The periodic table is the table arrangement of the chemical elements, commanded by their atomic number, electron configuration, and repetitive chemical properties, whose structures show periodic trends . Generally, in one line (period) elements are metal to the left, and non-metallic to the right, with elements having similar chemical behavior placed in the same column. The table rows are usually called periods and columns are called groups. Six groups have received the names and numbers assigned: for example, elements of group 17 are halogens; and group 18 is a noble gas. Also shown are four plots of simple rectangles or blocks associated with the filling of different atomic orbitals.
Organization of the periodic table can be used to derive relationships between various element properties, but also the predicted chemical and behavioral properties of undiscovered or newly synthesized elements. Russian Chemist Dmitri Mendeleev was the first to publish a recognizable periodic table in 1869, developed primarily to illustrate the periodic trends of the later elements. He also foresaw some properties of unidentified elements that are expected to fill the void in the table. Most of his predictions proved to be true. Mendeleev's ideas have been slowly expanded and perfected with the discovery or synthesis of further new elements and the development of new theoretical models to explain chemical behavior. The modern periodic table now provides a useful framework for analyzing chemical reactions, and continues to be widely used in chemistry, nuclear physics and other sciences.
All elements of the atomic number 1 (hydrogen) to 118 (oganesson) have been discovered or synthesized, completing the first seven lines of the periodic table. The first 98 elements exist in nature, although some are only found in small quantities and others are synthesized in laboratories before they are found in nature. The atomic numbers for elements 99 through 118 are only synthesized in laboratories or nuclear reactors. The synthesis of elements that have higher atomic numbers is currently being pursued: these elements will start the eighth row, and theoretical work has been done to suggest possible candidates for this extension. Many synthetic radionuclides of naturally occurring elements have also been produced in the laboratory.
Video Periodic table
Ikhtisar
Each chemical element has a unique atomic number ( Z ) representing the number of protons in its nucleus. Most elements have different numbers of neutrons between different atoms, with these variants being called isotopes. For example, carbon has three natural isotopes: all of its atoms have six protons and most have six neutrons as well, but about one percent has seven neutrons, and a very small fraction has eight neutrons. Isotopes are never separated in the periodic table; they are always grouped together under one element. Elements without stable isotopes have the atomic mass of their most stable isotopes, where such mass is shown, listed in parentheses.
In the standard periodic table, the elements are listed in the order of increasing the atomic number Z (the number of protons in the atomic nucleus). The new line ( period ) starts when the new electron shell has the first electron. Columns ( groups ) are determined by the atomic electron configuration; elements with the same number of electrons in a particular subshell fall into the same column (eg oxygen and selenium are in the same column because they have four electrons in the outer subshell). Elements with similar chemical properties generally fall into the same group in the periodic table, although in the f-block, and for some respect in the d-block, elements in the same period tend to have the same properties. Thus, it is relatively easy to predict the chemical properties of an element if one knows the properties of the elements around it.
In 2016, the periodic table has 118 confirmed elements, ranging from element 1 (hydrogen) to 118 (oganesson). Elements 113, 115, 117 and 118, the latest findings, formally confirmed by the International Union of Applied Pure and Applied Chemistry (IUPAC) in December 2015. Their proposed names, nihonium (Nh), moscovium (Mc), tennessine (Ts) and oganesson (Og) respectively, announced by IUPAC in June 2016 and made official in November 2016.
The first 94 elements occur naturally; the remaining 24, americium to oganesson (95-118), occurs only when synthesized in the laboratory. Of 94 natural elements, 83 are primordial and 11 occur only in the primordial element decay chains. No heavier element of einsteinium (element 99) ever observed in macroscopic amounts in its pure form, nor does it have astatin (element 85); franium (element 87) is only photographed in the form of light emitted from the microscopic number (300,000 atoms).
Maps Periodic table
Grouping method
Groups
Group or family is a vertical column in the periodic table. Groups usually have more significant periodic trends than periods and blocks, described below. The modern quantum mechanics theory of atomic structure explains group trends by proposing that elements in the same group generally have the same electron configuration in their valence shells. As a result, elements in the same group tend to have chemical similarities and show a clear tendency in properties with increasing atomic numbers. In some parts of the periodic table, such as d-block and f-block, horizontal similarity can be as important as, or more clearly than, vertical similarity.
Under international naming conventions, groups are numerically numbered from 1 to 18 from the leftmost column (alkali metal) to the far right column (noble gas). Previously, they were known by roman numerals. In America, the Roman numerals are followed by "A" if the group is in block s or p, or "B" if the group is in block d. The Roman numerals used correspond to the last digits of today's naming conventions (eg group element 4 is group IVB, and group of 14 elements is group IVA). In Europe, the writing is similar except that "A" is used if the group is before group 10, and "B" is used for the group included and after group 10. In addition, groups 8, 9 and 10 are used to be treated as a group of three , known collectively in both notations as group VIII. In 1988, the new IUPAC naming system began to be used, and the old group names were no longer in use.
