The electrons surrounding an atom’s nucleus exist within specific, defined regions of space referred to as electron shells or energy levels. These levels represent particular orbits where the electron possesses a fixed amount of energy. The concept of energy levels is fundamentally quantized, meaning an electron can only occupy one of these specific energy states and cannot exist in the space or energy value between them. When an electron absorbs energy, it jumps from a lower energy level to a higher one, entering an excited state. Conversely, when it falls back to a lower level, it releases the energy difference as a photon of light, a process that gives each element a unique spectral signature.
The Basics of Electron Shells
Scientists use the principal quantum number, designated by the letter $n$, to identify these main energy levels, starting with $n=1$ for the shell closest to the nucleus. These shells are also sometimes labeled alphabetically using capital letters, with $n=1$ being the K shell, $n=2$ the L shell, $n=3$ the M shell, and so on. As the principal quantum number $n$ increases, the electron shells are located progressively farther from the nucleus, and the electrons within them possess a higher overall energy.
Electrons fill these shells sequentially, beginning with the lowest energy level ($n=1$) and moving outward. Each shell has a theoretical maximum capacity for electrons, which is calculated using the formula $2n^2$. For example, the K shell ($n=1$) can hold a maximum of 2 electrons, while the L shell ($n=2$) can hold 8 electrons.
Mapping Energy Levels to Periodic Rows
The arrangement of the modern periodic table directly reflects how atoms fill these electron shells. The horizontal rows, known as periods, correspond precisely to the highest principal energy level ($n$) currently occupied by electrons in an element. Moving from left to right across any given period indicates the progressive filling of electrons into the outermost energy level.
For instance, all elements in Period 1 only contain electrons in the $n=1$ shell. The elements in Period 2 are building up their electrons in the $n=2$ shell. This pattern means that an atom of potassium, located in Period 4, has electrons occupying up to the fourth principal energy level ($n=4$).
Sublevels and the Periodic Table Blocks
Each principal energy level is further divided into one or more sublevels, which define the shapes of the electron clouds, or orbitals. These sublevels are labeled $s$, $p$, $d$, and $f$, and the number of sublevels available within a main shell is equal to the principal quantum number $n$. For example, the $n=1$ shell has only an $s$ sublevel, while the $n=2$ shell contains both $s$ and $p$ sublevels.
The existence and filling order of these sublevels account for the distinct rectangular blocks that shape the periodic table. The $s$-block elements are on the far left, filling their $s$ sublevels, which can hold a maximum of 2 electrons. The $p$-block elements are on the far right, filling their $p$ sublevels, which can hold up to 6 electrons. The $d$-block, containing the transition metals, is in the middle, filling $d$ sublevels (10 electrons), and the $f$-block, situated below the main body, is filling $f$ sublevels (14 electrons).
How Valence Shells Determine Chemical Behavior
The outermost electron shell of an atom is called the valence shell, and the electrons residing there are the valence electrons. These electrons are the farthest from the nucleus and are the least tightly bound, making them the primary participants in chemical reactions and bonding. The number of valence electrons an atom possesses dictates its chemical properties, including how readily it will form compounds and the type of bonds it will create.
For main-group elements, the number of valence electrons corresponds directly to the atom’s column, or group, number on the periodic table (Groups 1, 2, and 13–18). Atoms tend to react in ways that allow them to achieve the stable configuration of a full valence shell, which usually means eight electrons, following the octet rule. Elements with one or two valence electrons often lose them to form positively charged ions, while those with six or seven tend to gain electrons to form negatively charged ions, driving the formation of ionic bonds.