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Chapter 5 “Electrons in Atoms”. Ernest Rutherford’s Model Discovered dense + piece at nucleus Electrons move around it, like planets around sun Mostly.

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Presentation on theme: "Chapter 5 “Electrons in Atoms”. Ernest Rutherford’s Model Discovered dense + piece at nucleus Electrons move around it, like planets around sun Mostly."— Presentation transcript:

1 Chapter 5 “Electrons in Atoms”

2 Ernest Rutherford’s Model Discovered dense + piece at nucleus Electrons move around it, like planets around sun Mostly empty space Didn’t explain chemical properties of elements

3 Niels Bohr’s Model Why don’t electrons fall into nucleus? Move like planets around sun.  specific circular orbits at different levels.  An amount of fixed energy separates one level from another.

4 The Bohr Model of the Atom Niels Bohr I pictured the electrons orbiting the nucleus much like planets orbiting the sun. However, electrons are found in specific circular paths around the nucleus, and can jump from one level to another.

5 Bohr’s model Energy level measure of fixed energy e- can have Like rungs of ladder e- can’t exist between energy levels can’t stand between rungs on ladder Unlike ladder, energy levels not evenly spaced Higher energy levels closer together – less energy needed for jump

6 The Quantum Mechanical Model Energy is “quantized” - in chunks. A quantum is exact energy needed to move e - one energy level to another. Since energy of atom is never “in between” quantum leap in energy must exist. Erwin Schrodinger (1926) derived equation described energy and position of e - in an atom Schrodinger

7 The Quantum Mechanical Model Very small things behave differently from big things quantum mechanical model is a mathematical solution not like anything you can see (like plum pudding!)

8 The Quantum Mechanical Model Has energy levels for e - Orbits not circular. only tells us probability of finding e - a certain distance from nucleus.

9 The Quantum Mechanical Model e- found inside blurry “electron cloud” (area where there’s chance of finding e - Like fan blades or bicycle wheel

10 Atomic Orbitals Principal Quantum Number (n) = the energy level of e- (1, 2, 3, etc.) Within each energy level, Schrodinger’s equation describes several shapes. regions of space w/ high probability called atomic orbitals - regions of space w/ high probability of finding e- Sublevels like rooms in a hotel letters s, p, d, and f

11 Principal Quantum Number (n) denotes the energy level (shell) e- is located. Maximum number of e- that can fit in energy level is: 2n 2 How many e - in level 2? level 3?

12 Individual orbitals s spherical p dumbell d clover leaf f complicated # of orbitals Maximum electrons First possible energy level 1 2 1st 3 6 2nd rd th

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15 By Energy Level Energy Level 1 (n=1) Has only s sublevel 1s (1 orbital) w/ only 2 e- 2 total e- Energy Level 2 (n=2) Has s and p sublevels available 2s (1 orbital) w/ 2 e- 2p (3 orbitals) w/ 6 e- 2s 2 2p 6 8 total e

16 By Energy Level Energy level 3 (n=3) s, p, and d sublevels 3s (1 orbital) w/ 2 e- 3p (3 orbitals) w/ 6 e- 3d (5 orbitals) w/ 10 e- 3s 2 3p 6 3d total e- Energy level 4 (n=4) s, p, d, and f sublevels 4s (1 orbital) w/ 2 e- 4p (3 orbitals) w/ 6 e- 4d (5 orbitals) w/ 10 e- 4f (7 orbitals) w/ 14 e- 4s 2 4p 6 4d 10 4f total e

17 By Energy Level  Any more than the fourth and not all orbitals fill up. You simply run out of e-’s orbitals do not fill up in neat order. energy levels overlap Lowest energy fill first. s and p orbitals 1:20 d orbitals 3:40

18 Summary of Principal Energy Levels, Sublevels, and Orbitals Principal energy level Number of sublevels Type of sublevel Max # of electrons Electron configuration n = 1 1 1s (1 orbital) 2 1s 2 n = 2 2 2s (1 orbital 2p (3 orbitals) 8 2s 2 2p 6 n = 3 3 3s (1 orbital) 3p (3 orbitals) 3d (5 orbitals) 18 3s 2 3p 6 3d 10 n = 4 4 4s (1 orbital) 4p (3 orbitals) 4d (5 orbitals) 4f (7 orbitals) 32 4s 2 4p 6 4d 10 4f 14

19 ENERGY Energy Level s 2s 2p 3s 3p 3d Electron distribution in an Atom NUCLEUS

20 Section 5.2 Electron Arrangement in Atoms OBJECTIVES: Describe how to write the electron configuration for an atom.

21 Section 5.2 Electron Arrangement in Atoms OBJECTIVES: Explain why the actual electron configurations for some elements differ from those predicted by the aufbau principle.

