Thesis Abstract

Thesis Overview

<p><b>1.0 INTRODUCTION&nbsp;</b></p><p> The study of inorganic chemistry involves interpreting, correlating, and predicting the properties and structures of an enormous range of materials. Sulfuric acid is the chemical produced in the largest tonnage of any compound. A greater number of tons of concrete is produced, but it is a mixture rather than a single compound. Accordingly, sulfuric acid is an inorganic compound of enormous importance. On the other hand, inorganic chemists study compounds such as hexaaminecobalt(III) chloride, [Co(NH 3 ) 6 ]Cl 3, and Zeise’s salt, K[Pt(C 2 H 4 )Cl 3]. Such compounds are known as coordination compounds or coordination complexes. Inorganic chemistry also includes areas of study such as nonaqueous solvents and acid-base chemistry. Organometallic compounds, structures and properties of solids, and the chemistry of elements other than carbon are areas of inorganic chemistry. However, even many compounds of carbon (e.g., CO 2 and Na 2CO 3) are also inorganic compounds. The range of materials studied in inorganic chemistry is enormous, and a great many of the compounds and processes are of industrial importance. Moreover, inorganic chemistry is a body of knowledge that is expanding at a very rapid rate, and a knowledge of the behavior of inorganic materials is fundamental to the study of the other areas of chemistry. Because inorganic chemistry is concerned with structures and properties as well as the synthesis of materials, the study of inorganic chemistry requires familiarity with a certain amount of information that is normally considered to be physical chemistry. As a result, physical chemistry is normally a prerequisite for taking a comprehensive course in inorganic chemistry. There is, of course, a great deal of overlap of some areas of inorganic chemistry with the related areas in other branches of chemistry. A knowledge of atomic structure and properties of atoms is essential for describing both ionic and covalent bonding. Because of the importance of atomic structure to several areas of inorganic chemistry, it is appropriate to begin our study of inorganic chemistry with a brief review of atomic structure and how our ideas about atoms were developed.</p><p><b>&nbsp;1.1 SOME EARLY EXPERIMENTS IN ATOMIC PHYSICS&nbsp;</b></p><p>It is appropriate at the beginning of a review of atomic structure to ask the question, “ How do we know what we know? ”In other words, “ What crucial experiments have been performed and what do the results tell us about the structure of atoms? ”Although it is not necessary to consider all of the early experiments in atomic physics, we should describe some of them and explain the results. The fi rst of these experiments was that of J. J. Thomson in 1898–1903, which dealt with cathode rays. In the experiment, an evacuated tube that contains two electrodes has a large potential difference generated between the electrodes as shown in Figure 1.1 . Under the infl uence of the high electric fi eld, the gas in the tube emits light. The glow is the result of electrons colliding with the molecules of gas that are still present in the tube even though the pressure has been reduced to a few torr. The light that is emitted is found to consist of the spectral lines characteristic of the gas inside the tube. Neutral molecules of the gas are ionized by the electrons streaming from the cathode, which is followed by recombination of electrons with charged species. Energy (in the form of light) is emitted as this process occurs. As a result of the high electric fi eld, negative ions are accelerated toward the anode and positive ions are accelerated toward the cathode. When the pressure inside the tube is very low (perhaps 0.001 torr), the mean free path is long enough that some of the positive ions strike the cathode, which emits rays. Rays emanating from the cathode stream toward the anode. Because they are emitted from the cathode, they are known as cathode rays . Cathode rays have some very interesting properties. First, their path can be bent by placing a magnet near the cathode ray tube. Second, placing an electric charge near the stream of rays also causes the path they follow to exhibit curvature. From these observations, we conclude that the rays are electrically charged. The cathode rays were shown to carry a negative charge because they were attracted to a positively charged plate and repelled by one that carried a negative charge. The behavior of cathode rays in a magnetic fi eld is explained by recalling that a moving beam of charged particles (they were not known to be electrons at the time) generates a magnetic fi eld. The same principle is illustrated by passing an electric current through a wire that is wound around a compass. In this case, the magnetic fi eld generated by the fl owing current interacts with the magnetized needle of the compass, causing it to point in a different direction. Because the cathode rays are negatively charged particles, their motion generates a magnetic fi eld that interacts with the external magnetic fi eld. In fact, some important information about the nature of the charged particles in cathode rays can be obtained from studying the curvature of their path in a magnetic fi eld of known strength. Consider the following situation. Suppose a cross wind of 10 miles/hour is blowing across a tennis court. If a tennis ball is moving perpendicular to the direction the wind is blowing, the ball will follow a curved