from the integer. And the fact that iodine, for example, has a higher atomic number but a smaller atomic mass than tellurium, was understandable as a consequence of the respective terrestrial isotope mixtures of the two elements.

Niels Bohr (construction principle) edit

When an excited atom emits the excitation energy again, the emitted radiation usually has a precisely defined wavelength that depends on the type of atom and the excited state. Rutherford's atomic model, in which the electrons moved around a central nucleus, made it clear why there was radiation at all, since the electrons had to send out electromagnetic waves as moving charges. However, it could not explain why only certain wavelengths were emitted. [153] In addition, the electrons, which are in constant motion, would have had to constantly emit energy and, because of this continuous loss of energy, would have had to fall into the nucleus in the shortest possible time. [154]

Niels Bohr took up Max Planck's discovery that the energy distribution of black body radiation can only be explained on the assumption that the energy radiation does not take place continuously but in the form of discrete "energy packets". In 1913 [153] he created an atomic model of hydrogen in which an electron orbits the nucleus not on any, but on one of several permitted orbits. The electron is more energetic on the higher orbits. If it falls back on a deeper orbit, it gives off the energy difference Δ E < displaystyle Delta E> in the form of radiation. The frequency f < displaystyle f> of this radiation is not (as would have been expected according to Maxwell's electrodynamics) the oscillation frequency of the orbiting electron, but according to another Bohr postulate given by Planck's condition hf = Δ E < displaystyle hf , = , Delta E> with the Planck constant h < displaystyle h>. Bohr succeeded in formulating the condition for the permitted orbits in such a way that the energy differences between each two orbits exactly corresponded to the observed frequencies of the spectral lines in the spectrum of hydrogen. Bohr's atomic model was thus able to describe the hydrogen spectrum successfully; it also provided good results for the spectral lines of other atoms with only one electron (H, He +, Li ++, etc.). [155]

Bohr also tried to describe the electron configurations of atoms with several electrons by distributing the electrons to the different orbits of his atomic model. [156] His “construction principle” assumed that the electron configuration of an element could be derived from the configuration of the previous element by adding another electron (mostly on the outermost orbit). [156] If a lane (in today's parlance a "bowl") was full, the next lane began to be filled. However, Bohr could not deduce from his model how many electrons a path could take up and distributed the electrons, as was suggested by chemical and spectroscopic points of view. [156]

Irving Langmuir (valence electron octet) edit

Gilbert Newton Lewis formulated the octet theory of the chemical bond in 1916 on the basis of the abundance of individual facts about the chemical and crystallographic behavior of substances that has now accumulated. [157] According to this theory, the atoms always strive for an octet of valence electrons as a particularly stable configuration and, in the case of non-ionic bonds, can reach this state by entering into a chemical bond with other atoms and their own octet with the valence electrons of these other atoms complete. Lewis called such a bond mediated by shared electrons a covalent bond. [157] He imagined the valence electrons to be arranged at the eight corners of a cube surrounding the nucleus. [158]

Irving Langmuir arranged the octets in bowls, the diameter of which corresponded to the orbits in Bohr's atomic model. He was able to explain the individual behavior of the chemical elements using the octet rule: The noble gases already have a complete valence electron octet and are not inclined to enter into chemical bonds. Elements with one or a few electrons in the valence shell tend to give up these electrons. Elements that are missing one or a few electrons to complete an octet tend to take up the missing number of electrons. Langmuir was even able to explain the different values ​​of the elements, i.e. their tendency to combine with a certain number of the respective partner atoms (Edward Frankland had introduced the concept in 1852). According to Langmuir, the valence is the number of electrons taken in or released to complete an octet. Chlorine, for example, accepts an electron, is thus monovalent and therefore combines, for example, with exactly one hydrogen atom. [159] (The current concept of valence is kept more general.) Langmuir was also able to make the nature of isotopes understandable: Since the chemical behavior is determined by the valence electrons, all isotopes of an element obviously have the same number of valence electrons and thus the same for this element's characteristic chemical behavior. Different numbers of particles in the core, on the other hand, result in different atomic masses, but do not influence the chemical behavior. [160]

