Undergraduate
  1. General chemistry  1.1. Introduction to Chemistry  1.2. Atomic Structure  1.2.1. Subatomic particles  1.2.2. Atomic Model  1.2.3. Quantum numbers  1.2.4. Electron Configuration  1.2.5. Periodic trends  1.2.6. Isotopes and their applications  1.3. Chemical bond  1.3.1. Ionic bond  1.3.2. Covalent bonds  1.3.3. Metal Bond  1.3.4. Hybridization  1.3.5. Molecular geometry and VSEPR theory  1.3.6. Bond polarity and dipole moment  1.4. Stoichiometry  1.4.1. Mole concept  1.4.2. Empirical and molecular formula  1.4.3. Limiting reactant and excess reactant  1.4.4. Percent Yield and Purity  1.4.5. Dilution calculation  1.5. States of matter  1.5.1. Gas Laws  1.5.2. Matter and Intermolecular Forces  1.5.3. Solid and Crystal Structures  1.5.4. Phase diagram  1.5.5. Plasma and supercritical fluid  1.6. Chemical reactions  1.6.1. Types of Chemical Reactions  1.6.2. Balancing Chemical Equations  1.6.3. Thermochemical reactions  1.6.4. Reaction energy profiles  1.7. Solutions and Mixtures  1.7.1. Concentration units  1.7.2. Syndrome properties  1.7.3. Solubility and Solubility Rules  1.7.4. Henry's Law and Raoult's Law  1.7.5. Partition coefficient  1.8. Chemical equilibrium  1.8.1. Law of mass action  1.8.2. Le Chatelier's principle  1.8.3. Reaction quotient  1.8.4. To use this online calculator for Equilibrium Constant, enter Equilibrium Constant (Eq.) and hit the calculate button. Here is how the Equilibrium Constant calculation can be explained with given input values -> 0.05 = 0.05/0.  1.8.5. Acid-base balance  1.9. Acids and bases  1.9.1. Arrhenius, Brønsted–Lowry, and Lewis definitions  1.9.2. pH and pOH  1.9.3. Acid-base titration  1.9.4. Buffer Solution  1.9.5. Acid-base indicator  1.9.6. Polyprotic Acids and Bases  1.10. Electrochemistry  1.10.1. redox reactions  1.10.2. Galvanic and Electrolytic Cells  1.10.3. Nernst equation  1.10.4. Electrolysis  1.10.5. Faraday's laws of electrolysis  1.11. Thermodynamics  1.11.1. Laws of Thermodynamics  1.11.2. Enthalpy, entropy and Gibbs free energy  1.11.3. Spontaneity of reactions  1.11.4. Thermodynamic cycles (Hess's law, Carnot cycle)  1.12. Kinetics  1.12.1. Rate law and reaction order  1.12.2. Activation Energy and Catalyst  1.12.3. Collision theory  1.12.4. Reaction mechanism  1.12.5. Transition state theory
