Question

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

Short answer

The transition metals in Periods 5 and 6 have nearly identical radii in each group because the increase in effective nuclear charge due to the higher number of protons is balanced by the increased electron shielding effect from the additional 4f electrons. This causes the effective nuclear charge experienced by the outermost electrons to remain almost constant, leading to nearly identical atomic radii for these elements in each group.

Step-by-step solution

Recall the Electron Configuration of Transition Metals

Transition metals are located in the d-block of the periodic table. They have their outermost electrons in the d-subshell. The electron configuration of these elements can be represented in the form: [Noble Gas] \(\mathrm{n-1}\) d \(\mathrm{x}\) n s \(\mathrm{y}\), where n is the principal quantum number (period), x and y are the number of electrons in the d and s subshells, respectively.

Understand the Effective Nuclear Charge

The effective nuclear charge is the net positive charge experienced by an electron in an atom. It is affected by the number of protons in the nucleus (atomic number) and the electron shielding effect caused by the inner electron shells. As we move down a group in the periodic table, the atomic number increases, which also increases the effective nuclear charge.

Explain the Electron Shielding Effect

The electron shielding effect occurs when the inner electron shells shield the outer electrons from the full attraction of the nucleus. This causes the outer electrons to experience a reduced effective nuclear charge. When transitioning from periods 5 to 6, the electron shielding in the 4f subshell increases.

Relate the Effective Nuclear Charge and Electron Shielding to Atomic Radii

As the effective nuclear charge increases, the attraction between the nucleus and the outer electron shell increases, causing the atomic radius to decrease. However, an increase in electron shielding counteracts this effect by reducing the effective nuclear charge, causing the atomic radius to remain fairly constant.

Conclude with the Reason for Nearly Identical Radii

In the case of transition metals in Periods 5 and 6, the increase in effective nuclear charge due to the higher number of protons is balanced by the increased electron shielding effect from the additional 4f electrons. As a result, the effective nuclear charge experienced by the outermost electrons remains almost constant, leading to nearly identical atomic radii for these elements in each group.

Key concepts

Electron Configuration

Understanding electron configuration is essential when studying transition metals. These metals are part of the d-block in the periodic table and have electrons filling the d-subshell. Electron configuration is often expressed in the notation:

• [Noble Gas] \(n-1\) d \(x\) n s \(y\)
• Here, 'n' signifies the principal quantum number corresponding to the element's period.
• 'x' and 'y' represent the number of electrons in the d and s orbitals, respectively.

For instance, when moving across a period, the number of electrons in the d orbital increases. In transition metals, it's this d-subshell that plays a pivotal role in their chemical behavior.
As electrons occupy these d-orbitals, they interact with each other and with other parts of the atom, influencing properties like color, magnetism, and atomic size.

Effective Nuclear Charge

Effective nuclear charge (often abbreviated as Z extsubscript{eff}) is the "apparent" positive charge felt by an electron in an atom. Unlike the full nuclear charge, this value takes into account shielding effects from other electrons. This concept is pivotal in understanding various atomic properties:

• Z extsubscript{eff} increases across a period because the number of protons in the nucleus increases.
• It influences the attraction between electrons and the nucleus, affecting parameters like ionization energy and atomic radius.

In periods 5 and 6 of the transition metals, although atomic number (and thus nuclear charge) increases, the effect on size due to Z extsubscript{eff} is more complex due to other factors such as electron shielding.

Electron Shielding Effect

The electron shielding effect is crucial for understanding atomic behavior. Inner-shell electrons effectively "shield" outer electrons from the full charge of the nucleus. This phenomenon influences how electrons contribute to the atom's effective nuclear charge.

• As electrons are added in inner shells, like the 4f subshell in transition metals, they provide substantial shielding.
• This increased shielding can diminish the Z extsubscript{eff} felt by the outermost electrons.

In the transition between Periods 5 and 6, the electron shielding from the 4f subshell becomes significant. Though additional protons increase the atomic number, the 4f electrons counterbalance some of this change, stabilizing the effective nuclear charge experienced by outer electrons.

Atomic Radii

Atomic radii refer to the size of an atom, often measured by the distance from the nucleus to the boundary of the surrounding cloud of electrons. Several factors affect atomic radii:

• The effective nuclear charge directly influences an atom's radius—higher Z extsubscript{eff} tends to pull electrons closer, reducing the radius.
• Electron shielding can counteract changes in the radius by providing a repulsive force against the nucleus's pull.

For transition metals in Periods 5 and 6, the atomic radius remains fairly stable within each group despite increased atomic numbers. This is because the balance between a strong nuclear charge and increased electron shielding, especially from the 4f subshell, results in comparable effective nuclear charges. Thus, the atomic radii in these periods do not vary much.