1. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Electromagnetic Radiation Radiant energy that exhibits wavelength-like behavior and travels through space at the speed of light in a vacuum.

2. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 2 Waves Waves have 3 primary characteristics: 1.Wavelength: distance between two peaks in a wave. 2.Frequency: number of waves per second that pass a given point in space. 3.Speed: speed of light is 2.9979  10 8 m/s.

3. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 3 Figure 7.1 The Nature of Waves

4. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 4 Wavelength and frequency can be interconverted. = c/ = frequency (s  1 ) = wavelength (m) c = speed of light (m s  1 )

5. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 5 Figure 7.2 Classification of Electromagnetic Radiation

6. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 6 Planck’s Constant  E = change in energy, in J h = Planck’s constant, 6.626  10  34 J s = frequency, in s  1 = wavelength, in m Transfer of energy is quantized, and can only occur in discrete units, called quanta.

7. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 7 Energy and Mass Energy has mass E = mc 2 E = energy m = mass c = speed of light

8. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 8 Energy and Mass (Hence the dual nature of light.)

9. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 9 Figure 7.4 Electromagnetic Radiation

10. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 10 Figure 7.5 The Constructive and Destructive Interference of Waves

11. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 11 Wavelength and Mass = wavelength, in m h = Planck’s constant, 6.626  10  34 J s = kg m 2 s  1 m = mass, in kg = frequency, in s  1 de Broglie’s Equation

12. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 12 Atomic Spectrum of Hydrogen Continuous spectrum: Contains all the wavelengths of light. Line (discrete) spectrum: Contains only some of the wavelengths of light.

13. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 13 Figure 7.6 A Continuous Spectrum (a) and A Hydrogen Line Spectrum (b)

14. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 14 The Bohr Model E = energy of the levels in the H-atom z = nuclear charge (for H, z = 1) n = an integer The electron in a hydrogen atom moves around the nucleus only in certain allowed circular orbits.

15. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 15 The Bohr Model Ground State: The lowest possible energy state for an atom (n = 1).

16. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 16 Energy Changes in the Hydrogen Atom  E = E final state  E initial state

17. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 17 Figure 7.7 A Change between Two Discrete Energy Levels

18. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 18 Figure 7.8 Electronic Transitions in the Bohr Model for the Hydrogen Atom

19. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 19 Quantum Mechanics Based on the wave properties of the atom  = wave function = mathematical operator E = total energy of the atom A specific wave function is often called an orbital.

20. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 20 Figure 7.10 The Hydrogen Electron Visualized as a Standing Wave Around the Nucleus

21. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 21 Heisenberg Uncertainty Principle x = position mv = momentum h = Planck’s constant The more accurately we know a particle’s position, the less accurately we can know its momentum.

22. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 22 Probability Distribution 4 square of the wave function 4 probability of finding an electron at a given position Radial probability distribution is the probability distribution in each spherical shell.

23. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 23 Figure 7.13 Two Representations of the Hydrogen 1s, 2s, and 3s Orbitals

24. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 24 Figure 7.14 The Boundary Surface Representations of All Three 2p Orbitals

25. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 25 Figure 7.16 The Boundary Surfaces of All of the 3d Orbitals

26. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 26 Figure 7.17 Representation of the 4f Orbitals in Terms of Their Boundary Surfaces

27. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 27 Figure 7.18 Orbital Energy Levels for the Hydrog en Atom

28. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 28 Quantum Numbers (QN) 1.Principal QN (n = 1, 2, 3,...) - related to size and energy of the orbital. 2.Angular Momentum QN (l = 0 to n  1) - relates to shape of the orbital. 3.Magnetic QN (m l = l to  l) - relates to orientation of the orbital in space relative to other orbitals. 4.Electron Spin QN (m s = + 1 / 2,  1 / 2 ) - relates to the spin states of the electrons.

29. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 29 Pauli Exclusion Principle In a given atom, no two electrons can have the same set of four quantum numbers (n, l, m l, m s ). Therefore, an orbital can hold only two electrons, and they must have opposite spins.

30. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 30 Figure 7.19 A Picture of the Spinning Electron

31. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 31 Figure 7.20 A Comparison of the Radial Probability Distributions of the 2s and 2p Orbitals

32. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 32 Aufbau Principle As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to these hydrogen-like orbitals.

33. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 33 Figure 7.21 The Radial Probability Distribution for the 3s, 3p, and 3d Orbitals

34. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 34 Hund’s Rule The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli principle in a particular set of degenerate orbitals.

35. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 35 Valence Electrons The electrons in the outermost principle quantum level of an atom. Inner electrons are called core electrons.

36. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 36 Broad Periodic Table Classifications Representative Elements (main group): filling s and p orbitals (Na, Al, Ne, O) Transition Elements: filling d orbitals (Fe, Co, Ni) Lanthanide and Actinide Series (inner transition elements): filling 4f and 5f orbitals (Eu, Am, Es)

37. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 37 Figure 7.23 Mendeleev’s Early Periodic Table, Published in 1872

38. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 38 Figure 7.24 The Electron Configurations in the Type of Orbital Occupied Last for the First 18 Elements

39. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 39 Figure 7.25 Electron Configurations for Potassium Through Krypton

40. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 40 Figure 7.26 The Orbitals Being Filled for Elements in Various Parts of the Periodic Table

41. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 41 Figure 7.27 The Periodic Table With Atomic Symbols, Atomic Numbers, and Partial Electron Configura tions

42. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 42 Figure 7.29 The Order in which the Orbitals Fill in Polyelectronic Atoms

43. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 43 Ionization Energy The quantity of energy required to remove an electron from the gaseous atom or ion.

44. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 44 Figure 7.31 The Values of First Ionization Energy for the Elements in the First Six Periods

45. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 45 Periodic Trends First ionization energy: increases from left to right across a period; decreases going down a group.

46. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 46 Figure 7.32 Trends in Ionization Energies for the Representative Elements

47. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 47 Electron Affinity The energy change associated with the addition of an electron to a gaseous atom. X(g) + e   X  (g)

48. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 48 Figure 7.33 The Electronic Affinity Values for Atoms Among the First 20 Elements that Form Stable, Isolated X - Ions

49. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 49 Periodic Trends Atomic Radii: decrease going from left to right across a period; increase going down a group.

50. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 50 Figure 7.34 The Radius of an Atom

51. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 51 Figure 7.35 Atomic Radii for Selected Atoms

52. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 52 Information Contained in the Periodic Table 1.Each group member has the same valence electron configuration (these electrons primarily determine an atom’s chemistry). 2.The electron configuration of any representative element. 3.Certain groups have special names (alkali metals, halogens, etc). 4.Metals and nonmetals are characterized by their chemical and physical properties.

53. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 53 Figure 7.36 Special Names for Groups in the Periodic Table

54. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 54 Isoelectronic species Two or more molecular entities (atoms, molecules, ions) are described as being isoelectronic with each other if they have the same number of electrons or the same electron configuration and the same structure (number and connectivity of atoms), regardless of the nature of the elements involved.molecular entitiesatoms moleculesions

55. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 55 Valence isoelectronic The term valence isoelectronic is used when the above conditions are met but with only the same number of valence electrons.valence electrons

56. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 56 What it means.The statement "These compounds or molecules are isoelectronic" is not just an implemetation of the above definition. It has significance by fact that calculations on molecules and electron density, and therefore capability of reaction, have been performed on quite a lot of common substances. By identifying a new, rare or odd compound as being isoelectronic with an already known one offers clues to possible properties and reactions.

57. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 57 Radius of Isoelectronic species Atoms with the same number of valence electrons are arranged in radius smallest charge to largest charge. Sc 3+ Sc 3+, Ca 2+, K +, Ar, Cl −, S 2-, and P 3− are all isoelectronic.Ca 2+K +ArCl −S 2-P 3− In such monatomic cases, there is a clear trend in the sizes of such species, with atomic radius decreasing as charge increases.monatomicatomic radiuscharge