Wolfgang Pauli (principle of exclusion) edit

In 1921 Bohr took up the problem of the shell configuration of multi-electron atoms again. Arnold Sommerfeld had expanded Bohr's atomic model to include elliptical orbits and introduced a second quantum number to describe them. Bohr set up a new table with electron configurations in which a certain number of Sommerfeld's quantum numbers (in today's language: secondary quantum numbers) was allowed for each orbit number (in today's parlance: each main quantum number). [161]

Edmund Clifton Stoner created a table of shell configurations in 1924, in which he used a third quantum number, meanwhile introduced by Sommerfeld, to count the possible electron states. He was able to reproduce the values ​​of the elements better than Bohr.[162] The problem remained, however, that the number of additional spectral lines into which a spectral line splits when the atom is brought into a magnetic field, suggests twice as many possible states of the electrons as previously considered was. [163] Wolfgang Pauli was able to explain Stoner's table by introducing a fourth quantum number (the "spin quantum number"), which can take on two different values ​​and thus doubles the number of possible states, [164] and by assuming that no two electrons are in of the atomic shell could agree in all four quantum numbers (Pauli's principle of exclusion). [165] With this the reason was found why the main shells can each take up 2, 8, 18, 32 etc. electrons. [166]

Hund's rule, established by Friedrich Hund in 1927, describes the order in which the individual orbitals of a subshell are filled with electrons: If several orbitals have the same energy level, they are first filled with individual electrons (with mutually parallel spins). Only then are the orbitals each occupied with a second electron (according to the Pauli principle with opposite spin).

Erwin Schrödinger (hydrogen problem) edit

Quantum mechanics was developed in the 1920s, beginning with de Broglie (1924), Heisenberg (1925) and Schrödinger (1926). [167] [168] It replaced the descriptive electron orbits of Bohr's atomic model with abstract, mathematically described "orbitals". [169] [170]

In the case of the simple hydrogen atom, the quantum mechanical Schrödinger equation can be solved exactly. There is not just a single solution for this so-called “hydrogen problem”, but a whole set of solution functions that describe the various possible states of the electron. It is a set of discrete, that is, individually countable mathematical functions, which can therefore be differentiated from one another by means of "code numbers". As it turns out, exactly four such indicators are necessary for the unambiguous identification of each state, which can be identified with the four quantum numbers already derived from the experiments. The relationship between the first three quantum numbers can be derived from the Schrödinger equation for the hydrogen atom: [171]

This now also provides a physical-mathematical justification for the number of possible electron states for a given main quantum number, i.e. for the maximum number of electrons that each main shell can hold.

In atoms with several electrons, the electrons do not take on the one-electron states of the hydrogen atom just described, but rather multiple-electron states, for which the quantum numbers just described are no longer valid, strictly speaking. But they have analogue quantum numbers for which the same designations are used. [172]

Glenn T. Seaborg (transuranium elements) edit

Rutherford discovered in 1919 that nitrogen bombarded with alpha particles emitted a new type of particle. [173] Together with James Chadwick, he identified these particles as positively charged hydrogen nuclei and called them "protons". [174] This identified the source of the positive charge on the atomic nucleus. [174] Blackett and Harkins observed in a cloud chamber in 1925 that in such cases the alpha particle was swallowed instead of just knocking a proton out of the nitrogen nucleus in passing. [174] From this it could be concluded that according to the equation

the nitrogen atom had become an oxygen atom, the first example of an artificial element conversion ("transmutation"). [174] Because of their lower charge, protons can overcome the electrical repulsion of heavier nuclei more easily than alpha particles and are therefore better suited as projectiles in bombardment experiments, since they can reach these nuclei more easily. However, since there were no natural sources for protons with the required energy, proton accelerators were developed, partly as linear accelerators, but especially in the form of the cyclotron (Lawrence and Livingston, 1931), [175] which made numerous new transmutations possible. [176]

When bombarded with alpha particles, beryllium, boron and lithium emitted a previously unknown, very penetrating radiation that Chadwick identified as uncharged particles with the mass of a proton. [177] This “neutron” explained why different isotopes of an element could have the same atomic number, but different masses: They had different numbers of neutrons in the nucleus. Since it is not repelled by the nuclei as an uncharged particle, the neutron is also suitable as a projectile in bombardment experiments. [177]

