Problem 13 - Chemical Principles: The Quest for Insight

Question

2.13 For cach of the following ground-state atoms, predict the type of orbital \((1 s, 2 p, 3 d, 4 f\), and so on) from which an electron will need to be removed to form the \(+1\) ions: (a) \(\mathrm{Zn}\); (b) \(\mathrm{Cl}\); (c) \(\mathrm{Al}\); (d) \(\mathrm{Cu}\).

Step-by-Step Solution

Answer

Electrons are removed from: (a) Zn - 4s orbital, (b) Cl - 3p orbital, (c) Al - 3p orbital, (d) Cu - 4s orbital.

Step 1: Understanding Electron Configuration Principles

Recognize that the electron removed to form a +1 ion is generally the one that was added last according to the Aufbau principle. Electrons fill orbitals starting with the lowest energy level and work their way up, following the order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, etc.

Step 2: Identify Electron Configuration of Zn

Write the ground-state electron configuration for Zn (Zinc), which has an atomic number of 30. The configuration is 1s2 2s2 2p6 3s2 3p6 4s2 3d10. Since 4s is filled after 3d, the electron removed will be from the 4s orbital.

Step 3: Identify Electron Configuration of Cl

Write the ground-state electron configuration for Cl (Chlorine), which has an atomic number of 17. The configuration is 1s2 2s2 2p6 3s2 3p5. The electron removed will be from the 3p orbital, as it is the highest energy electron.

Step 4: Identify Electron Configuration of Al

Write the ground-state electron configuration for Al (Aluminum), which has an atomic number of 13. The configuration is 1s2 2s2 2p6 3s2 3p1. The electron removed will be from the 3p orbital.

Step 5: Identify Electron Configuration of Cu

Write the ground-state electron configuration for Cu (Copper), which has an atomic number of 29. The configuration is 1s2 2s2 2p6 3s2 3p6 4s1 3d10. Despite the 3d orbital being of higher energy than 4s, the 4s electron will be removed as copper has a special stability due to a completely filled d subshell.

Key Concepts

Aufbau Principle

The Aufbau principle serves as a guide for determining the electron configuration of atoms, ensuring that electrons populate orbitals in order of increasing energy.
This principle gets its name from the German word 'Aufbau', meaning 'building up' or 'construction'. It reflects the process of building up the electron structure of an element. According to this principle, electrons first fill the lowest available energy levels before moving to higher levels.
For instance, the 1s orbital is filled before the 2s orbital, adhering to the energy sequence of 1s, 2s, 2p, 3s, 3p, and so forth. A quick tip for students is to remember the order of the orbitals, which can often be visualized using a diagonal rule diagram or an electron configuration chart.
However, there are exceptions, particularly when dealing with transition metals where electrons in 4s are removed before 3d, even though 4s is filled before 3d, as shown in the example of Copper (Cu). This specific behavior is due to electron-electron interactions and increased stability from a full or half-full d orbital.

Orbital Types

Introduction to Orbitals

The various types of orbitals—s, p, d, and f—signify regions in an atom where electrons are likely to be found. They each have unique shapes, with the s being spherical and the p shaped like dumbbells. For d and f orbitals, the shapes become more complex.

Filling Up the Orbitals

Each type of orbital has a different capacity for holding electrons: s orbitals can hold 2, p orbitals can hold 6, d orbitals have room for 10, and f orbitals can accommodate 14 electrons. When predicting which orbital an electron will be removed from to form a positive ion, it's essential to look at the highest energy electrons within the configuration.
In atomic structures, the s orbital fills first, followed by p, then d, and finally f. This sequence is crucial when learning about the electronic structure of elements and for understanding the removal of electrons during ion formation.

Formation of Ions

When atoms form ions, they either gain or lose electrons to achieve a more stable arrangement, often resembling that of noble gases. In the context of the discussed exercise, creating a +1 ion involves the loss of one electron.
Electrons are generally removed from the outermost shell—an ion's valence shell—taking away those with the highest energy first. The element’s position in the periodic table often influences this due to the differentiating energy levels of s, p, d, and f orbitals.
Understanding the concept of ion formation is vital in chemistry, particularly when balanced equations, the prediction of product charges in reactions, or the properties of elements are being considered. Comprehending how ions form from neutral atoms can be a stepping stone to mastery in chemistry, as it applies to various topics like ionic bonding, electrical conductivity, and electrochemistry.