who determined the mass to charge ratio of an electron
Mass-to-charge ratio In the 19th century, the mass-to-charge ratios of some ions were measured by electrochemical methods. In 1897, the mass-to-charge ratio of the electron was first measured by J. J. Thomson.[6] By doing this, he showed that the electron was in fact a particle with a mass and a charge, and that its mass-to-charge ratio was much smaller than that of the hydrogen ion H+. In 1898, Wilhelm Wien separated ions (canal rays) according to their mass-to-charge ratio with an ion optical device with superimposed electric and magnetic fields (Wien filter). In 1901 Walter Kaufman measured the increase of electromagnetic mass of fast electrons (Kaufmann–Bucherer–Neumann experiments), or relativistic mass increase in modern terms. In 1913, Thomson measured the mass-to-charge ratio of ions with an instrument he called a parabola spectrograph.[7] Today, an instrument that measures the mass-to-charge ratio of charged particles is called a mass spectrometer.
Prafulla Chandra Ray Acharya Sir Prafulla Chandra Ray also spelled Prafulla Chandra Rây CIE (Bengali: প্রফুল্ল চন্দ্র রায় Praphulla Chandra Rāy; 2 August 1861 – 16 June 1944)[1] was a Bengali chemist, educator and entrepreneur.[2]
Ä Ä (lower case ä) is a character that represents either a letter from several extended Latin alphabets, or the letter A with an umlaut mark or diaeresis.
Atom In 1924, Louis de Broglie proposed that all particles behave to an extent like waves. In 1926, Erwin Schrödinger used this idea to develop a mathematical model of the atom that described the electrons as three-dimensional waveforms rather than point particles. A consequence of using waveforms to describe particles is that it is mathematically impossible to obtain precise values for both the position and momentum of a particle at a given point in time; this became known as the uncertainty principle, formulated by Werner Heisenberg in 1926. In this concept, for a given accuracy in measuring a position one could only obtain a range of probable values for momentum, and vice versa.[22] This model was able to explain observations of atomic behavior that previous models could not, such as certain structural and spectral patterns of atoms larger than hydrogen. Thus, the planetary model of the atom was discarded in favor of one that described atomic orbital zones around the nucleus where a given electron is most likely to be observed.[23][24]