  2. Organic chemistry  2.1. Structure and relationships  2.1.1. Hybridisation in Carbon Compounds  2.1.2. Resonance and aromatization  2.1.3. Molecular orbitals  2.1.4. Inductive and mesomeric effects  2.2. Hydrocarbons  2.2.1. Hydrocarbons  2.2.2. Alkene  2.2.3. Alkynes  2.2.4. Aromatic compounds  2.2.5. Cycloalkanes and Conformation  2.3. Functional Group  2.3.1. Alcohols and phenols  2.3.2. Ethers and epoxides  2.3.3. Aldehyde and ketone  2.3.4. Carboxylic acids and derivatives  2.3.5. Amin  2.3.6. Esters and amides  2.3.7. Thiols and sulfides  2.4. Stereoscopic  2.4.1. Isomerism in Stereochemistry  2.4.2. Chirality and optical activity  2.4.3. Structural analysis  2.4.4. Diastereomers and Enantiomers  2.5.1. Substitution reactions  2.5.2. Elimination reactions  2.5.3. Addition reactions  2.5.4. Radical reactions  2.5.5. Rearrangement reactions  2.5.6. Pericyclic Reactions  2.6. Spectroscopy and structural analysis  2.6.1. Infrared Spectroscopy  2.6.2. Nuclear magnetic resonance spectroscopy  2.6.3. Mass Spectrometry  2.6.4. UV-visible spectroscopy  2.6.5. X-ray crystallography  2.7. Polymer chemistry  2.7.1. Addition and Condensation Polymers  2.7.2. Polymerization mechanism  2.7.3. Biodegradable Polymer  2.7.4. Conducting and smart polymers
  3. Inorganic chemistry  3.1. Coordination chemistry  3.1.1. Ligands and coordination compounds  3.1.2. Crystal field theory  3.1.3. Molecular Orbital Theory in Coordination Compounds  3.1.4. Jahn–Taylor distortion  3.2. Main Group Chemistry  3.2.1. Alkali and alkaline earth metals  3.2.2. Halogens and Noble Gases  3.2.3. Transition metals and their complexes  3.2.4. Lanthanides and Actinides  3.3. Solid state chemistry  3.3.1. Crystal lattices and unit cells  3.3.2. Defects in solids  3.3.3. Electrical and magnetic properties  3.3.4. Superconductivity  3.4. Bio-inorganic chemistry  3.4.1. Metal ions in biology  3.4.2. Enzymatic reactions with metals  3.4.3. Metal poisoning and detoxification  3.4.4. Metalloproteins and Metalloenzymes
  4. Physical chemistry  4.1. Quantum Chemistry  4.1.1. Wave–particle duality  4.1.2. Schrödinger Equation  4.1.3. Quantum Numbers and Orbitals  4.1.4. Atomic and molecular spectroscopy  4.2. Statistical mechanics  4.2.1. Maxwell–Boltzmann distribution  4.2.2. Partitioning functions  4.2.3. Bose–Einstein and Fermi–Dirac statistics  4.3. Chemical thermodynamics  4.3.1. Gibbs Free Energy  4.3.2. Phase transition  4.3.3. Thermodynamic efficiency  4.4. Surface chemistry  4.4.1. Absorption and Catalysis  4.4.2. Colloids and Emulsions
  5. Analytical Chemistry  5.1. Classical methods  5.1.1. Gravimetric Analysis  5.1.2. titrimeter  5.2. Instrumental methods  5.2.1. Chromatography  5.2.2. Spectroscopy  5.2.3. Electroanalytical methods  5.2.4. X-ray diffraction
  6. Industrial Chemistry  6.1. Petrochemistry  6.2. Chemical engineering principles  6.3. Green Chemistry  6.4. Pharmaceuticals and Drug Synthesis
  7. Biochemistry  7.1. Carbohydrates  7.2. Proteins and enzymes  7.3. Lipids and membranes  7.4. Nucleic acids  7.5. Metabolism and Bioenergetics  7.6. Enzyme kinetics and mechanism

