Atomic Structure Part II Electrons in Atoms

Transcription

1 Atomic Structure Part II Electrons in Atoms

2 Radiant energy travels in the form of waves that have both electrical and magnetic properties. These electromagnetic waves can travel through empty space, as you know from the fact that radiant energy from the sun travels to Earth every day.

3 Electrons in Atoms I. Properties of Waves 1. Definition: Energy that exhibits wave-like (or oscillating) behavior as it travels through space

4 Electrons in Atoms 2. Wavelength (λ) distance from peak to peak, length of one complete wave 3. Frequency (ν) a. number of peaks that pass at a given point each sec b. can be called cycles per second (peak/sec) c. cps now called 1 Hertz (Hz)

5 Electrons in Atoms - Cont. 4. Velocity (C = speed of light) a. distance a given peak moves in a unit of time b. velocity (m/s) = frequency x wavelength c = ν x λ

6 II. Behavior of Light A. Newton (1600) thought light consisted of particles (beam of light is a stream of particles) B. Maxwell (1864) thought light was a wave phenomenon. Calculated the velocity of the propagation of an electromagnetic wave and found it was the same for light

7 II. Behavior of Light 1. some say light is like waves, some say its like particles 2. modern theory says that it behaves as both "wave/particle duality"

8 II. Behavior of Light 3. Max Planck (early 1900's) said: a. light is made up of bundles of energy called photons (or quanta) b. the energy of each photon is proportional to the frequency of the light (Quantum Theory) Quantum is the minimum amount of energy that can be gained or lost by an atom.

9 example: CONTINUOUS SPECTRUM *** when white light is passed through a prism, it is separated into a band of colors from red violet. It's called a continuous spectrum

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11 c. the work of Planck & Einstein led to E = h x ν E=energy, ν = frequency, h=planks constants (6.6262x10-34J /sec) J is the symbol for joule the SI unit for energy Energy of a quantum is related to the frequency of the emitted radiation by this equation

12 c. the work of Planck & Einstein led to E = h x ν According to Planck s theory, for a given frequency matter can emit or absorb energy only in whole-number mulitples of hv, that is 1hv, 2hv, 3hv, and so on.

13 c. The photoelectric effect In the photoelectric effect, electrons, called photoelectrons, are emitted from a metal s surface when light of a certain frequency shines on the surface. (example: solar calculator)

14 c. The photoelectric effect Einstein said light can both wavelike and particle like natures. That is, while a beam of light has many wavelike characteristics, it also can be thought of as a stream of tiny particles, or bundles or energy, called photons. A Photon is a particle of electromagnetic radiation with no mass that carries a quantum of energy.

15 III. Bright line spectrum A. a spectrum that shows separate bright lines, each with a specific wavelength B. bright-line spectra occur when an element is heated and the colored light given off is viewed through a spectroscope. Each element has a unique set of lines, characteristic of that element (like a fingerprint)

16 Line-Emission Spectrum excited state ENERGY IN PHOTON OUT ground state

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18 Fireworks? Hmmm

19 IV. Electromagnetic Spectrum A. visible light (like the continuous spectrum) is only one type of radiation. All other types are not visible to the human eye.

20 H I G H E N E R G Y L O W E N E R G Y

21 Electromagnetic Spectrum

22 L O W E N E R G Y Electromagnetic Spectrum R O Y G. B I V red orange yellow green blue indigo violet H I G H E N E R G Y

23 B. all forms of electromagnetic radiation travels at the speed of light. 1. speed of light = 3.00 x 10 8 meters/sec 2. use formula: c = ν x λ 3. each line spectrum has a particular frequency (ν ). If know wavelength (λ), we can find ν using c as a constant.

24 C. The energy in a photon of light is directly proportional to the frequency of the light. 1. frequency, energy 2. can find the energy of a single quantum (photon) of radiation at any given frequency.

25 C. The energy in a photon of light is directly proportional to the frequency of the light. 3. proportionality constant that relates the two is called Planck's constant (h). 4. formula: E = h x f

26 example: a spectral line has frequency of 3.5x10 12 hertz. What is the energy of a photon of radiation of this frequency? E = h x f (h=6.6262x10-34 J/sec) E = (3.5x10 12 Hz) (6.6262x10-34 J sec) E = (2.3x10-21 J)

27 V. Electron energy levels in Bohr's Model A. There are certain different orbits in which an electron can travel around a nucleus. 1. each circular orbit (or shell) is at a fixed distance from the nucleus

28 V. Electron energy levels in Bohr's Model 2. the greater the radius of that shell, the greater the energy of the electron in that shell. 3. these electron orbits are known as energy levels

29 B. When electrons absorb energy firm an outside source, they jump from lower to higher energy levels. when they fall back to their original levels, energy is emitted (light); the same amount as was absorbed.

