1. Electrons in Atoms

2.  Wavelength ( ) - length of one complete wave measured in m, cm, or nm  In light it tells us which color it is  Frequency ( ) - # of waves that pass a point during a certain time period,  hertz (Hz) = 1/s  Amplitude (A) - distance from the origin to the trough (dip) or crest (peak  Amount of energy the wave is carrying - height of wave. It is measured in meters.  In SOUND it tells us how LOUD it is. In LIGHT it tells how BRIGHT it is.

3. A greater amplitude greater frequency crest origin trough A

4.  Understanding electronic structure of atoms  Must understand light  Emitted and absorbed by substances.  Visible light - type of Electromagnetic Radiation (EM)  Carries (radiant) energy through space  Travels at speed of light  Exhibits wavelike behavior.  Think of light as particle  help understand how EM radiation and atoms interact

5. LOWENERGYLOWENERGY HIGHENERGYHIGHENERGY

6.  Move through a vacuum at the ‘speed of light’ 3.00 x 10 8 m/s  Behaves like waves that move through water  Result of energy transferred to the water (from a stone)  Expressed as up and down movement  Both electric and magnetic properties

7.  Wave Speed = (distance between wave peaks) x (frequency) = (wavelength) x (frequency)  EM radiation moves through a vacuum at the “speed of light” 3.00 x 10 8 m/s also called c.  A lower energy wave (infrared and red) has a longer wavelength( ) and lower frequency(f)  A higher energy wave (blue - violet) has a shorter wavelength( ) and higher frequency(f).

8.  Frequency & wavelength are inversely proportional c = c:speed of light (3.00  10 8 m/s) :wavelength (m, nm, etc.) :frequency (Hz)

9.  EX: Find the frequency of a photon with a wavelength of 434 nm. GIVEN: = ? = 434 nm = 4.34  10 -7 m c = 3.00  10 8 m/s WORK : = c = 3.00  10 8 m/s 4.34  10 -7 m = 6.91  10 14 Hz

10.  Planck (1900)  Observed - emission of light from hot objects  Concluded - energy is emitted (absorbed or released) in small, specific amounts (quanta)  Quantum - smallest energy packet that can be emitted or absorbed as EM radiation by an atom.

11. E:energy (J, joules) h:Planck’s constant (6.6262  10 -34 J·s) :frequency (Hz) E = h zPlanck proposed that the energy, E, of a single quantum energy packet equals a constant (h) times its frequency zThe energy of a photon is proportional to its frequency.

12.  EX: Find the energy of a red photon with a frequency of 4.57  10 14 Hz. GIVEN: E = ? = 4.57  10 14 Hz h = 6.6262  10 -34 J·s WORK : E = h E = ( 6.6262  10 -34 J·s ) ( 4.57  10 14 Hz ) E = 3.03  10 -19 J

13.  Energy is always emitted or absorbed in whole number multiples of hv, such as hv, 2 hv, 3 hv, 4hv, ….  The allowed energies are quantized  values are restricted to certain quantities.  The notion of quantized rather than continuous energies is strange.  Ramp vs Staircase  Ramp - vary the length your steps and energy used on the walk up.  Stairs - must exert exactly the specific amount of energy needed to reach the next step.  Your steps on steps are quantized, you cannot step between them.

14.  Planck (1900) vs. Classical TheoryQuantum Theory

15.  Einstein (1905)  Observed – photoelectric effect  Dispersed light falls on metal samples, the different frequencies produce different energetic photoelectrons

16.  Einstein (1905)  Concluded - light has properties of both waves and particles (photons) “wave-particle duality”  Photon - particle of light that carries a quantum of energy  Used planck’s quantum theory to deduced that: E photon = hv

17. Electrons in Atoms

18. ground state excited state ENERGY IN PHOTON OUT yElements’ atoms absorb electrical energy ye- get excited, become unstable, and release energy yEnergy is in form of light ySet of frequencies of EM waves

19.  e - exist only in orbits with specific amounts of energy called energy levels  Therefore…  e - can only gain or lose certain amounts of energy  only certain photons are produced  Ground state = lowest allowable atomic electron energy state  Excited state = any higher energy state

