6.5 Periodic Variations in Element Properties

Our goal in this lesson is to help you understand how the shape and organization of the modern periodic table are direct consequences of the atomic electronic structure of the elements.

 We begin with the image you saw in the preceding lesson, showing the long form of the table with the “block” structure emphasized. You will recall that the two f  blocks are written at the bottom merely to keep the table from becoming inconveniently wide; these two blocks actually go in between La-Hf and Ac-Db, respectively, in the d block.

Periods, groups and blocks

To construct the table, we place each sequence (denoted by the vertical red bar above) in a separate row, which we call a period. The rows are aligned in such a way that the elements in each vertical column possess certain similarities. 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. Notice that the first period is split in order to reflect these chemical similarities.

Figure 6.44 Periods, Groups, Blocks

Take a moment to see how the above image relates to the complete periodic table, which we reproduce below. Each row that begins with H down through Fr corresponds to a period; in this table, there are seven periods.

Each column, labeled with small blue numbers 1-18 along the top of the table corresponds to a group.

In the past, two different systems of Roman numerals and letters were used to denote the various groups. 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. But the rest of the world used A for the d-block elements and B for the others.

In 1985, a new international system was adopted in which the columns were simply labeled 1-18.

            Figure 6.45 Periodic blocks of Elements

Now look at the section of the table containing groups 3 through 12, labeled
d-block. At the bottom of Group 3 in this block, notice the elements Lanthanum and Actinium, which are shaded in light green. You will see a downward-pointing arrow that points to rwo more rows, also shaded in green.

These two f-blocks, as they are called, are not considered separate periods in this “short form” of the table, but have been squeezed in, beginning at the third period, where Scandium Sc is the first element at which the 3d shell begins to fill.

The “block” nomenclature of the periodic table refers to the sub-orbital type (quantum number ℓ, or s-p-d-f classification) of the highest-energy orbitals that are occupied in a given element. Notice that:

• 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.
• For the second period (n=2) there is a p block but no d block; in the usual “long form” of the periodic table it is customary to leave a gap between these two blocks in order to accommodate the d blocks that occur at n=3 and above.
• 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.

Periodic table families

Chemists have long found it convenient to refer to certain categories of elements by the by special names, commonly known as families. The labels superimposed on the periodic table below are widely used, and worth knowing.

Figure 6.46 Group names of Periodic Table

Of the families shown in the image, it is particularly important that you be able to recognize the alkali metals which begin each row, transition metals (all elements in the d-block Groups 3-12) and the noble gases of Group 18; these families relate directly to the electron configurations of the elements.

Other element families can be entirely arbitrary, such as the elements present in living organisms, the precious or coinage metals, the structural metals such as iron, aluminum and titanium, the elements that are commercially mined in a given country, etc.

2  How electron structures shape the periodic table

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 shell model of the atom

In order to relate the properties of the elements to their locations in the periodic table, it is often convenient to make use of a simplified view of the atom in which the nucleus is surrounded by one or more concentric spherical “shells”, each of which consists of the highest-principal quantum number orbitals (always s– and p-orbitals) that contain at least one electron.

As with any scientific model, the shell model offers a simplified view that helps us to understand and correlate diverse phenomena. 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.

Figure 6.47 Valence electrons

This diagram shows the first three rows of what are known as the representative elements— that is, the s– and p-block elements only. 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. This effect is especially noticeable in the transition-metal elements, and is the reason for not including the d-block with the representative elements.

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. Notice the use of noble-gas symbols to simplify the electron-configuration notation.

Valence electrons and the periodic table

In particular, the number of outer-shell electrons (which is given by the rightmost digit in the group number) is a major determinant of an element’s “combining power”, or valence. 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. This configuration, known as an octet, corresponds to that of one of the noble-gas elements of Group 18.

• the elements in Groups 1, 2 and 13 tend to give up their valence electrons to form positive ions such as Na⁺, Mg²⁺ and Al³⁺, as well as compounds NaH, MgH₂ and AlH₃. The outer-shell configurations of the metal atoms in these species correspond to that of neon.
• elements in Groups 15-17 tend to acquire electrons, forming ions such as P^3–, S^2– and Cl^– or compounds such as PH₃, H₂S and HCl. The outer-shell configurations of these elements correspond to that of argon.
• the Group 14 elements do not normally form ions at all, but share electrons with other elements in tetravalent compounds such as CH₄.

Effective nuclear charge

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. 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. All that matters is the net force, the difference between the nuclear attraction and the totality of the electron-electron repulsions.

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. Because the number of electrons in this model is less than the atomic number Z, the required nuclear charge will also be smaller, and is known as the effective nuclear charge. Effective nuclear charge is essentially the positive charge that a valence electron “sees”.

Part of the difference between Z and Z_effective 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 main actors here are the electrons in the much more compact inner shells which surround the nucleus and exert what is often called a shielding or “screening” effect on the valence electrons.

Figure 6.47 Effective Nuclear Charge

The formula for calculating effective nuclear charge is not very complicated, but we will skip a discussion of it here. 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.

Figure 6.48 Shielding effect