Complete the exercises below. Explain why the transition metals in periods 5 and 6 have nearly identical radii within each group.

Verified step by step guidance

1. Step 1: Understand the concept of atomic radius. The atomic radius is the distance from the nucleus of an atom to the outermost shell of electrons. It generally decreases across a period and increases down a group in the periodic table.
2. Step 2: Recognize the role of electron shielding. As you move down a group, additional electron shells are added, which typically increases the atomic radius. However, the effect of electron shielding can counteract this increase.
3. Step 3: Consider the lanthanide contraction. In period 6, the presence of the lanthanide series (elements 57-71) causes a contraction in atomic size due to poor shielding by the f-electrons. This results in a smaller than expected increase in atomic radius for period 6 transition metals.
4. Step 4: Compare periods 5 and 6 transition metals. Due to the lanthanide contraction, the atomic radii of period 6 transition metals are similar to those of period 5, despite being in a lower period.
5. Step 5: Conclude with the impact on transition metals. The nearly identical radii of transition metals in periods 5 and 6 within each group are primarily due to the lanthanide contraction, which offsets the expected increase in size from additional electron shells.