22 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f aufbau diagram - page 133 Aufbau is German for “building up”

23 Electron Configurations… …the way electrons arranged in various orbitals around nuclei of atoms. Three rules tell us how: 1) Aufbau principle - electrons enter the lowest energy first. follow diagram!causes difficulties b/c overlap of orbitals of different energies – follow diagram! 2) Pauli Exclusion Principle - 2 electrons max per orbital (hotel room) - different spins

24 Pauli Exclusion Principle No two electrons in an atom can have the same four quantum numbers. Wolfgang Pauli To show the different direction of spin, a pair in the same orbital is written as:

25 Quantum Numbers Each electron in an atom has a unique set of 4 quantum numbers which describe it. 1) Principal quantum number 2) Angular momentum quantum number 3) Magnetic quantum number 4) Spin quantum number

26 Electron Configurations 3) Hund’s Rule- When electrons occupy orbitals of equal energy, they don’t pair up until they have to. write e- configuration for Phosphorus  We need to account for all 15 e-’s in phosphorus

27 The first two electrons go into the 1s orbital Notice the opposite direction of the spins only 13 more to go... Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

28 The next electrons go into the 2s orbital Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

29 The next electrons go into the 2p orbital only 5 more... Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

30 The next electrons go into the 3s orbital only 3 more... Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

31 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f The last three electrons go into the 3p orbitals. They each go into separate shapes (Hund’s) 3 unpaired electrons = 1s 2 2s 2 2p 6 3s 2 3p 3 Orbital notation

32 Stairstep pattern

33 An internet program about electron configurations is: Electron Configurations (Just click on the above link) I electron config 3:24

34 Orbitals fill in an order Lowest energy  higher energy. Adding e-’s changes energy of orbital. Full orbitals are best situation. half filled orbitals  lower energy, next best more stable. Changes filling order

35 Write the electron configurations for these elements: Titanium - 22 electrons 11s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 Vanadium - 23 electrons 11s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 Chromium - 24 electrons 11s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 (expected) BBut this is not what happens!!

36 Chromium is actually: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 Why? This gives us two half filled orbitals (the others are all still full) Half full is slightly lower in energy. The same principal applies to copper.

37 Copper’s electron configuration Copper has 29 electrons so we expect: 1s 2 2s 2 2p 6 3s 2 3p 6 3d 9 4s 2 But the actual configuration is: 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 1 This change gives one more filled orbital and one that is half filled. Remember these exceptions: d 4, d 9

38 Irregular configurations of Cr and Cu Chromium steals a 4s electron to make its 3d sublevel HALF FULL Copper steals a 4s electron to FILL its 3d sublevel

39 Section 5.3 Physics and the Quantum Mechanical Model OBJECTIVES: Describe the relationship between the wavelength and frequency of light.

40 Section 5.3 Physics and the Quantum Mechanical Model OBJECTIVES: Identify the source of atomic emission spectra.

41 Section 5.3 Physics and the Quantum Mechanical Model OBJECTIVES: Explain how the frequencies of emitted light are related to changes in electron energies.

42 Section 5.3 Physics and the Quantum Mechanical Model OBJECTIVES: Distinguish between quantum mechanics and classical mechanics.

43 Light Study of light led to development of quantum mechanical model. Light is electromagnetic radiation. EM radiation: gamma rays, x-rays, radio waves, microwaves Speed of light = x 10 8 m/s abbreviated “c” “celeritas“ Latin for speed All EM radiation travels same in vacuum

44 - Page 139 “R O Y G B I V” Frequency Increases Wavelength Longer

45 Parts of a wave Wavelength Amplitude Origin Crest Trough

46 Equation: c = c = speed of light, a constant (2.998 x 10 8 m/s) (nu) = frequency, in units of hertz (hz or sec -1 ) (lambda) = wavelength, in meters Electromagnetic radiation propagates through space as a wave moving at the speed of light.

