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Rutherford model - Wikipedia, the free encyclopedia

The Rutherford model or planetary model is a model of the atom devised by Ernest Rutherford. Rutherford directed the famous Geiger-Marsden experiment in ...

Bohr model

In atomic physics, the Bohr model, devised by Niels Bohr, depicts the atom as a small, positively charged nucleus surrounded by electrons that travel in circular orbits around the nucleus—similar in structure to the solar system, but with electrostatic forces providing attraction, rather than gravity. This was an improvement on the earlier cubic model (1902), the plum-pudding model (1904), the Saturnian model (1904), and the Rutherford model (1911). Since the Bohr model is a quantum physics-based modification of the Rutherford model, many sources combine the two, referring to the Rutherford–Bohr model.

Introduced by Niels Bohr in 1913, the model's key success lay in explaining the Rydberg formula for the spectral emission lines of atomic hydrogen. While the Rydberg formula had been known experimentally, it did not gain a theoretical underpinning until the Bohr model was introduced. Not only did the Bohr model explain the reason for the structure of the Rydberg formula, it also provided a justification for its empirical results in terms of fundamental physical constants.

The Bohr model is a primitive model of the hydrogen atom. As a theory, it can be derived as a first-order approximation of the hydrogen atom using the broader and much more accurate quantum mechanics, and thus may be considered to be an obsolete scientific theory. However, because of its simplicity, and its correct results for selected systems (see below for application), the Bohr model is still commonly taught to introduce students to quantum mechanics, before moving on to the more accurate but more complex valence shell atom. A related model was originally proposed by Arthur Erich Haas in 1910, but was rejected. The quantum theory of the period between Planck's discovery of the quantum (1900) and the advent of a full-blown quantum mechanics (1925) is often referred to as the old quantum theory.

Origin

In the early 20th century, experiments by Ernest Rutherford established that atoms consisted of a diffuse cloud of negatively charged electrons surrounding a small, dense, positively charged nucleus. Given this experimental data, Rutherford naturally considered a planetary-model atom, the Rutherford model of 1911 – electrons orbiting a solar nucleus – however, said planetary-model atom has a technical difficulty. The laws of classical mechanics (i.e. the Larmor formula), predict that the electron will release electromagnetic radiation while orbiting a nucleus. Because the electron would lose energy, it would gradually spiral inwards, collapsing into the nucleus. This atom model is disastrous, because it predicts that all atoms are unstable.

Also, as the electron spirals inward, the emission would gradually increase in frequency as the orbit got smaller and faster. This would produce a continuous smear, in frequency, of electromagnetic radiation. However, late 19th century experiments with electric discharges through various low-pressure gases in evacuated glass tubes had shown that atoms will only emit light (that is, electromagnetic radiation) at certain discrete frequencies.

To overcome this difficulty, Niels Bohr proposed, in 1913, what is now called the Bohr model of the atom. He suggested that electrons could only have certain classical motions:

  1. The electrons can only travel in special orbits: at a certain discrete set of distances from the nucleus with specific energies.
  2. The electrons of an atom revolve around the nucleus in orbits. These orbits are associated with definite energies and are also called energy shells or energy levels. Thus, the electrons do not continuously lose energy as they travel in a particular orbit. They can only gain and lose energy by jumping from one allowed orbit to another, absorbing or emitting electromagnetic radiation with a frequency ν determined by the energy difference of the levels according to the Planck relation:\Delta{E} = E_2-E_1=h\nu \ , where h is Planck's constant.
  3. The frequency of the radiation emitted at an orbit of period T is as it would be in classical mechanics; it is the reciprocal of the classical orbit period: \nu = {1\over T}

The significance of the Bohr model is that the laws of classical mechanics apply to the motion of the electron about the nucleus only when restricted by a quantum rule. Although rule 3 is not completely well defined for small orbits, because the emission process involves two orbits with two different periods, Bohr could determine the energy spacing between levels using rule 3 and come to an exactly correct quantum rule: the angular momentum L is restricted to be an integer multiple of a fixed unit:

L = n{h \over 2\pi} = n\hbar

where n = 1, 2, 3, ... is called the principal quantum number, and ħ = h/2Ï€. The lowest value of n is 1; this gives a smallest possible orbital radius of 0.0529 nm known as the Bohr radius. Once an electron is in this lowest orbit, it can get no closer to the proton. Starting from the angular momentum quantum rule Bohr was a

