What are typical elements 1

Element groups in the periodic table



 
The periodic table is laid out in such a way that the elements are in vertical columns in 18 groups are divided. This division follows the current IUPAC convention. Depending on its meaning, the term element in the periodic table means an elementary substance or an elementary atom.

The previously used division into "main groups" and "subgroups" using the Roman numerals according to the CAS Convention has not been valid since 1986. The Roman numerals I to VIII are occasionally used in German-language textbooks for the “main groups”, internationally and also in the digitally available work they are no longer used. The main group can be determined from the final digit of the Arabic group number: Boron, for example, is in the 13th group, which, according to the old representation, would be III. Main group.

The 7 horizontal lines in the periodic table are called Periods. The distribution of electrons in their atomic shells repeats itself periodically in the elements: Elements standing one below the other have a similar electron distribution in the valence shells in their atomic structure. This also explains why they have similar substance properties. The last element of a period has the noble gas configuration. The occupation of the s orbital of a new shell leads to a jump into a new period.
 


1st group alkali metals
 
In the 1st group, all alkali metals are one below the other, hydrogen not counting among these metals. The name is derived from the Arabic word al kali which describes the potassium carbonate contained in the plant ash, which forms an alkaline solution with water. Lithium, sodium, potassium, rubidium, cesium and francium react exothermically with hissing water to form hydrogen and alkali metal hydroxides. These form an alkaline solution with water. The reactivity increases within the group from top to bottom with decreasing electronegativity, rubidium, cesium and francium react explosively with water. In the case of alkali metals, the oxidation state +1 generally occurs. Typical reactions occur, for example:

Alkali metal + water Hydrogen + alkali metal hydroxide
Alkali metal + hydrogen Alkali metal hydride
Alkali metal + oxygen Alkali metal oxide (oxide, peroxide, hyperoxide)
 
The high reactivity of the alkali metals is explained by the single valence electron. This can easily be given away. The atoms of the alkali metals have a very low electronegativity. Only a small amount of ionization energy is required to withdraw this single electron.



2nd group of alkaline earth metals
 
The second group contains the alkaline earth metals beryllium, magnesium, calcium, strontium, barium and radium. With the exception of beryllium, they like to react with water. This creates hydrogen and alkaline earth hydroxides, which form an alkaline solution with water. The oxides of these elements were formerly known as "alkaline earths".

Alkaline earth metal + water Hydrogen + alkaline earth metal hydroxide
Alkaline earth metal + hydrogen Alkali metal dihydride
Alkaline earth metal + oxygen Alkaline earth metal oxide (also barium peroxide)
 
The atoms of the alkaline earth metals with their two valence electrons gladly give up these electrons so that they reach the noble gas configuration. Therefore, the oxidation state +2 generally occurs.



3rd-12th Transition elements group
 
In the transition elements of the 3rd to the 12th group, the d orbitals are occupied by electrons. This explains in part why there are a large number of possible oxidation states associated with these elements. The transition elements used to be called "subgroup elements". This subdivision is no longer common internationally today.
 
 
3
4
5
6
7
8
9
10
11
12
Sc
Ti
V.
Cr
Mn
Fe
Co
Ni
Cu
Zn
Y
Zr
Nb
Mon
Tc
Ru
Rh
Pd
Ag
CD
La / Lu
Hf
Ta
W.
re
Os
Ir
Pt
Au
Ed
Ac / Lr
Rf
Db
Sg
Bra
Hs
Mt
Ds
Rg
Cn
 

Lanthanides, lanthanides
 
 
Correct representation in the present periodystem:IUPAC assigns all lanthanides (including lanthanides) to group 3 in the periodic table. According to IUPAC, the lanthanides begin with No. 57 Lanthanum, they end with No. 71 Lutetium. In the 15 lanthanides, predominantly the 4f orbitals of the atoms are gradually occupied with electrons. Lanthanum (5d1) and lutetium (5d1 with filled 6s2) make an exception. There are also several exceptions for actinides.

As elementary substances, all lanthanides are very similar to lanthanum: They are silvery, shiny and reactive metals that oxidize in moist air and react with water or dilute mineral acids to generate hydrogen and hydroxide. Some lanthanides are pyrophoric in their finely divided state, they can ignite by themselves. The chemical stability tends to increase to the right. All lanthanides have paramagnetic properties.

Problematic representation in a long period system:
According to the d-orbital occupation, the elementary atoms of lanthanum and lutetium can be assigned to the 3rd group. If one were to define the lanthanides only in terms of the occupation of the f orbitals, then lanthanum and lutetium would not be included, in contradiction to IUPAC. They would both be assigned to group 3 and that would be difficult to represent in a printed periodic table. In a two-dimensional long-period system, either lanthanum or lutetium would have to be assigned to group 3. For this reason, the representation in a long period system is problematic. The same problem occurs with the actinides.