path. It is easy to rationalize that if a second ball had a cross-sectional area that was twice that of a tennis ball but the same mass, it would follow a more curved path because the wind pressure on it would be greater. On the other hand, if a third ball having twice the cross-sectional area and twice the mass of the tennis ball were moving perpendicular to the wind direction, it would follow a path with the same curvature as the tennis ball. The third ball would experience twice as much wind pressure as the tennis ball, but it would have twice the mass, which tends to cause the ball to move in a straight line (inertia). Therefore, if the path of a ball is being studied when it is subjected to wind pressure applied perpendicular to its motion, an analysis of the curvature of the path could be used to determine the ratio of the cross-sectional area to the mass of a ball, but neither property alone. A similar situation exists for a charged particle moving under the infl uence of a magnetic fi eld. The greater the mass, the greater the tendency of the particle to travel in a straight line. On the other hand, the higher its charge, the greater its tendency to travel in a curved path in the magnetic fi eld. If a particle has two units of charge and two units of mass, it will follow the same path as one that has one unit of charge and one unit of mass. From the study of the behavior of cathode rays in a magnetic fi eld, Thomson was able to determine the charge-to-mass ratio for cathode rays, but not the charge or the mass alone. The negative particles in cathode rays are electrons, and Thomson is credited with the discovery of the electron. From his experiments with cathode rays, Thomson determined the charge-tomass ratio of the electron to be  1.76  10 8 coulomb/gram. It was apparent to Thomson that if atoms in the metal electrode contained negative particles (electrons), they must also contain positive charges because atoms are electrically neutral. Thomson proposed a model for the atom in which positive and negative particles were embedded in some sort of matrix. The model became known as the plum pudding model because it resembled plums embedded in a pudding. Somehow, an equal number of positive and negative particles were held in this material. Of course we now know that this is an incorrect view of the atom, but the model did account for several features of atomic structure. <br></p><p> The second experiment in atomic physics that increased our understanding of atomic structure was conducted by Robert A. Millikan in 1908. This experiment has become known as the Millikan oil drop experiment because of the way in which oil droplets were used. In the experiment, oil droplets (made up of organic molecules) were sprayed into a chamber where a beam of x-rays was directed on them. The x-rays ionized molecules by removing one or more electrons producing cations. As a result, some of the oil droplets carried an overall positive charge. The entire apparatus was arranged in such a way that a negative metal plate, the charge of which could be varied, was at the top of the chamber. By varying the (known) charge on the plate, the attraction between the plate and a specifi c droplet could be varied until it exactly equaled the gravitational force on the droplet. Under this condition, the droplet could be suspended with an electrostatic force pulling the drop upward that equaled the gravitational force pulling downward on the droplet. Knowing the density of the oil and having measured the diameter of the droplet gave the mass of the droplet. It was a simple matter to calculate the charge on the droplet, because the charge on the negative plate with which the droplet interacted was known. Although some droplets may have had two or three electrons removed, the calculated charges on the oil droplets were always a multiple of the smallest charge measured. Assuming that the smallest measured charge corresponded to that of a single electron, the charge on the electron was determined. That charge is  1.602  10  19 coulombs or  4.80  10 10 esu (electrostatic units: 1esu  1g 1/2 cm 3/2 sec  1 ). Because the charge-to-mass ratio was already known, it was now possible to calculate the mass of the electron, which is 9.11  10 31 kg or 9.11  10  28 g. The third experiment that is crucial to understanding atomic structure was carried out by Ernest Rutherford in 1911 and is known as Rutherford’s experiment. It consists of bombarding a thin metal foil with alpha ( α) particles. Thin foils of metals, especially gold, can be made so thin that the thickness of the foil represents only a few atomic diameters. The experiment is shown diagrammatically in Figure 1.2 . It is reasonable to ask why such an experiment would be informative in this case. The answer lies in understanding what the Thomson plum pudding model implies. If atoms consist of equal numbers of positive and negative particles embedded in a neutral material, a charged particle such as an α particle (which is a helium nucleus) would be expected to travel near an equal number of positive and negative charges when it passes through an atom.