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Ch. 6 Electronic Structure and the Periodic Table Part 1: Light, Photon Energies, and Emission Spectra.

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Presentation on theme: "Ch. 6 Electronic Structure and the Periodic Table Part 1: Light, Photon Energies, and Emission Spectra."— Presentation transcript:


2 Ch. 6 Electronic Structure and the Periodic Table Part 1: Light, Photon Energies, and Emission Spectra

3 Compare the wave and particle natures of light. radiation: the rays and particles —alpha particles, beta particles, and gamma rays—that are emitted by radioactive material Define a quantum of energy, and explain how it is related to an energy change of matter. Contrast continuous electromagnetic spectra and atomic emission spectra.

4 electromagnetic radiation wavelength frequency amplitude electromagnetic spectrum Light, a form of electronic radiation, has characteristics of both a wave and a particle. quantum Planck's constant photoelectric effect photon atomic emission spectrum

5 Review We left off with Rutherford and Chadwick discovering nucleus and neutrons This proved that J.J. Thomson’s “plum pudding” model of the atom was incorrect. So where do we go next?

6 Rutherford’s model of the atom

7 Rutherford Model: Problems Nucleus surrounded by electrons.  How did e- fill up space surrounding a (+) nucleus? What prevented the electrons from being pulled right into the nucleus? To answer this, must understand relationship of light and electrons

8 The Atom and Unanswered Questions (cont.) In the early 1900s, scientists observed certain elements emitted visible light when heated in a flame. Analysis of the emitted light revealed that an element’s chemical behavior is related to the arrangement of the electrons in its atoms.

9 At Rutherford Model Electrons pictured as particles Light pictured as waves Discovered electrons have wave-like properties, and light has particle-like properties

10 What do we do? Describe electrons as having dual wave-particle nature (or properties)  Sometimes it acts like a particle  Sometimes it acts like a wave Has stood up against many experiments to prove it wrong Explains how electron isn’t pulled into nucleus.

11 Electromagnetic (EM) radiation Light as a wave  Form of Energy that exhibits wavelike behavior as it travels  Speed = 3.00 x 10 8 m/s (speed of light through air)

12 The Wave Nature of Light (cont.) The wavelength (λ) is the shortest distance between equivalent points on a continuous wave.wavelength The frequency (ν) is the number of waves that pass a given point per second.frequency The amplitude is the wave’s height from the origin to a crest.amplitude

13 The Wave Nature of Light (cont.)

14 Wavelength Wavelength : distance between corresponding points on a wave  λ = wavelength λ is the Greek letter lambda Wavelength is usually in nanometers (nm) or meters

15 Wavelength

16 Frequency (ν), the Greek letter nu (not Vee) Number of waves that pass a given point in a specific amount of time Frequency units are in Hertz (Hz) or 1/ seconds ( /s )

17 Relationship between wavelength and frequency c = λν  Where c = speed of light Correlation?  As wavelength decreases, frequency increases  As wavelength increases, frequency decreases  This is an inverse relationship

18 The Wave Nature of Light (cont.) The speed of light (3.00  10 8 m/s) is the product of it’s wavelength and frequency c = λν.

19 What is the frequency of a wave with wavelength of 100 nm?

20 The Particle Nature of Light The wave model of light cannot explain all of light’s characteristics. Matter can gain or lose energy only in small, specific amounts called quanta. A quantum is the minimum amount of energy that can be gained or lost by an atom.quantum

21 Max Planck (1900) Described light as having particle-like properties When hot object loses energy, it doesn’t do it continuously as it would if it were a wave Loses energy in form of a quanta

22 Quanta? Quantum – finite quantity of energy that can be gained or lost by an atom  Specific: if it costs $1.25 to get a soda from machine, and you give it $1.00, do you get a a soda? Photon – individual quantum of light

23 The Particle Nature of Light (cont.) The photoelectric effect is when electrons are emitted from a metal’s surface when light of a certain frequency shines on it.photoelectric effect

24 Einstein In 1905, said Planck’s work applied to all EM. Explains photoelectric effect –  must absorb photon with specific energy to dislodge an electron  When electron is dislodged, it must be in the form of a particle  But as it moves, (we see it as color), it is in the form of a wave Shows dual nature of light (wave and particle)

25 The Particle Nature of Light (cont.) Albert Einstein proposed in 1905 that light has a dual nature. A beam of light has wavelike and particle-like properties. A photon is a particle of electromagnetic radiation with no mass that carries a quantum of energy.photon E photon = h E photon represents energy. h is Planck's constant. represents frequency.

26 Energy of a Photon (or any wave of energy) E = h ν  E = energy ( in joules, j)  v = frequency  h = Planck's constant 6.63 x 10 -34 J * s (Joule Seconds)  As frequency goes up, what happens to the energy?

27 What is the energy with a wave of frequency 1 x 10 16 Hz?

28 Light through a prism Continuous spectrum  All wavelengths in a given range are included  Why we see rainbows  Separated by wavelength Electromagnetic spectrum  Consists of all electromagnetic radiation, arranged by increasing wavelengths

29 Light through a prism

30 Electromagnetic Spectrum

31 Hydrogen Atom Spectrum Pass high voltage through Hydrogen gas Gas glows, and you can pass this light through prism Creates a bright line spectrum or atomic emission spectra

32 Atomic Emission Spectra (cont.)

33 The atomic emission spectrum of an element is the set of frequencies of the electromagnetic waves emitted by the atoms of the element.atomic emission spectrum Each element’s atomic emission spectrum is unique.

34 Hydrogen’s atomic emission spectra

35 Each line caused by light of a different wavelength What causes the light?

36 Electrons, man Electrons get boosted by voltage from ground state (or normal state) to excited state. When they relax back down to ground (almost immediately), they give off certain amounts of energy Line spectrum: produced when an electron drops from a higher energy orbit to a lower energy orbit

37 Ground State vs. Excited State Ground State  The state of lowest energy of an atom Excited State  A state in which an atom has a higher potential energy than its ground state

38 What causes the lines? Each line = energy from electron as it drops from excited state to ground state Energy of photon = difference in energy between ground and excited state.

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