Arrange these elements in order of increasing atomic size: \(\mathrm{Cs}, \mathrm{Sb}, \mathrm{S}, \mathrm{Pb}, \mathrm{Se}\).

The periodic table is a fundamental tool in chemistry, providing a systematic layout of all known chemical elements in order of increasing atomic number. It's arranged in rows (periods) and columns (groups), reflecting recurring patterns in properties. The key takeaway for students exploring atomic size trends is to understand the significance of an element's position on the table. Just by knowing the location, you can predict numerous characteristics, including atomic size.

Elements are aligned into groups based on similar electronic arrangements, which translates into similar chemical properties. As you glance across a period, each element has one more proton and is generally one electron heavier than its neighbor to the left. This additional positive charge pulls the electrons closer, affecting the atomic size.

Imagine atoms like tiny solar systems, with electron shells as orbits around the nucleus. An electron shell is essentially a group of atomic orbitals within a certain range of distances from the atom's nucleus. These shells are like layers of an onion, and each succeeding shell is further out from the nucleus.

Each shell can only hold a fixed number of electrons before additional electrons start to fill the next shell. As you move down the groups of the periodic table, new shells are added, making atoms larger because the outermost electrons are farther from the nucleus. However, the more electrons there are, the more they repel each other, leading to a larger overall size.

The atomic radius is a measurable concept which represents the size of an atom. It can be thought of as the distance from the nucleus to the boundary where the electron cloud is no longer influenced by the atom's positive charge. However, placing exact figures on atomic radii can be challenging due to the nebulous nature of electron clouds.

When comparing atomic radii, we often look at the trend of size increasing down a group as new shells are added. Conversely, the radius decreases across a period as more protons in the nucleus create a stronger attraction to electrons, pulling them in closer. This contraction of the electron cloud is why sulfur (S) is smaller than selenium (Se), for example.

Understanding trends in groups and periods is essential to predict an atom's relative size and many other chemical properties. Within a group (the vertical columns), atomic size increases from top to bottom. This is because each element down the group has an additional electron shell. On the other hand, within a period (the horizontal rows), atomic size decreases from left to right. As you navigate rightwards along a period, the nuclear charge substantially increases due to more protons being added, which attracts the electron clouds inward making the atoms smaller.

It's beneficial for students to relate the concept of atomic size trends with the visual layout of the periodic table and with the concept of electron shells. With this combined understanding, explaining why cesium (Cs) has a larger atomic radius than lead (Pb), despite being in the same period, becomes intuitive - it's due to cesium having fewer protons pulling on its electron cloud, so its electrons are less tightly held and the atom is larger.