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CHEM103 Lecture Notes - Lecture 3: Uncertainty Principle, Molecular Electronic Transition, Emission Spectrum

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12 Sep 2018

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CHEM103 verified notes

CHEM103 Lecture Notes - Lecture 2: Ion, Electromagnetic Spectrum, Electric Field

CHEM103 Lecture Notes - Lecture 3: Uncertainty Principle, Molecular Electronic Transition, Emission Spectrum

CHEM103 Lecture Notes - Lecture 4: Quantum Mechanics, Electron Density, Angular Momentum

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Chem 103 lecture 3 bohr"s model//the wave particle duality of matter and energy//quantum mechanics and the hydrogen atom. H atom is the simplest atom, it has 1e-, 1p+ and no neutrons. There were 2 issues with the rutherford model: atom is unstable according to classical physics, failed to explain the observation of line spectra. Bohr"s model blends classical and quantum theory: the electron occupies only certain energy levels and these are called stationary states. (cid:3041)= (cid:3027)2(cid:3041)2. Where: rh = rydberg"s constant for h atom = 2. 179 * 10-18 j. Energy scale n=1, e1 = (cid:3027)2(cid:3041)2 = rh = 2. 179 * 10-18 j n=2, e2 = - 5. 448 * 10-19 j n=3, e3 = - 2. 421 * 10-19 j n= , e = 0 j. In absorption, nf > ni e = positive. Problem: what is the frequency and the wavelength of the transition between ni = 5 and nf = 2 for hydrogen.

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Related Questions

CHEM 1430 LIGHT LAB SPRING 2018 A Danish physicist, Neils Bohr, a post-doctoral student in J. J. Thomson's lab and friend of Rutherford, took up the problem of electron arrangement within the atom. Bohr's breakthrough was recognizing Balmer's fixed emission wavelengths, when taken with Planck's Relationship, led to the conclusion the shells, like the radiation emitted, must be quantized (have set, specific energles). Bohr proposed the electrons circulated about the madeus, much like planets cirde surãin seres of concentric shells of increasing size. Each shell stood at a fixed (set) distance from the nucleus with the shell nearest the nucleus taking the m and a shell further away taking the n in the B-R equation. The integral values of m and n taken with the discrete wavelength led Bohr to propose the shells were "quantized," meaning electrons could only exist at the surface of the shells. When an electron moved from a high shell, n, to a low shell, m, light of a foxed wavelength was emitted Bohr reasoned the shells were special due to this quantized nature. An electron in a shell occupied what Bohr described as a "steady state," meaning the electron could circulate about the nucleus without loss or gain of energy. This steady state proposition, which conflicted with classical physics, could not be explained by then available data but was required for Bohr's proposed structure to be stable. Later work revealed the wave-particle duality of the electron which lead to acceptance of the Steady State Proposition. The energy of an electron in any shell, represented by n, can be calculated using an equation derived from combining the Balmer-Rydberg and Planck equations. The energy of an electron in the no shell is given by é¦¯.h RL where hcR,-2.18x10-*) where Ru, is the Rydberg's Constant for the hydrogen atom, h is Planck's Constant and c is the speed of light NOTE: The energy is given a nexative value to refect the stabillting, attractive force between the negatively charged electron and the positively charged nucleus. As the shell number n increases, the attractive force decreases leading to a rising shell energies (E, increases, values less negative). ACTIVITY Calculate the energles for an electron in the first 7 shells of the hydrogen atom using the above relationship, recording the values in the center column of the Hydrogen the right-hand column in the table, convert the shell with the exponent of 10 Improper sclentific notation is typically used to plot data covering several orders of magnitude (powers of ten). Shell Energy Table. For energies to improper scientific notation Hydrogen Shell Energy Table Shell Value, n Calculated Shell Energy.ã Express Shell Energy in improper Scientific Notation, x 10-1 Page 9 of 12