The Atom, Research Paper Example

The atom is considered to be the smallest particle of an element. However, it is a complex particle which has not been studied completely yet. Moreover, the outlook on the atom is constantly changing and there are several theories that mean to prove various points of view on the atomic structure.

For a long time there was no definite idea of an actual atomic structure. In the end of XIX – beginning of XX centuries the atom was proved to be a complex article consisting of simpler (elementary) particles. In 1911 on the basis of experimental data the English physicist Ernst Rutherford offered a nuclear atomic model in which a relatively small volume appeared to comprise almost total mass concentration. According to the physician, the atomic nucleus consisting of protons and neutron is charged positively (Salzman). At the same time, it is surrounded by electrons which are bearing negative charge.

Thus, Rutherford’s findings laid the foundation for the nuclear atomic theory. Since that time the physicists got a lot of vital information concerning chemical properties of the atom. So, hydrogen (?) has the lightest atom. While almost the whole atomic mass is concentrated in its nucleus, it is quite natural that the hydrogen nucleus represents a positively-charged elementary particle which was called proton (which means “first”). Thus, the proton has the mass almost equal to the one of the hydrogen nucleus (to be more exact 1,00728 carbon units) and the electrical charge of  +1. The atoms of other heavier elements have a larger charge and, probably, mass.

The measurements of nuclear charges showed that the charge of the atomic nuclear is equal to the atomic number of the element. However, it was impossible to suppose, because the latter, being like-charged would inevitably start out from each other and, consequently, the nucleus would appear unsteady. Moreover, the mass of atomic nuclei would turn out twice as much as the total mass of protons which determine the atomic charge of the corresponding elements.

Then, the atomic nucleus was supposed to contain protons which quantity exceeded the atomic number of the element, and the excessive positive charge was compensated by electrons which were also a part of the nucleus. These electrons might have to keep protons in the nucleus. However, this theory was to be rejected, as it appeared impossible to suppose that heavy (protons) and light (electrons) particles were able to co-exist in a compact atomic nucleus.

In 1932 James Chadwick discovered an elementary particle which was not charged at all and, therefore, was called a neutron. The neutron has a mass a little exceeding the proton mass (1,008665 carbon units) (Carpi). After that Dmitry Ivanenko and Werner Heisenberg independently of each other offered the theory of the atomic structure which soon became generally accepted.

According to this theory the atomic nuclei of all the elements (except for hydrogen) consist of protons and neutrons. The amount of protons in the nucleus determines the value of its positive charge and the summary number of protons and neutron becomes its mass value. Nuclear particles – protons and neutrons – are combined under the joint name “nucleon”. Thus, the amount of protons corresponds to the atomic number of the element and the total amount of nucleons – to its mass number (A). So, the number of neutrons (N) in the atomic nucleus can be found with the help of the following formula:

N = A – Z where Z is the atomic number.To conclude, the proton-neutron theory helped to adjust the risen earlier differences in theories on the atomic nucleus structure and the connection between its serial number and atomic mass.

The electronic structure of the atom determines its chemical properties including capability of creating chemical compounds – the most important function in chemistry. Because of a small size and large mass the atomic nucleus may be considered point like and resting at the center of the mass. Usually the chemistry studies in detail the system of electrons moving around the nucleus.

It is impossible to describe the electron motion in the atom from the point of view of classical mechanics and electrodynamics, because a charged particle moving in a circle in this case must emit electromagnetic waves and, therefore, waste energy and fall on the nucleus. In 1912 the Danish physicist Niels Bohr suggested that this problem could be solved by distinguishing so called fixed orbits moving along which electrons radiate no energy. Radiation may occur only in case an electron is moving from one orbit to another. Eventually new theories were offered which enabled to conceive the process of electron movement more clearly. Matrix mechanics of the German theoretical physicist Werner Heisenberg described the electron as a particle and wave mechanics of the Austrian theoretical physicist Erwin Schrödinger – as a wave. In the long run these two theories were combined into quantum mechanics which in reference to the chemistry developed into the quantum chemistry.

The essence of the theory of quanta consists in the fact that the energy is radiated not continuously but in definite small portions. Consequently, the energy store is changing in discrete steps – the fractional number of quanta can be neither radiated, no consumed by the object. The more distance from the orbit where the electron is situated to the one the electron means to transit, the more radiation frequency is. Bohr made an attempt to find out the radii of possible orbits of the only hydrogen electron. It appeared that that they relate as the squares of the natural numbers: 1 : 2 : 3 : …: n. The value n was called a principle quantum number (Farabee).

