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. 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.[citation needed] 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. Under the international naming system, the groups are numbered numerically 1 through 18 from the left most column (the alkali metals) to the right most column (the noble gases).[7] The older naming systems differed slightly between Europe and America (the table shown in this section shows the old American Naming System).[8] Some of these groups have been given trivial (unsystematic) names, such as the alkali metals, alkaline earth metals, halogens, pnictogens, chalcogens, and noble gases. However, some other groups, such as group 7, have no trivial names and are referred to simply by their group numbers, since they display fewer similarities and/or vertical trends.[7] Modern quantum mechanical theories of atomic structure explain group trends by proposing that elements within the same group generally have the same electron configurations in their valence shell, which is the most important factor in accounting for their similar properties.[1] Elements in the same group show patterns in 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 has a top to bottom decrease in electronegativity due to an increasing distance between valence electrons and the nucleus.[9] 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 lanthanides and actinides form two substantial horizontal series of elements.
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.[10] 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. Electronegativity increases in the same manner as ionization energy because of the pull exerted on the electrons by the nucleus.[9] 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.[11] Blocks
Because of the importance of the outermost electron shell, the different regions of the periodic table are sometimes referred to asperiodic 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 which are groups 13 through 18 in IUPAC (3A through 8A in American) and contains, among others, all of the semimetals. The d-block comprises groups 3 through 12 in IUPAC (or 3B through 8B in American group numbering) and contains all of the transition metals. The f-block, usually offset below the rest of the periodic table, comprises the lanthanides and actinides.[12] Effective nuclear charge
Effective Nuclear Charge Diagram
The effective nuclear charge is the net positive charge experienced by an electron in a multi-electron atom. The term "effective" is used because the shielding effect of negatively charged electrons prevents higher orbital electrons from experiencing the fullnuclear charge by the repelling effect of inner-layer electrons. The effective nuclear charge experienced by the outer shell electron is also called the core charge. It is possible to determine the strength of the nuclear charge by looking at the oxidation number of the atom.
Calculating the effective nuclear charge In an atom with one electron, that electron experiences the full charge of the positive nucleus. In this case, the effective nuclear charge can be calculated from Coulomb's law. However, in an atom with many electrons the outer electrons are simultaneously attracted to the positive nucleus and repelled by the negatively charged electrons. The effective nuclear charge on such an electron is given by the following equation:
Zeff = Z − S where Z is the number of protons in the nucleus (atomic number), and S is the average number of electrons between the nucleus and the electron in question (the number of nonvalence electrons). S can be found by the systematic application of various rule sets, the simplest of which is known as "Slater's rules" (named after John C. Slater). Douglas Hartree defined the effective Z of a Hartree-Fock orbital to be:
where
Note: Zeff is also often written Z*. Example Consider a sodium cation, a fluorine anion, and a neutral neon atom. Each has 10 electrons, and the number of nonvalence electrons is 2 (10 total electrons - 8 valence) but the effective nuclear charge varies because each has a different atomic number: - Zeff(F ) = 9 − 2 = 7 + Zeff(Ne) = 10 − 2 = 8 + + Zeff(Na ) = 11 − 2 = 9 + So, the sodium cation has the largest effective nuclear charge, and thus the smallest atomic radius.
Values Shielding effect The shielding effect describes the decrease in attraction between an electron and the nucleus in any atom with more than one electron shell. It is also referred to as the screening effect or atomic shielding. Slater's rules In quantum chemistry, Slater's rules provide numerical values for the effective nuclear charge concept. In a many-electron atom, each electron is said to experience less than the actual nuclear chargeowning to shielding or screening by the other electrons. For each electron in an atom, Slater's rules provide a value for the screening constant, denoted by s, S, or σ, which relates the effective and actual nuclear charges as
The rules were devised semi-empirically by John C. Slater and published in 1930.[1] Rules Firstly,[1][4] the electrons are arranged in to a sequence of groups in order of increasing principal quantum number n, and for equal n in order of increasing azimuthal quantum number l, except that s- and p- orbitals are kept together. [1s] [2s,2p] [3s,3p] [3d] [4s,4p] [4d] [4f] [5s, 5p] [5d] etc. Each group is given a different shielding constant which depends upon the number and types of electrons in those groups preceding it. The shielding constant for each group is formed as the sum of the following contributions:
1. An amount of 0.35 from each other electron within the same group except for the [1s] group, where the other electron contributes only 0.30. 2. If the group is of the [s p] type, an amount of 0.85 from each electron with principal quantum number (n) one less and an amount of 1.00 for each electron with an even smaller principal quantum number 3. If the group is of the [d] or [f], type, an amount of 1.00 for each electron inside it. This includes i) electrons with a smaller principal quantum number and ii) electrons with an equal principal quantum number and a smaller azimuthal quantum number (l) In tabular form, the rules are summarized as:
Electrons in group(s) Electrons in Electrons in all Other with principal quantum group(s) group(s) electrons in Group number n with principal with principal the same andazimuthal quantum quantum quantum number < group number < l number n-1 n-1
[1s] 0.3 N/A N/A N/A
[ns,np] 0.35 N/A 0.85 1
[nd] or [nf] 0.35 1 1 1
Atomic Radius
The atomic radius of an element is half of the distance between the centers of two atoms of that element that are just touching each other. Generally, the atomic radius decreases across a period from left to right and increases down a given group. The atoms with the largest atomic radii are located in Group I and at the bottom of groups.
