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Electronic Structures of Atoms / Periodic Trends / Ionic Bonding / Solids / Phase Changes

Electronic Structures of Atoms / Periodic Trends / Ionic Bonding / Solids / Phase Changes. H Advanced Chemistry Unit 3. Objectives #1-3 Atomic Theory. *review of electromagnetic radiation characteristics: (diagrams). Examples of Electromagnetic Radiation. Objectives #1-3 Atomic Theory.

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Electronic Structures of Atoms / Periodic Trends / Ionic Bonding / Solids / Phase Changes

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  1. Electronic Structures of Atoms / Periodic Trends / Ionic Bonding / Solids / Phase Changes H Advanced Chemistry Unit 3

  2. Objectives #1-3 Atomic Theory *review of electromagnetic radiation characteristics: (diagrams)

  3. Examples of Electromagnetic Radiation

  4. Objectives #1-3 Atomic Theory frequency, wavelength, energy frequency vs. wavelength (inverse relationship) frequency vs. energy (direct relationship) wavelength vs. energy (inverse relationship) c=fl(c = speed of light in m/s, f = frequency in Hz (1/s), l = wavelength in m)

  5. Max Planck (1858-1947)

  6. Objectives #1-3 Atomic Theory E = hf or hc/l h = Planck’s Constant (energy for waves) 6.626 X 10-34 Js

  7. Albert Einstein (1879-1955)

  8. Objectives #1-3 Atomic Theory E = mc2 (energy for particles) *Wave particle-duality Matter has wave and particle characteristics; acts as particle when interacting with matter; acts as wave when traveling through space (clip)

  9. Louis de Broglie (1892-1987)

  10. Derivation of de Broglie’s Equation: Ewaves = Eparticles hc / l = mc2 l (hc / l) = (mc2)l h = mcl l = h / mv (examples)

  11. #1 Calculate the frequency of light having a wavelength of 6.50 X 102 nm. c = fl (convert nm to meters) 6.50 X 102 nm X 1m / 1 X 109 nm = 6.50 X 10-7 m f = (3.00 X 108 m/s / 6.5 X 10-7 m) = 4.62 X 1014 Hz

  12. #2 Calculate the energy of the blue color of wavelength 4.50 X 102 nm emitted by an atom of copper. E = hc / l = (6.626 X 10-34 Js) (3.00 X 108 m/s) / 4.5 X 10-7 m = 4.42 X 10-19 J

  13. #3 Calculate the wavelength for an electron having a mass of 9.11 X 10-31 kg and traveling at a speed of 1.0 X 107 m/s. Calculate the wavelength for a ball having a mass of .10 kg and traveling at a speed of 35 m/s. • = h/mv = (6.626 X 10-34 Js) / (9.11 X 10-31 kg) (1.0 X 107 m/s) 7.27 X 10-11 m

  14. l = h / mv = (6.626 X 10-34 Js) / (.10 kg) (35 m/s) = 1.9 X 10-34 m

  15. Objectives #1-3 Atomic Theory *Work Function (Photoelectric Effect) Φ = hfo Φ = work function minimum energy required to remove electron from surface of metal fo = threshold frequency minimum frequency required to remove electrons from surface of metal (clip, diagram and examples to follow)

  16. Photoelectric Effect (Albert Einstein Nobel Prize 1921)

  17. #1 A gold strip is irradiated with radiation of a frequency of 6.9 X 1012 Hz. Calculate the energy of this radiation and determine if it is sufficient to cause electrons to be released from the metal. The work function of gold is 7.7 X 10-19 J. E = hf = 6.626 X 10-34 Js X 6.9 X 1012 Hz = 4.57 X 10-21 J; no electron ejection

  18. #2 The ionization energy of gold is 890.1 kJ/mole of electrons. Calculate the threshold frequency required to cause the photoelectric effect and eject an electron. Φ = hfo (convert I.E. into energy per electron) 890.1 kJ/mole X 1000 J / 1 kJ X 1 mole electrons / 6.02 X 1023 electrons = 1.479 X 10-18 J fo = 1.47 X 10-18 J / 6.626 X 10 -34 Js = 2.232 X 1015 Hz

  19. Niels Bohr (1892-1962)

  20. Objectives #1-3 Atomic Theory *Bohr’s Equation: E = -2.178 X 10-18 J (z2/n2) OR ∆E = -2.178 X 10-18 J (z2) X (1/n2final – 1/n2initial) used for: determining energy changes when electrons change energy levels for hydrogen; z = 1 (clip and examples)

  21. #1 Calculate the energy required to excite an electron from energy level 1 to energy level 3. ΔE = -2.178 X 10-18 J (1/9 – 1/1) = -2.178 X 10-18 J (-.8889) = 1.936 X 10-18 J

  22. #2 Calculate the energy required to remove an electron from a hydrogen atom. *this energy represents the ionization energy for hydrogen ΔE = -2.178 X 10-18 J (-1) =2.178 X 10-18 J

