Chapter 7: Electron Structure of the Atom

1. Chapter 7: Electron Structure of the Atom

2. Questions for Consideration • What is light and how can we describe it? • How can we describe the behavior of the electron in a hydrogen atom? • How can we describe electrons in all atoms? • How does the arrangement of electrons in atoms relate to the arrangement of elements in the periodic table? • Which electrons are chemically important in an atom? • How does the electron arrangement in ions differ from that in atoms? • What properties of atoms are related to their electron arrangements?

3. Chapter 7 Topics: • Electromagnetic Radiation and Energy • The Bohr Model of the Hydrogen Atom • The Modern Model of the Atom • Periodicity of Electron Configurations • Valence Electrons for the Main-Group Elements • Electron Configurations for Ions • Periodic Properties of Atoms

4. Light • What is light, and what does it have to do with the electrons in an atom? Figure 7.3 Figure 7.1 Figure 7.2

5. White Light • White light is composed of different colors that can be separated by a prism • Water acts as a prism for sunlight, giving the effect of a rainbow. • Sources of white light • Sun • Regular (incandescent) light bulbs Figure 7.1

6. 7.1 Electromagnetic Radiation and Energy • Light is electromagnetic radiation. • All forms of electromagnetic radiation travel through space as waves that move at the speed of light: 3.0×108 m/s. • Light is characterized by wavelength and frequency. Figure 7.5

7. Differentiating the Kinds of Electromagnetic Radiation • Wavelength (λ) • The distance between two corresponding points on a wave • Units are meters (m) or commonly nanometers (nm = 10-9 m) Figure 7.5

8. Differentiating the Kinds of Electromagnetic Radiation • Frequency (ν) • A measure of the number of wave cycles that move through a point in space in 1 s • Units are hertz (Hz) which are the same as inverse seconds (1/s) Figure 7.6

9. Electromagnetic Spectrum • Forms of electromagnetic radiation include: • gamma rays, X-rays, ultraviolet, visible, infrared, microwaves, and radio waves. Figure 7.6

10. Wavelength and Frequency • As wavelength increases, frequency decreases. • Frequency is in units of cycles per second, 1/s or s1. Figure 7.6

11. Wavelengths of Visible Light Figure from p. 263 ROY G. BIV

12. Wavelength (), Frequency (), and Energy (Ephoton) h = 6.626 × 10-34 J.s Figure 7.7

13. Duality of Light • Light exists as both waves and particles (or packets of light called photons) • Characteristics of waves c = λν • Frequency (ν) • Wavelength (λ) • Characteristic of photons • Energy of a photon is directly proportional to the frequency and inversely proportional to the wavelength. Ephoton = energy of the photon (in Joules), h = Planck’s constant (6.626 x 10-34 Js), ν= frequency (in Hz or s1)

14. λ, ν, and Ephoton • For c =λν, • we can solve for λ: λ = c/ν • we solve for ν: ν = c/λ • For Ephoton = hν • We can substitute c / λ for ν, giving us the equation: Ephoton = hc/λ • This equation shows the inverse proportionality between Ephotonand λ (wavelength). • The longer the wavelength the smaller the photon energy

15. Activity: Frequency, Wavelength, and the Electromagnetic Spectrum • Which has a greater frequency, the red light or the green light? Greater wavelength? Greater photon energy? Figure 7.7

16. Activity: Calculating Wavelength, Frequency, and Photon Energy • What is the frequency and photon energy of the laser light used for eye surgery if its wavelength is 710 nm? (watch your units) • What category does this light fall into? ν = 4.23 × 1014 s-1 Ephoton= 2.80 × 10-19 J visible light

17. Activity: λ, ν, and Ephoton • If the wavelength of a microwave beam is 11.5 cm, then what are the frequency of the radiation and the energy of its photons?

18. Activity Solutions: λ, ν, and Ephoton • If the wavelength of a microwave beam is 11.5 cm, then what are the frequency of the radiation and the energy of its photons? λ = 11.5 cm First, convert to meters: λ = 11.5 cm x (1 m/100 cm) = 1.15 x 10-3 m ν = c / λ ν = (3.0 x 108 m/s)/1.15 x 10-3 m = 2.6 x 1011 s-1

19. Activity Solutions: λ, ν, and Ephoton • If the wavelength of a microwave beam is 11.5 cm, then what are the frequency of the radiation and the energy of its photons? Then use the frequency to calculate photon energy: E = hν = (6.626 x 10-34 Js)(2.6 x 1011 s-1) = 1.7 x 10-23 J

20. Atomic Spectra • When visible light passes through a prism, its components separate into a spectrum. • White light, such as sun light or light from a regular light bulb, gives a continuous spectrum: Figure 7.8A

21. Atomic Spectra • Colored light, such as that from a neon sign, gives only specific colors in what is called a line spectrum. • Each line represents a specific wavelength of light: Figure 7.8B This spectrum is produced when an element has been heated or subjected to an electric charge. Why do only certain lines appear and what causes these specific colors of light to be emitted?

