Chemistry Time

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The Importance of Chemistry October 8, 2008

Filed under: Uncategorized — raaztarun @ 1:28 pm

How important is chemistry? For many, chemistry is considered the central science due to its significant connections and overlaps with other sciences. If a scientific discipline involves matter, chances are that chemistry is playing an important role.

A large amount of research breakthroughs in physics would not have been possible without the use of principles and methods of chemistry. The development of specialised drugs to cure diseases and understanding of biological and geological systems would be unimaginable if it weren’t for the contribution from chemistry. The list of examples is potentially endless (see the your examples section).

However, despite being readily accepted as such an important and fundamental science by the scientific community, especially those who have depended on chemistry in their work, the natural science of matter is rarely seen as interesting or important by the general public. Some university chemistry departments have faced the threat of closure, including Exeter’s late chemistry department and Sussex’s recent survival. These threats have not been due to poor department standards but due to failure of attracting the interest of enough new students. In response to this, the RSC recently acted as chemistry’s PR team, successfully getting the issue mentioned across in the media nation-wide.

I can remember once, when I told a friend of my choice of degree, being asked, “Are you sure you want to work in a chemist all your life?”…he was serious…true story.

On this site I will talk about an example of an area of chemical research that has immense potential applications in all areas of science, and everyday life – nanotechnology. The your examples section of the site is an interactive, message board-like environment where visitors can post brief examples of important areas of study/applications of chemistry.



Filed under: J.J.Thomson — raaztarun @ 4:27 pm

Sir Joseph John “J.J.” Thomson, OM, FRS (18 December 1856 – 30 August 1940) was a British physicist and Nobel laureate, credited for the discovery of the electron and of isotopes, and the invention of the mass spectrometer. He was awarded the 1906 Nobel Prize in Physics for the discovery of the electron and his work on the conduction of electricity in gases.


Joseph J. Thomson was born in 1856 in Cheetham Hill, Manchester in England, of Scottish parentage. His father died when he was only 16 years old.[1] In 1870 he studied engineering at University of Manchester known as Owens College at that time, and moved on to Trinity College, Cambridge in 1876. In 1880, he obtained his BA in mathematics (Second Wrangler and 2nd Smith’s prize) and MA (with Adams Prize) in 1883. In 1884 he became Cavendish Professor of Physics. One of his students was Ernest Rutherford, who would later succeed him in the post. In 1890 he married Rose Elisabeth Paget, daughter of Sir George Edward Paget, KCB, a physician and then Regius Professor of Physic at Cambridge. He fathered one son, George Paget Thomson, and one daughter, Joan Paget Thomson, with her. One of Thomson’s greatest contributions to modern science was in his role as a highly gifted teacher, as seven of his research assistants and his aforementioned son won Nobel Prizes in physics. His son won the Nobel Prize in 1937 for proving the wavelike properties of electrons. He was awarded a Nobel Prize in 1906, “in recognition of the great merits of his theoretical and experimental investigations on the conduction of electricity by gases.” He was knighted in 1908 and appointed to the Order of Merit in 1912. In 1914 he gave the Romanes Lecture in Oxford on “The atomic theory”. In 1918 he became Master of Trinity College, Cambridge, where he remained until his death. He died on 30 August 1940 and was buried in Westminster Abbey, close to Sir Isaac Newton. Thomson was elected a Fellow of the Royal Society on 12 June 1884 and was subsequently President of the Royal Society from 1915 to 1920.


Cathode rays

Thomson conducted a series of experiments with cathode rays and cathode ray tubes leading him to the discovery of electrons and subatomic particles. Thomson used the cathode ray tube in three different experiments.

First experiment

In his first experiment, he investigated whether or not the negative charge could be separated from the cathode rays by means of magnetism. He constructed a cathode ray tube ending in a pair of cylinders with slits in them. These slits were in turn connected to an electrometer. Thomson found that if the rays were magnetically bent such that they could not enter the slit, the electrometer registered little charge. Thomson concluded that the negative charge was inseparable from the rays.

Second experiment
Thomson's second experiment.

Thomson’s second experiment.

In his second experiment, he investigated whether or not the rays could be deflected by an electric field (something that is characteristic of charged particles).[2] Previous experimenters had failed to observe this, but Thomson believed their experiments were flawed because they contained trace amounts of gas. Thomson constructed a cathode ray tube with a practically perfect vacuum, and coated one end with phosphorescent paint. Thomson found that the rays did indeed bend under the influence of an electric field, in a direction indicating a negative charge.

Third experiment
Thomson's third experiment.

Thomson’s third experiment.