In the course of the bombardment experiments it was possible to produce new, non-naturally occurring isotopes through transmutation (the first in 1934 was the phosphorus isotope with a mass number of 30 by Irène and Frédéric Joliot-Curie). [178] The artificially produced isotopes are radioactive ("artificial radioactivity") and were quickly used for scientific and practical purposes because of their specific manufacturability. The bombardment of uranium atoms with neutrons led to the discovery of nuclear fission in 1938. [179]

In 1940, Edwin Mattison McMillan and Philip Hauge Abelson identified the new element neptunium in the products produced by bombarding uranium with neutrons. [180] [181] With the ordinal number 93 it was the first Transuran. A working group led by Glenn T. Seaborg examined the chemical properties of the new element (45 micrograms of it could be produced in a cyclotron). [180] It was to be expected that a beta decay of the isotope neptunium-239 would lead to the formation of the element with atomic number 94. [180] Targeted production of this isotope by bombarding uranium-238 with neutrons allowed Seaborg and colleagues in 1941 to produce the new element, plutonium. [182] [181]

I. II III IV V VI VII VIII .
5 Rb Sr Y Zr Nb Mon Tc Ru .
6 Cs Ba S.E. Hf Ta W. re Os .
7 Fr. Ra Ac Th Pa U .
S.E. = La Ce Pr Nd Pm Sm .
5 Rb Sr Y Zr Nb Mon Tc Ru .
6 Cs Ba LAN. Hf Ta W. re Os .
7 Fr. Ra ACT. Rf Db Sg Bra Hs .
LAN. = La Ce Pr Nd Pm Sm .
ACT. = Ac Th Pa U Np Pooh .

Until then it was established that the rare earths ("S.E." in the adjacent periodic table) represent an insert in the sixth period of the periodic table, in the seventh period, however, the possible existence of a similar insert had not yet been recognized. Francium, radium and actinium clearly belonged to the first, second and third groups of the seventh period. It was assumed that the subsequent elements thorium, protactinium and uranium had to belong to the fourth, fifth and sixth group, i.e., in the periodic table, they would come under the transition metals hafnium, tantalum and tungsten. [181] Some similarities with these groups, such as the value 4 of thorium or the value 6 of uranium, seemed to confirm the classification. [183] ​​This error delayed the progress of research, because in the identification and separation of new elements, their chemical similarity to known elements was often used. The new substances were mostly produced in too small quantities to be able to isolate them. However, if they were allowed to take part in a chemical reaction together with a known element, the product of which was then precipitated, and the easily measurable radioactivity was found in the precipitate, then the chemical similarity to the known element was shown. [184] (Marie and Pierre Curie had used this technique to concentrate the discovered radium in a barium chloride precipitate. [185]) However, if the radioactivity remained in the solution, the presumed similarity was refuted. Incorrectly assumed chemical similarities had delayed work on protactinium and uranium, among other things. Seaborg recognized in 1944 that the newly created transuranic elements were not transition metals, [181] but belong to an insert in the seventh period (the actinides), which corresponds to the insert (now called lanthanides) in the sixth period.

Seaborg et al. Produced element 96 (curium) in 1944 by bombarding plutonium-239 with helium ions and shortly afterwards element 95 (americium) by bombarding plutonium-239 with neutrons. [186] The bombardment of americium-241 with helium ions produced element 97 (Berkelium) in 1949, followed by element 98 (Californium) by bombarding Curium-242 with helium ions. [187]

Several working groups identified the elements 99 (Einsteinium) and 100 (Fermium) in the fallout of the nuclear weapons test "Mike" (1952). [187] With ever larger accelerators, heavier and heavier atoms could be used as projectiles, so that the generation of heavier and heavier transuranic elements was also possible. The heaviest Transuran produced so far (as of 2021) is Element 118 (Oganesson).

Ongoing discussions on positioning edit

Even today there are discussions about the position of some elements in the periodic table.

