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The Evolution of Atomic Models: Rutherford to Schrodinger

Explore the journey from Rutherford’s dense nucleus to Schrodinger’s quantum mechanical model, unravelling the mysteries of electron arrangement in atoms.

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The Evolution of Atomic Models: Rutherford to Schrodinger

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  1. Chapter 5Electrons in Atoms

  2. Ernest Rutherford’s Model • Discovered dense + nucleus • e-s move like planets around sun • Mostly empty space • Didn’t explain chemical properties of elements

  3. The Bohr Model • Neils Bohr (1885-1962) • Danish physicist • Student of Rutherford “Why don’t e-s fall into nucleus”? • Why don’t electrons fall into nucleus? Niels Bohr

  4. The Bohr Model I pictured the electrons found in specific circular paths around the nucleus, and can jump from one level to another. Furthermore, each level has a fixed amount of energy different from other levels Niels Bohr

  5. Bohr’s model • fixed energy e- have called Energy levels • Like rungsof ladder • e- can’t exist btwn energy levels • energy levels not evenly spaced • High levels closer (less energy needed to jump) Bohr’s model of the atom 5:17

  6. The Quantum Mechanical Model • e-’s don’t move like big objects • Rutherford & Bohr model • Energy - “quantized” (in chunks) • exact energy needed to move e- 1 energy level called a quantum • energy never “in btwn” • quantum leap in energy must exist • Erwin Schrodinger (1926) mathematically described energy & position of e- in atom Schrodinger Quantum Leap TV intro

  7. The Quantum Mechanical Model • energy levels for e- • Orbits not circular • Based on probabilityof finding e- certain distance from nucleus • electron cloud

  8. Atomic Orbitals • Principal Quantum # (n) - energy level of e- (1, 2, 3 etc.) • atomic orbitals - regions of space w/high probability of finding e- (not a true “orbit”) • within each energy level • Sublevels like rooms in a hotel • s, p, d, and f • Different shapes How many e- in level 2? level 3? Max # of e- that fit in energy level is: 2n2

  9. s and p orbitals 1:20

  10. d orbitals 3:40 atomic orbitals review (14:28)

  11. atomic orbitals First possible energy level # of orbitals (regions of space) Maximum electrons s spherical 2 1 1st p dumbell 6 3 2nd 10 d clover leaf 5 3rd 14 f complicated 7 4th

  12. ORDER OF ELECTRON SUBSHELL FILLING:NOT “IN ORDER” • energy levels overlap • Lowest energy fill first 1s2 2s2 2p6 3s2 3p6 3d10 Increasing energy 4s2 4p6 4d10 4f14 5s2 5p6 5d10 5f14 6s2 6p6 6d10 7s2 7p6 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6

  13. ELECTRON CONFIGURATION group # # valence e- s: 1 or 2 p: 1-6 d: 1-10 f: 1-14 Total e- = Atomic # 1s1 Principal energy level row # 1-7 7 rows sublevel s, p, d, or f 4 sublevels

  14. 1A 8A 1 7A 4A 6A 2A 3A 5A 2 8B 3 1B 2B 6B 5B 4B 7B 3B 4 d 5 6 7 f 6 7 Sublevels d and f are “special” group # = # valence e- 3d period # = # e- energy levels 4d 5d 6d 4f 5f

  15. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s Section 5.2 Electron Arrangement in Atoms aufbau diagram - page 133 Aufbau - German for “building up”

  16. Electron Configurations… • ….3 rules explain how e-’s fill their orbitals: • Aufbau principle– e-’s enter lowest energy level first. • Pauli Exclusion Principle- 2 e-’s max/orbital (hotel room) - different spins

  17. Pauli Exclusion Principle No 2 electrons in an atom can have the same four quantum numbers. To show different direction of spin, a pair in the same orbital is written as: Wolfgang Pauli

  18. Quantum Numbers Each e- has unique set of 4 quantum #’s describing it 1) Principal quantum # 2) Angular momentum quantum # 3) Magnetic quantum # 4) Spin quantum #