47 Wavelength and Frequency inversely related one gets bigger, other smaller Different frequencies of light different colors wide variety of frequencies (spectrum)

48 - Page 140 Use Equation: c =

49 Radio waves Micro waves Infrared. Ultra- violet X- Rays Gamma Rays Low Frequency High Frequency Long Wavelength Short Wavelength Visible Light Low Energy High Energy

50 Long Wavelength = Low Frequency = Low ENERGY Short Wavelength = High Frequency = High ENERGY

51 Atomic Spectra White light all colors of visible spectrum. prism prism separates it.

52 If the light is not white heating gas with electricity will emit colors. Passing this light through prism different.

53 Atomic Spectrum Each element gives off own characteristic colors. This is how we know composition of stars

54 atomic emission spectrum Unique to each element, like fingerprints! ID’s elements

55 Light is a Particle? Energy is quantized. Light is energy….. light must be quantized photons photons smallest pieces of light Photoelectric effect – Matter emits e- when it absorbs energy Albert Einstein Nobel Prize in chem Energy & frequency: directly related.

56 Planck-Einstein Equation: E = h E = Energy, in units of Joules (kg·m 2 /s 2 ) (Joule is metric unit of energy) (Joule is metric unit of energy) h = Planck’s constant (6.626 x J·s) (reflecting sizes of energy quanta) quanta = frequency, in units of hertz (hz, sec -1 ) = frequency, in units of hertz (hz, sec -1 ) energy (E ) of electromagnetic radiation is directly proportional to frequency ( ) of radiation.

57 The Math in Chapter 5 There are 2 equations: 1) c = 2) E = h Put these on your 3 x 5 notecard!

58 Examples 1) What is the wavelength of blue light with a frequency of 8.3 x hz? 2) What is the frequency of red light with a wavelength of 4.2 x m? 3) What is the energy of a photon of each of the above?

59 Explanation of atomic spectra electron configurations written in lowest energy. ground state energy level, and where electron starts from, called it’s ground state - lowest energy level.

60 Changing the energy Let’s look at a hydrogen atom, with only one electron, and in the first energy level.

61 Changing the energy “ excited ” Heat, electricity, or light can move electron up to different energy levels. The electron is now said to be “ excited ”

62 Changing the energy light As electron falls back to ground state, it gives energy back as light

63 Experiment #6, page 49-

64 may fall down in specific steps Each step has different energy Changing the energy

65 { { { Lyman series (UV) Balmer series (visible) Paschen series (infrared)

66 further they fall, more energy released = higher frequency orbitals also have different energies inside energy levels All electrons can move around. Ultraviolet Visible Infrared

67 What is light? Light is a particle - it comes in chunks. Light is a wave - we can measure its wavelength and it behaves as a wave combine E=mc 2, c=, E = 1/2 mv 2 and E = h, then we can get: = h/mv (from Louis de Broglie) Calculates wavelength of a particle. called de Broglie’s equation He said particles exhibit properties of waves

68 Wave-Particle Duality J.J. Thomson won the Nobel prize for describing the electron as a particle. His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron. The electron is a particle! The electron is an energy wave!

69 Confused? You’ve Got Company! “No familiar conceptions can be woven around the electron; something unknown is doing we don’t know what.” Physicist Sir Arthur Eddington The Nature of the Physical World 1934

70 The physics of the very small Quantum mechanics explains how very small particles behave Quantum mechanics is an explanation for subatomic particles and atoms as waves Classical mechanics describes the motions of bodies much larger than atoms

71 Heisenberg Uncertainty Principle impossible to know exact location and velocity of particle better we know one, less we know other Measuring changes properties. True in quantum mechanics, but not classical mechanics

72 Heisenberg Uncertainty Principle You can find out where the electron is, but not where it is going. OR… You can find out where the electron is going, but not where it is! “One cannot simultaneously determine both the position and momentum of an electron.” Werner Heisenberg

73 It is more obvious with the very small objects To measure where e-, we use light But light energy (photon) moves e- due to small mass And hitting e- changes frequency of light

74 Moving Electron Photon Before Electron velocity changes Photon wavelength changes After Fig. 5.16, p. 145

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