Question

Which of the series of elements listed below would have nearly same atomic radii (1) \(\mathrm{F}, \mathrm{Cl}, \mathrm{Br}, \mathrm{I}\) (2) \(\mathrm{Na}, \mathrm{K}, \mathrm{Rb}, \mathrm{Cs}\) (3) \(\mathrm{Li}, \mathrm{Be}, \mathrm{B}, \mathrm{C}\) (4) \(\mathrm{Fe}, \mathrm{Co}, \mathrm{Ni}, \mathrm{Cu}\)

Short answer

(4) Fe, Co, Ni, Cu

Step-by-step solution

Understand Atomic Radius

Atomic radius refers to the distance from the nucleus of an atom to the outermost electron. Generally, atomic radius increases as you go down a group in the periodic table and decreases as you move across a period.

Analyze Each Series

Look at each series of elements and consider if they belong to the same period or group. Elements in the same period that are close to each other tend to have nearly the same atomic radii due to similar nuclear charge and shielding effect.

Set 1: \( \mathrm{F}, \mathrm{Cl}, \mathrm{Br}, \mathrm{I} \)

Fluorine, Chlorine, Bromine, and Iodine are all halogens in Group 17. They are in different periods, so their atomic radii vary significantly.

Set 2: \( \mathrm{Na}, \mathrm{K}, \mathrm{Rb}, \mathrm{Cs} \)

Sodium, Potassium, Rubidium, and Cesium are alkali metals in Group 1. They are in different periods, so their atomic radii increase as you go down the group.

Set 3: \( \mathrm{Li}, \mathrm{Be}, \mathrm{B}, \mathrm{C} \)

Lithium, Beryllium, Boron, and Carbon are in the same period (Period 2) but are in different groups. Their atomic radii decrease as you move from left to right across the period.

Set 4: \( \mathrm{Fe}, \mathrm{Co}, \mathrm{Ni}, \mathrm{Cu} \)

Iron, Cobalt, Nickel, and Copper are all transition metals in the same period (Period 4). Transition metals in the same period tend to have similar atomic radii due to their d-electron shielding.

Conclusion

Among the given series, the transition metals Iron (Fe), Cobalt (Co), Nickel (Ni), and Copper (Cu) have nearly the same atomic radii because they are in the same period and their atomic radii vary very slightly.

Key concepts

Periodic Table Trends

Understanding the trends in the periodic table is fundamental in predicting and explaining the properties of elements. One of the key trends is atomic radius, which is the distance from the center of an atom's nucleus to its outermost electron. Based on periodic table trends:

• Atomic radius increases as we move down a group (or column) in the periodic table because each element down a group has an additional electron shell, increasing the size.
• Atomic radius decreases as we move from left to right across a period (or row). This is due to the increasing positive charge in the nucleus, which pulls electrons closer, decreasing the size.

These trends help explain why elements in the same group (like the alkali metals in Group 1) show increasing atomic radii down the group and why elements close to each other in a period (like period 4 transition metals) have similar atomic radii.

Transition Metals

Transition metals, found in the d-block of the periodic table, exhibit unique properties due to their electron configurations. These elements are found in groups 3-12 and are known for having multiple oxidation states, forming colorful compounds, and being good conductors of electricity. A crucial feature of transition metals is their relatively similar atomic radii, especially within the same period. This phenomenon occurs because:

• The additional electrons go into the inner d-subshell, rather than the outer valence shell. This results in a weaker impact on the overall size of the atom.
• The effective nuclear charge is somewhat balanced by the electron shielding from these inner d-electrons, making the change in atomic radius less pronounced than in s- or p-block elements.

Hence, elements like Iron (Fe), Cobalt (Co), Nickel (Ni), and Copper (Cu) in period 4 have nearly identical atomic radii.

Electron Shielding

Electron shielding, also known as screening, is a concept crucial for understanding atomic size and ionization energy. It refers to the phenomenon where inner-shell electrons blur the effect of the positive nuclear charge on the outer-shell electrons. Here’s how it influences atomic properties:

• As the number of inner electron shells increases in an atom, the outermost electrons experience less nuclear attraction. This increases the atomic radius.
• In transition metals, the d-electrons provide significant shielding, which stabilizes the size of the atom even when nuclear charge increases across the period.

This concept helps explain why even elements with the same number of outer electrons but differing inner shell configurations can exhibit varying atomic radii. For example, despite increasing nuclear charge from Iron (Fe) to Copper (Cu), the d-electron shielding minimizes the effect, resulting in nearly uniform atomic radii among these transition metals.