Chemical Element and Atomic Weights

The history of the periodic table reflects over a century of growth in the understanding of chemical properties, and culminates with the publication of the first actual periodic table by Dmitri Mendeleev in 1869. [1] While Mendeleev built upon earlier discoveries by such scientists as Antoine-Laurent de Lavoisier and Stanislao Cannizzaro, the Russian scientist is generally given sole credit for development of the actualperiodic table itself. The table itself is a visual representation of the periodic law which states that certain properties of elements repeat periodically when arranged by atomic number.

The table arranges elements into vertical columns (groups) and horizontal rows (periods) to display these commonalities. ————————————————- Elemental ideas from ancient times People have known about some chemical elements like gold, silver and copper from antiquity, as these can all be discovered in nature in native form and are relatively simple to mine with primitive tools. [2] However, the notion that there were a limited number of elements from which everything was composed originated with the Greek philosopher Aristotle.

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About 330 B. C Aristotle proposed that everything is made up of a mixture of one or more of four “roots” (originally put forth by the Sicilian philosopher Empedocles), but later renamed elements by Plato. The four elements were earth, water, air and fire. While the concept of an element was thus introduced, Aristotle’s and Plato’s ideas did nothing to advance the understanding of the nature of matter. ————————————————- Age of Enlightenment Hennig Brand was the first person recorded to have discovered a new element.

Brand was a bankrupt German merchant who was trying to discover the Philosopher’s Stone — a mythical object that was supposed to turn inexpensive base metals into gold. He experimented with distilling human urine until in 1649[3] he finally obtained a glowing white substance which he named phosphorus. He kept his discovery secret, until 1680 when Robert Boyle rediscovered it and it became public. This and related discoveries raised the question of what it means for a substance to be an “element”.

In 1661 Boyle defined an element as a substance that cannot be broken down into a simpler substance by a chemical reaction. This simple definition actually served for nearly 300 years (until the development of the notion of subatomic particles), and even today is taught in introductory chemistry classes. In 1817, Johann Wolfgang Dobereiner began to formulate one of the earliest attempts to classify the elements. He found that some elements formed groups of three with related properties. He termed these groups “triads”.

Some triads classified by Dobereiner are: 1. chlorine, bromine, and iodine 2. calcium, strontium, and barium 3. sulfur, selenium, and tellurium 4. lithium, sodium, and potassium In all of the triads, the atomic weight of the second element was almost exactly the average of the atomic weights of the first and third element. Dimitri Mendeleev, a Russian chemist, was the first scientist to make a periodic table much like the one we use today. Mendeleev arranged the elements in a table ordered byatomic weight, corresponding to relative molar mass as defined today.

It is sometimes said that he played “chemical solitaire” on long train rides using cards with various facts of known elements. [10] On March 6, 1869, a formal presentation was made to the Russian Chemical Society, entitled The Dependence Between the Properties of the Atomic Weights of the Elements. His table was published in an obscure Russian journal but quickly republished in a German journal,Zeitschrift fur Chemie (Eng. , “Chemistry Magazine”), in 1869. It stated: 1. The elements, if arranged according to their atomic weights, exhibit an apparent periodicity of properties. . Elements which are similar as regards to their chemical properties have atomic weights which are either of nearly the same value (e. g. , Pt, Ir, Os) or which increase regularly (e. g. , K, Rb, Cs). 3. The arrangement of the elements, or of groups of elements in the order of their atomic weights, corresponds to their so-called valencies, as well as, to some extent, to their distinctive chemical properties; as is apparent among other series in that of Li, Be, Ba, C, N, O, and Sn. 4. The elements which are the most widely diffused have small atomic weights. 5.

The magnitude of the atomic weight determines the character of the element, just as the magnitude of the molecule determines the character of a compound body. 6. We must expect the discovery of many yet unknown elements–for example, elements analogous to aluminium and silicon–whose atomic weight would be between 65 and 75. 7. The atomic weight of an element may sometimes be amended by a knowledge of those of its contiguous elements. Thus the atomic weight of tellurium must lie between 123 and 126, and cannot be 128. (This was based on the position of tellurium between antimonyand iodine whose atomic weight is 127.

However Moseley later explained the position of these elements without revising the atomic weight values — see below. ) 8. Certain characteristic properties of elements can be foretold from their atomic weights. Isotopes are variants of atoms of a particular chemical element, which have differing numbers of neutrons. Atoms of a particular element by definition must contain the same number of protons but may have a distinct number of neutrons which differs from atom to atom, without changing the designation of the atom as a particular element.