In his third experiment, Thomson measured the mass-to-charge ratio of the cathode rays by measuring how much they were deflected by a magnetic field and how much energy they carried. He found that the mass to charge ratio was over a thousand times lower than that of a hydrogen ion (H+), suggesting either that the particles were very light or very highly charged. Thomson’s conclusions were bold: cathode rays were indeed made of particles which he called “corpuscles“, and these corpuscles came from within the atoms of the electrodes themselves, meaning that atoms are in fact divisible. The “corpuscles” discovered by Thomson are identified with the electrons which had been proposed by G. Johnstone Stoney. He conducted this experiment in 1897. Thomson imagined the atom as being made up of these corpuscles swarming in a sea of positive charge; this was his plum pudding model. This model was later proved incorrect when Ernest Rutherford showed that the positive charge is concentrated in the nucleus.

Nobel Prize

Thomson’s discovery was made known in 1897, and caused a sensation in scientific circles, eventually resulting in him being awarded a Nobel Prize in Physics in 1906.[3] He notes that prior to his work: (1) the (negatively charged) cathode was known to be the source of the cathode rays; (2) the cathode rays were known to have the particle-like property of charge; (3) were deflected by a magnetic field like a negatively charged particle; (4) had the wave-like property of being able to penetrate thin metal foils; (5) had not yet been subject to deflection by an electric field. Thomson succeeded in causing electric deflection because his cathode ray tubes were sufficiently evacuated that they developed only a low density of ions (produced by collisions of the cathode rays with the gas remaining in the tube). Their ion densities were low enough that the gas was a poor conductor, unlike the tubes of previous workers, where the ion density was high enough that the ions could screen out the electric field. He found that the cathode rays (which he called corpuscles) were deflected by an electric field in the same direction as negatively charged particles would deflect. With the electrons moving along, say, the x-direction, the electric field E pointing along the y-direction, and the magnetic field B pointing along the z-direction, by adjusting the ratio of the magnetic field B to the electric field E he found that the cathode rays moved in a nearly straight line, an indication of a nearly uniform velocity v=E/B for the cathode rays emitted by the cathode. He then removed the magnetic field and measured the deflection of the cathode rays, and from this determined the charge-to-mass ratio e/m for the cathode rays. He writes: “however the cathode rays are produced, we always get the same value of e/m for all the particles in the rays. We may…produce great changes in the velocity of the particles, but unless the velocity of the particles becomes so great that they are moving nearly as fast as light, when other considerations have to be taken into account, the value of e/m is constant. The value of e/m is not merely independent of the velocity…it is independent of the kind of electrodes we use and also of the kind of gas in the tube.” Thomson notes that “corpuscles” are emitted by hot metals and “Corpuscles are also given out by metals and other bodies, but especially by the alkali metals, when these are exposed to light. They are being continually given out in large quantities and with very great velocities by radioactive substances such as uranium and radium; they are produced in large quantities when salts are put into flames, and there is good reason to suppose that corpuscles reach us from the sun.” Thomson also describes water drop experiments that enabled him to obtain a value for e that is about twice the modern value, and close to the then current value for the charge on a hydrogen ion in an electrolyte.

Isotopes and mass spectrometry

neon-20 and neon-22.

In the bottom right corner of this photographic plate are markings for the two isotopes of neon: neon-20 and neon-22.

In 1913, as part of his exploration into the composition of canal rays, Thomson channelled a stream of ionized neon through a magnetic and an electric field and measured its deflection by placing a photographic plate in its path. Thomson observed two patches of light on the photographic plate (see image on right), which suggested two different parabolas of deflection. Thomson concluded that the neon gas was composed of atoms of two different atomic masses (neon-20 and neon-22). This separation of neon isotopes by their mass was the first example of mass spectrometry, which was subsequently improved and developed into a general method by Thomson’s student F. W. Aston and by A. J. Dempster.

Other work

In 1906 Thomson demonstrated that hydrogen had only a single electron per atom. Previous theories allowed various numbers of electrons.[4][5]



Ernest Rutherford

Filed under: Ernest Rutherford — raaztarun @ 2:56 pm

Rutherford Model

The Rutherford model or planetary model was a model of the atom devised by Ernest Rutherford. Rutherford directed the famous Geiger-Marsden experiment in (1909), which suggested to Rutherford’s analysis (1911) that the Plum pudding model (of J. J. Thomson) of the atom was incorrect. Rutherford’s new model for theatom, based on the experimental results, had a number of essential modern features, including a relatively high central charge concentrated into a very small volume in comparison to the rest of the atom.

Experimental basis for the model

In the Geiger-Marsden experiment at Rutherford’s laboratory, alpha particles were used as a probe into atomic structure by being allowed to pass through a thin piece of gold foil, then detected. Rutherford predicted that all of the particles would pass through the foil, or be deflected slightly. This is indeed what happened most of the time, but a few particles, 1 in 8000, bounced unexpectedly nearly straight back toward the source. This supported the hypothesis that atoms have a dense region containing most of their mass, and associated with a highly concentrated electric field (probably positive in nature), instead of spread-out positive or negative field. Rutherford thought it likely, on purely symmetric and aesthetic grounds, that such a region of dense charge and mass would be located in the atom’s center. Such a region would then form a sort of atomic core.