Plum pudding model

The plum pudding model of the atom by J. J. Thomson, who discovered the electron in 1897, was proposed in 1904 before the discovery of the atomic nucleus. In this model, the atom is composed of electrons (which Thomson still called "corpuscles", though G. J. Stoney had proposed that atoms of electricity be called electrons in 1894) surrounded by a soup of positive charge to balance the electrons' negative charges, like negatively-charged "plums" surrounded by positively-charged "pudding". The electrons (as we know them today) were thought to be positioned throughout the atom, but with many structures possible for positioning multiple electrons, particularly rotating rings of electrons (see below). Instead of a soup, the atom was also sometimes said to have had a "cloud" of positive charge.

With this model, Thomson abandoned his earlier "nebular atom" hypothesis in which the atom was composed of immaterial vorticies. Now, at least part of the atom was to be composed of Thomson's particulate negative corpuscles, although the rest of the positively-charged part of the atom remained somewhat nebulous and ill-defined.

The 1904 Thomson model was disproved by the 1909 gold foil experiment, which was interpreted by Ernest Rutherford in 1911 to imply a very small nucleus of the atom containing a very high positive charge (in the case of gold, enough to balance about 100 electrons), thus leading to the Rutherford model of the atom. Finally, after Henry Moseley's work showed in 1913 that the nuclear charge was very close to the atomic number, Antonius Van den Broek suggested that atomic number is nuclear charge. This work had culminated in the solar-system-like (but quantum-limited) Bohr model of the atom in the same year, in which a nucleus containing an atomic number of positive charge is surrounded by an equal number of electrons in orbital shells.

Thomson's model was compared (though not by Thomson) to a British dessert called plum pudding, hence the name. Thomson's paper was published in the March 1904 edition of the Philosophical Magazine, the leading British science journal of the day. In Thomson's view: ... the atoms of the elements consist of a number of negatively electrified corpuscles enclosed in a sphere of uniform positive electrification, ...

In this model, the electrons were free to rotate within the blob or cloud of positive substance. These orbits were stabilized in the model by the fact that when an electron moved farther from the center of the positive cloud, it felt a larger net positive inward force, because there was more material of opposite charge, inside its orbit (see Gauss's law). In Thomson's model, electrons were free to rotate in rings which were further stabilized by interactions between the electrons, and spectra were to be accounted for by energy differences of different ring orbits. Thomson attempted to make his model account for some of the major spectral lines known for some elements, but was not notably successful at this. Still, Thomson's model (along with a similar Saturnian ring model for atomic electrons, also put forward in 1904 by Nagaoka after James C. Maxwell's model of Saturn's rings), were earlier harbingers of the later and more successful solar-system-like Bohr model of the atom.



