Arrangement of Electrons in Atoms

1. Arrangement of Electrons in Atoms Chapter 4

2. Historical Models • Democritus—indivisible particles • Thomson—electrons—plum pudding model • Rutherford—dense central core—nuclear model

3. New Model • Based on observations of absorption and emission of electromagnetic radiation by atoms • ***Video Clip***

4. EM Radiation • Wave-like properties • Wavelength—l—the distance between two corresponding points on successive waves • Velocity—c—speed of light = 3.00 x 108 m/s

5. EM Radiation • Frequency—v—the number of waves that pass a given point in one second • For any wave: c = lv

6. Proportionality • Since speed is constant (speed of light), frequency and wavelength vary inversely • Frequency increases; wavelength decreases • Frequency decreases; wavelength increases

8. Which has higher energy? • Wavelength of 400 nm or 700 nm? • Frequency of 108 or 1018? • Wavelength of 1 m or 1 cm? • Frequency of 106 or 1012?

9. Calculations • What is the wavelength of EM radiation with a frequency of 1016 Hz? • What type of radiation is it? (See p. 92 in textbook.)

10. Calculations • What is the frequency of light with a wavelength of 400 nm? (400 nm = 4 x 10-7 m) • What color is this light? (See p. 92 in textbook.)

11. Particle Behavior of Light • Max Planck • Objects do not emit light at continuous frequencies; rather the energy comes out in specific amounts or quanta • Quantum—the minimum amount of energy an atom can gain or lose

12. Energy of a Quantum • E = hv • E = energy (J) • h = Planck’s Constant = 6.623 x 10-34 J s • v = frequency (1/s)

13. Photoelectric Effect • Emission of electrons from a metal when light shines on the metal • If light is a wave, it should have enough energy to remove electrons at any frequency, but only certain frequencies of light will remove electrons.

14. Einstein’s Contribution • Light behaves as a particle and as a wave--Wave particle duality • Photon—a particle of light that carries a quantum of energy • Energy of photon = energy of quantum—use same equation

15. Quantum A specific “packet” of energy

16. Excited Electrons • Electrons that have absorbed energy become excited and move to higher energy states in an atom • When they return to lower energy states, they emit the energy

17. Line Emission Spectra • H always and only emits radiation at specific wavelengths • Other elements behave the same way, but patterns are more complex due to larger # of electrons

18. Line Emission Spectra • Very specific—Act like fingerprints by which an element can be identified ***Emission Tubes/ Spectrometers***

19. Do the Math • A mathematical relationship exists between the wavelengths of the emission spectra of hydrogen. • The spectral emission of elements is always the same.

21. Bohr Model • Energy of emission equals the difference in energy levels • Added orbits for electrons to the existing atomic model

22. Energy Levels • Electrons may only go to certain places (energy levels) within the atom. • Evidence: wavelengths in emission spectra of atoms

23. Don’t Write This!! It requires the same amount of energy every time you • walk from your seat to the café. • walk from your seat to your locker. • walk from your seat to the door. If these were the only three places you could go, you would only use energy in those amounts.

24. Limitations of Bohr’s Model • Since hydrogen has only one electron, Bohr’s model worked well for it. • The model did not work as well when applied to atoms with more electrons.

25. Bohr Model • Only certain numbers of electrons are allowed in each level • Level 1: 2 • Level 2: 8 • Level 3: 18 • Level 4: 32

26. The Quantum Model • De Broglie: electrons have properties of both waves & particles—interference & diffraction. • Wave-particle duality • Waves could only exist with with wavelengths that corresponded to Bohr’s orbitals

27. Heisenberg Uncertainty Principle • If electrons are like both waves and particles, where are they in an atom? • Electrons are located by their interaction with photons, but that interaction knocks them off course—as soon as you locate it, it’s somewhere else!

28. Heisenberg Uncertainty Principle It is impossible to know both the position and velocity of an electron or any other particle at the same time.

29. Schrodinger’s Wave Equation • Only certain frequencies of waves solve the equation • These frequencies correspond to quanta of energy & Bohr’s orbitals

30. Schrodinger’s Wave Equation • Schrodinger’s wave function describes a probable space in which electrons can be found • Orbital—a 3-dimensional region around the nucleus that indicates the probable location of an electron

31. s orbitals—1 per level beginning at level 1

32. p orbitals—”peanut” shape—3 per level beginning at level 2

33. d orbitals—”double dumbbell” shape—5 per level beginning at level 3

34. f orbitals—”flower” shape—7 per level beginning at level 4

35. Orbitals/Sublevels • Sets of orbitals (s, p, d & f) are sometimes called sublevels. • Each orbital holds 2 electrons • Different numbers of orbitals for each sublevel • s holds 2 (1x2) p holds 6 (3x2) • d holds 10 (5x2) f holds 14 (7x2)

36. Aufbau Principle • “Building up” • Electrons fill orbitals from lowest energy to highest energy. • Order is given by the diagonal rule.

37. Diagonal Rule Begin with 1s and follow the diagonal lines. Remember, each sublevel can only hold a specific # of electrons.

38. Orbital Diagrams • Represents electrons in their orbitals. • Arrows indicate spin. • Hund’s rule—electrons at a sublevel will not pair up until each orbital has one electron. • Pauli Exclusion Principle—electrons in the same orbital must have opposite spin.

39. Orbital Notation(Orbital Diagram) • Page 158 • __ __ __ = unoccupied orbitals • __ __ __ = orbitals occupied with electrons of the same spin • __ __ __ = orbitals occupied with electrons with different spins

40. Orbital Diagrams for C, N & Fe

41. Electron Configuration • Represents electrons in their sublevels with numbers and letters. • Follow Diagonal Rule • Does not show spin or pairing of electrons.

42. Examples • Si—Z = 14 1s2, 2s2, 2p6, 3s2, 3p2

43. Location on the periodic table also tells electron configuration

44. Noble Gas Notation • Shortcut for electron configurations • Use the Noble gas from the line above the element. • Continue configuration from the next s orbital. (The row number s)