1. Chapter 6 Electronic Structure of Atoms

2. Waves To understand the electronic structure of atoms, one must understand the nature of electromagnetic radiation. The distance between corresponding points on adjacent waves is the wavelength ( ).

3. Waves The number of waves passing a given point per unit of time is the frequency ( ). For waves traveling at the same velocity, the longer the wavelength, the smaller the frequency.

4. Electromagnetic Radiation All electromagnetic radiation travels at the same velocity: the speed of light (c), 3.00  10 8 m/s. Therefore, c =

5. The Nature of Energy The wave nature of light does not explain how an object can glow when its temperature increases. Max Planck explained it by assuming that energy comes in packets called quanta.

6. The Nature of Energy Einstein used this assumption to explain the photoelectric effect. He concluded that energy is proportional to frequency: E = h where h is Planck’s constant, 6.63  10 −34 J-s.

7. The Nature of Energy Therefore, if one knows the wavelength of light, one can calculate the energy in one photon, or packet, of that light: c = E = h

8. The Nature of Energy Another mystery involved the emission spectra observed from energy emitted by atoms and molecules.

9. The Nature of Energy One does not observe a continuous spectrum, as one gets from a white light source. Only a line spectrum of discrete wavelengths is observed.

10. The Nature of Energy Niels Bohr adopted Planck’s assumption and explained these phenomena in this way: 1.Electrons in an atom can only occupy certain orbits (corresponding to certain energies).

11. The Nature of Energy Niels Bohr adopted Planck’s assumption and explained these phenomena in this way: 2.Electrons in permitted orbits have specific, “allowed” energies; these energies will not be radiated from the atom.

12. The Nature of Energy Niels Bohr adopted Planck’s assumption and explained these phenomena in this way: 3.Energy is only absorbed or emitted in such a way as to move an electron from one “allowed” energy state to another; the energy is defined by E = h

13. The Nature of Energy The energy absorbed or emitted from the process of electron promotion or demotion can be calculated by the equation:  E = −R H ( ) 1nf21nf2 1ni21ni2 - where R H is the Rydberg constant, 2.18  10 −18 J, and n i and n f are the initial and final energy levels of the electron.

14. The Wave Nature of Matter Louis de Broglie posited that if light can have material properties, matter should exhibit wave properties. He demonstrated that the relationship between mass and wavelength was = h mv

15. The Uncertainty Principle Heisenberg showed that the more precisely the momentum of a particle is known, the less precisely is its position known: In many cases, our uncertainty of the whereabouts of an electron is greater than the size of the atom itself! (  x) (  mv)  h4h4

16. Quantum Mechanics Erwin Schrödinger developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated. It is known as quantum mechanics.

17. Quantum Mechanics The wave equation is designated with a lower case Greek psi (  ). The square of the wave equation,  2, gives a probability density map of where an electron has a certain statistical likelihood of being at any given instant in time.

18. Quantum Numbers Solving the wave equation gives a set of wave functions, or orbitals, and their corresponding energies. Each orbital describes a spatial distribution of electron density. The principal quantum number, n, describes the energy level on which the orbital resides.  The values of n are integers ≥ 0. We will not deal with the other three quantum numbers at this time.

19. s Orbitals Spherical in shape. Radius of sphere increases with increasing value of n.

20. p Orbitals Have two lobes with a node between them.

21. d Orbitals Four of the five orbitals have 4 lobes; the other resembles a p orbital with a doughnut around the center.

22. Energies of Orbitals For a one-electron hydrogen atom, orbitals on the same energy level have the same energy. That is, they are degenerate.

23. Energies of Orbitals As the number of electrons increases, though, so does the repulsion between them. Therefore, in many- electron atoms, orbitals on the same energy level are no longer degenerate.

24. Pauli Exclusion Principle No two electrons in the same atom can have exactly the same energy. We say that they have opposite spins, and represent them with arrows pointed up or down in orbital diagrams.