Some of these groups have been given trivial names (not systematic), as seen in the table below, although some are rarely used. Groups 3-10 do not have a trivial name and are only referenced by their group number or by the names of the first members of their group (such as "group scandium" for group 3), because they display fewer trend similarities and/or verticals.
Elements in the same group tend to exhibit patterns in atomic radius, ionization energies, and electronegativity. From top to bottom in groups, the radius of atomic elements increases. Because there are more energy levels, valence electrons are found farther from the nucleus. From the top, each sequential element has a lower ionisation energy because it is easier to remove electrons because the atoms are bound less tightly. Similarly, a group has a decrease in electronegativity upwards due to the increasing distance between valence and nuclear electrons. There are exceptions to this trend: for example, in Group 11, electronegativeness increases further down the group.
Period
Period is the horizontal row in the periodic table. Although the group generally has a more significant periodic trend, there are areas where horizontal trends are more significant than vertical group trends, such as f-blocks, where lanthanides and actinides form two substantial horizontal element sequences.
Elements in the same period showed trends in atomic radius, ionisation energy, electron affinity, and electronegativity. Moving left to right across a period, the radius of the atom usually decreases. This happens because each sequential element has additional protons and electrons, which causes the electrons to draw closer to the nucleus. The decrease in radius of these atoms also causes the ionization energy to increase as it moves from left to right across a period. The more bonded an element is, the more energy it will take to remove the electrons. Electronegativity increases in the same way as ionization energy because of the pull given to electrons by the nucleus. Electron affinity also shows a slight trend in one period. The metal (left side of a period) generally has a lower electron affinity than nonmetals (right side of a period), with the exception of the noble gases.
Block
The specific area of ​​the periodic table can be referred to as the block in recognition of the order in which the electron shell of the element is filled. Each block is named according to the subshell where the "last" electron is accidentally located. Block s consists of the first two groups (alkali metal and alkaline earth metals) as well as hydrogen and helium. P-block consists of the last six groups, namely groups of 13 to 18 in IUPAC group numbering (3A to 8A in American group numbering) and contain, among other elements, all metaloids. Block d consists of groups 3 to 12 (or 3B to 2B in American group numbering) and contains all transition metals. The f-block, often offset under the rest of the periodic table, has no group number and consists of lanthanides and actinides.
Metal, metaloid, and nonmetal
According to their physical and chemical properties together, the elements can be classified into major categories of metals, metalloids and nonmetals. The metals are generally shiny, producing highly solids that form alloys with each other and the non-metallic salt-like ionic compounds (other than noble gases). The majority of nonmetals are colorless or colorless insulation gases; not metals which form other non-metallic compounds, have covalent bonds. Among the metals and nonmetals are metalloids, which have medium or mixed properties.
Metals and non-metals can be further classified into subcategories showing gradients from metallic to non-metallic properties, as they move from left to right in rows. The metal can be subdivided into highly reactive alkali metals, via less reactive alkaline earth metals, lanthanides and actinides, via arcetipal transition metals, and ends in physically and chemically weak post-transition metals. Nonmetals can be easily divided into poliatomic nonmetals, which are closer to metalloids and show some new metallic characters so; non-metatomically substantially non-metallic, non-metallic, and monoatomically precious metals and gases that are almost inertomic. Specific groupings such as refractory metals and precious metals, are examples of subsets of transition metals, also known and sometimes denoted.
Placing elements into categories and sub-categories based solely on imperfectly distributed properties. There is a large difference in properties in each of the overlapping categories that can be seen on the border, as is the case with most classification schemes. Beryllium, for example, is classified as an alkaline earth metal although amphoteric chemistry and its tendency to form covalent compounds are attributes of weak or post-transition metal. Radon is classified as a nonmetallic noble gas but has some cationic chemistry which is a characteristic of metals. Other classification schemes are possible such as the division of elements into categories of mineralogical occurrences, or crystal structures. Categorizing the elements in this mode dates back to at least 1869 when Hinrichs wrote that a simple boundary line can be placed on the periodic table to show elements that share a common property, such as a metal, not a metal, or a gas element.
