Thursday, July 30, 2009

Atomic Theory Research Paper

Throughout the ages the study of what we call today chemistry has evolved into a highly developed point of study. One distinct element of chemistry is atomic theory. Throughout the ages atomic theory has been developed and extended by many different men who were all well-known chemists and physicists in their day. They developed the study of atoms from pure conjecture into known facts.

John Dalton originally was a schoolteacher in England. He thought about atoms as particles, which could make up the elements of our universe. Dalton reasoned about the nature of compounds, his theory became known as the “law of multiple proportions: when two elements form a series of compounds, the ratios of the masses of the second element that combine with one gram of the first element can always be reduced to small whole numbers” (43). Dalton created his theory on atoms and their nature. It was made up of four parts. First, “each element is made up of tiny particles called atoms.” Second, “The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways.” Third, “chemical compounds are formed when atoms of different elements combine with each other. A given compound always has the same relative numbers and types of atoms.” Fourth, “chemical reactions involve reorganization of the atoms-changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction.” Dalton, using the information that he had deduced and assumptions that he had made, created a table of the atomic masses, also known as atomic weights. Despite the fact that many of the masses on Dalton’s table were proved incorrect the creation of a table was a major step in the field of chemistry. (45)

The ideas and information that Dalton produced and showed to the scientific community was built upon in many ways by his successors. The aspect that was most developed by the scientists that followed Dalton was the periodic laws and the table that is organized based on those laws.

After [. . .] Dalton had developed the idea of atomic weights, chemists sought arithmetic connections between them, partly to see whether there was any likelihood of all elements being composed of a simple, common substance and partly to see whether occasional similarities in their properties pointed to similarities in structure. [. . .] Mendeleyev [. . .] formulated the periodic law, according to which, when all known elements are arranged in order of increasing atomic weight, the resulting table shows a periodicity of properties and allows one to observe the many types of chemical relation hitherto studied only in isolation. (Britannica)

At first, as any other scientific revolution, the scientific community did not accept Mendeleyev’s table when his laws were published in his book Principles of Chemistry (1868-1871). However, his periodic laws stood true to all elements of his time. Another interesting point is that since then many other elements have been discovered or created which all follow the same laws. Mendeleyev even left spaces for elements that were not known at the time foretelling that there were elements that had not been discovered when he died. Mendeleyev went so far as to predict the properties of many of the elements that he expected to be discovered. Today, the Periodic Table of the Elements looks very different and has many more elements on it but it still is based on the same laws, which is quite remarkable when one thinks about it. The periodic table is made up of transition metals, the representative elements, and the noble gases. The representative elements and the noble gases are organized into eight groups. These groups are organized based upon the number of valence electrons in the atom of that element. For example, oxygen has an electron configuration of [O] 1s2, 2s2, 2p4. The outer most orbital is the second, and there are six electrons in that orbital. The eighth group is Helium and the noble gases. The noble gases are very stable elements because they fill their subshells completely. They all end in a s2, p6, which fills the subshells. The other elements, the transition metals are set in the middle and the lanthanide series and actinide series are pulled out from the table for space reasons.

Joseph Gay-Lussac, a French chemist and physicist, used his experiments to decipher the means of determining the absolute formulas of compounds. The experiments that Gay-Lussac performed were mainly with gases. He worked with hydrogen and oxygen specifically. Gay-Lussac discovered that two quantities of hydrogen reacted with one quantity of oxygen, which formed two quantities of water vapor. Avogadro later used Gay-Lussac’s data when Avogadro formed his hypothesis.

Amodeo Avogadro, an Italian, worked towards discovering the means of determining the absolute formulas of compounds, as Gay-Lussac did. He used the data from Gay-Lussac’s experiments to postulate a hypothesis, known as Avogadro’s hypothesis. It read: “at the same temperature and pressure, equal volumes of different gases contain the same number of particles” (46). However, Avogadro’s deductions were ignored for fifty years before the chemistry community accepted his hypotheses.

William Crookes was another British scientist who helped the development of atomic theory. He developed atomic theory further through his cathode ray studies. Crookes’ experiments with cathode-ray tubes led him to show that when high-voltage electric current through the tube the tube glows. He also discovered that the cathode rays travel in straight lines. J.J. Thomson developed Crookes’ studies further a few years later.

