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The color of the light emitted depends upon the D E as the electron(s) move from higher to lower energy levels.

Chapter 6. The color of the light emitted depends upon the D E as the electron(s) move from higher to lower energy levels. . Ne. He. The shorter the wavelength ( ) , the higher the frequency ( ) . Energy of the wave increases as frequency increases. Electromagnetic spectrum.

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The color of the light emitted depends upon the D E as the electron(s) move from higher to lower energy levels.

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  1. Chapter 6 The color of the light emitted depends upon the DE as the electron(s) move from higher to lower energy levels. Ne He

  2. The shorter the wavelength (), the higher the frequency (). • Energy of the wave increases as frequency increases.

  3. Electromagnetic spectrum c = ln Converting from wavelength (l) to frequency (n) Ex: What is the frequency of a 400 nm EM wave? n = c/l = (3.00 x 108 m/s) / (4.00 x 10-7 m) = 7.50 x 1014 Hz

  4. 1 H 1 n = 1 n = 2 n = 3 n = 4 n = 5 Excited state: electron is above the lowest possible energy level Photon absorbed + Photon emitted Ground state: electron is at its lowest possible energy level

  5. A few definitions: Quantum: The smallest quantity of energy that can be absorbed or emitted as electromagnetic radiation Photon: (1) A whole number multiple of a quantum; (2) a ‘particle’ of light (from Einstein’s photoelectric effect) Monochromatic radiation: Electromagnetic radiation that consists of a single wavelength Spectrum: EM radiation that has been split into its separate wavelengths Continuous spectrum: A spectrum containing all possible wavelengths of EM radiation Line spectrum: A spectrum that does not contain all possible wavelengths of EM radiation

  6. Hydrogen Atomic emission (bright line) spectrum of hydrogen

  7. For emission: E = Ehigher – Elower = Ephoton = h Ephoton= hc/ • = frequency (Hz) h = Planck’s constant (6.626 x 10-34 J·s)  = wavelength (m) c = speed of light (3.00 x 108 m/s)

  8. n2 > n1 and both are positive RH = 1.096776 x 107 m-1 The Rydberg equation: used to calculate the wavelength of a photon emitted or absorbed by a hydrogen atom for specific values of n. Bohr modified Rydberg’s equation to calculate the energy levels available to an electron around a hydrogen atom: The more negative the energy value, the lower the energy of the electron in the orbit.

  9. The DE for an electronic transition can be found using: Notice that DE is negative when the electron goes from a higher to a lower energy level. Ex: What is the energy of the photon associated with an n = 4 to n = 2 electronic transition? = -4.09 x 10-19 J Q: Is energy absorbed or emitted? emitted Q: What is the wavelength of this photon? 4.86 x 10-7 m = 486 nm = 4860 Å

  10. If light can have characteristics of BOTH particles and waves… …can particles of matter also behave like waves? De Broglie’s equation h = 6.626 x 10-34 J·s m = mass in kg v = velocity in m/s Ex: What is the wavelength of an electron traveling at 5.97 x 106 m/s? (The mass of an electron is 9.109 x 10-31 kg) = 1.22 Å

  11. Heisenberg’s Uncertainty Principle: it is impossible to know precisely both the velocity (or energy) and the electron’s position at the same time. Given: me = 9.109 x 10-31 kg ve = 5 x 106 1% m/s Q: What is the uncertainty in the location of the electron? Dv = 1%(5 x 106 m/s) = 5 x 104 m/s = 1 x 10-9 m That’s an order of magnitude bigger than a hydrogen atom!

  12. The Four Quantum Numbers In the quantum mechanical model of the atom, each electron is described by four quantum numbers, and no two electrons in an atom can have the same numbers. The quantum numbers are: n: the principle quantum number. n = 1, 2, 3, 4… l: the angular quantum number. It describes the ‘shape’ of the orbital…the electronic distribution about the nucleus. l = n-1, n-2,… to 0 If l = 0, it’s an s orbital If l = 1, it’s a p orbital If l = 2, it’s a d orbital If l = 3, it’s an f orbital (and so on) ml: the magnetic quantum number. It describes the orientation of the orbital on an xyz-coordinate axis. ml = -l…0…+l in integral values. ms: the spin quantum number. It equals +½ or -½. Conventionally, the +½ is always given first.

  13. 1 1 4 4 2 3 3 2 n Orbital types One s-orbital One s-orbital Three p-orbitals One s-orbital Three p-orbitals Five d-orbitals One s-orbital Three p-orbitals Five d-orbitals Seven f-orbitals Within an energy level (n = 1, 2, 3, 4…), there exists n types of orbitals and n2 sublevels.

  14. A few examples: (all ms’s are ± ½, so won’t be shown here) n l ml Notes 0 0 The 1s orbital 1 2 1 -1, 0, 1 The three 2p orbitals. The 2s orbital 2 0 0 3 2 -2, -1, 0, +1, +2 The five 3d orbitals. 3 1 -1, 0, +1 The three 3p orbitals. 0 0 The 3s orbital 3 Notice that n = # of subshells present and that n2 = total number of orbitals for a particular energy level.

