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Chapter 11

Chapter 11. Modern Atomic Theory. Light. Made up of electromagnetic radiation. Waves of electric and magnetic fields are at right angles to each other. Parts of a wave. Wavelength. l. Frequency = number of cycles in one second

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Chapter 11

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  1. Chapter 11 Modern Atomic Theory

  2. Light • Made up of electromagnetic radiation. • Waves of electric and magnetic fields are at right angles to each other.

  3. Parts of a wave Wavelength l Frequency = number of cycles in one second Measured in hertz 1 hertz = 1 cycle/second

  4. Frequency = n

  5. The Nature of Waves

  6. Kinds of EM waves • There are many different l and n • Radio waves, microwaves, x rays and gamma rays are all examples. • Light is only the part our eyes can detect. G a m m a R a d i o w a v e s R a y s

  7. Electromagenetic Spectrum

  8. The speed of light • In a vacuum is 2.998 x 108 m/s = c • c = ln Ex 1: What is the wavelength of light with a frequency 5.89 x 105 Hz? Ex 2: What is the frequency of blue light with a wavelength of 484 nm?

  9. In 1900 • Matter and energy were seen as different from each other in fundamental ways. • Matter was particles. • Energy could come in waves, with any frequency. • Max Planck found that as the cooling of hot objects couldn’t be explained by viewing energy as a wave.

  10. Energy is Quantized • Planck found DE came in chunks with size hn DE = nhn • where n is an integer (which we will ignore for our calculations at this point). • h is Planck’s constant=6.626 x 10-34 J•s • These packets of hn are called quantum

  11. Einstein is next • Said electromagnetic radiation is quantized in particles called photons. • Each photon has energy E= hn = hc/l

  12. Fireworks in Washington D.C.

  13. What are electrons doing in an atom?

  14. Which is it? • Is energy a wave like light, or a particle? Yes • Concept is called the Wave -Particle duality. • What about the other way, is matter a wave? Yes

  15. Electrons move as waves • A wave has three characteristics: wavelength, frequency, and speed. • Wavelength-distance between wave peaks • Frequency-how many peaks pass a point per a certain amount of time • Speed-how fast a peak travels through water.

  16. Spectrum • The range of frequencies present in light. • White light has a continuous spectrum. • All the colors are possible. • A rainbow.

  17. Hydrogen spectrum • Emission spectrum because these are the colors it gives off or emits. • Called a line spectrum. • There are just a few discrete lines showing 656 nm 434 nm 410 nm 486 nm

  18. What this means • Only certain energies are allowed for the hydrogen atom. • Can only give off certain energies. • Use DE = hn = hc / l • Energy in the in the atom is quantized.

  19. What will an electron do? • It has mass, so it is matter. • A particle can only go through one hole. • A wave through both holes. • An electron does go though both, and makes an interference pattern. • It behaves like a wave. • Other matter has wavelengths too short to notice.

  20. The Bohr Ring Atom-1st quantum model n = 4 n = 3 n = 2 n = 1

  21. The Bohr Model • Doesn’t work. • Only works for hydrogen atoms. • Electrons don’t move in circles. • The quantization of energy is right, but not because they are circling like planets.

  22. We are worried about the change • When the electron moves from one energy level to another. DE = Efinal – Einitial

  23. Electrons also travel as particles • n is the principal energy level • n = 1 is called the ground state and is the lowest energy state of an atom • Excited state-atom has a higher potential • When an excited atom returns to its ground state, it gives off energy in the form of a photon. • Photon-stream of tiny energy packets.

  24. Electromagnetic Radiation Exhibits Wave Properties and Particulate Properties

  25. (a)The Probability Distribution (b) The Probability Along a Line Drawn From the Nucleus Outward in Any Direction

  26. Louis deBroglie Louis deBroglie said electrons can be considered as waves that are fixed around a nucleus instead of orbits as Bohr stated. This is known as the wave mechanical model.

  27. Heisenburg The Heisenburg Uncertainty Principle states that it is impossible to determine the position and velocity (speed) of an electron.

  28. Quantum numbers • Angular momentum quantum number l , gives the shape of the orbital. • integer values from 0 to n-1 • l = 0 is called s • l = 1 is called p • l =2 is called d • l =3 is called f • l =4 is called g

  29. The Angular Momentum Quantum Numbers and Corresponding Letters Used to Designate Atomic Orbitals

  30. Electron Shapes • s (groups 1 and 2 including H and He) holds up to 2 electrons • p (groups 13-18) holds up to 6 electrons • d (groups 3-12) holds up to 10 electrons • f (lanthanides and actinides) holds up to 14 electrons

  31. The Orbitals Being Filled for Elements in Various Parts of the Periodic Table

  32. Representation of the 2p Orbitals (a) The Electron Probability Distribution for a 2p Oribtal (b) The Boundary Surface Representations of all Three 2p Orbitals

  33. Representation of the 3d Orbitals (a) Electron Density Plots of Selected 3d Orbitals (b) The Boundary Surfaces of All of the 3d Orbitals

  34. Quantum numbers • Magnetic quantum number (m) • Tells direction in each shape. • Can have 2 values; either +1/2 or -1/2

  35. A Picture of the Spinning Electron

  36. The Periodic Table • Developed independently by German Julius Lothar Meyer and Russian Dmitri Mendeleev (1870”s). • Didn’t know much about atom. • Put in columns by similar properties and mass. • Predicted properties of missing elements.

  37. Aufbau Principle • Aufbau is German for building up. • As the protons are added one by one, the electrons fill up hydrogen-like orbitals. • Fill up in order of energy levels.

  38. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s He with 2 electrons

  39. Details • Valence electrons- the electrons in the outermost energy levels (not d). • Core electrons- the inner electrons. • Hund’s Rule- The lowest energy configuration for an atom is the one have the maximum number of unpairedelectrons in the orbital. • C 1s2 2s2 2p2

  40. Exceptions • Cr =[Ar] 4s1 3d5 • Cu=[Ar] 4s1 3d10 • Half filled orbitals. • Scientists aren’t sure of why it happens

  41. Electron Configurations for Potassium Through Krypton

  42. Ionization Energy • Ionization Energy- The energy necessary to remove an electron from a gaseous atom. • Highest energy electron removed first. • It takes much more energy to remove a core electron than a valence electron because there is less shielding.

  43. Trends in Ionization Energies (kj/mol) for the Representative Elements

  44. Shielding • Electrons on the higher energy levels tend to be farther out. • Have to look through the other electrons to see the nucleus. • They are less effected by the nucleus.

  45. Across a Period • Generally from left to right, IE increases because there is a greater nuclear charge with the same shielding. • As you go down a group IE decreases because electrons are farther away.

  46. The Radius of an Atom (r) is Defined as Half the Distance Between the Nuclei in a Molecule Consisting of Identical Atoms

  47. Atomic Radii (in Picometers) for Selected Atoms

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