25. Electron Configurations Distribution of all electrons in an atom Consist of  Number denoting the energy level

26. Electron Configurations Distribution of all electrons in an atom Consist of  Number denoting the energy level  Letter denoting the type of sublevel/orbitals

27. Electron Configurations Distribution of all electrons in an atom. Consist of  Number denoting the energy level.  Letter denoting the type of sublevel/orbitals.  Superscript denoting the number of electrons in those orbitals.

28. Orbital Diagrams Each box represents one orbital. Half-arrows represent the electrons. The direction of the arrow represents the spin of the electron.

29. Hund’s Rule “For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.”

30. Periodic Table We fill orbitals in increasing order of energy. Different blocks on the periodic table, then correspond to different types of orbitals.

31. Some Anomalies Some irregularities occur when there are enough electrons to half- fill s and d orbitals on a given row.

32. Some Anomalies For instance, the electron configuration for copper is [Ar] 4s 1 3d 5 rather than the expected [Ar] 4s 2 3d 4.

33. Some Anomalies This occurs because the 4s and 3d orbitals are very close in energy. These anomalies occur in f-block atoms, as well.

34. Periodic Properties of the Elements

35. Development of Periodic Table Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped.

36. Development of Periodic Table Mendeleev, for instance, predicted the discovery of germanium (which he called eka- silicon) as an element with an atomic weight between that of zinc and arsenic, but with chemical properties similar to those of silicon.

37. Periodic Trends In this section, we will rationalize observed trends in  Sizes of atoms  Sizes of ions.  Ionization energy.  Electronegativity.

38. Effective Nuclear Charge In a many-electron atom, electrons are both attracted to the nucleus and repelled by other electrons. The nuclear charge that an electron experiences depends on both factors. The effective nuclear charge, Z eff, is found this way: Z eff = Z − S where Z is the atomic number and S is a screening constant, usually close to the number of inner electrons.

39. Sizes of Atoms The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.

40. Sizes of Atoms Bonding atomic radius tends to… …decrease from left to right across a row due to increasing Z eff. …increase from top to bottom of a column due to increasing value of n

41. Sizes of Ions Ionic size depends upon:  Nuclear charge.  Number of electrons.  Orbitals in which electrons reside.

42. Sizes of Ions Cations are smaller than their parent atoms.  The outermost electron is removed and repulsions are reduced.

43. Sizes of Ions Anions are larger than their parent atoms.  Electrons are added and repulsions are increased.

44. Sizes of Ions Ions increase in size as you go down a column.  Due to increasing value of n.

45. Sizes of Ions In an isoelectronic series, ions have the same number of electrons. Ionic size decreases with an increasing nuclear charge.

46. Ionization Energy Amount of energy required to remove an electron from the ground state of a gaseous atom or ion.  First ionization energy is that energy required to remove first electron.  Second ionization energy is that energy required to remove second electron, etc.

47. Ionization Energy It requires more energy to remove each successive electron. When all valence electrons have been removed, the ionization energy takes a quantum leap.

48. Trends in First Ionization Energies As one goes down a column, less energy is required to remove the first electron.  For atoms in the same group, Z eff is essentially the same, but the valence electrons are farther from the nucleus.

49. Trends in First Ionization Energies Generally, as one goes across a row, it gets harder to remove an electron.  As you go from left to right, Z eff increases.

50. Trends in First Ionization Energies However, there are two apparent discontinuities in this trend.

51. Trends in First Ionization Energies The first occurs between Groups IIA and IIIA. Electron removed from p-orbital rather than s- orbital  Electron farther from nucleus  Small amount of repulsion by s electrons.

52. Trends in First Ionization Energies The second occurs between Groups VA and VIA.  Electron removed comes from doubly occupied orbital.  Repulsion from other electron in orbital helps in its removal.

53. Electronegativity: The ability of atoms in a molecule to attract electrons to itself. On the periodic chart, electronegativity increases as you go…  …from left to right across a row.  …from the bottom to the top of a column.