  19. Electron Configurations • Hund’s Rule- When e-’s occupy orbitals of same energy, they won’t pair up until they must • write e- configuration for Phosphorus • all 15 e-’s must be accounted for

  20. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The first 2 e-’s go into the 1s orbital Notice opposite direction of spins

  21. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next e-’s go in 2s orbital

  22. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next e-’s go in 2p orbital

  23. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next e-’s go in 3s orbital

  24. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The last 3 e-’s go in 3p orbitals They each go into separate shapes (Hund’s) • 3 unpaired e-’s = 1s22s22p63s23p3 Orbital notation

  25. Filling orbitals • Lowest  higher energy • Adding e-’s changes energy of orbital • Full orbitals best situation • half filled orbitals next best • more stable • Changes filling order

  26. Write the electron configurations for these elements: • Titanium - 22 electrons • 1s22s22p63s23p64s23d2 • Vanadium - 23 electrons • 1s22s22p63s23p64s23d3 • Chromium - 24 electrons • 1s22s22p63s23p64s23d4 (expected) • But this is not what happens!!

  27. Chromium is actually: • 1s22s22p63s23p64s13d5 Why? • 2 half filled orbitals • Half full slightly lower in energy • Same applies to copper

  28. Copper’s e- configuration • Copper has 29 e-s so expect: 1s22s22p63s23p63d94s2 • actual configuration is: • 1s22s22p63s23p63d104s1 • 1 more full orbital & 1 half filled • Exceptions • d4 • d9

  29. Irregular configurations of Chromium and Copper Chromium steals a 4s e- tomakeits 3d sublevel HALF FULL Copper steals a 4s electron to FILL its 3d sublevel

  30. Section 5.3Physics and the Quantum Mechanical Modelp. 138 Light • Study of light led to quantum mechanical model • Light is electromagnetic radiation • EM radiation: gamma rays, x-rays, radio waves, microwaves • Speed of light = 2.998 x 108 m/s • “c” - celeritas (Latin for speed) • All EM radiation travels same in vacuum

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

  32. Crest Wavelength Amplitude Trough Parts of a wave

  33. Electromagnetic radiation propagates through space as a wave moving at the speed of light. Equation: c = c = is a constant (2.998 x 108 m/s)  (lambda) = wavelength, in meters (nu) = frequency, in units of hertz (hz or sec-1)

  34. Wavelength and Frequency • inversely related • one gets bigger, other smaller • Different frequencies = different colors • wide range of frequencies (spectrum)

  35. - Page 140 Use Equation: c =

  36. Low Frequency High Frequency High Energy Low Energy X-Rays Radiowaves Microwaves Ultra-violet GammaRays Infrared . Long Wavelength Short Wavelength Visible Light

  37. Long  = Low Frequency = Low ENERGY Short  = High Frequency = High ENERGY

  38. Atomic Spectra • White light all colors of visible spectrum • prism separates it according to λ

  39. If the light is not white • heating gas with electricity will emit colors • this light thru prism is different

  40. Atomic Spectrum • elements emit own characteristic colors • composition of stars determined thru spectral analysis

  41. atomic emission spectrum • Unique to each element, like fingerprints! • ID’s elements

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

  43. energy (E ) of electromagnetic radiation directly proportional to frequency () of radiation. Planck-Einstein Equation:E = h E = Energy, in units of Joules (kg·m2/s2) (Joule…metric unit of energy) h = Planck’s constant (6.626 x 10-34 J·s) (reflecting sizes of energy quanta) = frequency, units of hertz (hz, sec-1)

  44. The Math in Chapter 5 • There are 2 equations: • c =  • E = h Put these on your 3 x 5 notecard!

  45. Examples • What is the wavelength of blue light with a frequency of 8.3 x 1015 hz? • What is the frequency of red light with a wavelength of 4.2 x 10-5 m? • What is the energy of a photon of each of the above?

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

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

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

  49. Changing the energy • As electron falls back to ground state, it gives energy back as light

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