30 6 5 Energy of photon 4 3 depends on the 2 difference in energy 1 levels Bohr s calculated energies matched the IR, visible, and UV lines for the H atom

31 C. In energy atom in its normal state, all electrons are in the lowest energy levels available (energetically stable)

32 VI. Atoms and Radiation A. When all of the lowest energy levels are occupied, the atom is in the ground state (unexcited).

33 VI. Atoms and Radiation B. When electron moves to higher energy level, atom is in the excited state, and is energetically unstable.

34 VI. Atoms and Radiation C. Bright line spectrum of an element represents the energy levels in its atoms.

35 problems with Bohr's Model: only explained some of the lines in the bright line spectrum really only worked for hydrogen need sublevels and electron cloud model to account for all of the lines.

36 VII. The Modern Model of the Atom A. Mechanics 1. Classical Mechanics - Newton's Laws of Motion (Newtonian Mechanics) Describes the behavior of visible objects traveling at ordinary velocities. Bohr s basis for his model, but couldn t explain why electrons would stay at on energy level or another. When looking at H-spectral lines, noticed more one (several closely spaced).

37 VII. The Modern Model of the Atom 2. Quantum Mechanics (wave mechanics) Describes the behavior of extremely small particles traveling at velocities at or near the speed of light

38 a. Louis de Broglie - particles could have properties of waves Planks quanta gave wave properties, debroglie said electron streams are like waves of light and have properties of both particles and waves (matter behaves as waves)

39 b. Schrodinger - described the behavior of electrons in terms of quantized energy changes "quantum mechanics" Describe a wave equation used to determine the probability of finding an electron in any give place or orbital Schrodinger s Cat

40 Orbital Radial Distribution Curve

41 c. Heisenberg - uncertainty principle "The more precisely the POSITION is determined, the less precisely the MOMENTUM is known" Region of space where there is a probability of finding an electron is called an orbital

42 B. Principal Energy Levels 1. Energy Levels Bohr - High Energy (outer level) Low Energy

43 Principal Quantum Numbers (N) Number of electrons 2 Corresponds to energy level

44 2. Sublevels Principal Quantum Numbers (N) Sublevel Present 1 1s 2 2s2p 3 3s3p3d 4 4s4p4d4f

45 Orbital - Electrons are represented by arrows Region of space where an electron is probably found Electron spin An orbital can hold 2 electrons that spin in opposite directions.

46 Rules: 1. Pauli Exclusion Principle Each orbital can hold TWO electrons with opposite spins. No two electrons in an atom can have the same 4 quantum numbers. Each e - has a unique address :

47 2. Aufbau Principle Electrons fill the lowest energy orbitals first. Electrons to be added must be placed in unfilled orbitals of lowest energy for stable configuration.

48 3. Hund s Rule Within a sublevel, place one e- per orbital before pairing them. Empty Bus Seat Rule WRONG RIGHT

49 Energy Level Diagram c Electrons have up and down spin orbital Orbital - Place where electrons are probably found

50 Shapes of electron orbitals

51 The s orbital

52 The p orbitals

53 p x The p orbitals p y p z

54 The d orbitals

55 The d orbitals

56 The f orbitals f

57 orbital views Click here for orbital viewer View the grand table

58 Energy Level C. Electron Configurations Sub Level # of Electrons 1s 2 = Helium 1s 2 2s 1 = Lithium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 = Krypton

59 Electron Configuration Notation Orbital Diagram O 8e - 1s 2s 2p Electron Configuration 1s 2 2s 2 2p 4

60 Notation Longhand Configuration S 16e - 1s 2 2s 2 2p 6 3s 2 3p 4 Core Electrons Valence Electrons Shorthand Configuration S 16e - [Ne] 3s 2 3p 4

61 s f (n-2) 6 7 Periodic Patterns p d (n-1) 1998 by Harcourt Brace & Company

62 Periodic Patterns Period # energy level (subtract for d & f) A/B Group # total # of valence e - Column within sublevel block # of e - in sublevel

63 Periodic Patterns Example - Hydrogen s 1 1st column of s-block 1st Period s-block

64 Periodic Patterns Shorthand Configuration Core e - : Go up one row and over to the Noble Gas. Valence e - : On the next row, fill in the # of e - in each sublevel.

65 Periodic Patterns Example - Germanium [Ar] 4s 2 3d 10 4p 2

66 Stability Full sublevel (s, p, d, f) Full energy level Half-full sublevel

67 Stability Ion Formation Atoms gain or lose electrons to become more stable. Isoelectronic with the Noble Gases.

68 Stability Ion Formation Atoms gain or lose electrons to become more stable. Isoelectronic with the Noble Gases.

69 Feeling overwhelmed?

70 Try a few! Mg = Fe = Ru = Ir = Ca +2 = Cl -1 =