&nbsp;</p><p>As a result, there should be no net effect on the α particle, and it should pass directly through the atom or a foil that is only a few atoms in thickness. A narrow beam of α particles impinging on a gold foil should pass directly through the foil because the particles have relatively high energies. What happened was that most of the α particles did just that, but some were defl ected at large angles and some came essentially back toward the source! Rutherford described this result in terms of fi ring a 16-inch shell at a piece of tissue paper and having it bounce back at you. How could an α particle experience a force of repulsion great enough to cause it to change directions? The answer is that such a repulsion could result only when all of the positive charge in a gold atom is concentrated in a very small region of space. Without going into the details, calculations showed that the small positive region was approximately 10  13 cm in size. This could be calculated because it is rather easy on the basis of electrostatics to determine what force would be required to change the direction of an α particle with a 2 charge traveling with a known energy. Because the overall positive charge on an atom of gold was known (the atomic number), it was possible to determine the approximate size of the positive region Rutherford’s experiment demonstrated that the total positive charge in an atom is localized in a very small region of space (the nucleus). The majority of α particles simply passed through the gold foil, indicating that they did not come near a nucleus. In other words, most of the atom is empty space.&nbsp;</p><p>The diffuse cloud of electrons (which has a size on the order of 10  8 cm) did not exert enough force on the α particles to defl ect them. The plum pudding model simply did not explain the observations from the experiment with α particles. Although the work of Thomson and Rutherford had provided a view of atoms that was essentially correct, there was still the problem of what made up the remainder of the mass of atoms. It had been postulated that there must be an additional ingredient in the atomic nucleus, and this was demonstrated in 1932 by James Chadwick. In his experiments a thin beryllium target was bombarded with α particles. Radiation having high penetrating power was emitted, and it was initially assumed that they were highenergy γ rays. From studies of the penetration of these rays in lead, it was concluded that the particles had an energy of approximately 7 MeV. Also, these rays were shown to eject protons having energies of approximately 5 MeV from paraffi n. However, in order to explain some of the observations, it was shown that if the radiation were γ rays, they must have an energy that is approximately 55 MeV. If an α particle interacts with a beryllium nucleus so that it becomes captured, it is possible to show that the energy (based on mass difference between the products and reactants) is only about 15 MeV. Chadwick studied the recoil of nuclei that were bombarded by the radiation emitted from beryllium when it was a target for α particles and showed that if the radiation consists of γ rays, the energy must be a function of the mass of the recoiling nucleus, which leads to a violation of the conservation of momentum and energy. However, if the radiation emitted from the beryllium target is presumed to carry no charge and consist of particles having a mass approximately that of a proton, the observations could be explained satisfactorily. Such particles were called neutrons, and they result from the reaction.</p><p> <b>1.2 THE NATURE OF LIGHT&nbsp;</b></p><p>From the early days of physics, a controversy had existed regarding the nature of light. Some prominent physicists, such as Isaac Newton, had believed that light consisted of particles or “ corpuscles. ” Other scientists of that time believed that light was wavelike in its character. In 1807, a crucial experiment was conducted by T. Young in which light showed a diffraction pattern when a beam of light was passed through two slits. Such behavior showed the wave character of light. Other work by A. Fresnel and F. Arago had dealt with interference, which also depends on light having a wave character. &nbsp; <br></p><p> The nature of light and the nature of matter are intimately related. It was from the study of light emitted when matter (atoms and molecules) was excited by some energy source or the absorption of light by matter that much information was obtained. In fact, most of what we know about the structure of atoms and molecules has been obtained by studying the interaction of electromagnetic radiation with matter or electromagnetic radiation emitted from matter. These types of interactions form the basis of several types of spectroscopy, techniques that are very important in studying atoms and molecules. In 1864, J. C. Maxwell showed that electromagnetic radiation consists of transverse electric and magnetic fi elds that travel through space at the speed of light (3.00  10 8 m/sec). The electromagnetic spectrum consists of the several types of waves (visible light, radio waves, infrared radiation, etc.) that form a continuum as shown in Figure 1.3 . In 1887, Hertz produced electromagnetic waves by means of an apparatus that generated an oscillating electric charge (an antenna). This discovery led to the development of radio. Although all of the developments that have been discussed are important to our understanding of the nature of matter, there are other phenomena that provide additional insight. One of them concerns the emission of light from a sample of hydrogen gas through which a high voltage is placed. The basic experiment is shown in Figure 1.4 . In 1885, J.J. Balmer studied the visible light emitted from the gas by passing it through a prism that separates the light into its components. <br></p>