Classification of the first period

Due to the electron configuration, but not due to the elemental properties, helium (electron configuration 1s 2) would have to be classified in the second main group, i.e. above beryllium in the periodic table. [188] Helium has only two electrons, in contrast to the other noble gases with eight electrons in the outermost shell. [189] Since helium behaves chemically like a noble gas, it is in the eighth main group with the other noble gases. When the noble gases were discovered around 1900, they were assigned to the zeroth main group, which no longer exists today. At that time, helium was at the top (i.e. in the first period) of the zeroth main group. Today the noble gases are positioned in the eighth main group according to IUPAC. [190]

Hydrogen can be positioned more clearly in the periodic table compared to helium, because it can assume the oxidation numbers 0 and +1 typical for the first main group and, like the lithium below, can form covalent bonds [191] [192] and is therefore counted among the alkali metals, [193] even if it is the only gaseous alkali metal and has a comparatively high electronegativity. Hydrogen forms alloy-like metal hydrides with some transition metals. [194] Nevertheless, due to its non-metallic chemical reactivity, hydrogen is occasionally classified in the seventh main group with the halogens. [195] Therefore, hydrogen is listed twice in some periodic tables, albeit rarely, in the first and seventh main group. [196] It was also proposed to sort hydrogen above carbon because its electronegativity, electron affinity and ionization potential correspond more closely to carbon, even if it can only react with one electron, in contrast to the representatives of the carbon group (fourth main group), that can react with four electrons. [197]

In order to take into account the different properties of hydrogen and helium, both are shown outside of the periodic table in rare cases. [198]

Lanthanoids and Actinides Edit

The classification of the lanthanides and actinides is relatively different compared to elements from other periods. Early attempts placed the lanthanoids and actinides between the main group elements. Bohuslav Brauner sorted the lanthanides and actinides below zirconium in 1902 - this arrangement was referred to as the "asteroid hypothesis" based on several asteroids in the same orbit, which Brauner described in a letter to Mendeleev in 1881. In 1922, Niels Bohr placed the lanthanides and actinides between the s-Block and the d-Block a. A periodic table was developed by Glenn T. Seaborg between 1944 and 1949, which shows the lanthanoids and actinides as footnotes below yttrium. [199] [200] However, it was also criticized that such a classification tears apart the representation of the periodic table. [201]

Scandium and yttrium are now relatively fixed, but the elements in the first subgroup below vary. Below yttrium are, depending on the representation, either the first representatives of the lanthanoids and actinides (lanthanum and actinium, i.e. in the order Sc-Y-La-Ac), or more rarely the last representatives of the lanthanoids and actinides (lutetium and lawrencium, i.e. in the order Sc-Y-Lu-Lr) or a gap with footnotes (i.e. in the order Sc-Y - * - *). These three variants depend on the discussion where the f-Block starts and ends. In a fourth variant, the third group is interrupted and an actinoid-lanthanide branch and a lutetium-lawrencium branch are inserted. [203] There are chemical and physical arguments for the variant with Lawrencium and Lutetium below yttrium, [204] [205] but this variant does not find a majority among experts on the subject of the periodic table. [206] Most chemists are unfamiliar with this discussion. [207] In 2015 the IUPAC set up a project group on the arrangement of the lanthanoids and actinides. [208] In January the project group published an interim report. In it she formulates three desiderata: 1) The order of the elements should follow their ordinal number. 2) The d-block is not intended to be split into two highly dissimilar parts. 3) The blocks should include two, six, ten and fourteen groups in accordance with the underlying quantum mechanical requirements. These are only possible with Sc-Y-Lu-Lr. [209]

Periodic table after discovering the elements

The dating of the discovery of such chemical elements, which have been known since prehistoric times or antiquity, is only imprecise and can vary by several centuries depending on the literature source. More reliable dating is only possible from the 18th century. Until then, only 15 elements were known and described as such: 12 metals (iron, copper, lead, bismuth, arsenic, zinc, tin, antimony, platinum, silver, mercury and gold) and three non-metals (carbon, sulfur and phosphorus). Most of the elements were discovered and scientifically described in the 19th century. At the beginning of the 20th century only ten of the natural elements were unknown. Since then, mainly difficult to access, often radioactive elements have been shown. Many of these elements do not occur naturally and are the product of man-made nuclear fusion processes. It was not until December 1994 that the two artificial elements Darmstadtium (Eka-Platin) and Roentgenium (Eka-Gold) were produced. Until the element names have been determined, new elements are identified with systematic element names.

This periodic table gives an overview of the discoverers or producers of the individual elements by clicking on the element identification. For the elements for which no discoverer / producer is known, the current historical state of knowledge is shown briefly under the overview plan.