Undergraduate → General chemistry → Atomic Structure ↓


Electron Configuration


Electron configuration is a fundamental concept in chemistry that describes the distribution of electrons in the electron shells and subshells of an atom. This topic is important for understanding the chemical properties of elements, their reactivity, and the structure of the periodic table.

Understanding electron configuration

To understand the concept of electron configuration, it is necessary to understand the structure of the atom. An atom consists of a nucleus containing protons and neutrons, with electrons orbiting outside the nucleus in electron clouds or shells. These electrons are arranged in a way to minimize their energy levels, resulting in specific patterns.

Basic principles

Aufbau principle

According to the Aufbau principle electrons occupy the lowest energy level available. The order of filling the subshells is determined by their increasing energy levels. The electrons will fill the orbitals starting from the lowest energy level to the highest energy level.

Pauli exclusion principle

The Pauli exclusion principle dictates that no two electrons in an atom can have the same set of four quantum numbers. In short, an orbital can hold a maximum of two electrons with opposite spins.

Hund's law

According to Hund's rule, electrons will fill degenerate orbitals - those of equal energy - first singly, then pairing up. This reduces electron-electron repulsion within the orbitals, leading to a more stable configuration.

Electron shells and subshells

Electrons are distributed around the nucleus in specific energy levels or shells, which can be divided into subshells called s, p, d, and f. Each subshell can hold a certain number of electrons:

  • s subshell: 2 electrons
  • p subshell: 6 electrons
  • d subshell: 10 electrons
  • f subshell: 14 electrons

Order of filling of electron subshells

The order of filling can be remembered using the periodic table or the diagonal rule, often expressed through an electron configuration chart. Here's a simplified version:

1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p

Let's consider a visual example of what the electron configuration might be like for some elements.

Examples of electron configurations

Hydrogen:

1s1

Hydrogen is the simplest element with one electron. That electron resides in the first shell and s subshell.

Helium:

1s2

Helium has two electrons, which completely fill the first shell in 1s subshell.

Carbon:

1s2 2s2 2p2

Carbon has six electrons, the first two of which fill 1s subshell, the next two fill 2s subshell, and the last two enter 2p subshell.

Neon:

1s2 2s2 2p6

Neon's second shell is filled, making it a very stable element due to its full outer electron shell, similar to the other noble gases.

Electron configuration notation

In chemical notation, electron configuration is represented by stating the number of electrons in each shell. For example, as mentioned, the configuration of neon can be described as 1s2 2s2 2p6 Alternatively, it can also be represented using noble gas notation:

[He] 2s2 2p6

The symbol [He] in brackets indicates that neon has the same inner shell configuration as helium, and the rest of the text describes the specific outer shell configuration.

Extended configuration and exceptions

While Aufbau theory generally accurately predicts the order of filling subshells, exceptions can occur, especially in the transition metals. Common exceptions include elements where the electron configurations will end with d4 or d9 In these cases, electrons from s orbitals can be promoted to d orbitals to obtain half-filled or completely filled d subshells, which are more stable.

Chromium:

Expected: 1s2 2s2 2p6 3s2 3p6 4s2 3d4 Actual: 1s2 2s2 2p6 3s2 3p6 4s1 3d5

Here, one electron from 4s is used to fill d orbital completely or at least halfway, providing greater stability.

Periodic table and electron configuration

Understanding electron configuration can help explain many aspects of the periodic table, such as why elements in the same group have similar chemical properties. Elements in the same group of the periodic table have the same number of electrons in their outermost energy level, or valence shell, which largely determines their chemical behavior.

Visual representation of electron configuration

Let us consider a diagram that shows the electron orbitals and their filling.

1s 2s 2P 2P 2P 3s

In this illustration, the different rectangles represent orbitals within subshells, and provide a diagrammatic representation of the positions of the electrons.

Important concepts related to electron configuration

Electron configurations play an important role in various chemical phenomena and applications.

Valence electrons

The electrons present in the outermost shell of an atom are known as valence electrons. These electrons are important in determining how elements interact and form chemical bonds.

Consider sodium (Na) with the configuration:

1s2 2s2 2p6 3s1

3s subshell has a single valence electron, which sodium can lose to attain a stable noble gas configuration, making it highly reactive.

Ion formation

Electron configuration can predict ion formation in both metals and nonmetals, as they gain or lose electrons to achieve a complete outer shell.

Cation example: Magnesium (Mg)

Original: 1s2 2s2 2p6 3s2 Ion (Mg2+): 1s2 2s2 2p6

By losing two electrons, magnesium attains a stable noble gas configuration. Therefore, its normal oxidation state is +2.

Anion example: Chlorine (Cl)

Original: 1s2 2s2 2p6 3s2 3p5 Ion (Cl-): 1s2 2s2 2p6 3s2 3p6

By gaining one electron, chlorine acquires a full outer shell, normally seen with -1 charge.

Conclusion

Understanding electron configuration in chemistry is essential for explaining and predicting atomic behavior, bonding, and the formation of cations and anions. Although the principles may seem complex at first, they provide valuable information about the reactivity and properties of elements. With practice and study, the patterns of electron configuration can become an essential tool in understanding the beauty of chemical interactions and the organization of the periodic table.


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Electron Configuration

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