1. Prentice Hall © 2003Chapter 6 Chapter 6 Electronic Structure of Atoms CHEMISTRY The Central Science 9th Edition David P. White

2. Prentice Hall © 2003Chapter 6 All waves have a characteristic wavelength,, and amplitude, A. The frequency,, of a wave is the number of cycles which pass a point in one second. The speed of a wave, v, is given by its frequency multiplied by its wavelength: For light, speed = c. The Wave Nature of Light

3. Prentice Hall © 2003Chapter 6 The Wave Nature of Light

5. Prentice Hall © 2003Chapter 6 Modern atomic theory arose out of studies of the interaction of radiation with matter. Electromagnetic radiation moves through a vacuum with a speed of 2.99792458  10 -8 m/s. Electromagnetic waves have characteristic wavelengths and frequencies. Example: visible radiation has wavelengths between 400 nm (violet) and 750 nm (red). The Wave Nature of Light

7. Prentice Hall © 2003Chapter 6 The Wave Nature of Light

8. Prentice Hall © 2003Chapter 6 Planck: energy can only be absorbed or released from atoms in certain amounts called quanta. The relationship between energy and frequency is where h is Planck’s constant (6.626  10 -34 J.s). To understand quantization consider walking up a ramp versus walking up stairs: For the ramp, there is a continuous change in height whereas up stairs there is a quantized change in height. Quantized Energy and Photons

9. Prentice Hall © 2003Chapter 6 The Photoelectric Effect and Photons The photoelectric effect provides evidence for the particle nature of light -- “quantization”. If light shines on the surface of a metal, there is a point at which electrons are ejected from the metal. The electrons will only be ejected once the threshold frequency is reached. Below the threshold frequency, no electrons are ejected. Above the threshold frequency, the number of electrons ejected depend on the intensity of the light. Quantized Energy and Photons

10. Prentice Hall © 2003Chapter 6 The Photoelectric Effect and Photons Einstein assumed that light traveled in energy packets called photons. The energy of one photon: Quantized Energy and Photons

12. Prentice Hall © 2003Chapter 6 Line Spectra Radiation composed of only one wavelength is called monochromatic. Radiation that spans a whole array of different wavelengths is called continuous. White light can be separated into a continuous spectrum of colors. Note that there are no dark spots on the continuous spectrum that would correspond to different lines. Line Spectra and the Bohr Model

14. Prentice Hall © 2003Chapter 6 Line Spectra Balmer: discovered that the lines in the visible line spectrum of hydrogen fit a simple equation. Later Rydberg generalized Balmer’s equation to: where R H is the Rydberg constant (1.096776  10 7 m -1 ), h is Planck’s constant (6.626  10 -34 J·s), n 1 and n 2 are integers (n 2 > n 1 ). Line Spectra and the Bohr Model

15. Prentice Hall © 2003Chapter 6 Bohr Model Rutherford assumed the electrons orbited the nucleus analogous to planets around the sun. However, a charged particle moving in a circular path should lose energy. This means that the atom should be unstable according to Rutherford’s theory. Bohr noted the line spectra of certain elements and assumed the electrons were confined to specific energy states. These were called orbits. Line Spectra and the Bohr Model

16. Prentice Hall © 2003Chapter 6 Bohr Model Colors from excited gases arise because electrons move between energy states in the atom. Line Spectra and the Bohr Model

17. Prentice Hall © 2003Chapter 6 Bohr Model Since the energy states are quantized, the light emitted from excited atoms must be quantized and appear as line spectra. After lots of math, Bohr showed that where n is the principal quantum number (i.e., n = 1, 2, 3, … and nothing else). Line Spectra and the Bohr Model

18. Prentice Hall © 2003Chapter 6 Bohr Model The first orbit in the Bohr model has n = 1, is closest to the nucleus, and has negative energy by convention. The furthest orbit in the Bohr model has n close to infinity and corresponds to zero energy. Electrons in the Bohr model can only move between orbits by absorbing and emitting energy in quanta (h ). The amount of energy absorbed or emitted on movement between states is given by Line Spectra and the Bohr Model