J.J. Thomson, yet another Englishman, also worked on atomic theory. Thomson studied Crookes’ experiments and tested them. Thomson went on to discover that because the ray went from the negative cathode in the tube the ray must be negative. He also deduced that the ray was a stream of particles. Combining this deduction and his study that the ray went from negative to positive, he theorized that the stream must be a stream of negatively charged particles, now known as electrons. Thomson’s further studies allowed him to determine the ratio between the charge and mass of an electron, -1.76 X 108 C/g. Also from his experiments, Thomson discovered that all atoms must contain electrons. This must be true, he thought, because the rays occurred with various metals. Thomson is also known for his ‘plum pudding’ model of the atom. He believed that an atom was an amorphous mass in which there was no defined structure.

The New Zealand scientist that is notable for his work on atomic theory is Ernest Rutherford. Just after the turn of the twentieth century Rutherford was working on many experiments expanding on his predecessors. One of Rutherford’s most famous experiments was the one in which he tested Thomson’s ‘plum pudding’ model of the atom. Rutherford set up a device that aimed a particles at a thin metal foil with a screen encircling the metal foil meant to detect the a particles. According to Thomson’s belief the particles would all go right through the foil and hit the screen behind the metal foil. However, in the experiment when Rutherford sent the a particles at the metal foil the particles went in more than one direction. Many of them went right through the foil, however, some also were redirected when they contacted the metal foil. Knowing that electrons were negatively charged, Rutherford was able to hypothesize that there was a positively charged mass in the center of an atom, which he named the nucleus. However, many a particles still went through the foil, which told Rutherford that an atom had a great amount of empty space in its structure. Therefore, the structure was a relatively massive, positively charged center with electrons spread out throughout a large volume containing the extremely small electrons. Based on Thomson and Rutherford’s experimentation today we have developed a great deal of knowledge on the structure of an atom. We know that the atom is made up of negatively charged electrons which have almost no mass, neutral neutrons with a mass of about one atomic mass unit, and positively charged protons with a mass of about one atomic mass unit.

The German scientist, Max Planck, is also known for his advances in atomic theory. Planck studied the radiation given off by solid bodies that were heated to incandescence. Through Planck’s experimentation he discovered that matter could absorb or radiate any amount of energy. Planck explained his observations through an assumption “that energy can be gained or lost only in whole-number multiples” (295). The equation that went along with this assumption was DE=nhv, often used as DE=hv. In this equation DE is the change in energy, n is an integer, h is Planck’s constant (6.626 X 10-34 J-s), and v is the frequency of the electromagnetic radiation that is either absorbed or radiated. Before Planck’s experiments the belief of the scientific community was that the energy of matter did not change. Today based on Planck’s studies, we know that energy can only be found in certain quantities, the hv of Planck’s equation, which are now known as quanta (singular – quantum) which indicates that energy has separate properties from any other substance; this leads to quantum physics.

A notable Danish scientist, Neils Bohr, worked on atomic theory and is known for his model of the atom. Bohr offered a new idea. He stated “that the electron in a hydrogen atom moves around the nucleus only in certain allowed circular orbits” (301). This is known as the quantum model for the hydrogen atom. Through his knowledge of classical physics and its theories and some assumptions of his own he calculated the radii of the ‘allowed’ orbits of the electrons in a hydrogen atom. Bohr developed many ideas, which supported his quantum model. One of which asserts the energy levels available to the electron in the hydrogen atom. This equation is as follows: E=(-2.178 X 10-18J)(Z2/n2); the n is an integer (the larger the integer the larger the radius of the orbit) and Z is the charge of the nucleus.

Werner Heisenberg, a German, Louis de Broglie, a Frenchman, and Erwin Schrodinger, an Austrian, all worked towards a new model of the atom knowing that the Bohr model was not correct. These three men adopted a new path in the study of the atom. It is known as the wave, or quantum, mechanics theory. This resulted in yet another model of the atom. De Broglie discovered that the electron has wave properties. Schrodinger continued in this thought by trying to emphasize the wave properties of an atom. De Broglie and Schrodinger believed that the connection between the nucleus and an electron was comparable to a standing wave. Examples of standing waves include things like string instruments and so they studied standing waves and compared their findings to atomic structure. Schrodinger came up with an atom model that had electrons that behaved as if they were standing waves.

