Question

Choose the larger atom in each pair. $$ \text { a. Al or } \mathrm{In} \quad \text { b. Si or } \mathrm{N} \quad \text { c. P or Pb } \quad \text { d. } \operatorname{C or F} $$

Short answer

The larger atoms for each pair are: a. In, b. Si, c. Pb, d. C.

Step-by-step solution

Understanding Atomic Size

The size of an atom generally increases as we move down a group in the periodic table, due to the addition of electron shells. It decreases as we move from left to right across a period because of the increasing nuclear charge, pulling the electrons closer to the nucleus.

Comparison a: Al vs In

Aluminum (Al) and Indium (In) are in the same group of the periodic table, with Indium being lower. Therefore, Indium has more electron shells and is the larger atom.

Comparison b: Si vs N

Silicon (Si) and Nitrogen (N) are in the same period of the periodic table, with Silicon being to the left. Therefore, Silicon has fewer protons in its nucleus compared to elements on its right, thus exerting less pull on its electron shells, making it the larger atom.

Comparison c: P vs Pb

Phosphorus (P) and Lead (Pb) are found in the same group with Lead being lower in the periodic table. This means Lead has more electron shells than Phosphorus, making Lead the larger atom.

Comparison d: C vs F

Carbon (C) and Fluorine (F) are in the same period with Carbon being to the left. This positions Carbon as having less nuclear charge than Fluorine, resulting in a larger atomic size for Carbon.

Key concepts

Periodic Table Trends

The periodic table is a comprehensive map of the elements, which provides a way to understand the relationships and properties of atoms. One of the fundamental trends you'll encounter is how atomic size changes within this structured framework.
As you begin to explore the vertical columns, known as groups, you'll notice that the atomic size increases as you descend. This is because each row down adds a new layer of electron shells, which are like concentric circles around the nucleus, increasing an atom's volume.

Horizontally, when examining periods from left to right, the trend reverses. Atomic size decreases across a period. This seems counterintuitive at first since more protons and electrons are added. But the increased nuclear charge pulls the electron shells closer to the nucleus due to stronger electromagnetic attraction, reducing the atom's radius.

Electron Shells

Imagine electron shells as the housing for an atom's electrons, with each shell equivalent to a different energy level. The shells closest to the nucleus have the lowest energy and those furthest away have the highest. When comparing atomic sizes, these shells are critical because they effectively dictate the distance of the outermost electrons from the nucleus.

As we move down a group in the periodic table, additional electron shells are added. This results in each successive element in a group having a larger size due to the new electron shell. To visualize: each row in the periodic table corresponds to a new larger shell, like a larger orbit that takes you further from the atomic core.

Nuclear Charge

Central to understanding these trends is the concept of nuclear charge, the net positive charge experienced by electrons in the valence shell of an atom. The charge is determined by the number of protons in an atom's nucleus since each proton bears a positive charge.

As elements on a period gain protons, their nuclear charge increases. This increased pull brings the electron shells tighter and tighter around the nucleus, thus shrinking the atom despite the additional electrons. The result is a more compact, smaller atom as you look across to the right on the periodic table. It's this delicate balance of nuclear charge that governs the atomic sizes and leads us to predict, for example, that Carbon (with lower nuclear charge) is larger than Fluorine.