Use shell notation to characterize rubidium's outermost electron.

Understanding the atomic number is fundamental in chemistry. It's like an atom's personal ID number, unique to each element, defined by the number of protons in an atom's nucleus. In a neutral atom, the atomic number also tells us how many electrons orbit around the nucleus because an atom in its neutral state has an equal number of protons and electrons.

For instance, rubidium (Rb) has an atomic number of 37, meaning it has 37 protons and, in a neutral state, 37 electrons. This atomic number not only identifies the element but also helps us determine its position in the periodic table and its chemical properties, including its electron configuration. Knowing the atomic number is the start point for many chemical calculations and understanding electron configurations.

Electrons reside in specific regions around an atom's nucleus known as energy levels or electron shells. These levels are like the floors in a hotel, with rooms (sub-shells) that can house electrons. The energy levels are organized by principal quantum numbers (\(n = 1, 2, 3, ...\)), which increase with distance from the nucleus. Moreover, each level has sub-levels designated as s, p, d, and f.

Each sub-shell has a maximum capacity of electrons it can hold: 2 for an s sub-shell, 6 for a p sub-shell, 10 for a d sub-shell, and 14 for an f sub-shell. These capacities are essential for building up an atom's electron configuration, as we place electrons in these sub-shells according to certain rules, such as the Aufbau principle. Understanding the arrangement of these shells and sub-shells is instrumental in figuring out how atoms will bond and interact with each other.

The Aufbau principle is a brilliant strategy that electrons use to organize themselves within an atom. It's like when you're packing a suitcase – you start with the essentials and then add in extra items where you find space. Similarly, in atomic terms, the Aufbau principle dictates that electrons fill lower-energy orbitals first before moving to higher-energy ones. This principle allows us to predict the electron configurations of atoms.

The rule gets its name from the German 'Aufbau,' meaning 'building up.' Following this principle, we fill the s sub-shell before p, then d, and finally f. This strategy guides us in writing the electron configuration of elements, as we've done for rubidium. As a result, we can determine the outermost electron and its probable chemical behavior, significantly aiding in the comprehension of atomic interactions and bonding patterns.