20.  Energy of photon depends on the difference in energy levels  Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom 1 2 3 4 5 6

21.  Each element has a unique bright-line emission spectrum.  “Atomic Fingerprint” Helium  Examples:  Iron  Now, we can calculate for all elements and their electrons – next section

23.  Louis de Broglie (1924)  Proposed eˉ in their orbits behave like a wave EVIDENCE: DIFFRACTION PATTERNS ELECTRONS VISIBLE LIGHT

24.  Heisenberg Uncertainty Principle  Impossible to know velocity and position of an electron at the same time  Trying to observe an electron’s position changes its momentum  Trying to observe an electron’s momentum changes its position  Electrons cannot be locked into well-defined circular orbits around the nucleus.

25.  Orbital (“electron cloud”)  a specific distribution of electron density in space.  Each orbital has a characteristic energy and shape. Orbital

26. Specify the “address” of each electron in an atom UPPER LEVEL

27. 1. Principal Quantum Number ( n = 1, 2, 3, …) (see periodic table left column)  Indicates the relative size and energy of atomic orbitals  As (n) increases, the orbital becomes larger, the electron spends more time farther from the nucleus  Each major energy level is called a principle energy level Ex: lowest level = 1 ground state, highest level = 7 excited state

28. 2. Energy Sublevel  Defines the shape of the orbital (s, p, d, f)  # of orbital related to each sublevel is always an odd # s = 1, p = 3, d = 5, f = 7  Each orbital can contain at most 2 electrons s p d f

29. Subscripts x, y, z designates orientation  Specifies the exact orbital within each sublevel

30. pxpx pypy pzpz

31. 4. Spin Quantum Number ( m s )  Electron spin  +½ or -½  An orbital can hold 2 electrons that spin in opposite directions.

32.  Pauli Exclusion Principle  A maximum of 2 electrons can occupy a single atomic orbital  Only if they have opposite spins 1. Principal #  2. Energy sublevel  3. Orientation  4. Spin #  energy level (s,p,d,f) x, y, z exact electron

33. Electron Configuration Electrons in Atoms

34. A. General Rules zAufbau Principle yElectrons fill the lowest energy orbitals first. y“Lazy Tenant Rule”

35. RIGHT WRONG A. General Rules zHund’s Rule yWithin a sublevel, place one e - per orbital before pairing them. y“Empty Bus Seat Rule”

36. O 8e - zOrbital Diagram zElectron Configuration 1s 2 2s 2 2p 4 B. Notation 1s 2s 2p

37. zShorthand Configuration S 16e - Valence Electrons Core Electrons S16e - [Ne] 3s 2 3p 4 1s 2 2s 2 2p 6 3s 2 3p 4 B. Notation zLonghand Configuration yValence electrons: determine chemical properties of that element & are the electrons in the atoms outermost orbital

38. © 1998 by Harcourt Brace & Company s p d (n-1) f(n-2) 12345671234567 6767 Notation

39. zShorthand Configuration yCore e - : Go up one row and over to the Noble Gas. yValence e - : On the next row, fill in the # of e - in each sublevel. Shorthand Notation

40. [Ar]4s 2 3d 10 4p 2 C. Periodic Patterns zExample - Germanium

41. zFull energy level zFull sublevel (s, p, d, f) zHalf-full sublevel D. Stability

42. zElectron Configuration Exceptions yCopper EXPECT :[Ar] 4s 2 3d 9 ACTUALLY :[Ar] 4s 1 3d 10 yCopper gains stability with a full d-sublevel. D. Stability

43. zElectron Configuration Exceptions yChromium EXPECT :[Ar] 4s 2 3d 4 ACTUALLY :[Ar] 4s 1 3d 5 yChromium gains stability with a half-full d-sublevel. D. Stability

44. zIon Formation yAtoms gain or lose electrons to become more stable. yIsoelectronic with the Noble Gases.

45. O 2- 10e - [He] 2s 2 2p 6 D. Stability zIon Electron Configuration yWrite the e - config for the closest Noble Gas yEX: Oxygen ion  O 2-  Ne