An electron is a subatomic particle with a negative elementary electric charge. It is a fundamental constituent of matter and has an intrinsic angular momentum (spin). Electrons are bound to an atomic nucleus by the electromagnetic force, while they are repelled from other electrons by the same force. This binding energy keeps atoms together and gives rise to the chemical properties of matter. The electron was first identified in 1897 as one of the components of cathode rays, and it was named after the Irish physicist George Johnstone Stoney who suggested that it be called an “electricon”.
The rest mass of an electron is 9.1093837015×10−31 kg, or about 1/1836 that of a proton. Its electric charge is −1.6021766208(98)×10−19 coulombs, which is approximately equal in magnitude to the elementary charge on a proton but has opposite sign; therefore, electrons are attracted to protons by their electrostatic force, but repel each other due to their identical charges. Like all elementary particles, electrons exhibit properties of both particles and waves: they can collide with other particles and can be diffracted like light. The wave properties of electrons are easier to observe with experiments than those of larger particles like atoms and molecules; for example, one famous experiment demonstrates that under certain circumstances electrons can behave like waves rather than like point-like particles—this phenomenon is known as electron diffraction.”
“Although scientists knew about electricity for centuries before they understood what it really was, they only began to investigate its nature in earnest in the late 1800s following William Crookes’ discovery (in 1879) that cathode rays could make objects glow.”
“In 1897 J. J. Thomson discovered that cathode rays were composed of previously unknown ‘negatively charged corpuscles’ which he later named ‘electrons’. In his famous plum pudding model he imagined them as being distributed throughout a sea of positively charged pudding.”
“It wasn’t until 1909 however that Robert Millikan conducted his famous oil drop experiment which measured the actual value of an electron’s charge for the first time (-1.59 x 10-19 C).”
“Today we know that electrons orbit around atomic nuclei much like planets orbit around stars except instead of being held in place by gravity they are held in place by electromagnetism -the attractive force between positive and negative charges.”
From this simple description you might think that atoms would fly apart since all their negatively charged electrons should repel each other -and you’d be right if it weren’t for two things: orbital angular momentum and Pauli’s exclusion principle.”
“Orbital angular momentum arises because just as Newton’s laws tell us planets will follow elliptical orbits so too will electrons follow elliptical orbits around nuclei (although these ellipses are much flatter than planetary ones).” This means there are regions close to the nucleus where there is a high probability finding an electron but as you move outwards from the nucleus this probability decreases until you reach what’s known as the Bohr radius -the average distance from nucleus at which an electron spends most its time.” At this radius there exists what’s known as an energy barrier because moving further away from the nucleus requires overcoming Coulombic electrostatic repulsion between protons in order for orbital stability to be maintained.” Consequently orbiting electrons can only exist stably at certain discrete distances from their nuclei -a bit like rungs on a ladder- with each rung corresponding to a different allowed energy state or shell.” As long as all these shells are filled then no net electrostatic forces act on any given atom meaning they won’t fall apart.” However if one or more shells aren’t completely full then electrostatic interactions will dominate over nuclear attraction leading to instability and possible loss or gain of outermost valence electrons via chemical reactions (ionic bonding etc.)”
“Pauli’s exclusion principle says no two identical fermions (particles with half integer spin such as 1/2 , 3/2 etc.) can occupy exactly same quantum state simultaneously i.e.. two identical fermions can never be found at same place doing same thing at same time!” This means within any given atom no two electrons can have exactly same set orbital parameters i..e.. n l m . Therefore neighbouring atoms must have slightly different values for at least one quantum number allowing themto avoid occupying exact same space -which would leadto their mutual annihilation! For example consider carbon atom which has 6 valence e⁻ arranged into 2s2 2p2 configuration…notice how next nearest neighbour nitrogen must have 5 valence e⁻ arranged into 2s2 2p3 configuration so even though both elements have complete p subshells available neitherof themcan fill these subshells identically thereby avoiding Pauliexclusion principle violation! If violation did occur then strong electromagnetic fields would quickly correct situation resulting lost energy visible light emission…”