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There is also a big increase in ionization energy for removal of an electron from an inner shell (lower n value). As the effective charge increases, or as the distance of the electron from the nucleus decreases, the greater the attraction between the nucleus and the electron. To understand how the periodic table is organized, imagine that we write down a long horizontal list of the elements in order of their increasing atomic number. Now if we look at the various physical and chemical properties of these elements, we would find that their values tend to increase or decrease with Z in a manner that reveals a repeating pattern— that is, a periodicity. In the past, two different systems of Roman numerals and letters were used to denote the various groups. The properties of an atom depend ultimately on the number of electrons in the various orbitals, and on the nuclear charge which determines the compactness of the orbitals. The electrons (denoted by the red dots) in the outer-most shell of an atom are the ones that interact most readily with other atoms, and thus play a major role in governing the chemistry of an element.
The above diagram shows the first three rows of what are known as the representative elements— that is, the s- and p-block elements only. Those electrons in the outmost or valence shell are especially important because they are the ones that can engage in the sharing and exchange that is responsible for chemical reactions; how tightly they are bound to the atom determines much of the chemistry of the element. We can simplify the shell model even further by imagining that the valence shell electrons are the only electrons in the atom, and that the nuclear charge has whatever value would be required to bind these electrons as tightly as is observed experimentally. Part of the difference between Z and Zeffective is due to other electrons in the valence shell, but this is usually only a minor contributor because these electrons tend to act as if they are spread out in a diffuse spherical shell of larger radius. The formula for calculating effective nuclear charge is not very complicated, but we will skip a discussion of it here.
When an atom is combined with other atoms in a solid element or compound, an effective radius can be determined by observing the distances between adjacent rows of atoms in these solids. A rough idea of the size of a metallic atom can be obtained simply by measuring the density of a sample of the metal.
Although the radius of an atom or ion cannot be measured directly, in most cases it can be inferred from measurements of the distance between adjacent nuclei in a crystalline solid. Covalent radius is half the distance between like atoms that are bonded together in a molecule. The lithium ion is sufficiently small that in LI, the iodide ions are in contact, so I-I distances are twice the ionic radius of I–. Many atoms have several different radii; for example, sodium forms a metallic solid and thus has a metallic radius, it forms a gaseous molecule Na2 in the vapor phase (covalent radius), and of course it forms ionic solids as mentioned above. The figure shows a periodic table in which the sizes of the atoms are represented graphically.
A positive ion is always smaller than the neutral atom, owing to the diminished electron-electron repulsion. Negative ions are always larger than the parent ion; the addition of one or more electrons to an existing shell increases electron-electron repulsion which results in a general expansion of the atom. An isoelectronic series is a sequence of species all having the same number of electrons (and thus the same amount of electron-electron repulsion) but differing in nuclear charge. Chemical reactions are based largely on the interactions between the most loosely bound electrons in atoms, so it is not surprising that the tendency of an atom to gain, lose or share electrons is one of its fundamental chemical properties. Note the very large jumps in the energies required to remove electrons from the 1s orbitals of atoms of the second-row elements Li-Ne. Ionization energies increase with the nuclear charge Z as we move across the periodic table.
This more detailed plot of the ionization energies of the atoms of the first ten elements reveals some interesting irregularities that can be related to the slightly lower energies (greater stabilities) of electrons in half-filled (spin-unpaired) relative to completely-filled subshells. Finally, a more comprehensive survey of the ionization energies of the main group elements is shown below. IE's (as well as many other properties) tend not to vary greatly amongst the d-block elements. Formation of a negative ion occurs when an electron from some external source enters the atom and become incorporated into the lowest energy orbital that possesses a vacancy. In general, electron affinities tend to be much smaller than ionization energies, suggesting that they are controlled by opposing factors having similar magnitudes. Because of these opposing effects, the periodic trends in electron affinities are not as clear as are those of ionization energies. In general, we can say that electron affinities become more exothermic as we move from left to right across a period (owing to increased nuclear charge and smaller atom size). In the Group 2 elements, the filled 2s orbital apparently shields the nucleus so effectively that the electron affinities are slightly endothermic.
The Group 15 elements have rather low values, due possibly to the need to place the added electron in a half-filled p orbital; why the electron affinity of nitrogen should be endothermic is not clear.
When two elements are joined in a chemical bond, the element that attracts the shared electrons more strongly is more electronegative.
