Potassium, symbolized as K, is the nineteenth element on the periodic table, meaning a neutral atom contains nineteen protons in its nucleus and nineteen electrons surrounding it. Determining the arrangement of these nineteen electrons is the first step toward understanding the element’s chemical behavior. The specific count of unpaired electrons is a direct consequence of fundamental rules that govern electron placement within atomic orbitals.
Understanding Electron Pairing Rules
The distribution of electrons within an atom’s energy levels and orbitals is governed by several established principles of quantum mechanics. An orbital is a specific region of space around the nucleus where an electron is most likely to be found, with different types designated as \(s\), \(p\), \(d\), and \(f\). Each orbital, regardless of its type, has a maximum capacity for two electrons. This capacity limit is established by the Pauli Exclusion Principle, which dictates that when two electrons occupy the same orbital, they must possess opposite spins. Electrons that exist in the same orbital with opposite spins are considered “paired.”
Orbitals of the same type and energy level are known as degenerate orbitals. When filling these sets of equal-energy orbitals, Hund’s Rule applies, stating that electrons will occupy each orbital singly before any orbital receives a second, pairing electron. This maximizes the number of electrons with parallel spins. An “unpaired electron” is, therefore, any electron that occupies an orbital by itself, without a second electron of opposite spin present.
Mapping Potassium’s Electron Configuration
To find the unpaired electrons in potassium, its nineteen electrons must be placed into the appropriate orbitals following the energy-filling sequence, often called the Aufbau principle. The lowest energy orbital, \(1s\), is filled first with two electrons (\(1s^2\)). Moving up in energy, the second shell is filled next, accommodating two electrons in the \(2s\) orbital (\(2s^2\)) and six electrons across the three \(2p\) orbitals (\(2p^6\)).
The third shell begins to fill with two electrons in the \(3s\) orbital (\(3s^2\)) and six electrons in the three \(3p\) orbitals (\(3p^6\)). These initial three principal quantum shells account for a total of eighteen electrons, all of which are paired. The final, nineteenth electron must occupy the next available energy level, which is the \(4s\) orbital, not the \(3d\) orbital, according to the energy sequence. The full electron configuration for a neutral potassium atom is \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^1\). The final electron is isolated in the \(4s\) orbital, making it the single unpaired electron.
The Final Count and Chemical Reactivity
Based on the electron configuration derived, the final count for the number of unpaired electrons in a neutral potassium atom is one. This single electron resides alone in the \(4s\) orbital, which represents the outermost principal quantum shell, or valence shell, of the atom. The presence of just one valence electron fundamentally dictates the chemical properties of potassium.
Potassium is classified as an alkali metal, a group characterized by having a single electron in their outermost shell. This lone electron is relatively far from the positively charged nucleus and is shielded by the eighteen inner-shell electrons. The resulting weak attraction from the nucleus means the electron is easily removed.
Potassium exhibits high metallic reactivity because it can readily lose this single \(4s\) electron to achieve a stable electron configuration identical to the noble gas argon (\(1s^2 2s^2 2p^6 3s^2 3p^6\)). Losing the unpaired electron requires a relatively small amount of energy, known as a low ionization energy. This tendency to lose an electron means potassium almost always forms a positively charged ion, \(\text{K}^+\), when it participates in chemical reactions.