  15. z z z y y y x x x The three p-orbitals are oriented along the x, y and z axes and can hold a maximum of 2 electrons each. px pz py It’s been suggested that the p-orbitals look like peanuts if that helps you to remember their shape.

  16. dz d 2 x2-y2 The d-orbitals dxy dxz dyz Even though the d-orbitals look like two p-orbitals, it is important to remember that each orbital can hold a maximum of 2 electrons regardless of how many lobes it has.

  17. The freaky (or flowery) f-orbitals

  18. Examples 1. (a) What is the designation for the subshell with n = 5 and l = 1? (b) How many orbitals are in this subshell? (c) Indicate the values of ml for each of these orbitals. (a) l = 1 is a p-orbital, so it is a 5p subshell (b) There are three p-orbitals (c) ml = -1, 0, 1 3. (6.54) Which of the following are permissible sets of quantum numbers for an electron in a hydrogen atom: (a)n = 2, l = 1, ml = 1; (b)n = 1, l = 0, ml = −1; (c)n = 4, l = 2, ml = −2; (d)n = 3, l = 3, ml = 0? For those combinations that are permissible, write the appropriate designation for the subshell to which the orbital belongs (that is, 1s, and so on). (a) 2px (c) 4dx2-y2

  19. Orbital Diagrams And Electron Configurations

  20. ms = spin quantum number Electrons behave as though they are spinning on their axis. A half-arrow is used to indicate if the direction of the spin is up ( , ms=+½) or down ( , ms = -½). Electron configuration H: 1 electron , 1s1

  21. Pauli Exclusion Principle: No two electrons in the same orbital can have identical quantum numbers. Hence, each orbital can contain 2 electrons iff the electron spins are in opposite directions. He: 2 electrons , 1s2

  22. Auf bau principle: the electrons fill the orbitals starting with the lowest energy level and working their way up the energy ‘ladder’ Li: 3 electrons , 1s2 2s1

  23. B: 5 electrons , 1s2 2s2 2p1

  24. Hund’s Rule: must fill energetically degenerate orbitals so that maximum multiplicity results  “Fill ‘em up with singles before you start pairing ‘em up.” C: 6 electrons , 1s2 2s2 2p2

  25. n p-block elements 1 d-block metals 2 n -1 3 4 5 6 7 s-block metals f-block metals n -2

  26. 1 1 2 1 s s Hydrogen Helium

  27. 2 2 1 2 2 2 1 2 2 2 1 1 1 2 s s s s s s p Lithium Beryllium Boron

  28. Lithium Beryllium Boron 2 1 2 2 2 2 1 2 2 2 1 1 1 2 s s s s s s p

  29. Boron 2 2 2 2 2 2 2 2 3 2 4 2 2 2 1 2 2 2 2 1 1 1 1 2 s s s s p p p s s s s p Carbon Nitrogen Oxygen

  30. Oxygen 2 2 2 2 2 2 5 4 6 2 2 2 2 2 2 1 1 1 s s s p p p s s s Fluorine Neon

  31. 2 2 2 2 2 2 5 6 6 6 6 6 6 2 2 3 2 4 3 3 2 2 1 2 2 2 2 2 2 1 3 3 1 1 5 4 3 4 8 10 3 3 s s s d d p p p p p p p s s s s s s s s s Chlorine (Cl) Nickel (Ni) Rubidium (Rb)

  32. Rules to Drawing Electron Dot Structures • Electrons 1-4 are added singly. • Exception: If there only two electrons, they’re shown as a pair. • Electrons 5-8 are paired with the single electrons that are already there. Electron Dot Structures Chemical reactivity depends upon the number of electrons in the highest energy level (the largest n). • Electron dot structures show only the outermost electrons so that reactivity can be easily predicted. • The outermost electrons are also called the valence electrons.

  33. Examples: 1. What is the electron configuration of strontium (Sr)? 1s2 2s2 2p6 3s2 3p6 4s2 3d10 3p6 5s2 The highest energy level is n=5 and there are 2 electrons in it. Therefore, Sr has 2 valence electrons. Sr 2. What is the electron configuration of nitrogen (N)? 1s2 2s2 2p3 N The highest energy level is n=2 and there are 5 electrons in it. Therefore, N has 5 valence electrons. 3. What is the electron configuration of xenon (Xe)? 1s2 2s2 2p6 3s2 3p6 4s2 3d10 3p6 5s2 4d10 5p6 The highest energy level is n=5 and there are 8 electrons in it. Therefore, Xe has 8 valence electrons. Xe

  34. 1 8 # Valence electrons 2 3 4 5 6 7

  35. Something to ponder… Excluding a few exceptions, how many valence electrons would all of the d-block and f-block elements have? 2

  36. Electron Configuration Exceptions Hund’s Rule (part 2): For n ≥ 4, the energy of the electrons decreases (and stability increases) when all of the boxes for a sublevel are either half-filled OR completely filled. [Ar] Cr All of the orbitals are half-filled 4s 3d [Ar] Cu All of the 3d orbitals are filled and the 4s orbital is half-filled 4s 3d

  37. Where would the exceptions be in the 5th through 7th periods?

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