The Properties of the Modern Periodic Table

Periodic trends are patterns in the characteristics of chemical elements in the periodic table. The significant trends are as follows:

• Ionisation energy
• Metallic character
• Atomic Radii
• Electronegativity
• Ionic radius
• Electron affinity
• Chemical reactivity
• Shielding effect

These patterns are caused by changes in the structure of the atoms of the elements within their groups and periods. There are a few exceptions, such as group 3 and 6 ionisation energies.

Periodic law

Periodic trends are founded on periodic law. “Chemical elements are listed in order of increasing atomic number, and their primary properties change in a cyclic pattern “in accordance with periodic law Elements with similar chemical properties repeat at regular intervals.”” Dmitri Mendeleev established this principle. He also claimed that the periodic table was based on several physical and chemical properties of elements rather than just atomic weights . The recurrence of features was later discovered to be attributable to the recurrence of comparable electronic configurations in the outer shells of atoms.

 Ionisation Energy

“Minimum energy required by an isolated atom to remove one electron in its neutral or gaseous state” is how the ionisation potential is defined. As one continues through the period, the ionisation energy increases. The reason for this is that as the nuclear charge increases over time, the nucleus tightens its grip on the electrons. The ionisation energy, on the other hand, decreases as one advances down the group. This is because as the nuclear charge decreases, the valence electrons move away from the nucleus. A number of things influence ionisation energy.

Different Factors Affecting Ionisation Energy Levels 

Nuclear charges – As the nuclear charge falls, the force of attraction between the nucleus and valence electrons reduces, resulting in lower ionisation energy.

 Shielding effect – As the nuclear charge increases, the shielding effect increases as well, resulting in an increase in ionisation energy.

Atomic radius – As the atomic radius increases, the force of attraction between the nucleus and the valence electrons decreases. As a result, the ionisation decreases as the atomic radius grows.

Half-Filled Valence Shells – In pseudo-filled or half-filled valence shells, ionisation energy is strong. The ionisation number for the electron in that shell will be high if the fundamental quantum number is low.

Exceptions – None of the elements in the oxygen and boron families follow the periodic trend. They use a little less energy than the conventional trend

Metallic Property

The metallic property of an element refers to its ability to conduct electricity. The metallic properties increase as the nuclear charge decreases in the group. Because the valence electrons are loosely bound by the nucleus, they can conduct electricity well. The metallic property, on the other hand, decreases over time as nuclear charge accumulates. As a result, the valence electrons’ force of attraction with the nucleus increases, making it impossible for them to conduct electricity or heat.

Atomic Radii

When an atom is in equilibrium, the atomic radius is the distance between its nucleus and its outermost stable electron orbital. The atomic radius shrinks as the nuclear charge increases. Because the force of attraction between the nucleus and the valence electrons increases as nuclear charge increases, the nucleus tightens its hold on the electron, resulting in smaller atomic radii. As a group advances, its atomic radius increases. The reason for this is that the nuclear charge decreases when more shells are added.. However, the atomic radii expand diagonally as well, resulting in rare exceptions.

Example: Along the Period – Li> Be > B > C > N > O > FDown the Grp – Li < Na < K < Rb < Cs.

 Electronegativity

Electronegativity is the ability of an atom or molecule to attract a pair of electrons. The bond formed as a result of this is determined by the difference in electronegativity of the atoms. With increasing nuclear charge, electronegativity rises. As you progress down a group, the electronegativity decreases as the nuclear charge decreases. The reason for this is that the distance between the nucleus of the atom and the valence electrons is large, making electron loss easy.

Example: Along the Period- Li < Be < B < C < N < O < F Down the Grp – Li > Na > K > Rb > Cs

Exception

The exception is group 13 elements, whose electronegativity climbs from aluminium to thallium as a result. In addition, tin has a higher electronegativity in group 14 than lead.

Electron Affinity

Electron affinity is the tendency of an atom to accept an electron or an electron pair. This is a hallmark attribute of nonmetals because they gain electrons and generate anions. As nuclear charge grows, the electron affinity rises with it. The group’s nuclear charge lowers as the nuclear charge decreases. Because fluorine has the highest electronegativity, noble gases are excluded. Because they have a full valence shell, they cannot gain or lose electrons.

Shielding Effect

Inner electrons resisting an outer electron can be described like this. It can also refer to the number of nuclei that can regulate electrons in the outer shell. The effective nuclear charge decreases along the group as the shielding effect increases. As the nuclear charge grows, the effective nuclear charge grows as well. We can draw the following conclusions to sum up everything.

Ionic Radius

The electrons in an ion’s many shells and the nucleus make up the ion. The ionic radius of an ion is the distance between the nucleus and the electron in the ion’s last outermost shell. There is a tendency that may be seen in the periodic table based on the ionic radius of different elements. This tendency can be summarised as follows: 

• As we proceed down the periodic table from top to bottom, the ionic radius of the elements increases in value. This happens because as we move down the periodic table, the number of layers or shells of electrons grows.
• As we proceed from left to right on the periodic table, the ionic radius tends to shorten. Although it appears strange that as additional protons, electrons, and neutrons are added, the ionic size decreases. However, when we move sideways on the periodic table, metals lose their outer electron layers, allowing cations to form. The ionic radius increases when the number of electrons in an ion exceeds the number of protons, resulting in a significant decrease in nuclear charge. This is true not just for ionic but also for atomic radius; nevertheless, the two are not the same.
• This is true not just for ionic but also for atomic radius; nevertheless, the two are not the same.

Chemical Reactivity

The reactivity of an atom refers to its ability to react with any other substance. Ionisation energy (how easily electrons are shed from the topmost layer) and electronegativity (how quickly an atom accepts another atom’s electrons) are frequently used to govern chemical reactivity. The periodic table’s chemical reactivity trend is based on this mechanism of electron transfer and interchanging.

Conclusion

One of the most significant achievements in the science of chemistry is the periodic table. It’s full of patterns that help us comprehend the world around us better. We wouldn’t have many of the items and medicines we have today if it weren’t for it. The knowledge gathered from the periodic table can open up a plethora of windows into the cosmos we live in. You should have a much better knowledge of the p after completing this activity.