Periodic Trends in Properties of Elements and Their Explanation

The periodic table is a tabular display of various chemical elements that are arranged according to their electronic configuration, atomic number, and common chemical properties. There are certain trends that are common throughout all groups and periods. 

The rows are termed as periods and the periodic table has 7 periods. The metals are present on the left rows and towards the right-hand side are the nonmetals. Whereas, the columns are termed as groups. In groups, elements have various chemical behaviors. There are 18 groups and the halogens are present under group 17 and noble gasses are present under group 18.

Periodic Trends

The specific patterns in the properties of chemical elements present in the periodic table are known as periodic trends. The important trends are,

1. Ionization energy

2. Metallic character

3. Atomic Radii

4. Electronegativity

5. Ionic radius

6. Electron affinity

7. Chemical reactivity

8. Shielding effect

These trends arise due to changes in the structure of atoms of the elements within their groups and periods. A few exceptions exist, for example, the ionization energy of groups 3 and 6.

Periodic Law

Periodic law forms the basis for periodic trends. According to periodic law, “the chemical elements are listed in an order of increasing atomic number, and main properties thus undergo cyclic changes. Elements having similar chemical properties re-occur in regular intervals”

This principle was given by Dmitri Mendeleev. He also stated that the periodic table was not just based on the atomic weights, but also based on various physical and chemical properties of elements.

Later it was also found that the recurrence of properties was due to the recurrence of similar electronic configurations in the outer shells of atoms.

1. Ionization Energy

The ionization potential can be defined as, 

“Minimum energy required by an isolated atom to remove one electron in its neutral or gaseous state”

As one goes across the period, the ionization energy increases. The reason behind this is that the nuclear charge across the period increases and thus the electrons are strongly held by the nucleus. 

But as one goes down the group, the ionization energy decreases down the group. The reason behind this is, down the group the valence electrons go farther away from the nucleus, thus the nuclear charge decreases.

Factors affecting ionization energy

Various Factors that Affect the Ionization Energy Levels 

Nuclear Charge

Lower the nuclear charge lower is the force of attraction between the nucleus and valence electrons, thus low ionization energy. 

Shielding Effect

Shielding effect increases as nuclear charge increases, thus with an increase in shielding effect the ionization energy also increases. 

Atomic Radius

As the atomic radius increases the force of attraction between the nucleus and valence electrons also decreases. Thus, with an increase in atomic radius the ionization decreases.

Half-Filled Valence Shells

Pseudo filled or half-filled valence shells have high ionization energy.

A simple principle that can be used is that, if the principal quantum number is low, then the ionization number will be high for the electron present in that shell.

Exceptions

All the elements in the oxygen and boron family are an exception to the above stated periodic trend. They require a little less energy than the usual trend.

2. Metallic Property

Metallic property of an element can be defined as their ability to conduct electricity. The metallic properties increase down the group as the nuclear charge decreases down the group. Since the valence electron is loosely bounded by the nuclei, they are able to conduct electricity well. 

But across a period, the metallic character decreases as nuclear charge increases. This causes the force of attraction between the valence electrons and the nuclei increases, thereby inhibiting them from conducting electricity or heat.

3. Atomic Radii

The atomic radius is the distance between the atomic nucleus and outermost stable electron orbital of an atom which is at equilibrium. Across a period the atomic radius decreases, as the nuclear charge increases. The reason for the decrease is as nuclear charge increases, the force of attraction between the nucleus and the valence electrons also increases, and the nucleus holds the electron tightly, thereby decreasing the atomic radii. 

In a group, the atomic radius increases down the group. The reason being, new shells are being added and thus the nuclear charge decreases. But the atomic radii also increase diagonally causing some exceptions. 

Example: 

Along the Period – Li> Be > B > C > N > O > F

Down the Grp – Li < Na < K < Rb < Cs

4. Electronegativity

Electronegativity can be defined as the ability of an atom or a molecule to attract a pair of electrons. The bond formed due to this is mainly determined by the difference between the electronegativity of the atoms. 

