Chapter 5 Electrons in Atoms p. 126

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

“Why don’t e-s fall into nucleus”? 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

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

fixed energy e- have called Energy levels Bohr’s model fixed energy e- have called Energy levels Like rungs of 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

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

The Quantum Mechanical Model energy levels for e- Orbits not circular Based on probability of finding e- certain distance from nucleus electron cloud

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

s and p orbitals 1:20

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

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

Summary of Principal Energy Levels, Sublevels, and Orbitals Number of sublevels Type of sublevel Max # of electrons Electron configuration n = 1 1 1s (1 orbital) 2 1s2 n = 2 2s (1 orbital 2p (3 orbitals) 8 2s2 2p6 n = 3 3 3s (1 orbital) 3p (3 orbitals) 3d (5 orbitals) 18 3s2 3p6 3d10 n = 4 4 4s (1 orbital) 4p (3 orbitals) 4d (5 orbitals) 4f (7 orbitals) 32 4s2 4p6 4d10 4f14

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

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

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

Section 5.2 Electron Arrangement in Atoms p. 133 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 - German for “building up”

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

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

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 #

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

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

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

Increasing energy The next e-’s go in 2p orbital 7p 6d 5f 7s 6p 5d 6s

Increasing energy The next e-’s go in 3s orbital 7p 6d 5f 7s 6p 5d 6s

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

An internet program about electron configurations is: I electron config (song) 3:24

Adding e-’s changes energy of orbital 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

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!!

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

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

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

Section 5.3 Physics and the Quantum Mechanical Model p. 138 Light Section 5.3 Physics and the Quantum Mechanical Model p. 138 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

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

Parts of a wave Crest Wavelength Amplitude Trough

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)

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

- Page 140 Use Equation: c =

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

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

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

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

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

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

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

E = Energy, in units of Joules (kg·m2/s2) 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)

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

What is the frequency of red light with a wavelength of 4.2 x 10-5 m? 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?

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.

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

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

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

Experiment #6, page 49-

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

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

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

= h/mv (from Louis de Broglie) 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=mc2 , c=, E = 1/2 mv2 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

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!

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

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

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

Heisenberg Uncertainty Principle “One cannot simultaneously determine both the position and momentum of an electron.” 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! Werner Heisenberg

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

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