Isotopes, sub-atomic particles and relative atomic mass

6 June 2017

Isotopes, sub-atomic particles and relative atomic mass Sub-atomic particles A subatomic particle Is a particle smaller than an atom: It may be elementary or composite. In 1905, Albert Einstein demonstrated the physical reality of the photons, hypothesized by Max Planck in 1900, in order to solve the problem of black body radiation In thermodynamics. In 1874, G. Johnstone Stoney postulated a rnlnlmum unit of electrical charge, for which he suggested the name electron in 1891. In 1897, J. J. Thomson confirmed Stoney’s conjecture by discovering the first subatomic particle, the electron (now denoted e-).

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Subsequent speculation about the structure of atoms was severely constrained by Ernest Rutherford’s 1907 gold foll experiment, showing that the atom is mainly empty space, with almost all its mass concentrated in a (relatively) tiny atomic nucleus. The development of the quantum theory led to the understanding of chemistry in terms of the arrangement of electrons in the mostly empty volume of atoms. Particle physics and nuclear physics concern themselves with the study of these particles, their interactions, and matter made up of them which do not aggregate Into atoms.

These particles include atomic constituents such s electrons, protons, and neutrons (protons and neutrons are actually composite particles, made up of quarks), as well as other particles such as photons and neutrinos which are produced copiously In the sun. However, most of the particles that have been discovered and studied are not encountered under normal earth conditions; they are produced in cosmic rays and during scattering processes in particle accelerators. Isotopes Isotopes are variants of a particular chemical element.

While all isotopes of a given element share the same number of protons, each isotope differs from the others In its number of neutrons. Radioactive isotopes The existence of isotopes was first suggested in 1912 by the radiochemist Frederick Soddy, based on studies of radioactive decay chains which indicated about 40 1 OF3 ditterent species described as radioelements (i. e. radioactive elements) between uranium and lead, although the periodic table only allowed for 1 1 elements from uranium to lead.

Several attempts to separate these new radioelements chemically had failed. For example, Soddy had shown in 1910 that mesothorium (later shown to be Ra-228), radium (Ra-226, the longest-lived isotope), and thorium X (Ra-224) are impossible to separate. Attempts to place the radioelements in the periodic table led Soddy and Kazimierz FaJans independently to propose their radioactive in 1913, to the effect that alpha decay produced an element two places to the left in the periodic table, while beta decay emission produced an element one place to the right.

Soddy recognized that emission of an alpha particle followed by two beta particles led to the formation of an element chemically identical to the initial element but with a mass four units lighter and with different radioactive properties. Soddy proposed that several types of atoms (differing in radioactive properties) could occupy the ame place in the table.

For example, the alpha-decay of uranium-235 forms thorium-231 , while the beta decay of actinium-230 forms thorium-230 The term “isotope”, Greek for “at the same place”, was suggested to Soddy by Margaret Todd, a Scottish physician and family friend, during a conversation in which he explained his ideas to her. In 1914 T. W. Richards found variations between the atomic weight of lead from different mineral sources, attributable to variations in isotopic composition due to different radioactive origins Stable isotopes The first evidence for isotopes of a stable (non-radioactive) element was found by J.

J. Thomson in 1913 as part of his exploration into the composition of canal (positive ions). Thomson channelled streams of neon ions through a magnetic and an electric field and measured their deflection by placing a photographic plate in their path. Each stream created a glowing patch on the plate at the point it struck. Thomson observed two separate patches of light on the photographic plate (see image), which suggested two different parabolas of deflection. Thomson eventually concluded that some of the atoms in the neon gas were of higher mass than the rest.

F. W. Aston subsequently discovered different stable isotopes for numerous elements sing a mass spectrograph. In 1919 Aston studied neon with sufficient resolution to show that the two isotopic masses are very close to the integers 20 and 22, and that neither is equal to the known molar mass (20. 2) of neon gas. This is an example of Aston’s whole number rule for isotopic masses, which states that large deviations of elemental molar masses from integers are primarily due to the fact that the element is a mixture of isotopes.

Aston similarly showed that the molar mass of chlorine (35. 45) is a weighted average of the almost integral masses for the two isotopes Cl-35 and Cl-37. Relative atomic mass An atomic weight (relative atomic mass) of an element from a specified source is the ratio of the average mass per atom of the element to 1/12 of the mass of an atom of 12C The relative masses of atoms are measured using an instrument called a mass spectrometer, invented by the English physicist Francis William Aston (1877-1945) when ne was working in Cambridge wit J J.

Thomson. It was in his use ot this instrument that the existence of isotopes of elements was discovered. Aston eventually discovered many of the naturally occurring isotopes of non-radioactive elements. He was awarded the Nobel Prize for Chemistry in 1922. Briefly, the mass pectrometer works by bombarding gaseous atoms with fast-moving electrons which knock out an electron from the atom. The cations formed are brought down on to a detector in turn according to their mass.

The instrument provides a measure of the relative mass (compared to 12C) and the relative number of each isotope. The diagrams below represent the mass spectrum of naturally occurring chlorine. The above right spectrum has been represented so that the most abundant isotope has a relative abundance of 100%, with the other mass peaks scaled in relation to this. The relative atomic mass of chlorine is now calculated as shown below:

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