2.  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 exhibits characteristics of waves  Wavelength (λ)  Frequency (ν)  amplitude

4.  Different types  Each type has characteristic λ and ν  All parts travel with the speed of light (c)… 3.00 x 10 8 m/s

7.  Microwaves are used to cook food and transmit information. What is the wavelength of a microwave that has a frequency of 3.44 x 10 9 Hz? Now You Try: Practice Problems 1 – 4 on page 140

8.  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.  Concluded that energy could only be gained or lost in small, specific amounts (like tiny packages)…called these amounts quanta

10. E = hν

11.  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)(video)

12.  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 10 14 s -1 ?

13.  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.  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.  Based on a hydrogen atom  Assigned a quantum number (n) to each orbit  As value of n increases, the amount of energy increases

17.  Worked well for hydrogen  Did not explain the atomic spectra produced by any other elements

18.  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

20.  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.  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.  Used to describe orbitals  Specify the properties of atomic orbitals and the properties of the electrons in the orbitals

23.  First three derived from the Schrodinger equation: › Main energy level (n) › Shape of orbital (l) › Orientation of orbital (m l )  Fourth is the spin quantum number (m s ) › Describes the fundamental state of the electron

24.  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 n 2

25.  Each main energy level (except the 1 st ) 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.  Depending on the value of l, the orbital is assigned a letter  0=s  1=p  2=d  3=f

27.  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.  Orbitals can have the same shape, but different orientations around the nucleus  Magnetic quantum numbers indicate the orientation (m l )  s orbitals have only one orientation  p orbitals can extend along the x, y, or z axis  3 p sublevels (p x, p y, or p z )

31.  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.  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.  Remember: › All electrons can be described by a set of quantum numbers › No two electrons can have the same set of quantum numbers

34.  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)

35. Energy LevelTypes of Orbitals 1s 2s, p 3s, p, d 4s, p, d, f 5 6s, p 7

36. Orbital Type# Sub-orbitals s1 p3 d5 f7

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

38. Type of OrbitalMax # electrons s2 p6 d10 f14

39. Energy LevelTypes of OrbitalMax # of e - 1s2 2s, p8 3s, p, d18 4s, p, d, f32 5s, p, d, f32 6s, p,8 7s, p8

40.  s, p, d, f blocks

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

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

43. 5d 3d 4d 6d

44. 4f 5f

45.  Carbon  Lithium  Sodium  Phosphorus  Neon

46.  Write the complete electron configuration for: › Helium › Sulfur › Magnesium › Silicon › Tin

47.  Carbon  Lithium  Sodium  Phosphorus  Neon

48.  Write the orbital notation for: › Helium › Sulfur › Magnesium › Silicon › Tin

49.  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.  Calcium: › Electron Configuration: › Noble Gas Notation:  You try these: a. Titanium b. Silicon c. Barium