PERIODIC TABLE The periodic table of the chemical elements (also Mendeleev's table, periodic table of the elements or just periodic table) is a tabular display of the chemical elements . Although precursors to this table exist, its invention is generally credited to Russian chemist Dmitri Mendeleev in 1869, who intended the table to illustrate recurring ("periodic") trends in the properties of the elements. The layout of the table has been refined and extended over time, as new elements have been discovered, and new theoretical models have been developed to explain chemical behavior. The periodic table is now ubiquitous within the academic discipline of chemistry , providing an extremely useful framework to classify, systematize, and compare all of the many different forms of chemical behavior. The table has found wide application in Dmitri Mendeleev chemical elements . tabular chemistr y chemical chemical engineering hemistry , physics , biology , and engineering , especially chemical engineering . The current standard table contains 117 elements as of July 2009 (elements 1 – 116 and element 118 ).
Structure of the periodic table Some versions of the table show a dark stair-step line along the metalloids. Metals are to the left of the line and non-metals to the right. The layout of the periodic table demonstrates recurring ("periodic") chemical properties. Elements are listed in order of increasing atomic number (i.e., the number of protons in the atomic nucleus ) . Rows are arranged so that elements with similar properties fall into the same columns ( groups or families ). According to quantum mechanical theories of electron configuration within atoms, each row ( period ) in the table corresponded to the filling of a quantum shell of electrons. There are progressively longer periods further down the table, grouping the elements into s- , p- , d- and f-blocks to reflect their electron configuration . In printed tables, each element is usually listed with its element symbol and atomic number ; many versions of the table also list the element's atomic mass and other information, such as its abbreviated electron configuration , electronegativity and most common valence numbers . As of 2006, the table contains 117 chemical elements whose discoveries have been confirmed. Ninety-four are found naturally on Earth, and the rest are synthetic elements that have been produced artificially in particle accelerators . Elements 43 (technetium), 61 (promethium) and all elements greater than 83 (bismuth), beginning with 84 (polonium) have no stable isotopes. The atomic mass of each of these element's isotope having the longest half-life is typically reported on periodic tables with parentheses . Isotopes of elements 43, 61, 93 (neptunium) and 94 (plutonium), first discovered synthetically, have since been discovered in trace amounts on Earth as products of natural radioactive decay
processes. The primary determinant of an element's chemical properties is its electron configuration , particularly the valence shell electrons. For instance, any atoms with four valence electrons occupying p orbital will exhibit some similarity. The type of orbital in which the atom's outermost electrons reside determines the "block" to which it belongs. The number of valence shell electrons determines the family, or group, to which the element belongs. The total number of electron shells an atom has determines the period to which it belongs. Each shell is divided into different sub-shells, which as atomic number increases are filled in roughly this order (the Aufbau principle ): Hence the structure of the table. Since the outermost electrons determine chemical properties, those with the same number of valence electrons are grouped together. Progressing through a group from lightest element to heaviest element, the outer-shell electrons (those most readily accessible for participation in chemical reactions) are all in the same type of orbital, with a similar shape, but with increasingly higher energy and average distance from the nucleus. For instance, the outer-shell (or "valence") electrons of the first group, headed by hydrogen , all have one electron in an s orbital. In hydrogen, that s orbital is in the lowest possible energy state of any atom, the first-shell orbital (and represented by hydrogen's position in the first period of the table). In francium , the heaviest element of the group, the outer-shell electron is in the seventh-shell orbital, significantly further out on average from the nucleus than those electrons filling all the shells below it in energy. As another example, both carbon and lead have four electrons in their outer shell orbitals.
Note that as atomic number (i.e., charge on the atomic nucleus ) increases, this leads to greater spin-orbit coupling between the nucleus and the electrons, reducing the validity of the quantum mechanical orbital approximation model, which considers each atomic orbital as a separate entity. The elements ununbium, ununtrium, ununquadium, etc. are elements that have been discovered, but so far have not received a trivial name yet. There is a system for naming them temporarily. Classification Groups A group or family is a vertical column in the periodic table. Groups are considered the most important method of classifying the elements. In some groups, the elements have very similar properties and exhibit a clear trend in properties down the group. These groups tend to be given trivial (unsystematic) names, e.g., the alkali metals , alkaline earth metals , halogens , pnictogens , chalcogens , and noble gases . Some other groups in the periodic table display fewer similarities and/or vertical trends (for example Group 14), and these have no trivial names and are referred to simply by their group numbers Periods A period is a horizontal row in the periodic table. Although groups are the most common way of classifying elements
there are some regions of the periodic table where the horizontal trends and similarities in properties are more significant than vertical group trends. This can be true in the d-block (or " transition metals " ), and especially for the f-block , where the lanthanoids and actinoids form two substantial horizontal series of elements. Blocks Because of the importance of the outermost shell, the different regions of the periodic table are sometimes referred to as periodic table blocks , named according to the subshell in which the "last" electron resides. The s-block comprises the first two groups ( alkali metals and alkaline earth metals ) as well as hydrogen and helium . The p-block comprises the last six groups (groups 13 through 18) and contains, among others, all of the semimetals . The d-block comprises groups 3 through 12 and contains all of the transition metals . The f-block , usually offset below the rest of the periodic table, comprises the rare earth metals . Other The chemical elements are also grouped together in other ways. Some of these groupings are often illustrated on the periodic table, such as transition metals , poor metals , and metalloids . Other informal groupings exist, such as the platinum group and the noble metals .