From Encyclopedia

Atomic Structure Atomic Structure

The ancient Greek philosophers Leucippus and Democritus believed that atoms existed, but they had no idea as to their nature. Centuries later, in 1803, the English chemist John Dalton, guided by the experimental fact that chemical elements cannot be decomposed chemically, was led to formulate his atomic theory. Dalton's atomic theory was based on the assumption that atoms are tiny indivisible entities, with each chemical element consisting of its own characteristic atoms.✶ ✶See Atoms article for further discussion of Dalton's atomic theory. The atom is now known to consist of three primary particles: protons, neutrons, and electrons, which make up the atoms of all matter. A series of experimental facts established the validity of the model. Radioactivity played an important part. Marie Curie suggested, in 1899, that when atoms disintegrate, they contradict Dalton's idea that atoms are indivisible. There must then be something smaller than the atom (subatomic particles) of which atoms were composed. Long before that, Michael Faraday's electrolysis experiments and laws suggested that, just as an atom is the fundamental particle of an element, a fundamental particle for electricity must exist. The "particle" of electricity was given the name electron. Experiments with cathode-ray tubes, conducted by the British physicist Joseph John Thomson, proved the existence of the electron and obtained the charge-to-mass ratio for it. The experiments suggested that electrons are present in all kinds of matter and that they presumably exist in all atoms of all elements. Efforts were then turned to measuring the charge on the electron, and these were eventually successful by the American physicist Robert Andrews Millikan through the famous oil drop experiment. The study of the so-called canal rays by the German physicist Eugen Goldstein, observed in a special cathode-ray tube with a perforated cathode, let to the recognition in 1902 that these rays were positively charged particles (protons ). Finally, years later in 1932 the British physicist James Chadwick discovered another particle in the nucleus that had no charge, and for this reason was named neutron. As a physical chemist, George Stoney made significant contributions to our understanding of molecular motion. However, this Irish scientist is better known for assigning a name to negative atomic charges, electrons, while addressing the Royal Society of Dublin in 1891. —Valerie Borek Joseph John Thomson had supposed that an atom was a uniform sphere of positively charged matter within which electrons were circulating (the "plum-pudding" model). Then, around the year 1910, Ernest Ruthorford (who had discovered earlier that alpha rays consisted of positively charged particles having the mass of helium atoms) was led to the following model for the atom: Protons and neutrons exist in a very small nucleus, which means that the tiny nucleus contains all the positive charge and most of the mass of the atom, while negatively charged electrons surround the nucleus and occupy most of the volume of the atom. In formulating his model, Rutherford was assisted by Hans Geiger and Ernest Marsden, who found that when alpha particles hit a thin gold foil, almost all passed straight through, but very few (only 1 in 20,000) were deflected at large angles, with some coming straight back. Rutherford remarked later that it was as if you fired a 15-inch artillery shell at a sheet of paper and it bounced back and hit you. The deflected particles suggested that the atom has a very tiny nucleus that is extremely dense and positive in charge. Also working with Rutherford was Henry G. J. Moseley who, in 1913, performed an important experiment. When various metals were bombarded with electrons in a cathode-ray tube, they emitted X rays, the wavelengths of which were related to the nuclear charge of the metal atoms. This led to the law of chemical periodicity, which provided refinement of the periodic table introduced by Mendeleev in 1869. According to this law, all atoms of an element have the same number of protons in the nucleus. It is called the atomic number and is given the symbol Z. Hydrogen is the simplest element and has Z = 1. Through Rutherford's work it was known that that electrons are arranged in the space surrounding the atomic nucleus. A planetary model of the atom, with the electrons moving in circular orbits around the nucleus seemed an acceptable model. However, such a "dynamic model" violated the laws of classical electrodynamics, according to which a charged particle, such as an electron, moving in the positive electric field of the nucleus, should lose energy by radiation and eventually spiral into the nucleus. To solve this contradiction, in 1913, the Danish physicist Neils Bohr (then studying under Rutherford) postulated that the electron orbiting the nucleus could move only in certain orbits, having in each a certain "quantized" energy. It turns out that the colors in fireworks would help prove him right. The colorful lights of fireworks are emitted by "excited" atoms; that is, by atoms that have absorbed extra energy. Light consists of electromagnetic waves, each (monochromatic) color with a characteristic wavelength λ and frequency v. Frequency is related to energy E through the famous Planck equation, E = hν, where h is Planck's constant (6.6256 x 10−34 J s). Note that white light, such as sunlight, is a mixture of light of all colors, so it does not have a characteristic wavelength. For this reason we say that white light has a "continuous spectrum." On the other hand, excited atoms emit a "line spectrum" consisting of a set of monochromatic visible radiations. Each element has a characteristic line spectrum that can be used to identify the element. Note that line emission spectra can also be obtained by heating a salt of a metal with a flame. For instance, common salt (sodium chloride) provides a strong yellow light to the flame coming from excited sodium, while copper salts emit a blue-green light and lithium salts a red light. The colors of fireworks are due to this phenomenon. Scientists in the late nineteenth century tried to quantify the line spectra of the elements. In 1885 the Swedish school teacher Johann Balmer discovered a series of lines in the visible spectrum of hydrogen, the wavelengths of which could be related with a simple equation: in which λ is wavelength, k is constant, a = 2, and b = 3, 4, 5, … This group of lines was called the Balmer series. For the red line b = 3, for the green line b = 4, and for the blue line b = 5. Similar series were further discovered: in the infrared region, the Paschen series (with a = 3 and b = 4, 5 … in the above equation), and much later in the ultraviolet region, the Lyman series (with a = 1 and b = 2, 3 …). In 1896 the Swedish spectroscopist Johannes Rydberg developed a general equation that allowed the calculation of the wavelength of the red, green, and blue lines in the atomic spectrum of hydrogen: where nL is the number of the lower energy level to which an electron falls and nH is the number of the higher energy level from which it falls. R is called the Rydberg constant (1.0974 x 10−7 m−1). R was later shown to be 2π 2me 4Z2/h 3c, where m is the mass of the electron, e is its charge, Z is the atomic number, h is Planck's constant, and c is the speed of light. As noted earlier, Bohr had suggested the quantization of Ruthford's model of the atom. Although he was not aware of the work of Balmer and Paschen when he wrote the first version of his 1913 article, he had incorporated Planck's constant h into his model, which turned out to be an important decision. Bohr assumed that the absorption or emission of radiation can occur only by "jumps" of the electron from one stationary orbit to another. (See Figure 1.) The energy differences between two such allowed orbits then provided the characteristic frequencies of the emitted light. ΔE = E n1 − E n2 = hν Planck's constant h was named by Bohr the "quantum of action." Bohr's theory was in close agreement with many experimental facts regarding one-electron atoms (the hydrogen