It is important to understand that electronegativities are properties of atoms that are chemically bound to each other; there is no way of measuring the electronegativity of an isolated atom. Because there is no single definition of electronegativity, any numerical scale for measuring it must of necessity be somewhat arbitrary.


The location of hydrogen on this scale reflects some of the significant chemical properties of this element. Make sure you thoroughly understand the following essential ideas which have been presented above.
You should be able to sketch out the general form of the periodic table and identify the various blocks and identify the groups corresponding to the alkali metals, the transition elements, the halogens, and the noble gases.
For the first eighteen elements, you should be able to predict the formulas of typical binary compounds they can be expected to form with hydrogen and with oxygen. WebElements (Mark Winters, Sheffield U., UK) The elements in this online periodic table are linked to an extensive variety of chemical and physical data as well as background, crystallographic, nuclear, electronic, biological and geological information. ChemiCool Periodic Table (MIT) is a nice-looking table in which you can click on individual elements to bring up more information. It's Elemental - this is not so much a periodic table as a set of links to a set of excellent articles that focus on the history and uses of the different elements, each written by someone having a special knowlege or interest in a particular element.
ChemistryCoach's Periodic Table links - a huge but well-organized list of every possible kind of specialized periodic table you can think of, as well as games, software, etc.
Main Group Elements - a somewhat more advanced comparison of these especially important elements. Mark Leach has assembled a very well-done and interesting page of historic and alternative periodic tables.
Chinese periodic table in simplified and traditional versions; see also this one which contains additional element information. Periodic tables to wear - Chemistry-related apparel including periodic-table T-shirts, neckties, etc. Its valence electrons have a higher principle quantum number, and are therefore further away from the nucleus, than C or B. For n=1 there is no p block, and the s block is split so that helium is placed in the same group as the other inert gases, which it resembles chemically. North Americans added the letter B to denote the d-block groups and A for the others; this is the system shown in the table above. Chemists have long found it convenient to refer to the elements of different groups, and in some cases of spans of groups by the names indicated in the table shown below.
The general trend is for an atom to gain or lose electrons, either directly (leading to formation of ions) or by sharing electrons with other atoms so as to achieve an outer-shell configuration of s2p6. The outer-shell configurations of the metal atoms in these species correspond to that of neon. As we move farther down (into the fourth row and below), the presence of d-electrons exerts a complicating influence which allows elements to exhibit multiple valances.
The degree of binding is the result of two opposing forces: the attraction between the electron and the nucleus, and the repulsions between the electron in question and all the other electrons in the atom. Because the number of electrons in this model is less than the atomic number Z, the required nuclear charge will also be smaller.
An even simpler although rather crude procedure is to just subtract the number of inner-shell electrons from the nuclear charge; the result is a form of effective nuclear charge which is called the core charge of the atom. An example would be the distance between the iodine atoms of adjacent I2 molecules in crystalline iodine. It is easy enough to measure the distance between adjacent rows of Na+ and Cl– ions in such a crystal, but there is no unambiguous way to decide what portions of this distance are attributable to each ion.
This is not true for KI, but in this solid, adjacent potassium and iodide ions are in contact, allowing estimation of the K+ ionic radius.
Since each row in the periodic table corresponds to an increment in n, atomic radius increases as we move down a column. The apparent discontinuities in this diagram reflect the difficulty of comparing the radii of atoms of metallic and nonmetallic bonding types. If a second electron is lost, the ion gets even smaller; for example, the ionic radius of Fe2+ is 76 pm, while that of Fe3+ is 65 pm.
Because the unit positive charge is concentrated into such a small volume of space, the charge density of the hydrogen ion is extremely high; it interacts very strongly with other matter, including water molecules, and in aqueous solution it exists only as the hydronium ion H3O+. Of course, only one member of such a sequence can be a neutral atom (neon in the series shown below.) The effect of increasing nuclear charge on the radius is clearly seen. They decrease as we move down the table because in each period the electron is being removed from a shell one step farther from the nucleus than in the atom immediately above it. This reflects the fact that as the more-compact d orbitals are being filled, they exert a screening effect that partly offsets that increasing nuclear charge on the outermost s orbitals of higher principal quantum number.
The reversal of the IE trend in this group is often attributed to the more easy removal of the single outer-shell p electron compared to that of electrons contained in filled (and thus spin-paired) s- and d-orbitals in the preceding elements. Because the entering electron is attracted to the positive nucleus, the formation of negative ions is usually exothermic.
