The region of space occupied by an atom's electrons (i.e.,
the atomic orbitals) is best thought of as a probability distribution of electron density,
therefore there is no well-defined "outer edge" of electron density. Still, the
size of an atom, even though it is difficult to define and accurately measure, often has
significant bearing chemical properties. There are several different ways to define the
size of an atom. Although we will focus on one particular definition, the others follow
the same periodic trend.
An atom's covalent radius is one-half the distance between atoms in the homonuclear substance. For instance, the I - I distance in I2 is 2.66 so the covalent radius of I is 1.33 . A remarkably good approximation of single bond distances in compounds can be obtained by taking the sum of covalent radii. For example, the covalent radius of carbon is 0.77 and the covalent radius of bromine is 1.14 . From these data we predict the C - Br distance in CBr4 (or any other compound for that matter) to be 1.91 which is quite close to the experimentally measured distance of 1.94 .
A plot of atomic radius versus atomic number (see the figure) reveals some interesting trends. Proceeding down a given group atoms generally get larger. This trend is expected since the value of the principal quantum number, n, is greater for the valence shell electrons of atoms lower in a group. Recall that the value of n determines the size of an orbital; the higher the value of n, the larger the orbitals. Shown below are the sizes, in angstroms, for some of the group 1, group 2, and group 13 elements. In parentheses are the atomic number and valence electron configuration.
|Li, 1.52 (Z = 3; 2s1)||Be, 1.13 (Z = 4; 2s2)||B, 0.88 (Z = 5; 2s22p1)|
|Na, 1.86 (Z = 11; 3s1)||Mg, 1.60 (Z = 12; 3s2)||Al, 1.43 (Z = 13; 3s23p1)|
|K, 2.27 (Z = 19; 4s1)||Ca, 1.97 (Z = 20; 4s2)||Ga, 1.22 (Z = 31; 4s23d104p1)|
|Rb, 2.47 (Z = 37; 5s1)||Sr, 2.15 (Z = 38; 5s2)||In, 1.63 (Z = 49, 5s24d105p1)|
For the alkali metals and alkaline earth metals, the predicted trend of size increasing as we progress down a group holds true. We can also see from these data (as well as from the graph) that atomic size decreases progressing from left to right across a period. Initially this trend may seem counter-intuitive, but the pattern makes sense if we remember the trend for effective nuclear charge, Zeff. As we progress from left to right across a period, Zeff increases. As Zeff increases, the valence electrons experience a greater nuclear charge and they are pulled closer to the nucleus, thus the trend is for decreasing size progressing to the right across a period.
In the data given for group 13 there is an interesting deviation
from the expected trend for atomic radius. At first, we might expect that an atom of Ga
would be larger than an atom of Al based on their positions in group. Gallium, however, is
considerably smaller than Al due to a phenomenon known as the d-block
contraction. As explained above, atoms get smaller as we progress from left to
right across a given period, thus in period three Na is larger than Mg, Mg is larger than
Al, Al is larger than Si, etc., until we arrive at Ar, the smallest atom in row. The next
element, K, is in period four and is, as expected, larger than any period three atom.
Again, as we progress to the right atomic radius decreases due to increasing Zeff.
It is in period four, however, that we first encounter the transition elements. Atomic
radius decreases as we move from Sc to Ti to V, etc., as electrons are added to the
d-subshell. By the time we completely fill the d-subshell and reach the p-block elements,
the atomic radius has decreased so much that Ga is actually smaller than Al. Because of
the d-block contraction, the post transition elements of period four are about the same
size as their period three congeners. Similarly, the occurrence of the f-block elements
between the period six s- and p-blocks causes the period six p-block elements to be about
the same size as their period five congeners. This effect is called the lanthanide
contraction, or f-block contraction.
Sizes of Ions
Cations are smaller than the corresponding neutral atoms; anions are larger than the neutral atoms from which they are formed. The Na1+ cation, for instance, with 11 protons attracting its 10 electrons, is smaller than an atom of Na which has 11 protons and 11 electrons. An atom of F, with 9 protons and 9 electrons, is smaller than the F1- anion. Similarly, dianions of a particular element are even larger than monoanions; dications are even smaller than monocations formed from the same atom.