From Yahoo Answers

Question:What can be said about the Rutherford model of an atom based on Newton's laws of motion, the laws of thermodynamics, and the nature of electromagnetic radiation? 1. All of these 2. The electrons are accelerating, so they would be giving off energy. 3. Continuous source of energy must be supplied to the atom. 4. Rutherford model of the atom could not work 5. None of these

Answers:Rutherfords model of the atom couldnt work.

Question:1) They show the exact location of electrons and protons 2)They show empty space between the nucleus and electrons 3)They contain electrons and protons 4)They contain energy levels for electrons which one?

Answers:2. Rutherford's model only showed that there was a nucleus, not protons. Thomson's just showed that electrons existed and that in between was empty space.

Question:a) Calculate the energy of an electron in the second Bohr orbit of a hydrogen atom. b) If the energy difference between the ground state of an atom and its excited stage is 5.4 x 10 (to the power of -19), what is the wavelength of the photon required producing this transition.

Answers:a) The energy of an electron in any orbit is given by En = -13.6/n^2 EV Therefore, energy in second orbit is E2 = - 13.6/4EV = - 3.4 EV b) delta E = 5.4 x 10^-19 (given) Using the formula delta E = hc/Wavelength Rearranging, we have wavelength = hc/deltaE =6.6x10^-34 x 3x10^8 /5.4x10^-19 = 3.66x10^-7m/s

Question:I need an IN-DETAIL description here. I need to make a model of both, but all of the pictures and descriptions I have found so far seem to look exactly the same. WHAT IS THE DIFFERENCE?

Answers:rutherford model just says that whole mass and positive charge of electron is concentrated at its nucleaus the electrons move around nucleaus in elliptical orbit in the same way as planets move around sun modern electron cloud theory says (in addition to above theory ) that trajectory of electron is impossible to determine this theory is completely based on schrodinger's equation. it tells about the most possible arrangements of electron in an atom

From Youtube

Structure of the Atom 3: The Rutherford Model :Another informative video from the Senior physics series describing the Rutherford model. More vids- www.shep.net More info- en.wikipedia.org

Chemistry: The Rutherford Experiment :www.mindbites.com In this lesson, Professor Harman explains the discovery of the nucleus and Neutrons. The nucleus of atoms was discovered using radioactivity, which is the spontaneous emissions of particles of radiation fron an atom. Prof. Harman talks about early experiments with radioactivity, the discovery of gamma, alpha, and beta particles and their characteristics. Alpha particles were instrumental in the discovery of the nucleus of atoms. The Rutherford Gold Foil Experiment used alpha particles aimed at gold foil. The vast majority of the particles went through the gold foil, but approximately 1 in every 8000 was deflected at a severe angle. Rutherford hypothesized that this was only possible if the majority of the atom's mass was held in one central location, which he deemed the 'nucleus.' It was detemined that the nucleus of an atom is very small, analogous to one lightbulb, if Las Vegas is an atom. This discovery led to a revision of the ""Plum Pudding"" model of an atom to the ""Planetary"" model. Professor Harman also talks about the Chadwick experiment that discovered Neutrons. There was mass in an atom still unaccounted for by protons and electrons. Commonly used experiments could not discover it, though, since Neutrons are neutral and all the experiments used a charge. After discussing the Chadwick experiment using Beryllium, Professor Harman revisits Dalton's Atomic Theory. Taught by Professor Harman, this lesson was selected from a broader, comprehensive ...