Problem 144 - A Complete Resource Book in Chemistry for JEE Main

Question

Lanthanoid contraction is caused due to (a) the appreciable shielding of outer electrons by \(4 \mathrm{f}\) electrons from the nuclear charge (b) the appreciable shielding on outer electrons by 5d-electrons from the nuclear charge (c) the same effective nuclear charge from \(\mathrm{Ce}\) to \(\mathrm{Lu}\) (d) the imperfect shielding of outer electrons by \(4 \mathrm{f}\) electrons from the nuclear charge

Step-by-Step Solution

Answer

The cause is the imperfect shielding of outer electrons by 4f electrons from the nuclear charge.

Step 1: Understand Lanthanoid Contraction

Lanthanoid contraction refers to the gradual decrease in atomic and ionic radii of the lanthanides with increasing atomic number, despite having the same general electronic configuration. This occurs as we move across the series from Ce to Lu.

Step 2: Examine the Shielding Effect

The outer electrons are affected by the shielding effect of the inner electrons. For lanthanides, the 4f electrons are poor at shielding the outer electrons from the nuclear charge due to their complex shape and penetration power.

Step 3: Analyze Each Option for Shielding Efficacy

Option (a): The 4f electrons shield the outer electrons well. Option (b): The 5d electrons do not significantly affect shielding in lanthanides. Option (c): Effective nuclear charge isn't constant throughout the series. Option (d): The 4f electrons shield imperfectly, leading to stronger nuclear attraction on outer electrons as atomic number increases.

Step 4: Identify Correct Option

Option (d) is correct because the imperfect shielding by 4f electrons allows the nuclear charge to be more effective in attracting outer electrons, leading to the lanthanoid contraction.

Key Concepts

Shielding Effect

The shielding effect is a phenomenon where inner shell electrons partially block the attractive force of the nucleus on the outer shell electrons. This essentially reduces the overall pull experienced by the outermost electrons due to the protons in the nucleus.

• Inner electrons act as a shield, decreasing the nucleus's direct impact on the valence electrons.
• The effectiveness of this shield depends on the orbitals' penetration power – how close to the nucleus these electrons can get.

Understanding the shielding effect is crucial because it influences an atom's effective nuclear charge – the net positive charge experienced by outer electrons. In the case of lanthanides, the 4f electrons, being buried deeper, do not shield effectively.
Since these electrons are poor at blocking, the shielding effect involved in lanthanoid contraction plays a pivotal role in the change in atomic size.

Lanthanides

Lanthanides are a series of elements in the periodic table, starting from cerium (Ce) and extending to lutetium (Lu). They are also known as rare earth elements because they typically occur together in natural sources.

• Their unique characteristics arise mostly from the filling of the 4f orbitals, which contain one to fourteen electrons in this series.
• Lanthanides share similar chemical properties due to their identical outer electron configuration, primarily differing in their 4f electron count.

Despite their similarity, the lanthanides demonstrate the lanthanoid contraction. This contraction is noteworthy because each subsequent element in the series is slightly smaller due to the imperfect shielding discussed in the earlier section. Understanding lanthanides helps in understanding the broader concept of how atomic structure impacts element properties.

Effective Nuclear Charge

The effective nuclear charge (\(Z_{eff}\)) is the net positive charge experienced by an electron in a multi-electron atom. This charge impacts how tightly an electron is bound to the nucleus and is calculated by considering both the protons in the nucleus and the shielding effect of the inner shell electrons.

• \(Z_{eff} = Z - S\), where \(Z\) is the atomic number and \(S\) is the shielding constant.
• The value of \(Z_{eff}\) increases across a period as electrons are added to the same shell, leading to less efficient shielding able to counteract the increasing nuclear charge.

In the lanthanides, this effect is crucial, as the 4f electrons don't shield effectively. Thus, the effective nuclear charge increases, pulling outer electrons more strongly towards the nucleus. This explains the reduction in atomic and ionic radii despite increasing atomic number.

Atomic and Ionic Radii

Atomic and ionic radii refer to the size of an atom or ion, determined by the space occupied by the electron cloud surrounding the nucleus.

• Atomic radius decreases across a period due to increasing effective nuclear charge, which pulls electrons closer to the nucleus.
• Ionic radius also varies, often depending on an element's electron configuration and how electrons are lost or gained during ion formation.

In lanthanides, the trend of atomic and ionic radii does not follow a straightforward increase as normally expected. Instead, due to lanthanoid contraction, there is a subtle but notable decrease in size across the series.
This results from the imperfect shielding by 4f electrons, which is compensated by an increase in effective nuclear charge, causing a net pull on the electron cloud and making the radii smaller. Understanding these variations is key to grasping how electron interactions within an atom dictate its physical dimensions.