1. Chapter 3 The Structure of the Atom In order to explain much of what is observed in chemistry, we need to adopt a model for the atom where the atom has components: Protons Neutrons Electrons (We do not need to account for any structure within these components.)

2. Rutherford’s Model of the Atom ● Mass of atom is concentrated in nucleus. ● Volume of atom is “empty.” ● Model is still used today.

3. Particles and Waves ● Light as a wave Wavelength, λ Speed, c Frequency, ν ● c=λν

4. Particles and Waves ● Particles have mass and volume. ● Waves have λ, ν and speed.

5. Light and Other Forms of Electromagnetic Radiation (EMR) ● Visible light is one form of EMR. ● Other forms include: Microwaves X rays Ultraviolet light ● In all cases c = νλ

6. Atomic Spectra ● Emission spectra are one type. ● Only discrete λ are emitted. ● Wavelengths of emitted light characteristic of element. ● Most interesting example was, and still is, hydrogen.

7. Atomic Spectra For hydrogen, it was discovered that the λ of the emitted light fit a simple equation: For hydrogen, it was discovered that the λ of the emitted light fit a simple equation:

8. The Wave-Packet Model of Electromagnetic Radiation ● Hydrogen spectrum raised many difficult questions. ● Solution to the problem required thinking about light as particles. ● These particles are called photons.

9. The Wave-Packet Model of Electromagnetic Radiation ● The energy of a photon is related to the frequency of the corresponding light wave. ● E = hν ● The constant of proportionality is called Planck’s constant, h. h has units of energy × time.

10. The Wave-Packet Model of Electromagnetic Radiation ● Energy levels of an atom are quantized. ● Emission occurs when an atom releases energy in the form of light as a photon: Atom(higher energy state) → Atom(lower energy state) + hν

11. The Bohr Model of the Atom ● Refinement of Rutherford Model. ● Mass at center. ● Electrons orbit the center like the moon around the Earth. ● The orbits are quantized.

12. The Bohr Model of the Atom ● Orbits ● E is proportional to radius ● Quantized orbits ● Absorption ● Emission

13. The Bohr Model of the Atom ● Fantastic achievement. ● Ushered in Quantum Mechanics. ● Only worked for 1 e - systems.

14. The Energy States of the Hydrogen Atom We already saw We now identify the two n i, n 1 and n 2, as labeling energy states of the hydrogen atom.

15. The Energy States of the Hydrogen Atom Figure 3.5

16. Electromagnetic Radiation and Color ● Primary additive colors RGB & color TVs Figure 3.6

17. Electromagnetic Radiation and Color ● Primary subtractive colors Absorbed (subtracted) color determines the observed color in transmission – A blue solution is absorbing yellow light (not blue light!). Figure 3.6

18. The First Ionization Energy ● Ionization: removing an electron. ● First Ionization: minimum amount of energy needed to accomplish this.

19. The First Ionization Energy Table 3.3

20. The First Ionization Energy Figure 3.7

21. The Shell Model ● Data in Figure 3.7 support a shell model of electrons surrounding a nucleus. ● Core charge

22. The Shell Model and Periodic Table ● Connection between Figure 3.7 and Periodic table. ● Valence electrons ● Core electrons

23. Photoelectron Spectroscopy (PES) and the Structure of Atoms ● PES uses light to ionize atoms, molecules or ions. ● Knowing the energy of the light, one can calculate the ionization energy IE = hv – KE.

24. Photoelectron Spectroscopy and the Structure of Atoms PES can remove an electron from any shell. Figure 3.14

25. Electron Configurations from PES ● PES reveals subshells. ● n for shell. ● s, p, d, f, … for subshell. ● Superscripts for occupation number.

26. Electron Configurations from PES Table 3.5

27. Allowed Combinations of Quantum Numbers ● Quantum numbers nl m l

28. Shells and Subshells of Orbitals ● The quantum numbers are used to label the shells and subshells. ● Three numbers are needed to specify an orbital. ● Each orbital can hold two electrons.

29. Orbitals and the Pauli Exclusion Principle ● There exists a fourth quantum number, m s. ● It can have one of two values: +½ or -½. ● Each electron in an atom has a set of 4 numbers: (n, l, m l, m s ). ● Pauli Exclusion Principle states that no two sets can be the same in a given atom.

30. Predicting Electron Configurations Figure 3.23

31. Electron Configurations and the Periodic Table ● Relationship between electron configurations and periodic table.

32. Electron Configurations and Hund’s Rule ● No effect on electron configuration. ● Important when drawing an orbital diagram. Which of these is correct for carbon and why?

33. Electron Configurations and Hund’s Rule ● The left one adheres to Hund’s rule.

34. The Sizes of Atoms: Metallic Radii Figure 3.25

35. The Sizes of Atoms: Covalent Radii Figure 3.26

36. The Relative Sizes of Atoms and Their Ions ● Cations are smaller than the atoms from which they originate. ● Anions are larger than the atoms from which they originate. Why?

37. Patterns in Ionic Radii ● Identical electron configurations come from isoelectronic species.

38. Second, Third, Fourth and Higher Ionization Energies ● Removal of an electron from a +1 cation is called the second ionization energy. ● Removal of an electron from a +2 cation is called the third ionization energy.

39. Average Valence Electron Energy ● AVEE ● Weighted average. ● Ionization energies used.

40. Average Valence Electron Energy ● Measure of attraction between electrons and nucleus. ● Measure of spacing of valence energy levels.

41. Average Valence Electron Energy AVEE can be used to characterize a material as metal or nonmetal. Metal: low AVEE, small separation. Nonmetal: high AVEE, large separation.

42. AVEE and Metallicity Figure 3.33