Electron configuration describes how an atom’s electrons are distributed among the various energy levels and subshells surrounding the nucleus. Understanding this arrangement is fundamental because it directly dictates an element’s chemical behavior, including how it forms bonds and reacts with other substances. Tin (Sn) is a metal with an atomic number of 50, meaning a neutral atom contains 50 protons and 50 electrons. Determining the specific configuration for these 50 electrons requires following a precise set of rules governing electron placement.
Foundational Principles of Electron Arrangement
The distribution of electrons is governed by the quantum mechanical model, which describes the space around the nucleus in terms of orbitals grouped into energy levels. These orbitals have distinct shapes and capacities, designated by the letters s, p, d, and f. An s-orbital holds two electrons, a p-subshell holds six electrons (three orbitals), a d-subshell holds ten electrons (five orbitals), and an f-subshell holds up to fourteen electrons (seven orbitals).
The Aufbau Principle
The Aufbau principle, derived from the German word for “building up,” dictates that electrons must fill the lowest energy orbitals first before occupying higher-energy ones. This sequential filling creates a predictable order for electron placement, starting with 1s, then moving to 2s, 2p, 3s, and so forth.
The Pauli Exclusion Principle
The Pauli exclusion principle imposes a limit on how many electrons can occupy any single orbital. It states that no two electrons in an atom can share the exact same set of quantum numbers, which practically limits each orbital to a maximum of two electrons. Furthermore, if an orbital contains two electrons, they must have opposite spins.
Hund’s Rule
A third rule, known as Hund’s rule, applies when multiple orbitals of the same energy level, called degenerate orbitals, are being filled. This rule specifies that electrons will occupy all empty degenerate orbitals singly before any orbital is filled with a second, paired electron. This maximizes the number of unpaired electrons within the subshell, leading to a more stable electron arrangement.
Step by Step Configuration of Tin
To determine the electron configuration for Tin (Sn), we start with its 50 electrons and follow the energy-based filling sequence established by the Aufbau principle. The process begins by filling the lowest-energy 1s orbital (two electrons), followed by 2s (two), and the 2p subshell (six), completing the first two principal energy levels. Next, the 3s and 3p subshells are filled with two and six electrons, respectively.
Next, the sequence requires filling the 4s orbital with two electrons before moving to the 3d subshell, which accommodates a full complement of ten electrons. The 4p subshell is then filled with six electrons, bringing the total count to 36 electrons, which is the configuration of the noble gas Krypton. At this point, we have filled the orbitals: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6\).
The remaining 14 electrons must be placed into the higher-energy orbitals, starting with the 5s orbital, which takes two electrons. Following the 5s, the 4d subshell is filled completely with ten electrons. This accounts for 48 of the 50 total electrons in the Tin atom, leaving two electrons left to place.
These final two electrons are placed into the 5p subshell, which can hold up to six electrons but only needs to accommodate the last two electrons for Tin. The filling order of the 5p orbitals would follow Hund’s rule, meaning the two electrons would occupy separate 5p orbitals before pairing up. The complete configuration is the result of this precise placement.
The Final Configuration and Shorthand Notation
The full electron configuration for a neutral Tin atom is written out as \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^2\). This notation confirms that the electrons fully complete the 5s and 4d subshells, and partially fill the 5p subshell with two electrons.
For practical purposes, chemists often use a condensed or shorthand notation to represent the configuration of heavier elements like Tin. This notation simplifies the expression by replacing the configuration of the inner, completed electron shells with the symbol of the preceding noble gas. The noble gas immediately before Tin is Krypton (Kr), which has an atomic number of 36 and the configuration \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6\).
The shorthand notation substitutes the first 36 electrons’ configuration with the bracketed symbol for Krypton, resulting in the concise expression: \([\text{Kr}] 5s^2 4d^{10} 5p^2\).
The Significance of Tin’s Outer Electrons
The chemical behavior of Tin is determined by its valence electrons, which are located in the outermost principal energy level. For Tin, these are the two electrons in the 5s subshell and the two in the 5p subshell, totaling four valence electrons (\(5s^2 5p^2\)). Having four valence electrons places Tin in Group 14 of the periodic table.
These four outer electrons enable Tin to exhibit two common oxidation states: \(+4\) and \(+2\). The \(+4\) state occurs when all four valence electrons (both \(5s\) and \(5p\)) are involved in chemical bonding. The \(+2\) state is explained by a phenomenon called the inert pair effect.
The inert pair effect describes the reluctance of the two \(5s\) electrons to participate in chemical reactions, especially in heavier p-block elements like Tin. Because the \(5s\) electrons are poorly shielded by the inner \(4d\) electrons, they are held more tightly by the nucleus. This stronger attraction means that only the two \(5p\) electrons are readily available for bonding, leading to the stable \(+2\) oxidation state.