Chapter 20, Problem 1RQ

What two first-row transition metals have unexpected electron configurations? A statement in the text says that first-row transition metal ions do not have 4s electrons. Why not? Why do transition metal ions often have several oxidation states, whereas representative metals generally have only one?

Interpretation Introduction

Interpretation: The first-row two elements that have unexpected electron configurations; the reason behind the absence of $4\text{s}$ electrons in first-row transition metal ions and the reason behind the exhibition of several oxidation states in transition metal ions but not only one oxidation state in representative elements is to be stated.

Concept introduction: The metal atoms that act as a bridge between the $\text{s}$ and $\text{p}$ block elements are called transition elements. They are called the $\text{d}$ block elements and these include the elements of group $3$ to group $12$. These elements forms a large number of coordination compounds.

To determine: The first-row two elements that have unexpected electron configurations; the reason behind the absence of $4\text{s}$ electrons in first-row transition metal ions and the reason behind the exhibition of several oxidation states in transition metal ions but only one oxidation state in representative elements.

Answer to Problem 1RQ

Answer

The two metals that have unexpected electron configuration are chromium and copper.

Due to less effect of effective nuclear charge over $4\text{s}$ electrons, first-row transition metals do not have $4\text{s}$ electrons.

Due to the less energy gap between the $\left(n-1\right)\text{d}$ orbital and $n\text{s}$ orbital, the electrons of both the orbital participates in the bond formation and leads to variable oxidation states of the transition metal atom.

Explanation of Solution

Explanation

The expected electronic configuration of chromium is represented as,

$\left[\text{Ar}\right]3{\text{d}}^{4}4{\text{s}}^2$

The actual electronic configuration of chromium is represented as,

$\left[\text{Ar}\right]3{\text{d}}^{5}4{\text{s}}^1$

The expected electronic configuration of copper is represented as,

$\left[\text{Ar}\right]3{\text{d}}^{9}4{\text{s}}^2$

The actual electronic configuration of chromium is represented as,

$\left[\text{Ar}\right]3{\text{d}}^{10}4{\text{s}}^1$

The half filled and fully filled orbital are more stable. Therefore to attain stability, these metals attain the unexpected configuration.

Due to less effect of effective nuclear charge over $4\text{s}$ electrons, first-row transition metals do not have $4\text{s}$ electrons.

The effect of effective nuclear charge over $4\text{s}$ electrons is lower than that is present over $3\text{d}$ electrons. So, it becomes easy to remove electrons from $4\text{s}$ orbital than $3\text{d}$ orbital. That’s why first-row transition metals do not have $4\text{s}$ electrons.

Due to the less energy gap between the $\left(n-1\right)\text{d}$ orbital and $n\text{s}$ orbital, the electrons of both the orbital participates in the bond formation and leads to variable oxidation states of the transition metal atom.

The electrons in transition metal atom are present in $\left(n-1\right)\text{d}$ and $n\text{s}$ orbital and any small effect can lead to anomalous electronic configuration. Therefore, the oxidation state varies from $+3$ to $+7$. But the number of electrons present in representative elements are lower than the transition metal atoms; therefore they exhibit only one oxidation state.

Conclusion

Conclusion

Transition metal atoms exhibit properties between $\text{s}$ and $\text{p}$ block elements and exhibit variable oxidation states due to participation of inner d orbital electrons and the $\text{s}$ orbital electrons in the bond formation.