Trends and periodic patterns
Electron configuration
Configurations of electrons or electron organizations that orbit neutral atoms show repetitive patterns or periodicity. Electrons occupy a series of electron shells (numbered 1, 2, and so on). Each shell consists of one or more sub-groups (named s, p, d, f, and g). As the atomic number increases, the electrons increasingly fill the shell and these subshells are more or less in line with Madelung's rules or energy ordering rules, as shown in the diagram. The electron configuration for the neon, for example, is 1s 2 2s 2 2p 6 . With the atomic number ten, the neon has two electrons in the first shell, and eight electrons in the second shell; there are two electrons in the subshell s and six in the p subshell. In terms of the periodic table, the first electron occupies a new shell according to the beginning of each new period, this position is occupied by hydrogen and alkali metals.
Since the properties of an element are largely determined by their electron configuration, the properties of the elements also exhibit repetitive patterns or periodic behavior, some examples shown in the diagram below for atomic radius, ionisation energy and electron affinity. This is the nature of periodicity, a manifestation noticed long before the underlying theory is developed, leading to the formation of periodic law (the properties of recurrent elements at various intervals) and the formulation of the first periodic table.
Atomic radius
The radius of the atom varies in a predictable way and can be described throughout the periodic table. For example, the radius is generally decreased throughout each period of the table, from alkali metals to noble gases; and up each group. The radius increases sharply between the noble gases at the end of each period and the alkali metal at the beginning of the next period. The tendency of atomic radius (and various other chemical and physical properties of the elements) can be explained by the electron shell theory of atoms; they provide important evidence for the development and confirmation of quantum theory.
The electrons in 4f-subshell, which progressively fill the lanthanide series, are not particularly effective in protecting the increased nuclear charge from further shells. The elements immediately following the lanthanides have atomic radii that are smaller than expected and which are almost identical to the radius of the atomic elements just above them. Hafnium therefore has almost the same atomic radius (and chemistry) as zirconium, and tantalum has atomic radii similar to niobium, and so on. This is known as lanthanide contraction. The effect of lanthanide contraction is seen to platinum (element 78), after which it is covered by a relativistic effect known as the inert pair effect. The d-block contraction, which is a similar effect between the d-block and the p-block, is less prominent than the lanthanide contraction but arises from a similar cause.
Energy ionization
The first ionisation energy is the energy needed to remove one electron from the atom, the second ionisation energy is the energy needed to remove the second electron from the atom, and so on. For certain atoms, successive ionization energies increase with the degree of ionization. For magnesium, for example, the first ionization energy is 738 kJ/mol and the second is 1450 kJ/mol. Electrons in closer orbitals experience greater electrostatic attraction; thus, their displacement requires more energy. The ionization energy becomes larger and to the right of the periodic table.
Large jumps in consecutive molar ionization energies occur when removing electrons from a noble gas configuration (complete electron electrons). For magnesium again, the first two ionization energies of magnesium given above correspond to the removal of two 3s electrons, and the third ionization energy is much greater than 7730 kJ/mol, to remove 2p electrons from a very stable neon. Mg 2 configuration. A similar jump occurs in the ionization energy of other third-line atoms.
electronegativity
Electronegativity is the tendency of atoms to attract the same pair of electrons. The electronegativity of the atom is affected by the atomic number and the distance between the valence electrons and the nucleus. The higher the electronegativity, the more elements attract the electrons. This was first proposed by Linus Pauling in 1932. In general, electronegativity increases as it moves from left to right over the period, and decreases in the declining group. Therefore, fluorine is the most electronegative of the elements, while cesium is the least, at least of those elements whose data are available.
There are some exceptions to this general rule. Gallium and germanium have higher electronegativity than aluminum and silicon due to d-block contraction. The elements of the fourth period immediately after the first line of the transition metal have very small radius of the atom because the 3d electrons are not effective in protecting the increased nuclear charge, and the smaller atomic size is correlated with the higher electronegativity. Very high lead electronegativity, especially when compared to thallium and bismuth, appears to be an artifact of data selection and data availability. Calculation methods other than the Pauling method show a normal periodic trend for these elements.
The electron equation
The atomic electron affinity is the amount of energy released when an electron is added to a neutral atom to form a negative ion. Although the affinity of electrons varies greatly, several patterns appear. Generally, nonmetals have more positive electron affinity values ​​than metals. Chlorine is very attractive to extra electrons. The similarities of noble gas electrons have not been measured conclusively, so they may or may not have few negative values.
Electron relations generally increase over the period. This is due to the filling of the atom's valence shell; the group 17 atom releases more energy than the group of 1 atoms to obtain the electrons as it obtains a charged valence shell and is therefore more stable.
The declining tendency of affinity groups to fall in electrons will be expected. The additional electrons will enter the orbital farther away from the nucleus. Thus this electron will be less attracted to the nucleus and will release less energy when added. In deriving a group, about one-third of the anomalous elements, with heavier elements have higher electron affinities than those that are lighter. Mostly, this is caused by a bad shield by electrons d and f. The uniform decrease in electron affinity applies only to group 1 atoms.
