Textbook Question
Draw the crystal-field energy-level diagrams and show the placement of electrons for each of the following complexes:
d. [NiCl4]2+ (tetrahedral),
Ch.23 - Transition Metals and Coordination Chemistry
Chapter 23, Problem 68
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 guidance1
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.
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.
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.
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.
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.
Key Concepts
Here are the essential concepts you must grasp in order to answer the question correctly.
Atomic Radius
The atomic radius is the distance from the nucleus of an atom to the outermost shell of electrons. In transition metals, the atomic radius is influenced by the number of electron shells and the effective nuclear charge experienced by the outer electrons. As you move down a group in the periodic table, additional electron shells are added, but the increase in nuclear charge is offset by electron shielding, leading to similar radii.
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02:02
Atomic Radius
Electron Shielding
Electron shielding occurs when inner-shell electrons repel outer-shell electrons, reducing the effective nuclear charge felt by the outer electrons. In transition metals of periods 5 and 6, the presence of d-electrons contributes to this shielding effect. As a result, even though period 6 elements have more protons, the increased shielding from the additional inner electrons leads to similar atomic radii compared to period 5.
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Effective Nuclear Charge (Z_eff)
Effective nuclear charge (Z_eff) is the net positive charge experienced by an electron in a multi-electron atom, accounting for both the total nuclear charge and the shielding effect of other electrons. In transition metals, Z_eff increases down a group, but the increase is not sufficient to significantly alter the atomic radius due to the compensating effects of electron shielding. This results in nearly identical radii for transition metals in periods 5 and 6.
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Related Practice
Textbook Question
Sketch the structure of the complex in each of the following compounds and give the full compound name:
a. cis-[Co(NH3)4(H2O)2] (NO3)2
Textbook Question
Sketch the structure of the complex in each of the following compounds and give the full compound name:
b. Na2[Ru(H2O)Cl5]
Textbook Question
Sketch the structure of the complex in each of the following compounds and give the full compound name:
c. trans-NH3[Co(C2O4)2(H2O)2]