This is particularly evident in the first few rows of the periodic table, in which small effects tend to be magnified anyway because an added electron produces a large percentage increase in the number of electrons in the atom. The vertical trend is for electron affinity to become less exothermic in successive periods owing to better shielding of the nucleus by more inner shells and the greater size of the atom, but here also there are some apparent anomalies. Elements with low electronegativities (the metallic elements) are said to be electropositive.


Nevertheless, electronegativity is eminently useful in summarizing the chemical behavior of an element.
Most such scales are themselves based on atomic properties that are directly measurable and which relate in one way or the other to electron-attracting propensity. Although it acts like a metallic element in many respects (forming a positive ion, for example), it can also form hydride-ion (H–) solids with the more electropositive elements, and of course its ability to share electrons with carbon and other p-block elements gives rise to a very rich chemistry, including of course the millions of organic compounds.
These articles appeared in an issue of Chemical & Engineering News that celebrated the 80th anniversary of that publication. This Phoenix College website is an interactive periodic table with a comprehensive database of element properties and isotopes, which can be searched and collated in novel and useful ways. The rows are aligned in such a way that the elements in each vertical column possess certain similarities.
The elements belonging to a given group bear a strong similarity in their chemical behaviors. The two of these that are most important for you to know are the noble gases and the transition metals.
The shell model (as with any scientific model) is less a description of the world than a simplified way of looking at it that helps us to understand and correlate diverse phenomena. This configuration, known as an octet, corresponds to that of one of the noble-gas elements of Group 18.
This effect is especially noticeable in the transition-metal elements, and is the reason for not including the d-block with the representative elements at all. All that matters is the net force, the difference between the nuclear attraction and the totality of the electron-electron repulsions.
It is not possible to specify a definite value for the radius of an isolated atom; the best we can do is to define a spherical shell within whose radius some arbitrary percentage of the electron density can be found. Because of the different ways in which atoms can aggregate together, several different kinds of atomic radii can be defined. The atoms are assumed to be spheres of radius r in contact with each other, each of which sits in a cubic box of edge length 2r.
Because such solids fall into several different classes, several kinds of atomic radius are defined. The best one can do is make estimates based on studies of several different ionic solids (LiI, KI, NaI, for example) that contain one ion in common. The other important factor is the nuclear charge; the higher the atomic number, the more strongly will the electrons be drawn toward the nucleus, and the smaller the atom. If formation of the ion involves complete emptying of the outer shell, then the decrease in radius is especially great. An examination of the successive ionization energies of the first ten elements (below) provides experimental confirmation that the binding of the two innermost electrons (1s orbital) is significantly different from that of the n=2 electrons.Successive ionization energies of an atom increase rapidly as reduced electron-electron repulsion causes the electron shells to contract, thus binding the electrons even more tightly to the nucleus. This results in the familiar zig-zag lines when the first ionization energies are plotted as a function of Z. But the latter, only a minor actor in positive ion formation, is now much more significant. You will make considerable use of electronegativity when you study chemical bonding and the chemistry of the individual elements. The most widely used of these scales was devised by Linus Pauling and is related to ionization energy and electron affinity.
Thus the first short-period elements H and He are chemically similar to the elements Li and Ne at the beginning and end of the second period.
At n=6 we introduce an f block, but in order to hold the table to reasonable dimensions the f blocks are placed below the main body of the table.
In 1985, a new international system was adopted in which the columns were simply labeled 1-18. The principal simplification here is that it deals only with the main group elements of the s- and p-blocks, omitting the d- and f-block elements whose properties tend to be less closely tied to their group numbers. The volume of each box is just the total volume of the solid divided by the number of atoms in that mass of the solid; the atomic radius is the cube root of r. This effect is responsible for the contraction we observe as we move across the periodic table from left to right.
For some atoms, the electron affinity appears to be slightly negative, suggesting that electron-electron repulsion is the dominant factor in these instances.
One reason for this is that the electrons contained in the inner shells of the atom exert a collective negative charge that partially cancels the charge of the nucleus, thus exerting a so-called shielding effect which diminishes the tendency for negative ions to form. Values less than about 2.2 are usually associated with electropositive, or metallic character. In the representation of the scale shown in figure, the elements are arranged in rows corresponding to their locations in the periodic table.
The correlation is obvious; electronegativity is associated with the higher rows and the rightmost columns.



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