Periodicity of chemical properties The main value of the periodic table is the ability to predict the chemical properties of an element based on its location on the table. It should be noted that the properties vary differently when moving vertically along the columns of the table than when moving horizontally along the rows Periodic trends of groups Modern quantum mechanical theories of atomic structure explain group trends by proposing that elements within the same group have the same electron configurations in their valence shell , which is the most important factor in accounting for their similar properties. Elements in the same group also show patterns in their atomic radius , ionization energy , and electronegativity . From top to bottom in a group, the atomic radii of the elements increase. Since there are more filled energy levels, valence electrons are found farther from the nucleus. From the top, each successive element has a lower ionization energy because it is easier to remove an electron since the atoms are less tightly bound. Similarly, a group will also see a top to bottom decrease in electronegativity due to an increasing distance between valence electrons and the nucleus
Periodic trends of periods Elements in the same period show trends in atomic radius , ionization energy , electron affinity , and electronegativity . Moving left to right across a period, atomic radius usually decreases. This occurs because each successive element has an added proton and electron which causes the electron to be drawn closer to the nucleus. This decrease in atomic radius also causes the ionization energy to increase when moving from left to right across a period. The more tightly bound an element is, the more energy is required to remove an electron. Similarly, electronegativity will increase in the same manner as ionization energy because of the amount of pull that is exerted on the electrons by the nucleus. Electron affinity also shows a slight trend across a period. Metals (left side of a period) generally have a lower electron affinity than nonmetals (right side of a period) with the exception of the noble gases.
I n 1817, Johann Wolfgang Döbereiner began to formulate one of the earliest attempts to classify the elements. He found that some elements formed groups of three with related properties. He termed these groups "triads". Some triads classified by Döbereiner:
chlorine , bromine , and iodine
calcium , strontium , and barium
sulfur , selenium , and tellurium
lithium , sodium , and potassium
In all of the triads, the atomic weight of the second element was almost exactly the average of the atomic weights of the first and third element.
John Newlands John Newlands was an English chemist who in 1865 classified the 56 elements that had been discovered at the time into 11 groups which were based on similar physical properties. Newlands noted that many pairs of similar elements existed which differed by some multiple of eight in atomic weight. However, his law of octaves , likening this periodicity of eights to the musical scale, was ridiculed by his contemporaries. It was not until the following century, with Gilbert N. Lewis ' valence bond theory (1916) and Irving Langmuir 's octet theory of chemical bonding (1919) that the importance of the periodicity of eight would be accepted.
Mendeleev Dmitri Mendeleev , a Siberian -born Russian chemist, was the first scientist to make a periodic table much like the one we use today. Mendeleev arranged the elements in a table ordered by atomic weight corresponding to relative molar mass as defined today. It is sometimes said that he played "chemical solitaire" on long train rides using cards with various facts of known elements. On March 6, 1869, a formal presentation was made to the Russian Chemical Society, entitled The Dependence Between the Properties of the Atomic Weights of the Elements . His table was published in an obscure Russian journal but quickly republished in a German journal, Zeitschrift für Chemie (Eng., "Chemistry Magazine"), in 1869. DMITIRI MENDEELEV
Mendeleev predicted the discovery of other elements and left space for these new elements, namely eka-silicon ( germanium ), eka- aluminium ( gallium ), and eka-boron ( scandium ) . Thus, there was no disturbance in the periodic table.
He pointed out that some of the then current atomic weights were incorrect.
He provided for variance from atomic weight order.
Shortcomings of Mendeleev's table
His table did not include any of the noble gases , which were discovered later. These were added by Sir William Ramsay as Group 0, without any disturbance to the basic concept of the periodic table.
There was no place for the isotopes of the various elements, which were discovered later.
Refinements to the periodic table Henry Moseley I n 1914 Henry Moseley found a relationship between an element's X-ray wavelength and its atomic number (Z), and therefore re-sequenced the table by nuclear charge rather than atomic weight. Before this discovery, atomic numbers were just sequential numbers based on an element's atomic weight. Moseley's discovery showed that atomic numbers had an experimentally measurable basis. Thus Moseley placed argon (Z=18) before potassium (Z=19) based on their X-ray wavelengths, despite the fact that argon has a greater atomic weight (39.9) than potassium (39.1). The new order agrees with the chemical properties of these elements, since argon is a noble gas and potassium an alkali metal . Similarly, Moseley placed cobalt before nickel , and was able to explain that tellurium occurs before iodine without revising the experimental atomic weight of tellurium (127.6) as proposed by Mendeleev. Moseley's research also showed that there were gaps in his table at atomic numbers 43 and 61 which are now known to be Technetium and Promethium , respectively, both radioactive and not naturally occurring. Following in the footsteps of Dmitri Mendeleev, Henry Moseley also predicted new elements.