Moving from left to right across a period, electrons are added one at a time to the outer energy shell. Electrons within a shell cannot shield each other from the attraction to protons. Since the number of protons is also increasing, the effective nuclear charge increases across a period. This causes the atomic radius to decrease.
Moving down a group in the periodic table, the number of electrons and filled electron shells increases, but the number of valence electrons remains the same. The outermost electrons in a group are exposed to the same effective nuclear charge, but electrons are found farther from the nucleus as the number of filled energy shells increases. Therefore, the atomic radii increase.
Ionization Energy
The ionization energy, or ionization potential, is the energy required to completely remove an electron from a gaseous atom or ion. The closer and more tightly bound an electron is to the nucleus, the more difficult it will be to remove, and the higher its ionization energy will be. The first ionization energy is the energy required to remove one electron from the parent atom. The second ionization energy is the energy required to remove a second valence electron from the univalent ion to form the divalent ion, and so on. Successive ionization energies increase. The second ionization energy is always greater than the first ionization energy. Ionization energies increase moving from left to right across a period (decreasing atomic radius). Ionization energy decreases moving down a group (increasing atomic radius). Group I elements have low ionization energies because the loss of an electron forms a stable octet.
The property is alternately still often called the ionization potential, measured in volts. In chemistry it is often referred to one mole of substance (molar ionization energy or enthalpy) and reported in kJ/mol. In atomic physics the ionization energy is typically measured in the unit electron volt (eV). The ionization energy is different for electrons of different atomic or molecular orbitals. More generally, the nth ionization energy is the energy required to strip off the nth electron after the first n − 1 electrons have been removed. It is considered a measure of the tendency of an atom or ion to surrender an electron, or the strength of the electron binding; the greater the ionization energy, the more difficult it is to remove an electron. The ionization energy may be an indicator of the reactivity of an element. Elements with a low ionization energy tend to be reducing agents and form cations, which in turn combine with anions to form salts.
Electron binding energy (BE), more accurately, is the energy required to release an electron from its atomic or molecular orbital when adsorbed to a surface rather than a free atom. Binding energy values are normally reported as positive values with units of eV. The binding energies of 1s electrons are roughly proportional to (Z-1)² (Moseley's law). Values and trends Main article: Molar ionization energies of the elements Generally the (n+1)th ionization energy is larger than the nth ionization energy. Always, the next ionization energy involves removing an electron from an orbital closer to the nucleus. Electrons in the closer orbitals experience greater forces of electrostatic attraction; thus, their removal requires increasingly more energy. Some values for elements of the third period are given in the following table: Successive molar ionization energies in kJ/mol (96.485 kJ/mol = 1 eV/particle)
Element First Second Third Fourth Fifth Sixth Seventh
Na 496 4,560
Mg 738 1,450 7,730
Al 577 1,816 2,881 11,600
Si 786 1,577 3,228 4,354 16,100
P 1,060 1,890 2,905 4,950 6,270 21,200
S 999.6 2,260 3,375 4,565 6,950 8,490 27,107
Cl 1,256 2,295 3,850 5,160 6,560 9,360 11,000
Ar 1,520 2,665 3,945 5,770 7,230 8,780 12,000
Large jumps in the successive molar ionization energies occur when passing noble gas configurations. For example, as can be seen in the table above, the first two molar ionization energies of magnesium (stripping the two 3s electrons from a magnesium atom) are much smaller than the third, which requires stripping off a 2p electron from the very stable neon configuration of Mg2+.
Periodic trend for ionization energy. Each period begins at a minimum for the alkali metals, and ends at a maximum for the noble gases. Ionization energy is also a periodic trend within the periodic table organization. Moving left to right within aperiod or upward within a group, the first ionization energy generally increases. As the atomic radiusdecreases, it becomes harder to remove an electron that is closer to a more positively charged nucleus. ionization enthalpy increases from left to right in a period and decreases from top to bottom in a group.
Electron Affinity
Electron affinity reflects the ability of an atom to accept an electron. It is the energy change that occurs when an electron is added to a gaseous atom. Atoms with stronger effective nuclear charge have greater electron affinity. Some generalizations can be made about the electron affinities of certain groups in the periodic table. The Group IIA elements, the alkaline earths, have low electron affinity values. These elements are relatively stable because they have filled s subshells. Group VIIA elements, the halogens, have high electron affinities because the addition of an electron to an atom results in a completely filled shell. Group VIII elements, noble gases, have electron affinities near zero, since each atom possesses a stable octet and will not accept an electron readily. Elements of other groups have low electron affinities.
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