19. Prentice Hall © 2003Chapter 6 Bohr Model We can show that When n i > n f, energy is emitted. When n f > n i, energy is absorbed Line Spectra and the Bohr Model

20. Bohr Model Line Spectra and the Bohr Model

21. Prentice Hall © 2003Chapter 6 Limitations of the Bohr Model Can only explain the line spectrum of hydrogen adequately. Electrons are not completely described as small particles. Line Spectra and the Bohr Model

22. Prentice Hall © 2003Chapter 6 Knowing that light has a particle nature, it seems reasonable to ask if matter has a wave nature. Using Einstein’s and Planck’s equations, de Broglie showed: The momentum, mv, is a particle property, whereas is a wave property. de Broglie summarized the concepts of waves and particles, with noticeable effects if the objects are small. The Wave Behavior of Matter

23. Prentice Hall © 2003Chapter 6 The Uncertainty Principle Heisenberg’s Uncertainty Principle: on the mass scale of atomic particles, we cannot determine exactly the position, direction of motion, and speed simultaneously. For electrons: we cannot determine their momentum and position simultaneously. If  x is the uncertainty in position and  mv is the uncertainty in momentum, then The Wave Behavior of Matter

24. Prentice Hall © 2003Chapter 6 Schrödinger proposed an equation that contains both wave and particle terms. Solving the equation leads to wave functions. The wave function gives the shape of the electronic orbital. The square of the wave function, gives the probability of finding the electron, that is, gives the electron density for the atom. Quantum Mechanics and Atomic Orbitals

25. Prentice Hall © 2003Chapter 6 Quantum Mechanics and Atomic Orbitals

26. Prentice Hall © 2003Chapter 6 Orbitals and Quantum Numbers If we solve the Schrödinger equation, we get wave functions and energies for the wave functions. We call wave functions orbitals. Schrödinger’s equation requires 3 quantum numbers: 1.Principal Quantum Number, n. This is the same as Bohr’s n. As n becomes larger, the atom becomes larger and the electron is further from the nucleus. Quantum Mechanics and Atomic Orbitals

27. Prentice Hall © 2003Chapter 6 Orbitals and Quantum Numbers 2.Azimuthal Quantum Number, l. This quantum number depends on the value of n. The values of l begin at 0 and increase to (n - 1). We usually use letters for l (s, p, d and f for l = 0, 1, 2, and 3). Usually we refer to the s, p, d and f- orbitals. 3.Magnetic Quantum Number, m l. This quantum number depends on l. The magnetic quantum number has integral values between -l and +l. Magnetic quantum numbers give the 3D orientation of each orbital. Quantum Mechanics and Atomic Orbitals

28. Prentice Hall © 2003Chapter 6 Orbitals and Quantum Numbers Quantum Mechanics and Atomic Orbitals

29. Prentice Hall © 2003Chapter 6 Orbitals and Quantum Numbers Orbitals can be ranked in terms of energy to yield an Aufbau diagram. Note that the following Aufbau diagram is for a single electron system. As n increases, note that the spacing between energy levels becomes smaller. Quantum Mechanics and Atomic Orbitals

30. Prentice Hall © 2003Chapter 6 Orbitals and Quantum Numbers Quantum Mechanics and Atomic Orbitals

31. Prentice Hall © 2003Chapter 6 Orbitals and Quantum Numbers Quantum Mechanics and Atomic Orbitals

32. Prentice Hall © 2003Chapter 6 The s-Orbitals All s-orbitals are spherical. As n increases, the s-orbitals get larger. As n increases, the number of nodes increase. A node is a region in space where the probability of finding an electron is zero. At a node,  2 = 0 For an s-orbital, the number of nodes is (n - 1). Representations of Orbitals

34. Prentice Hall © 2003Chapter 6 The s-Orbitals Representations of Orbitals

35. Prentice Hall © 2003Chapter 6 The p-Orbitals There are three p-orbitals p x, p y, and p z. The three p-orbitals lie along the x-, y- and z- axes of a Cartesian system. The letters correspond to allowed values of m l of -1, 0, and +1. The orbitals are dumbbell shaped. As n increases, the p-orbitals get larger. All p-orbitals have a node at the nucleus. Representations of Orbitals