  • C: Known since prehistoric times.
  • S: Known since prehistoric times, its elemental character was probably first recognized by Lavoisier.
  • Zn: Since around 1300 BC Processed in China.
  • Sb: Recent discoveries show the use of antimony by the Mesopotamian peoples for about 4000 years.
  • Hg: Known for about 3000 years.
  • Np - Og: The people named here as discoverers of the transuranic elements represent the participating research groups at the United Institute for Nuclear Research in Dubna, at the Lawrence Berkeley National Laboratory in Berkeley, at CERN in Geneva and at the GSI Helmholtz Center for Heavy Ion Research in Darmstadt.

2019: International Year of the Periodic Table Edit

The United Nations (UN) has declared 2019 the "International Year of the Periodic Table of the Chemical Elements" (IYPT 2019): With this, they want to arouse awareness worldwide of how chemistry can promote sustainable development and solutions to global challenges in energy, education, agriculture or Health can offer. In this way, the most recent discoveries and names of four "super-heavy" elements of the periodic table with the ordinal numbers 113 (Nihonium), 115 (Moscovium), 117 (Tenness) and 118 (Oganesson) are to be made better known. The dedication also coincides with the 150th anniversary of the development of the periodic table. [210] Events in Paris, Murcia and Tokyo will commemorate the event. [211]

Future expansions of the periodic table edit

Experiments to create synthetic elements continue and are expected to result in the creation of elements with atomic numbers above 118 as well.If these fit into the previous scheme, there will be a for the first time in the eighth period G- Lower shell (namely that of the fifth main shell) filled up. Since the G-Subshell contains nine orbitals that can hold 18 electrons, the eighth period will include a total of fifty (2 · 5 2) elements: Eight main group elements (filling the 8s- and the 8pLower shells), ten outer transition elements (filling the 7dLower shell), fourteen inner transition elements in which the 6f- Lower shell is filled, and another eighteen inner transition elements in which the 5g- Lower shell is filled up. [55] [212] The same would apply to the ninth period.

It is possible that the relativistic effects that increase with the ordinal number (see above) will blur the periodicities more and more. They influence the behavior of the electrons and thus the chemical properties so that they no longer necessarily have to correspond to the position of the element in the periodic table. [55] This is already indicated by the known elements: This is how the transition metals should 104Rutherfordium and 105Dubnium are similar in their behavior to the transition metals hafnium and tantalum, respectively, which are above in the periodic table. However, experiments show a behavior that is more similar to the actinides plutonium or protactinium. [55] The following elements 106Seaborgium and 107Bohrium, on the other hand, again show the behavior corresponding to their position in the periodic table. [55] 114As an element of the fourth main group, Flerovium should be similar to lead, but seems to behave more like a precious metal. [55]

In general, the heavier the atoms produced, the shorter their lifespan. Theoretical estimates suggest that from ordinal numbers of a little over 170 the lifetime of the generated atoms approaches zero, [55] [213] so that one can no longer speak of generated atoms at all. If applicable, this would be the theoretical upper limit for the scope of the periodic table.

Long period system edit

The medium-length form of the periodic table most commonly used nowadays (with 18 columns and outsourced to save space f-Block) has already been explained in detail. If you do not outsource the f-Blocks, which includes the lanthanoids and actinides, you get the so-called long form of the periodic table with 32 columns. In this representation, in contrast to the medium-long form, there are no interruptions in the sequence of the ordinal numbers. [214]

A first long-period system was proposed by Alfred Werner in 1905. [78] William B. Jensen recommended the long-period table because the lanthanides and actinides presented separately in shorter periodic tables would appear unimportant and boring to the students. [215] Despite the complete representation, the long-period system is seldom used because of its unwieldy format for letterpress printing. [216] A periodic table that goes beyond the atomic number 118 can be found under Extended Periodic Table.

