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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.


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

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Answers:The Bohr model puts electrons in circular orbits around the nucleus that correspond only to the principle quantum #. It does not distinguish the subshells (s, p, d, etc). Carbon is 1s2 2s2 2p2 So in the Bohr model there would be two concentric circular orbits for n = 1 and n = 2 shells. The smaller n = 1 orbit would have two electrons and the larger n = 2 orbit 4 electrons. oxygen is 1s2 2s2 2p4 again two orbits; n = 1 with 2 electrons and n = 2 with 6 Oxygen has the greater electronegativity because it has a greater effective nuclear charge (Zeff)

Question:I need to do a bohr model of carbon and i have to label a valence electron, except I have no idea what it is/ means. I've tried searching in google images for a diagram labeled where it is but I cant seem to find one. Can anyone explain what it is or just post a link to a diagram labeled where the valence electron is. thanks you!

Answers:Here's a Bohr Model of Carbon: http://www.chemicalelements.com/elements/c.html Valence electrons are the electrons on the outermost shell. As you can see, Carbon has 4 valence electrons.

Question:http://physics.uoregon.edu/~jimbrau/BrauImNew/Chap04/FG04_11.jpg the second picture is that of carbon....i have to make a 3d project for science class is that the right model to follow after if not please give me ideas yeah i have to use pipecleaners and styrophome balls i sont think that picture is gonna worl

Answers:yeah that would be the right model to use. http://www.mbe.doe.gov/me70/manhattan/images/AtomLabeledLarge.gif is one thats more familiar. Are you going to do the styrofoam ball and pipecleaner route or otherwise?

Question:Which of the following transitions in the Bohr hydrogen atom leads to the absorption of a photon with the highest frequency? A) n = 1 n = 3 B) n = 5 n = 2 C) n = 4 n = 3 D) n = 2 n = 4

Answers:The transition that will release the most energy (highest frequency of light) is the jump from n=5 to n=2 as this transition is from a high orbital to a low orbital. You can think of the electron as being attached to a spring. When you pull it out to a high orbital like n=5 you have a lot of potential energy then when you let go and let it fall back to the lower n=2 orbital, you release a lot of kinetic energy in the form of a high frequency photon of light. Now, this question asks which transition leads to the absorption of a photon of high frequency. Well, this is the opposite of the above. Now you are putting energy in to force an electron to jump from a low orbit to a high orbit. Therefore, based on the available choices, the transition from n=1 to n=3 would require the most energy (if you look at the Bohr model you'll see that the distance between n=1 and n=3 is greater than the distance between n=2 and n=4). In order for the electron to make the jump, it would need to absorb a higher-frequency photon of light. Answer: n=1 to n=3 Hope that helps.

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