Trends in the Periodic Table

TRENDS IN THE PERIODIC TABLE

            Though the Periodic Law states that the periodicity of elements is based on increasing atomic number, the physical and chemical properties of elements are ultimately based on the electron configuration of its atoms. Some of the key properties of elements that can be studied in the periodic table are Electronegativity, Ionization energy, Electron Affinity, and Atomic radius.

                Two important factors must be discussed first before venturing into the periodic trend. These factors should be understood first in order for us to better understand why such trends occur. These are the (1) Principal Quantum Number (n) of the valence electrons and the (2) Effective Nuclear Charge (Z_eff) “felt” by the valence electrons.

                The Principal Quantum Number (n) tells us the average distance of the electron from the nucleus. As n increases, the outermost electrons called valence electrons gets farther away from the nucleus. The Effective Nuclear Charge (Z_eff) is the attraction of the valence electrons to the nucleus. This attraction is affected by the inner electrons that screens or shields the valence electrons from the nucleus. This is called the Shielding Effect. Because of the inner electrons, the valence electrons do not feel the full nuclear charge.

                                Z_eff< Z

              Z_eff = Z – S

                where: S is the Screening or Shielding factor;

                                 Z is the Nuclear Charge; and

                                Z_eff is the effective Nuclear Charge

The valence electron does not feel the full nuclear charge because the inner electrons screen or shield it from the nucleus

               The Effective Nuclear Charge can be likened to Family A with one child and to Family B with seven or eight children. The child in Family A can have more attention coming from his parents compared to the children in Family B.

                Electronegativity is defined as the ability of an atom to attract electrons towards itself. In other words, it is the electron appeal of an atom. The higher the electronegativity of an element, the more it can pull its own valence electrons and the valence electrons of other elements towards its nucleus. The trends for electronegativity of elements in the periodic table are as follows:

§  Increases from Left to Right (across the period)

Ø  This is due to the fact that the average distance of the valence electrons from the nucleus (n) are the same while the number of charge of the elements increases as you go along the period. In effect, the nuclear charge becomes stronger from left to right of the periodic table, making it attract more electrons towards itself.

§  Decreases from Top to Bottom (down the group)

Ø  This is because the average distance of the valence electrons from the nucleus (n) becomes greater, hence, making the effective nuclear charge of the element weaker. Because the effective nuclear charge is weak, its electronegativity becomes weaker also.

                 Noble gases have weak or no electronegativity values because they have completely filled s- and p-sublevels and attracting electrons toward itself would make it unstable. This explains why Noble Gases are unreactive to most substances. This also means that the most electronegative element is Fluorine and the least electronegative is Francium. Furthermore, this makes Fluorine the most reactive nonmetal and why most nonmetals are negatively charged.

                Ionization Energy is the energy absorbed by an atom to dislodge or remove its valence electrons. The trends for ionization energy of elements in the periodic table are as follows:

§  Increases from Left to Right (across the period)

Ø  Because the effective nuclear charge increases as you move from left to right of the periodic table, it would require less energy to dislodge electrons for the elements in the Alkali Metal Group while more energy is needed for elements in the Halogen Group.

§  Decreases from Top to Bottom (down the Group)

Ø  As n increases, the effective nuclear charge decreases. As effective nuclear charge decreases down the group, the lesser energy is required to dislodge electrons

This trend explains why metals are positively charged, they tend to give away their valence electrons because of weak effective nuclear charge and large atomic size. This also makes Francium the most reactive metal.

As more electrons are removed or dislodged from the atom, the greater the ionization energy becomes. This is because the number of negative charge becomes lesser, thus, the nucleus has a greater pull on the electrons, like in a tug-o-war where one team loses a member and the whole team gets pulled.

   

  

                Electron Affinity is defined as the energy released by an atom after gaining an electron. Generally, the greater the electron affinity of an atom, the more stable it becomes after gaining an electron. Electron Affinity can be likened to the excitement of atoms to have the same electron configuration as the noble gases, which are stable atoms. The greater their excitement, the more energy they release after gaining an electron. The trends for electron affinity of elements in the periodic table are as follows:

§  Increases from Left to Right (across the period)

Ø  A metal’s tendency is to give away their electrons so their excitement to receive electrons is low. A nonmetal’s tendency is to accept electrons (due to high effective nuclear charge), thus, it has high excitement to receive electrons, hence, high energy is release after gaining an electron.

§  Decreases from Top to Bottom (down the group)

Ø  Effective nuclear charge decreases down the group, thus, lesser excitement to receive electrons an lesser energy is released after gaining an electron.

Ø  Effective nuclear charge decreases down the group, thus, lesser

excitement to receive electrons and lesser energy is released after gaining an electron

This trend is only applicable for the first electron gained by a nonmetal. For the

second electron, electron affinity is positive, thus, energy is absorbed because, like in the case of O^-, energy would be needed to attach another electron to a negatively charged ion due to the electrostatic repulsion between the two.

                Atomic radius is the average distance between the nucleus and the valence electrons.

The trends for atomic radius of elements in the periodic table are as follow:

§  Decreases from left to right (across the period)

Ø  Though the principal quantum number (n) of the valence electrons across the group is the same, the effective nuclear charge increases from left to right thus, the valence electrons get attracted towards the nucleus making the atomic radius smaller

§  lncreases from top to bottom - (down the group)

Ø  The principal quantum number (n) increases from top to bottom, thus, the distance between the nucleus and the valence electrons becomes greater.

 Atomic Radius vs Ionic Radius

Atomic radius, as studied earlier, is the average distance between the nucleus and the valence electrons. When atoms lose or gain electrons, they become positive or negatively charged and become ions. The ionic radius is an estimated radius of an ion in an ionic compound. When an atom loses or gains electrons, the size of the particle suddenly changes.

When metals lose their valence electrons, the number of protons becomes greater than the number of electrons, thus, the effective nuclear charge becomes greater and the electrons get attracted towards the nucleus, making the ionic radius smaller. The tendency of metals is to accept electrons due to their strong electronegativity. When a nonmetal accepts an electron, the electron reduces the effective nuclear charge. The additional electron causes the other electrons to move away from each to relieve the repulsive force between them, making the ionic radius larger. As more electrons are added to the nonmetal ion, the larger the ion becomes.

To summarize the main points in the trend in ionic radius:

§  Increases from top to bottom

§  Decreases with increasing positive charge for metals along a period

§  Increases a charge becomes more negative for nonmetals along a period

Enhancing Skills:

Using your knowledge on the Periodic Trends, arrange the following elements:

A. ln lncreasing Electronegativity:

1. Ca, Ba, Sr, Be, Ra

2. Sb, P, N, Bi, As

3. Br, K, Ge, Ca, Se

B. ln Decreasing lonization Energy:

1. K, Na, Li, Rb, Cs

2. Br, F, Cl, l, At

3. Po, Bi, Pb, At Tl

C. ln Decreasing Atomic Radius

1, Mo, Nb, Tc, Ru, Zr

2. Cu, Fe, Zn, Ni, Co

3. B, ln, Tl, Ga, Al