Periodic table of elements

History of the periodic table of elements

 * First attempt to divide the elements - German chemist Johann Döbereiner - 19th century - triad rule - in a triad of elements (e.g. Li, Na, K) the middle element has the average properties of the outermost elements.
 * 2nd half of the 19th century - English chemist John Newlands - rule of octaves - first arrangement of elements according to their atomic weight (he used the analogy of the similarity of elements with the similarity of tones in music).
 * Dmitry Ivanovich Mendeleev published his Periodic Table in the Journal of the Russian Chemical Society in 1869.
 * He ranked the elements according to their relative atomic weight (63 elements were known at the time). He formed groups of elements with similar properties, breaking the list into a series of rows and columns, thus creating the periodic table of elements.
 * The original definition of the Periodic Law: The properties of the elements are a periodic function of their atomic weights.
 * Mendeleev left spaces in the table for elements yet to be discovered and also predicted their chemical properties with high accuracy.
 * The gradual discovery of further knowledge proved that the Periodic Law is one of the fundamental laws of nature.
 * Today, based on the knowledge of the internal structure of atoms, the periodic law is formulated as follows: The properties of the elements are a periodic function of their proton numbers.

Periodic table of elements

 * The short and long forms (the most commonly used, it also contains lanthanoids and actinoids), of the non-traditional modifications, the main example is the table of elements according to their occurrence on Earth.
 * The elements in the periodic table are arranged in 7 horizontal series - periods, which are denoted by Arabic numerals, and in 18 vertical groups denoted by Roman numerals. The groups are further divided into major subgroups (I.A...VIII.A) and minor groups (I.B...VIII.B).
 * The period number is the same as the maximum principal quantum number, i.e. the number of the valence layer of electrons.
 * The first period begins with hydrogen H, which has an electron configuration of 1s1, and ends with helium, which has an electron configuration of 1s2
 * The element of each subsequent period begins with the electron configuration ns1
 * According to the filling of certain types of orbitals with valence electrons, the elements can be divided into:
 * non-transitional elements:
 * s-elements - valence electrons are in ns orbitals
 * p-elements – valence electrons are in ns np orbitals
 * transitional elements:
 * d-elements – valence electrons are in ns (n – 1)d
 * internally transient elements
 * f-elements - lanthanoids and actinoids - valence electrons complement the (n - 2)f orbitals
 * the number of valence electrons is identical to the number of the group in which the element is located (exception - some elements of group VIII.B)
 * From the position of the element in the table we can determine:
 * the electron configuration of the atom
 * the physical and chemical properties of the element (they are periodic in nature)
 * many properties and possible compounds can be predicted

First ionization energy of atoms of elements

 * the energy required to remove one electron from each atom in 1 mole of gaseous atoms in the ground state
 * the ionisation energy value characterises the ability of an atom to give up electrons
 * the main factors affecting the 1st ionization energy:
 * the size of the positive charge on the nucleus of an atom - the increase in the nucleus of an atom (with increasing proton number) will also cause an increase in the 1st ionization energy - in layman's terms: the larger the nucleus, the larger the ionization energy
 * the distance of the electron from the nucleus - the further the valence layer of electrons is from the nucleus, the less ionization energy is required to detach the electron
 * electron capping effect - refers to atoms that contain inner electron layers. The inner electrons "dampen" the effect of the attractive nucleus. The greater the covering effect, the less ionization energy is needed to strip the electron
 * Hence: the first ionization energy of atoms of elements increases from left to right in periods and decreases downwards in groups
 * while the combined effect of distance and overlap weakens the effect of increasing the positive charge of the nucleus

Atomic radii of elements

 * atomic radii of the elements in the period decrease from left to right and in groups increase downwards

Elektronegativity

 * Electronegativity is the ability of an element to attract binding electrons to itself (a measure of its ability to attract its binding electrons)
 * it rises from left to right in a period and falls downward in a group
 * electronegative elements: trying to reach the electron configuration of the nearest noble gas - the so-called octet rule
 * electropositive elements: trying to reach the configuration of the preceding noble gas

Metals, non-metals and semi-metals

 * metals - electropositive elements, mainly s-elements and p-elements with a small number of electrons in the valence layer, d and f-elements. They easily form cations, in the solid state they form a metallic lattice
 * non-metals - electronegative elements, especially p-elements with a larger number of electrons in the valence layer (for example: halogens, chalcogens, hydrogen, carbon, nitrogen,...). They easily form anions.
 * Semi-metals - have some properties of metals and some of non-metals.
 * Metallic properties of elements increase in groups downwards and in periods from right to left

Related articles

 * Atom
 * Atomic nucleus

Literature used


Categoy:Nuclear and atomic physics