Electrons are fundamental subatomic particles with a negative electric charge and a very small rest mass compared with protons and neutrons. Their electric charge is exactly one elementary negative charge, while their mass is approximately 1 over 1836 of the proton mass. Electrons belong to the family of leptons and, according to the Standard Model of particle physics, are elementary particles rather than composites made from smaller known constituents.
They have intrinsic angular momentum called spin one half and possess a corresponding magnetic moment. Electrons are central to chemistry because their arrangement around atomic nuclei determines bonding, ion formation, spectroscopy and many physical properties of matter. In metals, electrons can move through a material and carry electrical current, while in semiconductors their behaviour underpins modern electronic devices.
The electron was identified experimentally in studies of cathode rays by J J Thomson in the late nineteenth century. Experiments showed that cathode rays consisted of particles with a consistent charge to mass ratio and that these particles were present regardless of the material used to produce the rays. This demonstrated that atoms were not indivisible entities but contained smaller charged constituents.
The discovery of the electron contributed to the development of early atomic models, including the Thomson model, Rutherford nuclear model, Bohr model and ultimately quantum mechanical descriptions. Modern physics no longer describes electrons as tiny classical spheres orbiting nuclei. Instead, electrons in atoms are described by quantum states and wavefunctions, with measurable probabilities for their positions and other properties.
Electrons have several properties that are fundamental to quantum physics. Their charge is negative, their spin is one half, and they obey Fermi Dirac statistics because they are fermions. The Pauli exclusion principle therefore prevents two identical electrons in the same atom from occupying exactly the same complete quantum state. Electrons also possess intrinsic magnetic moments associated with their spin and magnetic properties.
They can occupy different energy levels and orbitals in atoms, but an orbital should not be interpreted as a physical path. In quantum mechanics, an orbital is a mathematical description of a one electron spatial state or probability amplitude within an atom or molecule. The observable distribution of electron position is related to the squared magnitude of the relevant wavefunction.
Electrons are responsible for most familiar chemical bonding. In covalent bonds, electrons are shared between atoms in molecular orbitals or related bonding states. In ionic bonding, electrons are transferred sufficiently to produce oppositely charged ions whose electrostatic attraction contributes to the structure. In metallic bonding, electrons are delocalised across a lattice and can move through the material. Electron configurations also explain why elements have characteristic valencies and reactivities.
Electrons in inner shells generally remain strongly bound to their nuclei, whereas outer or valence electrons are more directly involved in chemical interactions. The energy required to remove an electron is described by ionisation energy, while electron affinity describes the energy change associated with adding an electron under defined conventions.
Electrons also interact with electromagnetic radiation. An electron in an atom can absorb or emit photons when transitions occur between allowed quantum states, provided the relevant selection rules are satisfied. These transitions generate characteristic atomic spectra and provide a powerful way to identify elements and investigate their electronic structures. In solids, electron energy levels broaden into bands, producing distinctions among conductors, semiconductors and insulators.
Electrons can also participate in photoelectric effects, electron emission, fluorescence processes, semiconductor carrier transport and many other phenomena. Their interaction with electromagnetic fields is fundamental to electricity, magnetism, optics and electronics. The behaviour of electrons is therefore not confined to chemistry; it is one of the central subjects connecting atomic physics, condensed matter physics, electrical engineering and materials science.
In an atom, electrons are bound to the positively charged nucleus by electromagnetic attraction. Their quantum mechanical energies are quantised, meaning that bound electrons can occupy specific allowed stationary states rather than an arbitrary continuous set of classical orbital energies. The precise states depend on the nuclear charge, electron electron interactions and other physical effects.
Hydrogen provides the simplest exact model because it contains one proton and one electron, whereas multi electron atoms require more sophisticated approximations because electrons repel one another and their motions are correlated. Electron configurations are often represented using principal shells and subshells such as 1s, 2s, 2p, 3s and 3p. These labels describe quantum states rather than classical trajectories.
Electrons can exist as bound particles in atoms and molecules or as mobile particles in many materials. In metals, some electrons occupy delocalised states extending through the solid and can respond to electric fields, producing electrical conductivity. In semiconductors, electrons and positively behaving holes can be controlled by doping, electric fields, light and other means, forming the basis of transistors, diodes, solar cells, integrated circuits and many sensors.
In vacuum systems, beams of electrons can be accelerated and manipulated using electric and magnetic fields. Electron microscopes exploit the short de Broglie wavelengths associated with energetic electrons to resolve structures much smaller than those readily accessible with ordinary visible light microscopy. These applications arise from the combination of electron charge, mass, wave properties and quantum statistics.
A common misconception is that electrons simply orbit atomic nuclei like planets orbiting the Sun. This classical picture is useful only as a very rough historical analogy and becomes misleading when applied literally. Quantum mechanics describes an electron in an atom using a quantum state, with position represented probabilistically rather than by a definite classical trajectory. The electron is not normally sitting at one fixed point inside an orbital, nor is an orbital a solid shell surrounding the nucleus.
Electron behaviour can involve superposition, interference, tunnelling and entanglement, and these phenomena become essential in appropriate experiments. The electron should therefore be understood as a quantum particle whose measurable properties are governed by quantum theory. Its behaviour simultaneously exhibits particle and wave characteristics without being reducible to either classical category.
