Wave Particle Duality & Electron Configurations

1. Wave Particle Duality &Electron Configurations

2. Remember Rutherford? • Proposed model of the atom had a nucleus of positive charge surrounded by a relatively large area of empty space where electrons orbited • Did not propose an arrangement for the electrons • Did not explain why the electrons were not pulled into the nucleus (attraction of opposite charges)

3. Light as a Wave • Light exhibits characteristics of waves • Wavelength (λ) • Frequency (ν) • amplitude

4. Electromagnetic Waves • Different types • Each type has characteristic λ and ν • All parts travel with the speed of light (c)… 3.00 x 108 m/s

6. Electromagnetic Wave Formula

7. EXAMPLE PROBLEM 5.1 • Microwaves are used to cook food and transmit information. What is the wavelength of a microwave that has a frequency of 3.44 x 109 Hz? Now You Try: Practice Problems 1 – 4 on page 140

8. The Particle Nature of Light • Objects that are heated often give off a characteristic color (red of stove burner, white of light bulb) • View of light as a wave did not provide an accurate explanation of why this occurs • So….

9. Max Planck (1858 – 1947) • Concluded that energy could only be gained or lost in small, specific amounts (like tiny packages)…called these amounts quanta

10. Energy of a Quantum E = hν

11. The Photoelectric Effect • Another phenomenon that could not be explained with light as a wave • When light of a certain minimum frequency shines on a metal’s surface, the metal will eject electrons (video)

12. Example 5.2 • Every object gets its color by reflecting a certain portion of incident light. The color is determined by the wavelength of the reflected photons, thus by their energy. What is the energy of a photon from the violet portion of the Sun’s light if it has a frequency of 7.230 x 1014 s-1?

13. Atomic Emission Spectra • Also called line spectra…not continuous • Set of frequencies of electromagnetic waves emitted by an element • Not continuous • Unique for each element (like a fingerprint)

14. Electron Configurations • Bohr • Model stated that atoms orbit the nucleus in definite paths (energy levels) • Patterned this model after planets orbiting the sun • Electrons in a particular path (energy level) have a fixed amount of energy…quantized • Energy levels are analogous to the rungs on a ladder

15. Bohr’s Model • Based on a hydrogen atom • Assigned a quantum number (n) to each orbit • As value of n increases, the amount of energy increases

16. Energy relationships…

17. Limitations of Bohr’s Model • Worked well for hydrogen • Did not explain the atomic spectra produced by any other elements

18. The Quantum Mechanical Model of the Atom • Louis de Broglie (1892 – 1987) • Recognized that electrons exhibited characteristics of waves • Recognized that light has properties of both waves and particles • Theorized that matter must be able to possess qualities of waves and particles as well

19. de Broglie Equation: • Predicts that all matter has wave characteristics

20. Heisenberg Uncertainty Principle • States that it is fundamentally impossible to know both the precise velocity (momentum) and location of a particle at the same time • Applies to all matter, but is useful only with really small particles…like electrons

21. Electron Configurations • Schrodinger • Revised Bohr’s model • Mathematical equation to determine the most likely place an electron would orbit the nucleus • Gives the probability of finding an electron in a particular place within the atom

22. Quantum Numbers • Used to describe orbitals • Specify the properties of atomic orbitals and the properties of the electrons in the orbitals

23. 4 Quantum Numbers • First three derived from the Schrodinger equation: • Main energy level (n) • Shape of orbital (l) • Orientation of orbital (ml) • Fourth is the spin quantum number (ms) • Describes the fundamental state of the electron

24. Principal Quantum number • Symbolized by n • Indicates the main energy level occupied by an electron • Values are positive integers • As n increases, so does the distance from the nucleus • More than one electron can have the same n • Total number of orbitals for a given energy level is given by n2

25. Angular Momentum Quantum Number • Each main energy level (except the 1st) has different orbitals of different shapes • Symbolized by l • Indicates the shape of the orbital • The number of orbital shapes possible for each energy level is equal to the value of n • The values of l allowed are zero through n-1

26. Angular Momentum Quantum Number • Depending on the value of l, the orbital is assigned a letter • 0=s • 1=p • 2=d • 3=f

27. Shapes of sub-orbitals • s = spherical • p = dumbbell shaped • d= complex • f = way to complicated to explain…see illustrations • Atomic orbitals are designated by the n followed by the letter of the sublevel

30. Magnetic Quantum Number • Orbitals can have the same shape, but different orientations around the nucleus • Magnetic quantum numbers indicate the orientation (ml) • s orbitals have only one orientation • p orbitals can extend along the x, y, or z axis • 3 p sublevels (px, py, or pz)

31. Magnetic Quantum Numbers • Values for m sublevels correspond values m = -1 m=0 and m=1 • 5 different d orbitals…therefore 5 different orientations • m=-2 m=-1 m=0 m=+1 m=+2 • 7 different f orbitals….7 orientations

32. Spin Quantum Numbers • Electrons spin on an internal axis • Can spin in one of two possible directions • Spin quantum numbers can be +1/2 or -1/2 • A single orbital can hold a maximum of 2 electrons • The electrons in a single orbital must have opposite spins

33. Electron Configurations • Remember: • All electrons can be described by a set of quantum numbers • No two electrons can have the same set of quantum numbers

34. Rules for Electron Configurations • Aufbau Principle • An electron will occupy the lowest energy orbital that can receive it • Pauli Exclusion Principle: • No 2 electrons in the same atom can have the same set of quantum numbers • Hund’s Rule: • Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron….All electrons in singly occupied orbitals have the same direction of spin (parallel spin)

37. Sub-orbitals • Each sub-orbital can hold two electrons • The electrons in a sub-orbital must have opposite spin

40. Electron Configurations & the Periodic Table • s, p, d, f blocks

41. 1 s 2 s 3s 4 s 5s 6 s 7 s s - block

42. 2p 3p 4p 5p 6p 7p p - block

43. 3d 4d 5d 6d d - block

44. f - block 4f 5f

45. Give the Electron Configurations: • Carbon • Lithium • Sodium • Phosphorus • Neon

46. You Try: • Write the complete electron configuration for: • Helium • Sulfur • Magnesium • Silicon • Tin

47. Show orbital notations for: • Carbon • Lithium • Sodium • Phosphorus • Neon

48. You try • Write the orbital notation for: • Helium • Sulfur • Magnesium • Silicon • Tin

49. Noble Gas Notations • Use the noble gas that comes before the element • Write the noble gas’s symbol in brackets • Continue with the electron configuration from there

50. Noble Gas Notation Example • Calcium: • Electron Configuration: • Noble Gas Notation: • You try these: a. Titanium b. Silicon c. Barium