History of the Atomic Model

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1 Chapter 5 Lecture Chapter 5 Electronic Structure and Periodic Trends 5.1 Electromagnetic Radiation Learning Goal Compare the wavelength, frequency, and energy of electromagnetic radiation. Fifth Edition Democritus (400 B.C.) Believed that matter was composed of invisible particles of matter he called atoms. Antoine Lavoisier (1700 s) Law of Conservation of Mass Matter is not created or destroyed. Joseph Proust (1700 s) Law of constant composition compounds are composed of atoms in definite ratios. John Dalton (Late 1700 s) First atomic theory explaining chemical reactions J.J. Thomson (1897) Discovered the electron using cathode ray tubes and suggested the Plum Pudding model Hantaro Nagaok (1903) Postulated a "Saturnian" model of the atom with flat rings of electrons revolving around a positively charged particle. Robert Millikan (1909) Found the charge and mass of the electron in his famous oil-can experiment. Ernest Rutherford (1911) Discovered the nucleus in his famous gold foil experiment. Data from his experiments led Rutherford to propose a planetary model in which a cloud of electrons surrounded a small, compact nucleus of positive charge. Only such a concentration of charge could produce the electric field strong enough to cause the heavy deflection of alpha particles observed. Electromagnetic Radiation Electromagnetic radiation is energy that travels as waves through space is described in terms of wavelength and frequency moves at the speed of light in a vacuum speed of light = 3.0 x 10 8 m/s = wavelength ( ) x frequency ( ) Wavelength Wavelength (, lamda) is the distance from a peak in one wave to the peak in the next is expressed in meters (m) or nanometers (nm) 5 6 LecturePLUS Timberlake 1

2 Frequency Wavelength and Frequency Frequency (, nu) is the number of waves that pass by each second is expressed in hertz (Hz), which is equal to cycles per second of one megahertz is equal to 1 x 10 6 Hz 3.5 MHz = 3.5 x 10 6 Hz White light passing through a prism separates into all colors, which is a continuous spectrum Waves with different wavelengths have different frequencies. 7 8 Inverse Relationship of and The inverse relationship of wavelength and frequency means that longer wavelengths have lower frequencies shorter wavelengths have higher frequencies different types of electromagnetic radiation have different wavelengths and frequencies Electromagnetic Spectrum The electromagnetic spectrum arranges forms of energy from lowest to highest arranges energy from longest to shortest wavelengths shows visible light with wavelengths from nm 9 10 The short wavelengths of the color blue are dispersed more by the molecules in the atmosphere than longer wavelengths of visible light, which is why we say the sky is blue. If blue light has a wavelength of 450 nm, is the wavelength of red light greater or less than 450 nm? If the frequency of blue light is 6.3 x Hz, what is its frequency in kilohertz (khz) and in megahertz (MHz)? Visible light LecturePLUS Timberlake 2

3 1. Which of the following has the shortest wavelength? A. microwaves B. blue light C. UV light 2. Which of the following has the lowest energy? A. red light B. blue light C. green light Chapter 5 Lecture Chapter 5 Electronic Structure and Periodic Trends 5.2 Atomic Spectra and Energy Levels Fifth Edition 3. Which of the following has the highest frequency? A. radio waves B. infrared C. X-rays 13 Learning Goal Explain how atomic spectra correlate with the energy levels in atoms. Max Plank (1900) Quanta = minimal amount of energy that can be emitted or absorbed by electromagnetic radiation. Studying energy absorbed and emitted by hot glowing matter, He noted that energy is only released or absorbed in chunks of some minimal size he called quanta. Albert Einstein (1905) By observing the Photoelectric effect, Einstein proposed that a beam of light is not a wave propagating through space, but rather a collection of discrete wave packets(photons), each with energy hf. This shed light on the previous discovery of the Planck relation (E = hν) linking energy (E) and frequency (ν) as arising from quantization of energy. The factor h is known as the Planck constant. The Photoelectric Effect In opposition to Maxwell's theory that EMR energy is proportional to intensity, Einstein concluded that EMR energy is proportional to wave frequency: E photon = h where h is Planck s constant ( J-s.) Einstein's Wave-Particle Duality of Light Albert Einstein determined that E.M.R. is composed of packets of quantized energy called photons; each having its own characteristic wavelength and traveling at a constant speed, the speed of light (c), 3.00 X 10 8 m/s. Therefore, if one knows the wavelength of light, one can calculate the energy of one photon-wave, or quanta, of that light: c = E = h Reminder kg m 1 J 1 2 s 2 LecturePLUS Timberlake 3

4 The Emission Spectra of Elements One property of the elements that really captured the attention of scientists is that one does not observe a continuous spectrum for hydrogen, as one gets from a white light source. Only a line spectrum of discrete wavelengths is observed. Atomic Spectrum An atomic spectrum consists of lines of different colors formed when light from a heated element passes through a prism photons emitted when electrons drop to lower energy levels 20 Johann Balmer (1885) Showed that the wavelengths of the four visible lines of hydrogen fit a simple formula relationship. λ = hm 2 /(m 2 - n 2 ) 1 λ Johannes Rydberg (1888) (R Expanded Balmer s relationship to a more general equation that could be used to calculate all spectral lines of hydrogen, not only the visible, known as the Rydberg equation: H 1 1 ) 2 2 n1 n2 Niels Bohr (1913) Solidified Rutherford s Planetary atomic model by using the work of Max Plank and Albert Einstein on the nature of Electromagnetic Radiation to predict the spectral lines of hydrogen described by the work of Johann Balmer and Johhanes Rydberg. Electron Energy Levels Electrons are arranged in specific energy levels that are labeled n = 1, n = 2, n = 3, and so on increase in energy as n increases have the electrons with the lowest energy in the first energy level (n=1) closest to the nucleus 24 LecturePLUS Timberlake 4

5 Energy Level Changes An electron absorbs photons of a specific energy to jump to a higher energy level falls to a lower energy level by emitting photons of a specific energy DEMO Changes in Energy Levels An electron in the n = 3 energy level falls to the n = 2 energy level by emitting photons with energy equal to the energy difference of the two energy levels. An electron in the n = 5 energy level moves to the n = 2 energy level by releasing the energy equal to the energy difference of the fifth and second energy levels Energy Emitted In each of the following energy level changes, indicate if energy is 1) absorbed 2) emitted 3) not changed A. An electron moves from the first energy level (n =1 ) to the third energy level (n = 3). B. An electron falls from the third energy level to the second energy level. C. An electron moves within the third energy level. When electrons drop from a higher level to the first level, second level, and third level, photons of ultraviolet light, visible light, and infrared are emitted (not to scale) LecturePLUS Timberlake 5

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