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CHEM 1300 Lecture Notes - Lecture 3: Photon, Junkers J 1, Niels Bohr

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3 Aug 2018

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= 1. 02 (cid:1005)(cid:1004) (cid:1005)(cid:1013)j (cid:3042)(cid:3047)(cid:3042)(cid:3041)= h(cid:2021) (cid:2019) = 7. 01 (cid:1005)(cid:1004) (cid:1011)m (cid:2021)(cid:2019)= c the electron. Sub in : e = (cid:1011). (cid:1004)(cid:1005) (cid:1005)(cid:1004) (cid:1011)(cid:3040) Binding energy = photon energy - kinetic energy. What is nu ? nu and mu sound really similar. E = h(cid:2021) (cid:2021)= 4. 3 (cid:1005)(cid:1004)(cid:1005)(cid:1008) (cid:1005) h = 6. 63 (cid:1005)(cid:1004) (cid:1007)(cid:1008)j s. 2. 8 (cid:1005)(cid:1004) (cid:1005)(cid:1013)j < 3. 7 (cid:1005)(cid:1004) (cid:1005)(cid:1013)j *so it does(cid:374)"t do a(cid:374)ythi(cid:374)g. E = h(cid:2021) (cid:2021)= 7. 5 (cid:1005)(cid:1004)(cid:1005)(cid:1008) (cid:1005) h = 6. 63 (cid:1005)(cid:1004) (cid:1007)(cid:1008)j s. Kinetic energy = 5. 0 (cid:1005)(cid:1004) (cid:1005)(cid:1013)j - 3. 7 (cid:1005)(cid:1004) (cid:1005)(cid:1013)j. The energy of electrons are described relative to a free, stationary electron which has 0 energy. The photoelectric effect focused on energies of a bound electron as compared to a free electron. Energy (cid:272)a(cid:374)(cid:374)ot disappea(cid:396) (cid:449)he(cid:374) it"s t(cid:396)a(cid:374)sfe(cid:396)(cid:396)ed to ato(cid:373). If the photon energy matches an energy difference in an atom, the photon can be absorbed.

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Related Questions

Titanium metal requires a photon with a minimum energy of 6.94 × 10^–19 J to emit electrons. (a) What is the minimum frequency of light necessary to emit electrons from titanium via the photoelectric effect? (b) What is the wavelength of this light? (c) Is it possible to eject electrons from titanium metal using visible light? (d) If titanium is irradiated with light of wavelength 233 nm, what is the maximum possible kinetic energy of the emitted electrons? (d) What is the maximum number of electrons that can be freed by a burst of light
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LIGHT LAB SPRING 2018 CHEM 1430 Ill. Quantization of Electromagnetic Radiation In 1894, the German physidist, Max Plank, was working for a utility company to maximize visible light output from incandescent lamps using minimum power. Based on theoretical considerations, Planck postulated that electromagnetic energy could only be emitted at discrete values, rather than the continuum of predicted by classical physics, as multiples of the frequency of the radiation. Planck's Postulate took the mathematical form: EÂ·hv where the energy emitted, in Joules, is the product of a constant, h, known as Planck's Constant, and the frequency of the radiation, v. Planck's Postulate was experimentally verified and h found to have the value of 6.626 x 1034 Js(Joule seconds). The idea that electromagnetic radiation energy had specific frequency dependent energy was termed a "quantum of action," a distressing finding for the dlassically trained Planck. Despite his objections, Planck adopted this quantization of energy as a last resort to reconcile experiment with theory. This discovery set the stage for the early understanding of atomic structure, led to the term "photon" to describe the packet of energy carried by light waves, led to an understanding of the photoelectric effect, eventually leading to the development of general quantum theory Planck's Relationship, in combination with the tight equation, c the energy of a photon from frequency or wavelength. Î» w, allow expressions to be written to calculate ACTIVITY 4: QUANTIZATION OF THE ENERGY OF RADIATION To practice cakculations with Planck's Relationship, complete the following table, showing work with units in the right mest column, Convert given frequency to energy or given wavelength to energy. Take the value for Planck's Constant, h, to be 6.626 x 10"Js, show a calculation string include dimensions, reporting the final value to the proper number of significant figures Spectral Region Wavelength Frequency Energy in J/photon Show Calculation String Î» in meters v in s or cps X-ray 6.0 x 10"m Blue light 4.5 x 10m Infrared 4.1 x 10" cps Microwave 7.3 x 101 cps Radio Wave 3.0x 10 m IV. Bohr's Quantum Model for the Hydrogen Atom Ernest Rutherford's Gold Foil Experiment of 1911 demonstrated the nuclear structure of the atom. Questions remained as to how electrons were arranged in the atom. The most common belief called for the negatively charged the late 19h century recognized a particle (an electron) traveling around a body (the positively charged nucleus) would continuously radiate energy, slowing over time, eventually falling into the nucleus. Matter would collapse. Page 8 of 12