22. Activity: Atomic Spectra • The hydrogen line spectrum contains four colored lines in the visible region: violet, blue, green, and red. • Which color has the longest wavelength? • Which color has the highest frequency? • Which color has the largest photon energy? red violet violet

23. 7.2 The Bohr Model of the Hydrogen Atom • The hydrogen line spectrum contains four visible lines: Figure 7.10

24. The Bohr Model of the Hydrogen Atom • Max Planck first hypothesized that energy produced by atoms can only have certain values and is therefore quantized. • That’s the reason why only distinct lines are seen in element line spectra. Energy is quantized and can only exist at certain wavelengths. Figure 7.8B

25. Bohr’s Model of the Hydrogen Atom • Niels Bohr, a student of Rutherford, studied the line spectra of the hydrogen atom to try to understand the electrons in the nuclear model of the atom. • Bohr came up with a planetary model, in which electrons orbit the nucleus in circular pathways. • His model was based on the idea that electrons and their energies are quantized: they can have only certain values.

26. Bohr’s Planetary Model • The electron in the H atom can occupy a specific orbit. Its energy increases as its distance from the nucleus increases. • Bohr labeled the electron orbits with a number, starting with 1 closest to the nucleus and increasing as the orbits get further away from the nucleus. Figure 7.9

27. Electron transitions to lower energy levels result in the emission of light Figure 7.10

28. Bohr Model • When the electron moves from a higher-energy orbit to a lower-energy orbit, it loses energy equal to the difference in energy between the orbits: ΔE = EfinalEinitial • The energy is released as a photon: Ephoton = ΔE

29. Activity: Electron Transitions and Light • Which transition results in the emission of light with the greatest photon energy? • Longest wavelength? • Highest frequency? n = 6 to n = 2 n = 3 to n = 2 Figure 7.10 n = 6 to n = 2

30. Deficiencies of the Bohr Model • The Bohr model explains only the hydrogen atom and its spectrum. • A more complex model was needed to describe all atoms in general. • However, the Bohr model did give insight into the quantized behavior of the atom that was later better understood.

31. 7.3 The Modern Model of the Atom • The modern model of the atom is based on Schrödinger’s mathematical model of waves • This model describes electrons as occupying orbitals, not orbits. • Orbitals • Three dimensional regions in space where electrons are likely to be found; not a circular pathway • Principal energy level • Orbitals of similar size Figure 7.11 Figure 7.11

32. Probability Map for lowest-energy state of the electron in a H atom: • A probability map shows where the electron is most likely to be found. The closer the dots, the higher the probability. Figure 7.11

33. A typical representation of an s-orbital: Figure 7.11

34. Principal Energy Levels • Orbitals of similar size exist in the same principal energy level (n=1, 2, 3 …). Figure 7.12

35. Orbitals • There are four general types of orbitals: • s, p, d, and f Figure 7.13

36. As the principal energy level increases, the orbital extends further from the nucleus, as shown here for the first three s orbitals. Figure 7.14

37. p Orbitals • There are three p orbitals in a sublevel. They lie perpendicular to one another. Figure 7.15A

38. p Orbitals • p-orbitals overlap with their centers at the nucleus. Figure 7.15B

39. d Orbitals • There are five d orbitals in a sublevel. Figure 7.16

40. f Orbitals • How many f orbitals are in a sublevel? Figure 7.13

41. What orbitals (sublevels) exist at each principal energy level? • The first principal energy level (n=1): • a single s orbital • The second principal energy level (n=2): • 2s orbital and the 2p sublevel (three 2p orbitals) • The third principal energy level (n=3): • 3s orbital, the 3p sublevel (three 3p orbitals), and the 3d sublevel (five 3d orbitals) • The fourth principal energy level (n=4): • 4s orbital, the 4p sublevel (three 4p orbitals), and the 4d sublevel (five 4d orbitals), and seven 4f orbitals.

42. Orbital Diagrams • Orbital diagrams • Show the sublevels and orbitals that can exist at each principal energy level. • Each box represents an orbital. • Groups of boxes represent sublevels. Figure 7.17

43. Orbital Diagram for the Hydrogen Atom • In the hydrogen atom only, the sublevels within a principal energy level all have the same energy. Each box represents an orbital. Figure 7.17

44. Orbital Diagram for a Multielectron Atom • In multielectron atoms, the sublevels within a principal energy level have different energies. • Sublevels within a principal energy level split so that • s < p < d < f Figure 7.18

45. Arrangement of Electrons • How are electrons arranged in an atom? • We will consider the ground state – the most stable arrangement of lowest energy.

46. Orbital Diagram for Carbon • What are some rules for arranging electrons in orbitals and sublevels? Electrons are represented by up and down arrows. Figure from p. 273 Carbon

47. Orbital Diagram for Carbon • Electrons occupy the lowest-energy orbitals first (aufbau principle). • No more than two electrons occupy each orbital (Pauli exclusion principle). Two electrons in the same orbital must have opposite spins, represented by the up and down arrows. Figure from p. 273

48. Helium’s Orbital Diagram Figure from p. 274

49. Lithium’s Orbital Diagram Figure from p. 274

50. Boron’s Orbital Diagram Figure from p. 274