Alternative periodic tables edit

The form of the periodic table by Dmitri Mendeleev has prevailed. Nevertheless, there were (and are) other suggestions for alternative arrangements of the elements according to their properties. In the first hundred years since Mendeleev's draft of 1869, an estimated 700 variants of the periodic table were published. [215] [217] [218] In addition to many rectangular variants, there were also circular, spherical, cube, cylinder, spiral, pyramidal, layered, flower, loop, [219] octagonal and triangular shapes Periodic tables. [220] The different forms mostly serve to emphasize certain properties. Most representations are two-dimensional. The first three-dimensional representation was published in 1862 by de Chancourtois before Mendeleev's periodic table. Another three-dimensional representation of several paper loops was published in 1925 by M. Courtines, [221], [222] and a layered representation was created by A. N. Wrigley in 1949. [223] [224] Paul-Antoine Giguère published in 1965 a periodic table compiled from several plates [225] and Fernando Dufour a tree-shaped representation in 1996. [226] [227] The periodic table by Tim Stowe [228] from 1989 was made including one color dimension is described as four-dimensional. [229]

In addition, there are more chemically and more physically oriented representations of the periodic table. [230] A chemically oriented periodic table was published in 2002 by Geoff Rayner-Canham for inorganic chemists, [231] in which tendencies and patterns as well as unusual chemical properties are emphasized. A physically oriented periodic table was published by Charles Janet in 1928, with a stronger focus on electron configuration and quantum mechanics, [232] as well as by Alper from 2010. [233] The latter was criticized, however, because of the lack of representation of the periodicity of the properties. [234] The mixed forms include the standard periodic system, which lists both chemical and physical properties such as oxidation numbers, electrical and thermal conductivities. [235] Its prevalence is attributed to the balance and practicality of the properties indicated. [236] [237] No alternative periodic table, but a clearly different looking representation is the short periodic table (see above), in which main and subgroups are nested in one another. Other classification methods are based on the natural occurrence of the elements in minerals (Goldschmidt classification) or on the crystal structure.

Periodic table of the elements with melting and boiling points, electronegativity and density

Element spiral by Benfey (1960) [238] [239]

Stowe's 3-D Periodic Table

Periodic table by Zmaczynski and Bayley

“In addition to predicting new elements and their expected properties, the periodic table has also proven to be invaluable when looking for promising research approaches in the production of new compounds. Chemists have internalized this way of thinking to such an extent that they are barely aware of how extremely difficult their task would be if they could not rely on periodic trends. Your work can be planned successfully because you can anticipate the effects of replacing an element or group in a compound. The prudent chemist always keeps an eye on the possibility that surprisingly new effects or unexpected factors may occur. "


Periodic table of the elements

Periodic table of the elements Periodic table.

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Manfred Weber, Frankfurt [MW1] (A) (28)
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- IUPAC group number (1989/1970)
- Name in German and Latin names
- Element symbol, atomic number (atomic number)
- relative atomic mass
- crystal structure etc.

- IUPAC group number (1989/1970)
- Name in German and English
- Element symbol, atomic number (atomic number)
- relative atomic mass
- main oxides and their chemical character
- atomic radius etc.

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Depiction

The following is the periodic table in its most well-known form today as the long-period table:

  • The elements are listed with their atomic number and their symbol.
  • as Periods the horizontal lines or rows are designated,
  • as groups the vertical columns.
  • the Peel refer to the shell model of atomic physics.

(A periodic table extended by atomic number 118 can be found under Extended Periodic Table.)

group
formerly (CAS group):
1
I & # 160A
2
II & # 160A
3
III & # 160B
4
IV & # 160B
5
V & # 160B
6
VI & # 160B
7
VII & # 160B
8
VIII & # 160B
9
VIII & # 160B
10
VIII & # 160B
11
I & # 160B
12
II & # 160B
13
III & # 160A
14
IV & # 160A
15
V & # 160A
16
VI & # 160A
17
VII & # 160A
18
VIII & # 160A
 
period peel
1 1
H
2
Hey
K
2 3
Li
4
Be
5
B.
6
C.
7
N
8
O
9
F.
10
No
L.
3 11
N / A
12
Mg
13
Al
14
Si
15
P.
16
S.
17
Cl
18
Ar
M.
4 19
K
20
Approx
21
Sc
22
Ti
23
V
24
Cr
25
Mn
26
Fe
27
Co
28
Ni
29
Cu
30
Zn
31
Ga

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With the most recently recognized element, Oganesson, the number of known elements increases to 118. But what is the most sensible way to sort this amount of basic chemical components? As early as 1869, the chemist Dmitri Ivanovich Mendeleev noted the elements known at the time in tabular form. This table, modified somewhat, is still one of the most important tools in chemistry. This production explains how it can be used and what information is in it.

