Understanding the arrangements of electrons in an atom is crucial for predicting chemical behaviour. Electron configurations provide this insight, illustrating the distribution of electrons among various atomic orbitals.
Defined Energy States and Electron Spin
- Every atomic orbital has a specific energy state depending on the electron configuration and the atom's chemical environment.
- Each orbital can hold a maximum of two electrons. Crucially, these electrons must have opposite spins. Representing electron spin, up-spin is typically symbolised as ↑ and down-spin as ↓.
Sublevels and Number of Orbitals
- Main energy levels are divided into sublevels: s, p, d, and f.
- Each sublevel contains a fixed number of orbitals. These orbitals are regions in space where there's a high likelihood of locating an electron.
- s sublevel – 1 orbital
- p sublevel – 3 orbitals
- d sublevel – 5 orbitals
- f sublevel – 7 orbitals
Image courtesy of trinset
Fundamental Principles of Electron Configurations
Aufbau Principle
- Electrons will fill the lowest energy orbitals first before moving to higher energy orbitals.
Image courtesy of Chemistry Learner.
Hund's Rule
- Within a sublevel, electrons prefer to occupy separate orbitals and will do so with parallel spins before pairing up in any single orbital.
Pauli Exclusion Principle
- No two electrons in an atom can have the same set of four quantum numbers. Essentially, this means an orbital can contain a maximum of two electrons, and they must have opposite spins.
Using these principles, you can deduce the electron configurations for atoms and ions up to Z = 36.
Image courtesy of sciencenotes.
Writing Electron Configurations
Full Electron Configurations
- This configuration details every electron in an atom.
- For instance, carbon (Z = 6) has a configuration of 1s² 2s² 2p².
Condensed Electron Configurations using Noble Gas Core
- Uses the previous noble gas's configuration followed by the electrons added beyond that noble gas.
- For chlorine (Z = 17), the condensed configuration is [Ne] 3s² 3p⁵.
Image courtesy of CK-12 Foundation (raster), Adrignola (vector)
Orbital Diagrams (Arrow-in-box diagrams)
- These diagrams visually represent electron placements within orbitals.
- Each box in the diagram stands for an orbital, with arrows representing electrons.
- For carbon (Z = 6):
- 1s: ↑↓
- 2s: ↑↓
- 2p: ↑ ↑
Exceptional Electron Configurations: Chromium and Copper
- While most atoms follow the rules set out above, a few, like chromium (Cr) and copper (Cu), are exceptions.
- Chromium (Z = 24)
- Expected: [Ar] 4s² 3d⁴
- Actual: [Ar] 4s¹ 3d⁵
- Half-filled d sublevel is more stable.
- Copper (Z = 29)
- Expected: [Ar] 4s² 3d⁹
- Actual: [Ar] 4s¹ 3d¹⁰
- Fully filled d sublevel is more stable.
FAQ
The number of orbitals in a sublevel is determined by the magnetic quantum number, which can have integer values between -l and +l, where l is the azimuthal quantum number associated with each sublevel. For the p sublevel, l=1, which gives magnetic quantum numbers of -1, 0, and 1. These three values correspond to the three p orbitals: px, py, and pz. For the d sublevel, l=2, yielding magnetic quantum numbers of -2, -1, 0, 1, and 2, corresponding to the five different d orbitals. The shape and orientation of these orbitals in space are distinct due to these magnetic quantum numbers.
Yes, there are instances, particularly among transition metals, where electrons from the s sublevel are removed before the d electrons during ionisation. This phenomenon can be attributed to the relatively close energy levels of the s and d orbitals in these elements. For example, in the ionisation of chromium, the 4s electrons are removed before the 3d electrons. When the 4s and 3d orbitals are both occupied, the 4s electrons are actually at a slightly higher energy than the 3d electrons, making them easier to remove. This illustrates the complexity and nuances of electron configurations and ionisation in transition metals.
Using the noble gas core in condensed electron configurations offers a shorthand method to represent electron configurations without writing out all the electrons. Noble gases have a full set of electrons in their outermost energy level, which makes them stable. By starting with the electron configuration of the closest noble gas with fewer electrons than the atom in question, chemists can save time and provide a clear picture of the valence electrons, which are most pertinent in chemical reactions. It simplifies the representation and allows for a more direct focus on the outermost, and chemically significant, electrons of the atom.
The Pauli Exclusion Principle states that no two electrons in an atom can have the same set of four quantum numbers, meaning two electrons in the same orbital must have opposite spins. On the other hand, Hund's Rule states that electrons will occupy degenerate (equal energy) orbitals singly before pairing up. When determining electron configurations, these two principles work hand-in-hand. For instance, when filling the p orbitals, an electron will go into each of the three p orbitals (px, py, pz) with parallel spins before any pairing occurs, abiding by Hund's Rule. However, when electrons do pair up in an orbital, they must have opposite spins, which is dictated by the Pauli Exclusion Principle.
The naming convention for orbitals, especially the d orbitals, can indeed be a bit perplexing. The primary reason for this lag is historical and stems from the way the energy levels of these orbitals were first observed and categorised. The 4s orbital is actually lower in energy than the 3d orbitals when they're both empty. Thus, the 4s fills before the 3d. However, once the 4s orbital starts filling, its energy increases, making the 3d orbitals more stable in comparison. The 3d name simply signifies that these orbitals are associated with the third main energy level (n=3), even though they fill during the fourth period of the periodic table.
Practice Questions
Electron configurations are determined by filling up the orbitals in order of increasing energy levels. According to the Aufbau principle, electrons will fill the orbitals starting from the lowest energy. In the case of bromine, after the 1s, 2s, 2p, 3s, and 3p orbitals have been filled, the 4s orbital is filled next. After 4s, the 3d orbitals have a lower energy than the 4p orbitals, which means they are filled before the 4p orbitals. The 3d orbitals can hold up to 10 electrons, and in bromine, they are fully occupied with 10 electrons. After the 3d orbitals, the 4p orbitals begin to fill, and in bromine, they contain 5 electrons.
The electron configuration for chromium is [Ar] 4s¹ 3d⁵. While one might expect it to be [Ar] 4s² 3d⁴ based on the Aufbau principle, chromium's configuration is an exception due to the increased stability offered by a half-filled d sublevel. The energy difference between the 4s and 3d orbitals is very small. By promoting one electron from the 4s orbital to the 3d orbital, chromium achieves a half-filled d sublevel (3d⁵). This configuration is more stable because electrons in the same sublevel with parallel spins tend to repel each other less. Therefore, the [Ar] 4s¹ 3d⁵ configuration provides a lower energy and greater stability for chromium.
