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Arrangement of Electrons in Atoms Part One Learning Objectives Read Pages 97-106 Asgn #16: 103/1-6 1

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Section 1 Light and Electrons Explain the mathematical relationship amount the speed, wavelength, and frequency of electromagnetic radiation. Discuss the dual wave-particle nature of light. Discuss the significance of the photoelecvtric effect and the line-emission spectrum of H to the development of the atomic model. Describe the Bohr model of the H atom. 2

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Section 2 Quantum Model Discuss De Broglie’s role in the development of the quantum model of the atom. Compare and contrast the Bohr model and the quantum model of the atom. Explain how the Heisenberg uncertainty principle and the Schrodinger wave equation led to the idea of atomic orbitals. 3

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Electromagnetic Radiation (EMR) or light is a form of energy that moves as a wave through space. Electromagnetic Spectrum is made up of many kinds of EMR: visible, X rays, ultraviolet, infrared, microwaves, and radiowaves. 4

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Electromagnetic (E-M) Waves (LIGHT!) Do not require a medium through which to travel Light travels at 3.0 x 10 8 m/s in a vacuum or air Its wavelength and frequency varies according to the type of E-M wave 5

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Higher frequency Greater energy More penetration What type has a.Greatest frequency? b.Less frequency than infrared light? 6

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c = c, the speed of light which is 3x10 8 m/s in a vacuum or air. Units: m/s wavelength or distance between corresponding points on adjacent waves – Units: m or nm frequency or number of waves passing a point in a given amount of time. Units: Hertz, Hz or 1/s or s -1 7

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Light Problems: What is the frequency of light whose wavelength is 600 nm? nm means 10 -9 m c = c 3x10 8 m = 5 x 10 14 s -1 600x10 -9 m s 8

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Photoelectric Effect This is the emission of electrons from a metal when electromagnetic radiation shines on the metal. P. E. shows that energy is emitted in small, specific packets called quanta. A quantum of energy is the minimum quantity of energy that can be lost or gained by an atom. The photoelectric effect showed that light behaves as particles, too! 9

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E = h E, energy of a quantum of radiation in joules, J h, Planck’s constant is 6.626 x 10 -34 Js, frequency in s -1 Problem: What is the frequency of a photon whose energy is 3.4 x 10 -19 J? = E/h = 3.4 x 10 -19 J / 6.626 x 10 -34 Js = 5.1 x 10 14 s -1 *Wavelength-frequency relationship was proposed by Planck in 1900. 10

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Einstein explained the photoelectric effect was due to metal absorbing energy in discrete amounts of photons. Ground state – lowest energy state of an atom Excited state – state where an atom has a higher potential energy than ground state. 11

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webexhibits.org 12

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Tutorvista.com 13

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Bohr Model 1913 – Niels Bohr proposed a hydrogen atom model where electrons circle the nucleus only in allowed paths or orbits with a definite amount of energy. If an electron absorbs energy, it can go to a higher level. If in a higher energy level, an electron can emit a certain amount of energy to move to a lower level. 14

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chemweb.ucc.ie 15

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Quantum Model of the Atom Questions were unanswered regarding how electrons could be particles yet they gave off waves of light. De Broglie suggested that electrons could be considered waves confined to space around a nucleus only at specific frequencies. Diffraction experiments proved that electron beams can interfere with each other and produce areas of low energy and high energy areas as a result of interference. 16

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Heisenberg Uncertainty Principle (1927) – it is impossible to determine simultaneously both the position and velocity of an electron or any other particle. Schrodinger Wave Equation (1926) - developed an equation that treated electrons in atoms as waves. Heisenberg and Schrodinger laid the foundation for mathematical descriptions of wave properties of very small particles such as electrons – the probable location of electrons around the nucleus. AKA: Quantum Theory and Quantum Numbers. 17

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End of Section 1 and part of Section 2 of Chapter 4 Arrangement of Electrons in Atoms 18

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