Learning goals:
Relation to curricula and educational standards
The students
- Can reproduce the historical development process of the periodic table
- learn the basic structure of the periodic table (main groups, periods, ordinal number)
- get to know Bohr's atomic model and apply the knowledge gained
- can use the periodic table to make statements about the number of protons, electrons, neutrons, valence electrons and the probability areas of an element
- learn the definitions of the following terms: isotope, octet rule, electronegativity
- Gather information on selected main groups and elements
- describe the trends in atomic radius and ionization energy and can apply this to selected examples
- can derive the electron configuration based on the position of an element in the periodic table and draw conclusions about its reaction behavior
- Can name and describe possibilities for achieving the noble gas configuration (electron pair bond, ion bond) and explain the meaning of the valence electrons in this context.

Production features:
1 film, 6 sequences, 17 graphics, 6 PDF worksheets, 6 Word worksheets, 1 use in class, 1 film commentary / film text, 1 accompanying information

Technical requirements:
A current web browser must be installed.

Information:
Media type: Didactic FWU online medium
Recommended addressees: Grammar School 7.-11. Great
Production year: 2020 (2020)
License rights: yes
This medium is GEMA-free
Language: German German subtitles
Running time: 19 min

Subject areas:
Chemistry - Physical Chemistry - Atomic Structure, Periodic Table
Chemistry - Inorganic Chemistry - Elements


Electronegativity in the periodic table

the Electronegativity indicates the aspiration of an atom within a molecule To attract binding electrons. Electronegativity increases from bottom to top and from left to right in the periodic table.

Francium is the element with the least electronegativity | Source: Visualhunt

The values ​​of the Electronegativity are entered in the periodic table for each element. For example, sodium has an electronegativity of 0.9 and chlorine has an electronegativity of 3.2. The difference, which must always be positive, is 2.3 in this case. There is an ionic bond in the case of NaCl table salt.

Results in the difference of Electronegativity has a value from 0.1 to 0.4, it is a weak polar atomic bond. One atom therefore demands electron pairs a little more than the other atom. At values ​​below 0.1, the bond is non-polar and there is a covalent atomic bond.

A strong polar atomic bond is present at values ​​from 0.4 to 1.7. Is the difference the Electronegativity of the elements involved greater than 1.7, one speaks of a ionic compound. The binding electrons are completely transferred to a binding partner.

the Electronegativity is an indication of polarity, i.e. for the centers of charge and the ionic bond character of a bond. The bond becomes more polar, the more the bonded elements differ in their electronegativity.


Periodic Table of the Elements - Chemistry and Physics

All chemical elements are arranged in the PSE.
A field is provided for each element. The elements are classified according to the number of protons (according to the size) of their atoms.
The periods (left on the side) indicate how many electron shells are occupied. The main group indicates how many external electrons the element has. Elements with similar properties are in one column.
The number of neutrons indicates the reactivity. These elements are classified in separate rows as lanthanoids and actinides.

Atomic number
Electronegativity
Electron configuration
Atomic mass in u
Name symbol

Bullet points

- Every element with the same number of protons and electrons has an atomic number
- in the periodic table every element has an atomic number
- the 7 horizontal rows are called periods
- Elements with similar properties and a certain number of external electrons are in a column
- if elements are in the 8 main group, i.e. have 8 outer electrons, this is a particularly stable form which is also called the octet rule. The noble gases are then in this main group


Periodic table

That Periodic table (Long version Periodic table of the elements, abbreviated PSE or PSdE) is a list of all chemical elements, sorted according to increasing nuclear charge (atomic number). The list is subdivided into rows (periods) so that elements with similar chemical properties are in each column (main group / subgroup) of the resulting table. The name Periodic table (from Greek περίοδος períodos, German, handling, circulation, circulation ‘) [1] [2] indicates that many properties of the elements repeat themselves periodically with increasing atomic number.

The periodic table was presented in 1869 independently and almost identically by two chemists, first by the Russian Dmitri Mendeleev (1834–1907) and a few months later by the German Lothar Meyer (1830–1895). Historically, the periodic table was of particular importance for the prediction of as yet undiscovered elements and their properties, since the properties of an element can be approximately predicted if the properties of the surrounding elements in the periodic table are known. Today it is mainly used as a clear organization scheme for the elements and to determine possible chemical reactions.


Video: PERIODIC TABLE OF ELEMENTS Animation (August 2022).