Question

Explain why atomic radius decreases as you move to the right across a period for main-group elements but not for transition elements.

Short answer

Atomic radius decreases across a period for main-group elements due to increasing effective nuclear charge that pulls electrons closer. For transition elements, despite an increase in nuclear charge, the screening effect from inner d electrons is stronger, leading to a less pronounced decrease in atomic radius.

Step-by-step solution

Understand Periodic Trends

In the periodic table, atomic size or atomic radius tends to decrease as we move from left to right across a period. This is due to the increasing number of protons in the nucleus, which leads to a greater effective nuclear charge. As the effective nuclear charge increases, the electrons are pulled closer to the nucleus, reducing the size of the atom.

Main-group vs Transition Elements

For main-group elements (s and p blocks), the electrons being added are to the same principal energy level, so the increase in effective nuclear charge has a more pronounced effect in shrinking the atomic radius. However, for transition elements (d block), electrons are added to the penultimate energy level (not the outermost), and the increase in nuclear charge is often screened by the additional electrons in the d orbital, leading to a less pronounced decrease in atomic radius.

Screening Effect

With transition metals, the inner d electrons do not shield the outermost electrons from the nucleus effectively. Despite this poor shielding, the addition of d electrons somewhat counteracts the increase in nuclear charge. As a result, the atomic radius does not decrease as significantly across the transition elements as it does across the main-group elements.

Key concepts

Periodic Trends

The periodic table showcases a recurring trend known as periodic trends, which are patterns seen in the properties of elements when organized by increasing atomic number. One such trend is the atomic radius—a measure of the size of an element's atoms.

As you move left to right across a period, generally, the atomic radius decreases. This occurs because with each step to the right, an additional proton is added to the nucleus, creating a stronger pull—known as effective nuclear charge—on the electrons in the same energy level. This increasing pull causes the electrons to orbit closer to the nucleus, and hence, the atom's radius shrinks.

Effective Nuclear Charge

Imagine the nucleus of an atom holding onto its electrons like a magnet. The strength of this magnetic pull is referred to as the effective nuclear charge (ENC). ENC is essentially the net positive charge experienced by an electron in a multi-electron atom.

The ENC increases from left to right across a period thereby pulling electrons closer and reducing the atomic radius. However, ENC is not simply the total number of protons in the nucleus; it's also affected by factors such as the screening effect of inner-shell electrons, which can shield outer electrons from the full charge of the nucleus.

Main-group Elements

Main-group elements are found in the s and p blocks on the periodic table and involve the addition of electrons to the outermost energy level as we proceed across a period.

For these elements, as new electrons are added, they enter into the same principal energy level while the nucleus gains more protons. Thus, a greater ENC is felt by the electrons in the outermost energy level, causing the atomic radius to shrink more evidently in main-group elements than in transition elements when moving across a period.

Transition Elements

The situation with transition elements, which occupy the d block of the periodic table, differs from the main-group elements. Here, electrons are added to an inner (penultimate) d subshell, while the outer s shell remains filled with just two electrons for a large portion of the series.

This phenomenon leads to an interesting behavior where despite the increasing nuclear charge, the rate at which the atomic radius decreases is slower due to the poor shielding of the d electrons. Thus, this interplay between increasing nuclear charge and a less effective screening effect explains why transition elements don't follow the typical trend of shrinking atomic radii across a period.

Screening Effect

The screening effect, also known as the shielding effect, is a pivotal factor affecting the atomic radius. It describes the protection or 'screening' of outer electrons from the nucleus by inner-shell electrons.

In transition elements, the d electrons, which are found in the penultimate shell, are less effective at screening due to their shape and energy compared to the s and p electrons in the outer shell. Therefore, although there is an increase in the number of protons, the screening effect causes a less significant decrease in the atomic radius across the transition metals, as the full positive charge of the nucleus is not as effectively felt by the outermost electrons.