Hint: Think about what the terms atom and ion mean. Consider the number of electrons relative to the number of protons and what it might mean for the effective nuclear charge. How does the effective nuclear charge affect the radius of the atom or the ion?
Complete step by step answer:
The radius of an atom or an ion is defined as the distance from the center of the nucleus of the atom to the electron present in the outermost orbit.
- Atomic radius:
The atomic radius is the standard radius of that particular element. The number of electrons in the outermost orbit does not change and neither does the orbit which is declared as the outermost one varies. Thus, the number of electrons will always be equal to the number of protons in the nucleus. The atomic radius decreases along a period as we go from left to right in the periodic table and increases as we go down the group
Before going to what the ionic radius means, let us first look at what the term effective nuclear charge means. Effective nuclear charge is the net positive charge that is experienced by an electron due to the protons in the nucleus. The decrease in atomic radius along a period is due to the increase in effective nuclear charge on each electron. They are attracted more strongly towards the nucleus as the number of protons go on increasing and the distance of the electron in the outermost shell from the nucleus remains similar.
- Ionic radius:
Ionic radius is when the atom either gains or loses an electron and turns into an anion or a cation respectively.
When any atom gains an electron, it is called an anion. The number of electrons increases but the number of protons in the nucleus remains the same. Thus, the effective nuclear charge on the electrons in the outermost orbit reduces. The nuclear force attracting the electrons becomes lesser and the atomic size increases.
When an atom loses an electron, it is called a cation. The number of electrons decreases but the number or protons in the nucleus remains the same. Thus, the effective nuclear charge on every electron in the outermost shell increases. The force attracting each electron in the outermost shell increases and the atomic size decreases.
Note: Remember that for larger atoms, you also need to take into account factors like shielding effect and electronegativity. The atomic size may be closer to the ionic size due to the shielding effect.