Metallic characters
The lower the value of ionisation energy, electronegativity and the affinity of electrons, the more metallic characters the element possesses. In contrast, non-metallic characters increase with a higher value than this property. Given the periodic trend of these three properties, the metal character tends to decrease over a period (or row) and, with some deviations (mostly) due to poor screening of the nucleus by d and f electrons, and relativistic effects, tends to increase. down to group (or column or family). Thus, the most metallic elements (such as cesium and franium) are found in the lower left of the traditional periodic table and the most nonmetallic elements (oxygen, fluorine, chlorine) in the upper right. The combination of horizontal and vertical trends in metallic characters describes the staircase-shaped dividing line between metal and non-metal found on some periodic tables, and the practice sometimes categorizes some elements adjacent to that line, or elements adjacent to those elements, as metaloid.
Connect or bridge the group
From left to right across four blocks of long periodic table forms or 32 columns is a set of groups of connecting or bridging elements, which lie roughly between each block. These groups, such as the metaloid, show properties in between, or which are a mixture of, groups on either side. Chemically, the elements of group 3, scandium, yttrium, lanthanum and actinium behave very much like alkaline earth metals or, more generally, block metal but have some physical properties d blocking transition metal. Lutetium and lawrencium, at the end of the block f , can form other bridging or bridging groups. Lutetium behaves chemically as lanthanides but exhibits a mixture of lanthanides and the physical properties of transition metals. Lawrensium, as an analogue of luteium, will feature such characteristics. Coin metals in clusters 11 (copper, silver, and gold) are chemically capable of acting as transition metals or primary group metals. Volatile groups of 12 metals, zinc, cadmium and mercury are sometimes regarded as a block link d to the p block. Logically they are d block elements but they have some transition metallic properties and are more like their neighboring blocks p in group 13. The relatively inertia of noble, in group 18, the bridge the most reactive group of elements in the periodic table - halogens in groups of 17 and alkali metals in group 1.
History
The first systemization attempt
In 1789, Antoine Lavoisier published a list of 33 chemical elements, grouping them into gases, metals, nonmetals, and earth. The chemist spent the next century looking for a more appropriate classification scheme. In 1829, Johann Wolfgang DÃÆ'¶bereiner observed that many elements can be grouped into triads based on their chemical properties. Lithium, sodium, and potassium, for example, are grouped together in triads as reactive soft metals. DÃÆ'¶bereiner also observes that, when arranged with atomic weights, the second member of each triad is approximately the average of the first and the third; this is known as the Triad Law. The German chemist Leopold Gmelin worked with this system, and in 1843 he had identified ten triads, three groups of four, and one group of five. Jean-Baptiste Dumas published a work in 1857 that explains the relationship between various metal groups. Although various chemists are able to identify relationships between small groups of elements, they have not yet developed a scheme that includes everything.
In 1857, German chemist August KekulÃÆ'Ã… © observed that carbon often has four other atoms attached to it. Methane, for example, has one carbon atom and four hydrogen atoms. This concept is finally known as valence; different elements are bound to different numbers of atoms.
In 1862, Alexandre-Emile BÃ © Â © guyer de Chancourtois, a French geologist, published the early form of the periodic table, which he called helical telluric or screw. He was the first to notice the periodicity of the elements. With elements arranged in a spiral on a cylinder with a sequence of increasing atomic weights, de Chancourtois suggests that elements of a similar nature seem to occur at regular intervals. His chart includes several ions and compounds other than elements. His paper also uses geological terms rather than chemical terms and does not include diagrams; as a result, it received little attention until the work of Dmitri Mendeleev.
In 1864, Julius Lothar Meyer, a German chemist, published a table with 44 elements composed by valence. The table shows that elements with similar properties often share the same value. Simultaneously, the English chemist William Odling published the arrangement of 57 elements, arranged according to the atomic weight. With some irregularities and loopholes, he sees what appears to be the atomicity of the atomic weights between the elements and that this corresponds to the "group they normally receive". Odling alluded to the idea of ​​periodic law but did not pursue it. He then proposed (in 1870) the classification of valence-based elements.
The British chemist John Newlands produced a series of papers from 1863 to 1866 noting that when they were listed in order of increasing atomic weights, similar physical and chemical properties were repeated at interval eight; he likened that periodicity to the musical octave. This Octave Law is called ridiculed by Newlands contemporaries, and the Chemical Society refuses to publish his work. However, Newlands is still able to compile tables of elements and use them to predict the presence of missing elements, such as germanium. The Chemical Society only recognizes the importance of its invention five years after they are credited Mendeleev.
