What is Electromagnetic radiation?

1. What is Electromagnetic radiation? Any type of radiation where light is adsorbed or emitted aka radiant energy Examples: Visible light, radio waves, X-rays, heat, etc.

2. Moves like 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 (l).

3. Waves • The number of waves passing a given point per unit of time is the frequency (u). • 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 × 108 m/s. • Therefore, c = lu

5. The Nature of Energy Max Planck explained it by assuming that energy comes in packets called quanta. Energy is some ratio of hv (2hv, 3hv) like stairs.

6. The Nature of Energy • Einstein used this assumption to explain the photoelectric effect. • He concluded that energy is proportional to frequency: E = hu where h is Planck’s constant, 6.626 × 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 = lu E = hu

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

9. The Nature of Energy • For atoms and molecules, 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: • 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: • 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: • Energy is only absorbed or emitted as to move an electron from one “allowed” energy state to another; the energy is defined by E = hu

13. The Nature of Energy The energy absorbed or emitted from the process of electron promotion or demotion can be calculated

14. The Uncertainty Principle Heisenberg showed that the more precisely the momentum of a particle is known, the less precisely is its position is known:

15. The Uncertainty Principle In many cases, our uncertainty of the whereabouts of an electron is greater than the size of the atom itself!

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

17. Quantum Mechanics • The wave equation is designated with a lowercase Greek psi (y). • The square of the wave equation, y2, 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 an area where the electron density is the highest. • An orbital is described by a set of three quantum numbers which are no longer a part of the AP syllabus.

19. s Orbitals • They are spherical in shape. • The radius of the sphere increases with the value of n.

20. s Orbitals Observing a graph of probabilities of finding an electron versus distance from the nucleus, we see that s orbitals possess n − 1 nodes, or regions where there is 0 probability of finding an electron.

21. p Orbitals • They have two lobes with a node between them.

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

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

24. 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.

25. Pauli Exclusion Principle • No two electrons in the same atom can have exactly the same energy. • Therefore, no two electrons in the same atom can have identical sets of quantum numbers.

26. Electron Configurations • This term shows the distribution of all electrons in an atom. • Each component consists of • A number denoting the energy level, 4p5

27. Electron Configurations • This term shows the distribution of all electrons in an atom • Each component consists of • A number denoting the energy level, • A letter denoting the type of orbital, 4p5

28. Electron Configurations • This term shows the distribution of all electrons in an atom. • Each component consists of • A number denoting the energy level, • A letter denoting the type of orbital, • A superscript denoting the number of electrons in those orbitals. 4p5

29. Orbital Diagrams • Each box in the diagram represents one orbital. • Half-arrows represent the electrons. • The direction of the arrow represents the relative spin of the electron.

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

31. Periodic Table • We fill orbitals in increasing order of energy. • Different blocks on the periodic table (shaded in different colors in this chart) correspond to different types of orbitals.

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

33. Some Anomalies For instance, the electron configuration for silver is [Ar] 5s1 4d10 rather than the expected [Ar] 5s2 4d9. Don’t be concerned about exceptions.

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

35. Periodic Trends • In this chapter, we will rationalize observed trends in • Sizes of atoms and ions. • Ionization energy. • Electron affinity.

36. 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.

37. Effective Nuclear Charge The effective nuclear charge, Zeff, is found this way: Zeff = Z − S where Z is the atomic number and S is a screening constant, usually close to the number of inner electrons.

38. What Is the Size of an Atom? The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei.

39. Sizes of Atoms The bonding atomic radius tends to — Decrease from left to right across a row (due to increasing Zeff). — Increase from top to bottom of a column (due to the increasing value of n).

40. Sizes of Ions • Ionic size depends upon • The nuclear charge. • The number of electrons. • The orbitals in which electrons reside.

41. Sizes of Ions • Cations are smaller than their parent atoms: • The outermost electron is removed and repulsions between electrons are reduced.

42. Sizes of Ions • Anions are larger than their parent atoms” • Electrons are added and repulsions between electrons are increased.

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

44. Ionization Energy • The ionization energy is the amount of energy required to remove an electron from the ground state of a gaseous atom or ion. • The first ionization energy is that energy required to remove the first electron. • The second ionization energy is that energy required to remove the second electron, etc.

45. 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.

46. 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, Zeff is essentially the same, but the valence electrons are farther from the nucleus.

47. 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, Zeff increases.

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

49. Trends in First Ionization Energies • The first occurs between Groups 2 and 3. • In this case the electron is removed from a p orbital rather than an s orbital. • The electron removed is farther from the nucleus. • There is also a small amount of repulsion by the s electrons.

50. Trends in First Ionization Energies • The second discontinuity occurs between Groups 15 and 16. • The electron removed comes from a doubly occupied orbital. • Repulsion from the other electron in the orbital aids in its removal.