Across the period, the electronegativity increases as nuclear charge increases. Moving down a group, the electronegativity decreases as nuclear charge decreases. The reason being the distance between the nucleus of the atom and the valence electrons is long and thus the electrons are easily lost. 

Example: 

Along the Period- Li < Be < B < C < N < O < F 

Down the Grp - Li > Na > K > Rb > Cs

Exception

The group 13 elements are an exception and thus the electronegativity increases from aluminum to thallium. Also, in group 14, the electronegativity of tin is higher than lead.

5. Electron Affinity

Electron affinity can be defined as the tendency of an atom to accept an electron or an electron pair. This is a characteristic feature of nonmetals as they gain electrons to become anions. Across a period, the electron affinity increases as nuclear charge increases.

Down the group, it decreases, as the nuclear charge decreases. Fluorine has the highest electronegativity and noble gasses are not included in this. The reason being they have a full valence shell and thus can neither gain nor lose electrons.

6. Shielding Effect

It can be defined as the repelling of an outer electron by the inner electrons. It can also be used to explain how many nuclei can control the outer electrons. The effective nuclear charge decreases down the group due to increased shielding effect. Across a period, the effective nuclear charge increases as nuclear charge increases.

To summarize the whole thing, we can make the following conclusions.

7. Ionic Radius

An ion consists of the electrons in its numerous shells and the nucleus. The distance between the nucleus and the electron in the last outermost shell of an ion is known as the ionic radius of an ion. Based on the ionic radius of different elements there’s a trend that can be identified in the periodic table. Basically, this trend can be seen as;

• If we move from the top of the periodic table down to its bottom the ionic radius of the elements will increase in value. This happens because as we move down the periodic table the number of layers or shells of electrons increase in number.

• If we move sideways from left to right on the periodic table then the ionic radius tends to decrease in size. Although it seems odd that the ionic size would decrease as more protons, electrons and neutrons are added. However, this happens because as we move sideways on the periodic table the metal shed their outer electrons layers in order to form cations. For non-metals the ionic radius increases as the number of electrons present in the ion exceeds the number of protons causing significant decrease in nuclear charge.

This trend applies to not only ionic radius but also to atomic radius; however, these are different from each other.

8. Chemical Reactivity

Reactivity of an element refers to the capacity at which an atom tends to react with any other substance. The chemical reactivity is often regulated by the ionization energy (how simply electrons are shed from the outermost layer) and electronegativity (how fast an atom takes another atom’s electrons). This process of transfer and interchanging of electrons is the principle on which the chemical reactivity trend occurs in the periodic table.

• In metals the chemical reactivity decreases as we move sideways from left to right on the periodic table. Whereas, the reactivity increases as we move from top to the bottom groups of the table. The farther downwards or towards the left we move the exchange of electrons becomes easier and more rapid, increasing the chemical reactivity of the elements.

• In non-metals it’s the opposite. The chemical reactivity increases as we move from left to the right of the table. And the reactivity decreases as we move towards the bottom groups from the top groups. The farther upwards or to the right the easier it becomes for atoms to shed their electrons in exchange of other electrons, increasing the electronegativity which in turn makes the chemical reactivity of the elements more rigorous.

Facts Based on the Period Trends and Periodic Tables

• Most noble gases- helium(He), neon(Ne), argon(Ar), krypton(Kr), xenon(Xe), and radon(Rn)- have zero electronegativity because they are extremely stable. They have full valence electronic layers due to which they do not lose or gain electrons easily.

• Out of the 118 elements of the periodic table, 90 can be found in nature while the rest 28 are completely man-made.

• While hydrogen(H) is the lightest element present in the periodic table (can be found in the top left corner), oganesson (Og) is the heaviest element (can be found in the lower right corner)

• Almost 75% of elements in the periodic table are metals. While there are only a few non-metals.

• The only two elements that are liquid at room temperature are bromine and mercury.

• If the periodic table is folded half along its group 4 elementsThe groups that lie on top of each other can be fused with each other perfectly because they have harmonizing electron structure. Hence, they fit together in complete stability.