In 1867, Gustavus Hinrichs, an American-born Danish-born academic chemist, published a spiral periodic system based on atomic and weight spectra, and chemical similarity. His work is regarded as idiosyncratic, plush and labyrinthine and this may have been against his confession and acceptance.
Mendeleev table
Russian chemistry professor Dmitri Mendeleev and German chemist Julius Lothar Meyer separately published their respective periodic tables in 1869 and 1870. Mendeleev's table is the first published version; that Meyer was an expanded version of its chart (Meyer) in 1864. They built their tables by listing the elements in rows or columns in the order of atomic weights and starting new rows or columns when the characteristic elements began to recur..
The admission and acceptance given to Mendeleev's desk stems from two decisions he made. The first is to leave the gap in the table when it seems the relevant element has not been found. Mendeleev was not the first chemist to do so, but he was the first to be recognized using trends in his periodic table to predict the properties of missing elements, such as gallium and germanium. The second decision is to sometimes ignore the order suggested by atomic weights and switch adjacent elements, such as tellurium and iodine, to further classify them into chemical families.
Mendeleev published in 1869, using atomic weights to organize elements, information that could be determined for a fair precision of his time. The atomic weight works well enough to allow Mendeleev to accurately predict the properties of missing elements.
After the discovery, in 1911, by Ernest Rutherford of the atomic nucleus, it was proposed that the number of integers of nuclear charge is identical to the sequential place of each element in the periodic table. In 1913, Henry Moseley used X-ray spectroscopy to confirm this proposal experimentally. Moseley determines the nuclear charge value of each element, and suggests that Mendeleev's reservations actually put the elements in sequential order with nuclear charges. The nuclear charge is identical to the number of protons, and determines the atomic number (Z) of each element. Using atomic numbers gives a definitive, sequence-based sequence for an element. Moseley predicted, in 1913, that the only missing element between aluminum (Z = 13) and gold (Z = 79) was Z = 43, 61, 72, and 75, all of which were later discovered. The atomic number is the absolute definition of an element, and provides a factual basis for ordering the periodic table. Periodic tables are used to predict the properties of new synthetic elements before they are produced and studied.
Second version and further development
In 1871, Mendeleev published his periodic table in a new form, with groups of similar elements arranged in columns rather than in rows, and the columns were numbered I to VIII corresponding to the oxidation status of the element. He also provides detailed predictions for the properties of elements that he has previously recorded missing, but must exist. This gap is then filled as chemists discover additional natural elements. It is often stated that the last natural element found was franium (referred to by Mendeleev as cesium cesium) in 1939. Plutonium, produced synthetically in 1940, was identified in small amounts as natural. elements that occurred in 1971.
The popular periodic table layout, also known as the general or standard form (as pointed out at various other points in this article), is caused by Horace Groves Deming. In 1923, Deming, an American chemist, published the periodic table of short forms (Mendeleev style) and medium (18-column). Merck and Company prepared a handout form from Deming's 18-column middle table, in 1928, which was widely circulated in American schools. In the 1930s, Deming's tables appeared in textbooks and chemical encyclopedias. It has also been distributed for years by the Sargent-Welch Scientific Company.
With the development of modern quantum mechanics theory of the electron configuration in atoms, it becomes clear that each period (row) in the table corresponds to the filling of quantum electrons. The larger atoms have more electron sub-shells, so the tables then require longer periods.
In 1945, Glenn Seaborg, an American scientist, made the suggestion that actinide elements, such as lanthanides, fill sub-levels. Prior to this time actinides were thought to form the fourth d-block block. Seaborg's colleagues advised him not to publicize such radical suggestions as likely to undermine his career. Because Seaborg thinks he does not have a career to smear, he keeps publishing it. Suggestions Seaborg proved true and he later won the 1951 Nobel Prize in chemistry for his work in synthesizing actinide elements.
Although minute quantities of some transuranic elements occur naturally, they are all first found in the laboratory. Their production has significantly expanded the periodic table, the first being neptunium, synthesized in 1939. Because many transuranic elements are so unstable and rapidly decaying, they are challenging to detect and characterize when produced. There is controversy over the acceptance of competing claim claims for some elements, requiring an independent review to determine which party has priority, and therefore to name the right. In 2010, a joint Russian-US joint venture in Dubna, the Moscow Oblast, Russia, claimed to have synthesized six atoms of tennessine (element 117), making it the most recent invention claimed. This, together with the nihonium (element 113), moscovium (element 115), and oganesson (element 118), are the four most recently named elements, whose names all become official on November 28, 2016.