 
  
Actinides, actinides
 
According to IUPAC, the actinides (also actinides) begin with the actinium and end with the lawrencium. All elements are very similar to actinium: They are silvery, shiny, reactive metals that are highly toxic and radioactive. The artificially produced elements beyond uranium are also called transuranium elements.
 
 
Lantha-
nide
La
Ce
Pr
Nd
Pm
Sm
Eu
Gd
Tb
Dy
Ho
He
Tm
Yb
Lu
Acti-
nide
Ac
Th
Pa
U
Np
Pooh
At the
Cm
Bk
Cf
It
Fm
Md
No
Lr


13. Group Boron Group
 
The 13th group includes all elements below boron such as aluminum, gallium, indium or thallium. In contrast to the alkaline earth metals, these elements hardly react with water. Pure, unprotected aluminum reacts with water, but a protective oxide layer forms immediately.

The atoms in this group have three valence electrons. Therefore the oxidation state +3 occurs mainly, in a few cases +1 or +2.


14. Group carbon group
 
In the 14th group with the elements carbon, silicon, germanium, tin or lead, the elements differ greatly in their chemical and physical properties, as the dividing line between the metals and the non-metals runs through this group. The metallic character increases in the direction of lead. One thing they have in common is the formation of a dioxide when it reacts completely with pure oxygen.

Due to the four valence electrons, the atoms mainly have the oxidation states +2, +4 and −4.
 


15. Group nitrogen group or pnicogens
 
The 15th group contains nitrogen, phosphorus, arsenic, antimony or bismuth. The term pnicogene recommended by the IUPAC is derived from the Greek word for suffocation. Nitrogen is a non-metal, while bismuth has typical metallic properties. In the case of phosphorus, arsenic and antimony, there are both non-metallic and metallic modifications. Red and white phosphorus are non-metals, black and purple phosphorus are metallic, the latter occur in crystalline form and conduct electricity.

The elementary atoms of the nitrogen group have five valence electrons. For the noble gas configuration, three electrons can be absorbed or five electrons can be emitted. The oxidation states therefore predominantly −3 or +5, but +2, +3 or +4 can also be found sporadically.
 


16. Chalcogen group
 
The 16th group includes the chalcogens. They include oxygen, sulfur, selenium, tellurium or polonium. The name is derived from the Greek and means something like "ore builder": Chalcogens sometimes react with each other and with the metals to form oxides and sulfides, which make up the largest proportion of the minerals for ore extraction. Chalcogen dioxides react with water to form acids with the empirical formula H.2XO3, For example, sulfur dioxide forms sulfurous acid with water, selenium dioxide forms selenious acid. The chalcogens form chalcogen hydrogen with hydrogen with the empirical formula H.2X.

Metal + oxygen Metal oxide
Metal + chalcogen Metal chalcogenide (sulfide, selenide, telluride)
Chalcogen dioxide + water Chalcogenic acid
Chalcogen + hydrogen Chalcogen hydrogen
 
To achieve the noble gas configuration, their atoms can donate six electrons or accept two. For this reason, the oxidation states -2 and +6 occur in the chalcogens, in isolated cases there are also −1, +2 and +4 and others. Typical reactions are:



17. Group halogens

 
In the 17th group are the halogens, which, according to the Greek, means "salt formers". Fluorine, chlorine, bromine and iodine sometimes react violently with the metals and form the corresponding salts. The halogens are soluble in water; fluorine even reacts with water to form hydrogen fluoride and oxygen. The halogens react with hydrogen in an exothermic reaction to form the hydrogen halides. The chlorine-oxyhydrogen gas reaction is an example of this. The resulting hydrogen halides dissolve in water as Brönsted acids: Hydrogen chloride, for example, dissolves in water to form hydrochloric acid.

Halogen + water Halogen solution (exception: fluorine reacts with water)
Metal + halogen Metal halide
Halogen + hydrogen Hydrogen halide
Hydrogen halide + water Hydrogen halide solution

The electronegativity of halogens is relatively high: they like to accept an electron in order to achieve the noble gas configuration. Compared to less electronegative elements, the halogens often occur in the −1 oxidation state. This explains the extreme reactivity and also the high toxicity of the halogens. Two atoms of the halogens immediately react with each other: They always form diatomic molecules.


18. Noble Gases Group

 
The noble gases in the 18th group form the end of each period in the periodic table. Helium, neon, argon, krypton, xenon or radon are gaseous non-metals that are extremely reluctant to react with other substances, they behave so “nobly”.

The noble gases already have the "ideal" noble gas configuration on the valence shell, so that a very high ionization energy has to be used to remove electrons from the shell. For this reason there are very few noble gas compounds.