  23. Johannes Rydberg (1854-1919)

  24. Objectives #1-3 Atomic Theory *Rydberg Equation: 1/λ = 1/91 nm (1/nL2 – 1/nH2) *used for: determining wavelength of photons released when electrons change energy levels (Examples)

  25. #1 Calculate the wavelength of light released when an electron drops from the following states: 2 to 1: 1 / l = 1 /91 nm (1/1 – 1/4) = .75/91 nm l = 91nm/.75 = 121.33 nm

  26. 4 to 1: 1 / l = 1 /91 nm (1/1 – 1/16) = .9375/91 nm l = 91nm/.9375 = 97.07 nm

  27. 6 to 1: 1 / l = 1 /91 nm (1/1 – 1/36) = .9722/91 nm l = 91nm/.9722 = 93.60 nm

  28. *relationships of answers: the greater the energy difference, the smaller the wavelength

  29. Erwin Schrodinger (1887-1961)(clip)

  30. Objectives #4-5 The Quantum Numbers and Quantum States *Review of Quantum Theory: • Quantum Numbers • Principle (n) *energy level of shell of electron *n = 1,2,3….. *(old system) n = K, L, M, …. *indicates the number of sublevels in energy level

  31. Illustration of Principle Quantum Number

  32. Objectives #4-5 The Quantum Numbers and Quantum States • Orbital (l) *indicates orbital shape *l = 0, n-1 *s, p, d, f

  33. Illustration of Orbital Quantum Number / Orbital Shapes

  34. Objectives #4-5 The Quantum Numbers and Quantum States • Magnetic (ml) *indicates orientation of orbital in space *ml = 0, +/-l *the number of ml values indicate the number of orbitals within sublevel

  35. Illustration of Magnetic Quantum Number

  36. Objectives #4-5 The Quantum Numbers and Quantum States • Spin (ms) *indicates spin of electron *+1/2 or -1/2 *allows for up to 2 electrons per orbital *s 2 electrons p 6 electrons d 10 electrons f 14 electrons

  37. Illustration of Spin Quantum Number

  38. Objectives #4-5 The Quantum Numbers and Quantum States *Quantum Number Sets for Electrons in Atoms:

  39. Illustration of Quantum States

  40. Objectives #4-5 The Quantum Numbers and Quantum States (examples of quantum number states problems)

  41. Objectives #7-9 Electron Configurations of Ions / Orbital Filling and Periodic Trends *valence electrons and occasionally the electrons contained within the d sublevel are involved in chemical bonding *atoms tend to lose or gain electrons in such a way to complete octets (s2p6) or to form similarly stable arrangements called pseudo noble-gas configurations (clip and examples)

  42. Objectives #7-9 Electron Configurations of Ions / Orbital Filling and Periodic Trends *Orbital Filling and Periodic Trends • Ionization Energy (from period 2) Group 1 Li 5.4 ev Group 2 Be 9.3 ev (spike) Group 15 N 14.5 ev (spike) Group 17 F 17.4 ev (spike) Group 18 Ne 21.6 ev (spike)

  43. Trends in Ionization Energy

  44. The effective nuclear charge on a particular electron in an atom is less than the actual nuclear charge of the atom due to the screening or shielding effect of the inner core electrons. For example, the effective nuclear charge on the 3s electrons in magnesium is lessthan the 2p electrons because there is less nuclear charge acting on the 3s electron. Since the 2p electrons are attracted more strongly, they are more stable, and have less energy than the 3s electron

  45. Objective #10-12 Formation of the Ionic Bond / Born-Haber Cycle and Lattice Energy *ionic bonds involve the transfer of valence electrons from a metal to a nonmetal *the tendency for a metal to lose electrons depends on its ionization energy and the tendency of a nonmetal to gain electrons depends on its electron affinity *the loss of an electron requires a gain of energy and is therefore an endothermic process example: Na + energy › Na+1 + e- *the gain of an electron releases energy and is therefore an exothermic process example: Cl + e- › Cl-1 + energy

  46. Formation of Sodium Chloride

  47. Formation of Crystal Lattice

  48. Objective #10-12 Formation of the Ionic Bond / Born-Haber Cycle and Lattice Energy *combinations of elements with low ionization energies and high electron affinities will cause an extremely exothermic reaction and generally be the most stable *example: Na(s) + Cl2(g) › NaCl(s) + energy *the energy produced when the ionic bond forms is referred to as the lattice energy; this energy is also equal to the energy required to break apart the ionic bond *chemical bonding not only involves a rearrangement of electrons but it also involves changes in energy (clip)

  49. Illustration of Born-Haber Cycle

  50. Objective #10-12 Formation of the Ionic Bond / Born-Haber Cycle and Lattice Energy *the formation of an ionic compound; such as the following reaction: Na(s) + 1/2Cl(2)(g) › NaCl(s) + ∆Hof = -410.9 kJ where ∆Hof refers to the standard heat of formation which is the energy change involved when a compound is formed from its elements, involves a series of energy changing steps known as the Born-Haber cycle *these steps are as follows:

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