Different periodic table
Long or 32-column table
Modern periodic tables are sometimes extended to their long or 32-column form by returning the f-block elements of footnotes to their natural position between block-and-block. Unlike the 18-column form, this setting produces "no interruptions in the order of increasing atomic numbers". The relationship of the f-block to the other blocks of the periodic table also becomes easier to see. Jensen supports a 32-column table with the argument that lanthanide and actinide are otherwise degraded in the students' minds as boring, non-essential elements that can be quarantined and ignored. Although this advantage of 32 columns is generally avoided by editors because of the undue rectangular ratios (compared to the book page ratios), and the familiarity of chemists with modern forms (as introduced by Seaborg).
Table with different structure
Within 100 years after the appearance of Mendeleev's table in 1869, it was estimated that about 700 different versions of periodic tables were published. As well as various rectangular variations, other periodic table formats have been formed, for example, such as circles, cubes, cylinders, buildings, spirals, lemniscates, octagonal prisms, pyramids, balls, or triangles. Such alternatives are often developed to highlight or emphasize the chemical or physical properties of elements not seen in the traditional periodic table.
A popular alternative structure is Theodor Benfey (1960). The elements are arranged in a continuous spiral, with hydrogen at the center and transition metals, lanthanides, and actinides occupying the peninsula.
The most periodic table is two dimensions; three dimensional tables are known so far at least 1862 (Mendeleev's two-dimensional table in 1869). More recent examples include the Courtines' Periodic Classification (1925), Lamina Systems Wringley (1949), the Periodic helix GiguÃÆ'¨re (1965) and Dufour's Periodic Tree (1996). Furthermore, the Periodic Table of Physical Stowe (1989) has been described as four dimensions (having three spatial dimensions and one color dimension).
The various forms of the periodic table can be regarded as lying on the chemistry-physics continuum. Toward the chemical ends of the continuum can be found, for example, the irregular Inorganic Chemical Periodic Table (2002), which emphasizes trends and patterns, as well as unusual chemical relations and properties. Near the tip of continuum physics is the Periodic Table of the Left-Step Janet (1928). It has a structure that shows a closer relationship to the electron-shell filling sequence and, by association, quantum mechanics. A somewhat similar approach has been taken by Alper, although criticized by Eric Scerri for neglecting the need to display chemical and physical periodicity. Somewhere in the middle of the continuum is the general or standard periodic table that is everywhere. It is considered as better to express empirical trends in physical, electrical and thermal conductivity, and oxidation numbers, and other easily inferable properties from traditional techniques of chemical laboratories. Its popularity is regarded as a result of this layout having a good feature balance in terms of ease of construction and size, and atomic sequence depiction and periodic trends.
Open questions and controversy
Placement of hydrogen and helium
Simply follow the electron configuration, hydrogen (electronic configuration 1s 1 ) and helium (1s 2 ) should be placed in groups 1 and 2, above lithium (1s 2 2s 1 ) and beryllium (1s 2 2s 2 ). While such placement is common to hydrogen, it is rarely used for helium outside the electron configuration context: When noble gases (then called "inert gases") were first discovered around 1900, they are known as "group 0", reflecting no chemical reactivity of elements known at that point, and helium is placed at the top of the group, as it shares the extreme chemical extremism seen throughout the group. When the group changed its official number, many authors continued to assign helium directly over the neon, in group 18; one example of such placement is the current IUPAC table.
The chemistry of hydrogen is not too close to the alkali metal, which occupies group 1. On this basis it is sometimes placed elsewhere. A common alternative is at the top of a group of 17 hydrogen-given strictly univalent and largely nonmetallic chemistry, and highly univalent and non-metallic fluorine chemistry (the opposite element at the top of group 17). Sometimes, to show hydrogen has properties corresponding to that of alkali and halogen metals, it is shown at the top of two columns simultaneously. Another suggestion is that on carbon in group 14: placed in that way, it fits with the trend of increasing the potential value of ionization and the value of electron affinity, and not too far from the trend of electronegativity, although hydrogen can not exhibit more severe tetravalent characteristics. groups of 14 elements. Finally, hydrogen is sometimes placed separately from any group; this is based on how the general properties of hydrogen differ from each group. The other 1 period element, helium, is sometimes placed separately from any group. The property that distinguishes helium from the rest of the noble gases (although the extraordinary scarcity of helium is very close to neon and argon gases) is that in its closed electron shell, helium has only two electrons in the outer electron orbital, while the rest of the noble gas has eight.
Group 3 and its elements in periods 6 and 7
Although scandium and yttrium are always the first two elements in group 3, the identity of the next two elements is not completely resolved. They are generally lanthanum and aktinium, and more rarely luteium and lawrensium. Both variants derive from the historical difficulty in placing lanthanides in the periodic table, and arguments about where the block elements of f begin and end. It has been argued that such an argument is evidence that, "it is an error to break the [periodic] system into sharply constrained blocks". The third variant shows two positions under yttrium as occupied by lanthanides and actinides.
