Thursday 17/11/2011 The Periodic Table The first periodic table was devised by Dmitri Mendeleev and published in 1869. Mendeleev found he could arrange the 65 elements that were then known in a grid or table so that each element had: 1. A higher atomic weight than the one on its left. 2. Similar chemical properties to other elements in the same column. He realized that the table in front of him lay at the very heart of chemistry. In his table he noted gaps - spaces where elements should be but none had yet been discovered. In 1913, Henry Moseley, who worked with Rutherford, showed that it is atomic number (electric charge) which is most fundamental to the chemical properties of any element. Mendeleev had believed chemical properties were determined by atomic weight. Moseley correctly predicted the existence of new elements based on atomic numbers. Today the chemical elements are still arranged in order of increasing atomic number (Z) as you go from left to right across the table. We call the horizontal rows periods and the vertical rows groups. We also know now that an element's chemistry is determined by the way its electrons are arranged - its electron configuration. The noble gases are found in group 18, on the far right of each period. The reluctance of the noble gases to undergo chemical reactions indicates that the atoms of these gases strongly prefer their own electron configurations - featuring a full outer shell of electrons - to any other. In contrast to the noble gases, the elements with the highest reactivity are those with the greatest need to gain or lose electrons in order to achieve a full outer shell of electrons. Elements that sit in the same group (e.g. the alkali metals in Group 1) all have the same number of outer electrons, leading to similar chemical properties. Likewise the halogens in Group 17 also have similar properties to one another. When halogens react, they gain an electron to form negatively charged ions. Each ion has the same electron configuration as the noble gas in the same period. The ions are therefore more chemically stable than the elements from which they formed. There is a progression from metals to non-metals across each period. The block of elements in groups 3 - 12 contains the transition metals. These are similar to one another in many ways: they produce colored compounds, have variable valency and are often used as catalysts. Then we come to the lanthanides (elements 58 - 71) and actinides (elements 90 - 103). The lanthanides are often called the rare earth elements, although in fact these elements are not rare. The actinides include most of the well-known elements that take part in or are produced by nuclear reactions. No element with atomic number higher than 92 occurs naturally in large quantities. Tiny amounts of plutonium and neptunium exist in nature as decay products of uranium. These elements, and higher elements, are also produced artificially in nuclear reactors or particle accelerators The periodic table of the chemical elements (also known as the periodic table or periodic table of the elements) is a tabular display of the 118 known chemical elements organized by selected properties of their atomic structures. Elements are presented by increasing atomic number, the number of protons in an atom's atomic nucleus. While rectangular in general outline, gaps are included in the horizontal rows (known as periods) as needed to keep elements with similar properties together in vertical columns (known as groups), e.g. alkali metals, alkali earths, halogens, noble gases. The following is the periodic table as defined by the International Union of Pure and Applied Chemistry (IUPAC): Group # 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 Period 1 2 1 H He 3 4 5 6 7 8 9 10 2 Li Be B C N O F Ne 11 12 13 14 15 16 17 18 3 Na Mg Al Si P S Cl Ar 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 4 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 5 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe 55 56 * 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 6 Cs Ba Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn 87 88 ** 104 105 106 107 108 109 110 111 112 113 114 115 116 117 118 7 Fr Ra Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo * Lanthanides 57 58 59 60 61 62 63 64 65 66 67 68 69 70 71 La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu (Lanthanoids) ** Actinides 89 90 91 92 93 94 95 96 97 98 99 100 101 102 103 Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr (Actinoids) Organizing principles 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. 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 Atomic number By definition, each chemical element has a unique atomic number, the number of protons in its nucleus. Different atoms of many elements have different numbers of neutrons, which differentiates between isotopes of an element. For example, all atoms of hydrogen have one proton, and no atoms of any other element have exactly one proton. On the other hand, a hydrogen atom can have one or two neutrons in its nucleus, or none at all, yet all of these cases are isotopes of hydrogen, not instances of some other element. (A hydrogen atom with no neutrons in addition to its sole proton is called protium, one with one neutron in addition to its proton is called deuterium, and one with two additional neutrons, tritium.) In the modern periodic table, the elements are placed progressively in each row (period) from left to right in the sequence of their atomic numbers, with each new row starting with the next atomic number following the last number in the previous row. No gaps or duplications exist. Since the elements can be uniquely sequenced by atomic number, conventionally from lowest to highest, sets of elements are sometimes specified by such notation as "through", "beyond", or "from ... through", as in "through iron", "beyond uranium", or "from lanthanum through lutetium". The terms "light" and "heavy" are sometimes also used informally to indicate relative atomic numbers (not densities), as in "lighter than carbon" or "heavier than lead", although technically the weight or mass of atoms of an element (their atomic weights or atomic masses) do not always increase monotonically with their atomic numbers. The significance of atomic numbers to the organization of the periodic table was not appreciated until the existence and properties of protons and neutrons became understood. Mendeleev's periodic tables instead used atomic weights, information determinable to fair precision in his time, which worked well enough in most cases to give a powerfully predictive presentation far better than any other comprehensive portrayal of the chemical elements' properties then possible. Substitution of atomic numbers, once understood, gave a definitive, integer-based sequence for the elements, still used today even as new synthetic elements are being produced and studied. Periodicity of chemical properties 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 orbitals 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. Subshell S G F D P Period 1 1s 2 2s 2p 3 3s 3p 4 4s 3d 4p 5 5s 4d 5p 6 6s 4f 5d 6p 7 7s 5f 6d 7p 8 8s 5g 6f 7d 8p The total number of electron shells an atom has determines the period to which it belongs. Each shell is divided into different subshells, which as atomic number increases are filled in roughly this order (the Aufbau principle) (see table).[5] Hence the structure of the periodic table. Since the outermost electrons determine chemical properties, those with the same number of valence electrons are generally 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.