The structure of an atom is visualized using electron configuration, which describes how electrons are distributed in atomic orbitals. The orbital diagram is a graphical representation showing the specific orbitals and the electron arrangement within them. We will explore this arrangement for Calcium (Ca), an alkaline earth metal with an atomic number of 20. A neutral Calcium atom possesses 20 electrons that must be placed precisely.
Understanding Orbitals and Sublevels
Electrons exist in distinct energy levels, visualized as shells surrounding the nucleus, designated by the principal quantum number \(n\). Within each shell are sublevels (\(s, p, d\), and \(f\)) that define the shape of the region where the electron is likely to be found. These sublevels contain the actual orbitals, which are the specific three-dimensional spaces holding the electrons.
The \(s\) sublevel contains just one spherical orbital, which can hold a maximum of two electrons. The \(p\) sublevel is comprised of three dumbbell-shaped orbitals, giving it a total capacity of six electrons. Further out, the \(d\) sublevel holds five orbitals, accommodating up to ten electrons in total. The \(f\) sublevel consists of seven orbitals that can house fourteen electrons.
Fundamental Rules for Electron Placement
Placing electrons into these defined orbitals requires adherence to three fundamental principles that ensure the atom achieves its lowest energy state. The Aufbau principle dictates the order of filling, stating that electrons must first occupy the lowest energy orbitals available before they move into higher energy orbitals.
The Pauli Exclusion Principle governs the behavior of electrons within a single orbital. It specifies that any given orbital can contain no more than two electrons, and these two electrons must exhibit opposite spins. This opposing spin is represented by arrows pointing up and down in the orbital diagram, reflecting their distinct quantum states.
When multiple orbitals have the exact same energy, such as the three \(p\) orbitals, Hund’s rule must be applied. This principle states that electrons will occupy separate orbitals within that sublevel first, maximizing the number of unpaired electrons. Only after each equal-energy orbital contains one electron will a second electron be added to pair up the first one, which helps to minimize natural electron repulsion.
Step-by-Step Derivation for Calcium
We can now apply these principles to position the 20 electrons of a neutral Calcium atom into the orbital structure. The process begins by filling the lowest energy level, the \(1s\) orbital, with two electrons according to the Pauli Exclusion Principle (\(1s^2\)). Next, we move to the second shell, filling the \(2s\) orbital with two electrons (\(2s^2\)), followed by the three \(2p\) orbitals.
The \(2p\) sublevel accommodates six electrons, which are placed singly into each of the three orbitals before pairing up, following Hund’s rule. This completes the second shell with eight electrons. The third shell begins with the \(3s\) orbital, which takes two electrons (\(3s^2\)), and then the three \(3p\) orbitals, which accept six electrons (\(3p^6\)).
At this point, 18 electrons have been placed, giving Calcium the electron configuration of \(1s^2 2s^2 2p^6 3s^2 3p^6\). The next available orbital would logically seem to be the \(3d\) sublevel; however, the Aufbau principle dictates a different path. The \(4s\) orbital possesses a slightly lower energy level than the \(3d\) orbitals, meaning the last two electrons enter the \(4s\) orbital first.
The complete and final electron configuration for Calcium is therefore \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2\). The orbital diagram visually represents this by showing ten boxes: one each for \(1s, 2s, 3s\), and \(4s\); three for \(2p\); and three for \(3p\). Each of these ten orbitals contains a pair of electrons with opposite spins, resulting in a diagram where all 20 electrons are fully paired.
How Calcium’s Electron Structure Dictates Reactivity
The chemical behavior of Calcium is directly determined by its outermost electrons, known as the valence electrons. Calcium’s structure shows two electrons residing in its highest energy level, the \(4s\) orbital. This placement identifies Calcium as an Alkaline Earth Metal, positioning it in Group 2 of the periodic table.
Atoms tend to react in ways that allow them to achieve a more stable, lower-energy electron arrangement, often mimicking the configuration of a noble gas. By losing these two valence electrons from the \(4s\) orbital, Calcium achieves the stable, filled-shell configuration of the noble gas Argon.
This loss results in the formation of a positively charged ion, \(\text{Ca}^{2+}\), which has 18 electrons and a full outer shell. The tendency to easily donate these two electrons makes Calcium highly reactive, as metals are characterized by their low ionization energy.