36. Prentice Hall © 2003Chapter 6 The p-Orbitals Representations of Orbitals

37. Prentice Hall © 2003Chapter 6 The d and f-Orbitals There are five d and seven f-orbitals. Three of the d-orbitals lie in a plane bisecting the x-, y- and z-axes. Two of the d-orbitals lie in a plane aligned along the x-, y- and z-axes. Four of the d-orbitals have four lobes each. One d-orbital has two lobes and a collar. Representations of Orbitals

39. Prentice Hall © 2003Chapter 6 Orbitals and Their Energies Orbitals of the same energy are said to be degenerate. For n  2, the s- and p-orbitals are no longer degenerate because the electrons interact with each other. Therefore, the Aufbau diagram looks slightly different for many-electron systems. Many-Electron Atoms

40. Orbitals and Their Energies

41. Prentice Hall © 2003Chapter 6 Electron Spin and the Pauli Exclusion Principle Line spectra of many electron atoms show each line as a closely spaced pair of lines. Stern and Gerlach designed an experiment to determine why. A beam of atoms was passed through a slit and into a magnetic field and the atoms were then detected. Two spots were found: one with the electrons spinning in one direction and one with the electrons spinning in the opposite direction. Many-Electron Atoms

42. Electron Spin and the Pauli Exclusion Principle Many-Electron Atoms

43. Prentice Hall © 2003Chapter 6 Electron Spin and the Pauli Exclusion Principle Since electron spin is quantized, we define m s = spin quantum number =  ½. :Pauli’s Exclusions Principle: no two electrons can have the same set of 4 quantum numbers. Therefore, two electrons in the same orbital must have opposite spins. Many-Electron Atoms

44. Prentice Hall © 2003Chapter 6 Electron Spin and the Pauli Exclusion Principle In the presence of a magnetic field, we can lift the degeneracy of the electrons. Many-Electron Atoms

46. Prentice Hall © 2003Chapter 6 Hund’s Rule Electron configurations tells us in which orbitals the electrons for an element are located. Three rules: electrons fill orbitals starting with lowest n and moving upwards; no two electrons can fill one orbital with the same spin (Pauli); for degenerate orbitals, electrons fill each orbital singly before any orbital gets a second electron (Hund’s rule). Electron Configurations

47. Prentice Hall © 2003Chapter 6 Condensed Electron Configurations Neon completes the 2p subshell. Sodium marks the beginning of a new row. So, we write the condensed electron configuration for sodium as Na: [Ne] 3s 1 [Ne] represents the electron configuration of neon. Core electrons: electrons in [Noble Gas]. Valence electrons: electrons outside of [Noble Gas]. Electron Configurations

48. Prentice Hall © 2003Chapter 6 Transition Metals After Ar the d orbitals begin to fill. After the 3d orbitals are full, the 4p orbitals being to fill. Transition metals: elements in which the d electrons are the valence electrons. Electron Configurations

49. Prentice Hall © 2003Chapter 6 Lanthanides and Actinides From Ce onwards the 4f orbitals begin to fill. Note: La: [Xe]6s 2 5d 1 4f 0 Elements Ce - Lu have the 4f orbitals filled and are called lanthanides or rare earth elements. Elements Th - Lr have the 5f orbitals filled and are called actinides. Most actinides are not found in nature. Electron Configurations

50. Prentice Hall © 2003Chapter 6 The periodic table can be used as a guide for electron configurations. The period number is the value of n. Groups 1A and 2A have the s-orbital filled. Groups 3A - 8A have the p-orbital filled. Groups 3B - 2B have the d-orbital filled. The lanthanides and actinides have the f-orbital filled. Electron Configurations and the Periodic Table

53. Prentice Hall © 2003Chapter 6 End of Chapter 6: Electronic Structure of Atoms