Schrodinger’s equation is: Hy=Ey; where y, called the wave function, is a function of the coordinates (x, y, and z) of the electron’s position in three-dimensional space and H represents a set of mathematical instructions called an operator. In this case, the operator contains mathematical terms that produce the total energy of the atom (the sum of the potential energy due to the attraction between the proton and electron and the kinetic energy of the moving electron). When this equation is analyzed, many solutions are found. Each solution consists of a wave function y that is characterized by a particular value of E. A specific wave function is often known as an orbital. (307)

When Schrodinger’s equation is solved the results show that there are many wave functions, orbitals, that are solutions. Every individual orbital is distinguished by quantum numbers, which can be deciphered so that the properties of the orbital are known. There are three kinds of quantum numbers: principal, angular momentum, and magnetic.

The principal quantum number (n) has integral values: 1, 2, and 3,…. The principal quantum number is related to the size and energy of the orbital. As n increases, the orbital becomes larger and the electron spends more time further from the nucleus. An increase in n also means higher energy, because the electron is less tightly bound to the nucleus and the energy is less negative.

The angular momentum quantum number (l) has integral values from 0 to (n – 1) for each value of n. This quantum number is related to the shape of atomic orbitals. The value of l for a particular orbital is commonly assigned a letter: l = 0 is called s; l = 1 is called p; l = 2 is called d; l = 3 is called f. This system arises from early spectral studies [. . .]. [l is also sometimes known as a subshell.]

The magnetic quantum number (ml) has integral values between l and –l, including 0. The value of ml is related to the orientation of the orbital in space relative to the other orbitals in the atom. (309-310)

Heisenberg’s discoveries added another twist to atomic theory, especially the wave, or quantum mechanics model. He concluded that “there is a fundamental limitation to just how precisely we can know both the position and momentum of a particle at a given time” (307). This is known as the Heisenberg uncertainty principle. Mathematically speaking it is shown as follows: “(Dx)[D(mv)]і[h/(4p)] where Dx is the uncertainty in a particle’s position, D(mv) is the uncertainty in a particle’s momentum, and h is Planck’s constant. Thus the minimum uncertainty in the equation is h/(4p)” (307). The only problem with the equation of the uncertainty principle is that the closer or more accurate a position measurement the more inaccurate the measurement of momentum is. When this principle is exercised in thinking about electrons it is shown that it is impossible for us to know the exact motion of an electron on its path around the nucleus. Therefore, Heisenberg disproves Bohr’s atomic model, a specified orbit, by his uncertainty principle.

The periodic table can be looked in blocks as well. The blocks are called the s-block, p-block, d-block, and f-block. These blocks have to do with the subshells of orbitals. Electrons fill up orbitals in order from 1 up. However, within these orbitals electrons fill up subshells. The subshells fill up in order from s, to p, to d, and then finally to f. The s subshell can hold up to two electrons, the p subshell can hold up to six electrons, the d subshell holds up to ten electrons, and the f subshell holds at maximum fourteen electrons. From left to right the first two vertical columns of elements and Helium are in the s-block because the last subshell in which electrons fall into is a s subshell. The p-block includes the elements in rest of the representative elements and the noble gases. The d-block includes all of the transition metals with exception of the lanthanide series and actinide series. The lanthanide series and actinide series are the elements in the f-block. These subshells fill in an order, which is shown in the aufbau principle.

From the work of the scientists who worked on atomic theory we know that electrons can be moved around between atoms. The atoms that lose an electron or electrons are called ions. Ionization energy is the energy needed to remove an electron from an atom so that the atom is at its considered ground state, forming a positive ion. An atom’s ground state is defined as the lowest possible energy state. There are two types of ionization energy: first and second ionization energy. First ionization energy is the energy need to remove the highest energy electron from an atom. Second ionization energy is much greater than first ionization energy and removes the next highest energy electron from the atom. This is true is because when an atom has a positive charge, after one electron has already been removed; the atom is bounded stronger than before. Therefore, more energy is necessary to remove a second electron from an atom. On the periodic table ionization energy reduces as one goes down a group of elements; for example, Helium has the highest ionization energy in the eighth group while Radon has the lowest. Also, as one moves from left to right ionization energy increases in most cases. For instance, Hydrogen has a lower ionization energy level than Helium. Electron affinity is the energy required for adding an electron to an atom, forming a negative ion. Generally, when going down a group on the periodic table electron affinity decreases, or gets closer to zero or more positive, but the differences are very small.

In conclusion, over many years scientists have developed atomic theory. However, it is all based on the things the early chemists did, including the experiments they performed and the information they passed on. In many respects, the early atomic theorists were the most important because they helped pave the way for the other scientists that continued to develop atomic theory.

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