Periodic table atomic radii

As we move down a column (or family) in the periodic table, atomic radius increases. 

Within a given group of the Periodic Table, the radius increases with increasing atomic number. This fact is... 

Figure 1 shows atomic radius as a function of atomic number for the first three periods of the periodic table. Within a period the atomic radius decreases as the atomic number increases, but the atomic radius increases as the period increases. 

As a general rule, the radius of an atom will increase down and to the left in the periodic table. The radius of an atom depends on its electron probability. Because of the inherent uncertainty in defining where the end of one atom s sphere of influence over the electron density ends and the next begins, the radius of an atom can be defined in a number of different ways, and the magnitude of an atom s radius will vary from compound to compound. Furthermore, even for the same molecule, the value might depend on the experimental technique employed. For instance, electron diffraction measures the distance between two nuclei, whereas X-ray crystallography measures the distance between peaks of maximum electron density. Other molecules cannot be crystallized and the radius can only be measured in the gas phase, typically using microwave spectroscopy... 

The electron configuration or orbital diagram of an atom of an element can be deduced from its position in the periodic table. Beyond that, position in the table can be used to predict (Section 6.8) the relative sizes of atoms and ions (atomic radius, ionic radius) and the relative tendencies of atoms to give up or acquire electrons (ionization energy, electronegativity). 

In this section we will consider how the periodic table can be used to correlate properties on an atomic scale. In particular, we will see how atomic radius, ionic radius, ionization energy, and electronegativity vary horizontally and vertically in the periodic table. 

The decrease in atomic radius moving across the periodic table can be explained in a similar manner. Consider, for example, the third period, where electrons are being added to the third principal energy level. The added electrons should be relatively poor shields for each other because they are all at about the same distance from the nucleus. Only the ten core electrons in inner, filled levels (n = 1, n = 2) are expected to shield the outer electrons from the nucleus. This means that the charge felt by an outer electron, called the effective nuclear charge, should increase steadily with atomic number as we move across the period. As effective nuclear charge increases, the outermost electrons are pulled in more tightly, and atomic radius decreases. 

The radii of cations and anions derived from atoms of the main-group elements are shown at the bottom of Figure 6.13. The trends referred to previously for atomic radii are dearly visible with ionic radius as well. Notice, for example, that ionic radius increases moving down a group in the periodic table. Moreover the radii of both cations (left) and anions (right) decrease from left to right across a period. 

Because carbon stands at the head of its group, we expect it to differ from the other members of the group. In fact, the differences between the element at the head of the group and the other elements are more pronounced in Group 14/IV than anywhere else in the periodic table. Some of the differences between carbon and silicon stem from the smaller atomic radius of carbon, which explains the wide occurrence of C=C and G=Q double bonds relative to the rarity of Si=Si and Si=0 double bonds. Silicon atoms are too large for the side-by-side overlap of p-orbitals necessary for -it-bonds to form between them. Carbon dioxide, which consists of discrete 0=C=0 molecules, is a gas that we exhale. Silicon dioxide (silica), which consists of networks of —O—Si- O - groups, is a mineral that we stand on. 

As the positive charge of the nucleus increases, the electrical force exerted by the nucleus on the negatively charged electrons increases, too, and electrons become more tightly bound. This in turn reduces the radius of the orbital. As a result, each orbital shrinks in size as atomic number increases. For example, the 2s orbital steadily decreases in size across the second row of the periodic table from Li (Z = 3 ) to Ne... 

One way that a solid metal can accommodate another is by substitution. For example, sterling silver is a solid solution containing 92.5% silver and 7.5% copper. Copper and silver occupy the same column of the periodic table, so they share many properties, but copper atoms (radius of 128 pm) are smaller than silver atoms (radius of 144 pm). Consequently, copper atoms can readily replace silver atoms in the solid crystalline state, as shown schematically in Figure 12-4. 

Figure 5.2 Atomic radius increases going down a column of the periodic table and generally decreases going across a row.

The Periodic Table forms one of the most remarkable, concise, and valuable tabulations of data in science. Its power lies in the regularities that it reveals, thus, in some respects, it has the same role as the SOM. Construct a SOM in which the input consists of a few properties of some elements, such as electronegativity, atomic mass, atomic radius, and electron affinity. Does the completed map show the kind of clustering of elements that you would expect What is the effect of varying the weight given to the different molecular properties that you are using ... 

Consider the element with atomic number 116 in Group 6A. Even though it has not been isolated, its atomic radius is expected to be somewhat larger than that of Po (1.68 A), probably about 1.9 - 2.0 A, since it lies just below Po on the periodic table. Its outer electrons would lie in the n=l shell, which would be further away from the nucleus than Po s outermost electrons in the n=6 shell. 

II. The general trend is for ionization energy to increase as one moves from left to right across the periodic table and to decrease as one moves down this is the inverse of the trend one finds in examining the atomic radius. 

As it is mentioned, electronegativity is dependent upon atomic radius. In the periodic table, as a period is crossed from left to right, atomic radius decreases, and hence the ability of an atom to attract valence electrons increases. However, as you descend a group, atomic radius increases and therefore ability of an atom to attract valence electrons decreases. So consequently, electronegativity decreases from top to bottom in a group and increases from left to right across a period. 

---