Chemical and physical arguments have been made to support luteium and lawrencium but most authors seem to be unsure. Most of the working chemists are unaware of any controversy. In December 2015, an IUPAC project was established to make recommendations on the issue.
Lanthanum and aktinium
Lanthanum and actinium are usually described as the 3 remaining group members. It has been argued that this layout dates from the 1940s, with the appearance of a periodic table that depends on the electron configuration of the distinguishing elements and ideas of electrons. Configuration of cesium, barium and lanthanum is [Xe] 6s 1 , [Xe] 6s 2 and [Xe] 5d 1 6s 2 . Lanthanum thus has a 5d distinguishing electron and this forms it "in group 3 as the first member of block-d for period 6". A consistent set of electron configurations is then seen in group 3: scandium [Ar] 3d 1 4s 2 , yttrium [Kr] 4d 1 5s < soup> 2 and lanthanum [Xe] 5d 1 6 s 2 . Still in period 6, ytterbium is given electron configuration [Xe] 4f 13 5d 1 6,0 2 and luteium [Xe] 4f 14 5d 1 6s 2 , "produces a 4f diffusing electron for luteium and strongly sets it as the last member of f-block for period 6". Then the work of spectroscopy found that the ytterbium electron configuration was actually [Xe] 4f 14 6s 2 . This means that ytterbium and lutetium - the latter with [Xe] 4f 14 5d 1 6.0 2 - both have 14 f-electrons, "produce d-not electron f-differentiating" for luteium and make it an "equally valid candidate" with lanthanum [Xe] 5d 1 6s 2 , for group 3 the position of the periodic table under yttrium. Lanthanum has the advantage of incumbency since 5d 1 electron appears for the first time in its structure while it appears for the third time in lutium, after also making a second brief appearance in gadolinium.
In terms of chemical behavior, and trends down group 3 for properties such as melting points, electronegativity and ionic radii, scandium, yttrium, lanthanum and aktinium are similar to those of their 1-2 counterparts. In this variant, the number of electrons f in the most common (trivalent) ions of the f-block element consistently corresponds to its position in the f-block. For example, the number of f-electrons for the trivalent ions of the first three f-block elements is Ce 1, Pr 2 and Nd 3.
Lutetium and lawrencium
In another table, lutesium and lawrensium are the remaining 3 members of the group. Initial techniques for separating chemically scandium, yttrium and lutetium depend on the fact that these elements occur together in the so-called "yttrium group" whereas La and Ac occur together in the "cerium group". Thus, luteium instead of lanthanum was assigned to group 3 by several chemists in the 1920s and 30s. Some physicists in the 1950s and 1960s favored luteium, given the comparison of some of its physical properties with lanthanum. This arrangement, in which the lanthanum is the first member of the f-block, is debated by some authors because lanthanum has no f-electrons. It has been argued that this is not a valid concern given the other periodic table anomalies - thorium, for example, has no f-electrons but is part of the f-block. For lawrensium, the gas phase phase electron configuration is confirmed in 2015 as [Rn] 5f 14 7s 2 7p 1 . Such configuration is another periodic table anomaly, regardless of whether lawrencium is located in block-f or d-block, since the only potentially p-block position has been provided for nihonium with a prediction configuration [Rn] 5f 14 6d 10 7s 2 7p 1 .
Chemically, scandium, yttrium and lutium (and possibly lawrensium) behave like a trivalent version of a group of 1-2 metals. On the other hand, the downtrend of groups for properties such as melting point, electronegativity and ionic radius, is similar to that found among groups of 4-8 â € <â € Lanthanides and actinides
Some authors position thirty lanthanides and actinides in two positions under yttrium (usually through footnote markers). This variant emphasizes the similarity in chemistry of 15 lanthanide elements (La-Lu), perhaps at the expense of ambiguity as which element occupies two groups of 3 positions below yttrium, and a width of 15 columns f blocks (there are only 14 elements in each block line f ).
Group included in transition metal
Transition metal definitions, such as those supplied by IUPAC, are elements whose atoms have an incomplete sub-shell, or which may cause cations with an incomplete sub-shell. By this definition all elements in groups 3-11 are transition metals. Therefore, the IUPAC definition does not include group 12, consisting of zinc, cadmium and mercury, from the transition metal category.