In 1911, Rutherford came forth with his own physical model for subatomic structure, as an interpretation for the unexpected experimental results. In it, the atom is made up of a central charge (this is the modern atomic nucleus, though Rutherford did not use the term “nucleus” in his paper) surrounded by a cloud of orbiting electrons. In this 1911 paper, Rutherford only commits himself to a small central region of very high positive or negative charge in the atom, but uses the following language for pictorial purposes:[verification needed]

“For concreteness, consider the passage of a high speed a[lpha] particle through an atom having a positive central charge N e, and surrounded by a compensating charge of N electrons.”

From purely energetic considerations of how far alpha particles of known speed would be able to penetrate toward a central charge of 100 e, Rutherford was able to calculate that the radius of his gold central charge would need to be less (how much less could not be told) than 3.4 x 10-14 metres (the modern value is only about a fifth of this). This was in a gold atom known to be 10-10 metres or so in radius— a very surprising finding, as it implied a strong central charge less than 1/3000th of the diameter of the atom.

The Rutherford model did not attribute any structure to the orbits of the electrons themselves, though it did mention the atomic model of Hantaro Nagaoka, in which the electrons are arranged in one or more rings (this is the ONLY previous atomic model mentioned in Rutherford’s 1911 paper).

The Rutherford paper suggested that the central charge of an atom might be “proportional” to its atomic mass in hydrogen mass units (roughly 1/2 of it, in Rutherford’s model). For gold, this mass number is 197 (not then known to great accuracy) and was therefore modeled by Rutherford to be possibly 196. However, Rutherford did not attempt to make the direct connection of central charge to atomic number, since gold’s place on the periodic table was known to be about 79, and Rutherford’s more tentative model for the structure of the gold nucleus was 49 helium nuclei, which would have given it a mass of 196 and charge of 98. This differed enough from gold’s “atomic number” (at that time merely its place number in the periodic table) that Rutherford did not formally suggest the two numbers might be exactly the same

Successor model

The Rutherford model of the atom was soon superseded by the Bohr model, which used some of the early quantum mechanical results to give locational structure to the behavior of the orbiting electrons, confining them to certain circular (and later elliptical) planar orbits. In the Bohr model, expanding on the work of Henry Moseley, the central charge was identified as being directly connected with the atomic number (that is, the element’s place on the periodic table). Since the Bohr model is an improvement on the Rutherford model in this and other ways, some sources combine the two, referring to the Bohr model as the Rutherford-Bohr model. However, even an atom with a core containing an atomic number of charges was the work of a number of men, including those mentioned, and also lesser known workers such as Antonius Van den Broek.

The Rutherford model was important because it essentially proposed the concept of the nucleus, although this word is not used in the paper. What Rutherford notes as the (probable) concomitant of this results, is a “concentrated central charge” in the atom: “Considering the evidence as a whole, it seems simplest to suppose that the atom contains a central charge distributed through a very small volume, and that the large single deflexions are due to the central charge as a whole, and not to its constituents.” The central charge containing most of the atom’s positive charge, invariably later become associated with a concrete structure, the atomic nucleus.

After the Rutherford model and its confirmation in the experiments of Henry Moseley and its theoretical description in the Bohr model of the atom, the study of the atom branched into two separate fields, nuclear physics, which studies the nucleus of the atom, and atomic physics which studies atom’s electronic structure.


Shield of the U.S. Atomic Energy Commission.

Shield of the U.S. Atomic Energy Commission.

Despite its inaccuracy, the Rutherford model caught the imagination of the public in a way that the more correct Bohr model did not, and has continually been used as a symbol for atoms and atomic energy. Examples of its use over the past century include:

Key points of Rutherford model

  • The electron clouds of the atom do not influence alpha scattering.
  • A large number of the atom’s charges, up to a number equal to about half the atomic mass in hydrogen units, are concentrated in very small volume at the center of the atom. These are responsible for deflecting both alpha and beta particles.
  • The mass of heavy atom s such as gold is mostly concentrated in the central charge region, since calculations show it is not deflected or moved by the high speed alpha particles, which have very high momentum in comp arison to electrons, but not with regard to heavy atoms (such as gold) on the whole. This suggests that much of the mass of atoms is concentrated in their centres.

The Rutherford model’s contribution to modern science

After this discovery, scientists started to realize that the atom is not ultimately a single particle, but is made up of far smaller subatomic particles. Later workers began research to figure out the exact atomic structure which led to Rutherford’s gold foil experiment. They eventually discovered that atoms have a positively-charged nucleus (with an exact atomic number of charges) in the center, with a radius of about 1.2 x 10-15 meters x [Atomic Mass Number]1/3. Since electrons were found to be even smaller, this meant that the atom consists of mostly empty space.

Later on, scientists found the expected number of electrons (the same as the atomic number) in an atom by using X-ray beams. When an X-ray passes through an atom, some of it is scattered, while the rest passes through the atom. Since (in many cases with X-rays of the proper frequency) the X-ray loses its intensity primarily due to electron scattering, by noting the rate of decrease in X-ray intensity, the number of electrons contained in an atom could be estimated accurately.