Later on the theory of Bohr was used in study of the atomic structure of other elements. In enabled to answer the question of the electron arrangement in atoms of different elements and find the relation between the properties of the elements and the structure of the atomic electron shells.  Nowadays the atomic structure of all elements is found. Though it should be taken into account that it is just another theory which may contribute to the explanation of many physical and chemical properties of the elements.

Soon after that the notion of valence was discovered. The chemical behavior of atoms is mostly influenced by the interactions between electrons. All the atomic electrons are thought to remain within definite electron configurations. As it has been already mentioned, the electrons are separated into shells which are situated on some distance from the nucleus. The electrons in the outermost shell are called the valence electrons. Electrons fill orbitals and shells from the inside.

The number of electrons in an atom’s outermost valence shell influences its bonding behavior. Consequently, elements with the same number of valence electrons are grouped together in the periodic table of the elements.

Group 1 elements contain one electron on their outer shell; Group 2, two electrons; Group 3, three electrons; etc. As a general rule, the fewer electrons in an atom’s valence shell, the more reactive it is. Group 1 metals are therefore very reactive.

Every atom is much more stable that is less energetic, if the shell is of full valence. This can be achieved in case the atom shares electrons with neighboring atoms. It will be called a covalent bond or if it removes electrons from other atoms – an ionic bond. Another form of ionic bonding involves an atom giving some of its electrons to another atom; this also works because it can end up with a full valence by giving up its entire outer shell. By moving electrons, the two atoms become linked. This is known as chemical bonding and contributes to the building molecules (Moring).

The quantum mechanical theory of the atomic structure views the atom as a system of microparticles which are not governed by laws of the classical mechanics. The first atomic structure models were similar to the structure of the planetary system. Nevertheless, it turned out to be impossible to describe the electron motion the way it is done with planets. From the point of view of the quantum mechanics one may speak of the definite state of the atom which may be characterized as charged with some amount of energy which according to the intermittency hypothesis can change only by the transit from one such state to another.

Furthermore, the quantum mechanics admits that electrons may behave in the atom like both particles and waves (it is called the principle of wave-corpuscle duality). And, finally, according to the Heisenberg indeterminacy principle it seems to be impossible to determine the motion trajectory of electrons in the atom (Moring). At present due to the new methods of quantum mechanics the electronic structure of all existing types of the atom is well-known. The atom of the chemical element is described with a definite electronic configuration (electronic formula) which enables to suggest the chemical properties of the given element.

To conclude, in our research we managed to view the history of development of the atomic theories. It becomes evident that at the very beginning they were very simple and failed to reflect the real picture of our chemical world. However, as far as the science was developing, it became possible to combine several theories of different scientists (including not only chemists, but also physicists and biologists) and select the most beneficial and authentic facts. Moreover, we got the opportunity to trace how exactly the outlook on the atomic structure was changing with time. Nowadays the quantum atomic theory is at its height and because of the fact that much is still unknown today; it seems that it will take much time and effort to bring this theory to its logical end. To add, as today the scientists are apt to combine several sciences to make the study more effective, the quantum atomic theory appears to be extremely important in our increase of knowledge about the world.

Works Cited

Carpi, A. Atomic Theory II. Ions, Isotopes and Electron Shells. 2003. 28 Mar. 2009. <http://www.visionlearning.com/library/module_viewer.php?mid=51>

Farabee, M. Chemistry I: Atoms and Molecules. 2007. 28 Mar. 2009. <http://www.estrellamountain.edu/faculty/farabee/biobk/BioBookCHEM1.html>

Hot Atom Chemistry 1968-1976. Jan. 31. 2008. 28 Mar. 2009. <http://www.bnl.gov/chemistry/History/HotAtomChemistry1968-1976.asp>

Moring, G. Bohr’s Atomic Theory. 2001. 28 Mar. 2009. <http://www.infoplease.com/cig/theories-universe/bohrs-atomic-theory.html>

Salzman, W. Atomic and Molecular Orbitals Page. May, 30.1997. 28 Mar. 2009. <http://chemistry.about.com/gi/dynamic/offsite.htm?zi=1/XJ/Ya&sdn=chemistry&zu=http%3A%2F%2Fwww.chem.arizona.edu%2F~salzmanr%2Forbitals.html>

The Atomic Theory. Aug. 1999. 28 Mar. 2009. <http://library.thinkquest.org/27948/bohr.html>