Some chemists treat the categories of "d-block elements" and "transition metals" interchangeably, thus including groups 3-12 between transition metals. In this case the group 12 elements are treated as special cases of the transition metal in which d electrons are not normally involved in chemical bonds. The 2007 report on mercury (IV) fluoride (HgF 4 ), a compound in which mercury will use electrons to bind, has prompted some commentators to suggest that mercury can be considered a transition metal. Other commentators, such as Jensen, argue that the formation of compounds such as HgF 4 can occur only under very abnormal conditions; indeed, its existence is currently disputed. Thus, mercury can not be considered a transition metal by a reasonable interpretation of the usual meaning of the term.
Still other chemists further exclude group 3 elements from the definition of transition metal. They do so on the basis that the group 3 elements do not form ions which have a half-filled skin and therefore do not exhibit the chemical characteristics of the transition metal. In this case, only groups 4-11 are considered transition metals. Although group 3 elements show some of the characteristic chemical properties of transition metals, they show some characteristics of their physical properties (due to their presence in every single electron atom). Element
with unknown chemical property
Although all elements up to oganesson have been found, from the elements above the hassium (element 108), only copernicium (element 112), nihonium (element 113), and flerovium (element 114) have known chemical properties, and only for copernicium there is. sufficient evidence for the current conclusive categorization. Other elements may behave differently than what would be predicted by extrapolation, because of the relativistic effect; for example, flerovium has been predicted to exhibit properties such as noble gases, although it is currently placed in a carbon group. The current experimental evidence still leaves the question of whether flerovium behaves more like a metal or a noble gas.
More table extensions
It is unclear whether new elements will continue the current periodic table pattern as period 8, or require further adjustment or adjustment. Seaborg expected the eighth period to follow a predetermined pattern, so it would include a two-element block for elements 119 and 120, a new g-block for the next 18 elements, and 30 additional elements continuing f -, d-, and p- blocks, culminating in element 168, the next noble gas. Recently, physicists such as Pekka PyykkÃÆ'¶ have theorized that these additional elements do not follow Madelung's rules, which predict how electron shells are filled and thus affect the appearance of the current periodic table. Currently there are several competing theoretical models for the placement of atomic number elements less than or equal to 172. In all this is element 172, not element 168, which emerges as the next noble gas after oganesson, although this must be regarded as speculative because it is not there is a complete calculation done outside element 122.
element with highest atomic number
The number of possible elements is unknown. The very first suggestion made by Elliot Adams in 1911, and based on the arrangement of elements in each row of the horizontal periodic table, is that the elements of atomic weight are greater than about 256 (which would equalize between elements 99 and 100 in modern terms)) There is no. The higher - more recent - estimate is that the periodic table can end soon after the island's stability, which is predicted to center around the element 126, as the extension of the periodic table and the nuclide is limited by the proton and neutron drip lines. Other predictions that end the periodic table include the element 128 by John Emsley, on element 137 by Richard Feynman, and on element 155 by Albert Khazan.
- Bohr Model
The Bohr model shows difficulties for atoms with atomic numbers greater than 137, since any element with an atomic number greater than 137 would require 1s electrons to travel faster than c , the speed of light. Therefore the non-relativistic Bohr model is inaccurate when applied to such elements.
- Dirac Relativistic Equation
The relativistic Dirac equation has problems for elements with more than 137 protons. For such elements, the wave function of Dirac's ground state is more oscillatory than bound, and there is no gap between the positive and negative energy spectrum, as in the Klein paradox. A more accurate calculation considering the limited effect of nucleus size suggests that the first binding energy exceeds the limit for elements with more than 173 protons. For heavier elements, if the deepest orbital (1s) is not filled, the core electric field will pull the electrons out of a vacuum, resulting in spontaneous emission from the positron. This does not happen if the deepest orbital is filled, so element 173 is not necessarily the end of the periodic table.
Optimal shape
The various forms of the periodic table have prompted the question of whether there is an optimal or definitive periodic table form. The answer to this question is deemed to depend on whether the apparent chemical periodicity between elements has an underlying truth, which is effectively programmed into the universe, or if such periodicity is not a product of subjective human interpretation, depending on the circumstances, beliefs and predilections of observers human. The objective basis for chemical periodicity will solve the question of the location of hydrogen and helium, and the composition of the group 3. The underlying truth, if any, is thought to be undiscovered. In its absence, the various forms of the periodic table can be regarded as variations on the theme of chemical periodicity, each of which explores and emphasizes different aspects, properties, perspectives, and relationships of and between elements.
See also
Note
References
Bibliography
External links
- Periodic table on the IUPAC website, notes no official IUPAC table
- Dynamic periodic table, with interactive layout
- Eric Scerri, a leading philosopher of science specializing in the history and philosophy of the periodic table
- The INTERNET Database of the Periodic Table
- Periodic table of endangered elements
- Periodic sample table
- Periodic video table
- WebElements
Source of the article : Wikipedia
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