The number of nucleons (protons and neutrons) in the nucleus, known as themass number, is not the same for two isotopes of any element. For example, carbon-12, carbon-13 and carbon-14 are three isotopes of the element carbon with mass numbers 12, 13 and 14 respectively. The atomic number of carbon is 6 (every carbon atom has 6 protons); therefore the neutron numbers in these isotopes are 6, 7 and 8 respectively. A nuclide is an atom with a specific number of protons and neutrons in the nucleus, for example carbon-13 with 6 protons and 7 neutrons.

The nuclide concept (referring to individual nuclear species) emphasizes nuclear properties over chemical properties, while the isotopeconcept (grouping all atoms of each element) emphasizes chemical over nuclear. The neutron number has drastic effects on nuclear properties, but its effect on chemical properties is negligible in most elements, and still quite small in the case of the very lightest elements, although it does matter in some circumstances (for hydrogen, the lightest of all elements, the isotope ffect is large enough to strongly affect biology). Since isotope is the older term, it is better known than nuclide, and is still sometimes used in contexts where nuclide might be more appropriate, such as nuclear technology and nuclear medicine. An isotope and/or nuclide is specified by the name of the particular element (this indicates the atomic number implicitly) followed by a hyphen and the mass number (e. g. helium-3, helium-4, carbon-12, carbon-14, uranium-235 and uranium-239). When a chemical symbol is used, e. . , “C” for carbon, standard notation is to indicate the number of nucleons with a superscript at the upper left of the chemical symbol and to indicate the atomic number with a subscript at the lower left (e. g. 3 2He, 4 2He, 12 6C, 14 6C, 235 92U, and 239 92U, respectively). Since the atomic number is implied by the element symbol, it is common to state only the mass number in the superscript and leave out the atomic number subscript (e. g. 3 He, 4 He, 12 C, 14 C, 235 U, and 239 U, respectively).

The letter m is sometimes appended after the mass number to indicate ametastable or energetically-excited nuclear state (rather than the lowest-energy ground state), for example 180m 73Ta (tantalum-180m). Some isotopes are radioactive and are therefore described as radioisotopes or radionuclides, while others have never been observed to undergo radioactive decay and are described as stable isotopes. For example, 14 C is a radioactive form of carbon while 12 C and 13 C are stable isotopes. There are about 339 naturally occurring nuclides on Earth,[1] of which 288 are primordial nuclides, meaning that they have xisted since the solar system’s formation. These include 33 nuclides with very long half-lives (over 80 million years) and 255 which are formally considered as “stable isotopes”,[1] since they have not been observed to decay. Many apparently “stable” isotopes are predicted by theory to be radioactive, with extremely long half-lives (this does not count the possibility of proton decay, which would make all nuclides ultimately unstable). Of the 255 nuclides never observed to decay, only 90 of these (all from the first 40 elements) are stable in theory to all known forms of decay.

Element 41 (niobium) is theoretically unstable via spontaneous fission, but this has never been detected. Many other stable nuclides are in theory energetically susceptible to other known forms of decay, such as alpha decay or double beta decay, but no decay has yet been observed. The half-lives for these processes often exceed a million times the estimated age of the universe, and in fact there are 27 known radionuclides (see primordial nuclide) with half-lives longer than the age of the universe.

Adding in the radioactive nuclides that have been created artificially, there are more than 3100 currently known nuclides. [2] These include 905 nuclides which are either stable, or have half-lives longer than 60 minutes. See list of nuclides for details. There are three primary types of radiation: * Alpha – these are fast moving helium atoms. They have high energy, typically in the MeV range, but due to their large mass, they are stopped by just a few inches of air, or a piece of paper. * Beta – these are fast moving electrons.

They typically have energies in the range of a few hundred keV to several MeV. Since electrons are might lighter than helium atoms, they are able to penetrate further, through several feet of air, or several millimeters of plastic or less of very light metals. * Gamma – these are photons, just like light, except of much higher energy, typically from several keV to several MeV. X-Rays and gamma rays are really the same thing, the difference is how they were produced. Depending on their energy, they can be stopped by a thin piece of aluminum foil, or they can penetrate several inches of lead.

In this experiment, we study the penetrating power of each type of radiation. For this test, we used an old radium coated watch hand, obtained from, where else,  The radium is an alpha source, and we also get several beta and gamma emissions as well from daughter products. This hand is about 50 years old, so we should get a nice mixture of radiations. This makes it an ideal source to use for these experiments. We first placed the radium watch hand about 1 cm from the face of a GM-45 detector. We obtained an average of 44307 CPM.

Next, we placed a single piece of paper in between the watch hand and detector. The rate dropped to 35111 CPM. The paper blocked all of the alpha rays. We then put a second piece of paper in between, the counts dropped to 31583 CPM. With a third piece of paper, 27977 CPM. We are now blocking beta rays with our pieces of paper, each piece blocks some more of the radiation. With 3 mm of steel (1/8 inch), the reading went down to 394 CPM. Background is about 50 CPM, so at this point we’re picking up some of the gammas from the source.



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