Chapter 3 Part II

How are theelectronsarranged

around thenucleus?

Normal or Rest Position

Crest

Wavelength Amplitude

Amplitude

Wavelength Trough

3

Parts of a Wave

Parts of a WaveA. Amplitude: Height of the wave. The higher

the wave the greater the intensity.

B. Wavelength: (λ , “lambda”) in nanometers (1 x 10-9 m) measured from crest to crest or from trough to trough.

C. Frequency: (ν , “nu”) The number of times a wave passes a fixed point. Measured in cycles/second (1/s) 1 cycle/second = Hertz (Hz)ex) Radio FM 93.3 megahertz (MHz) is 93.3 x 106

cycles/sec.

Visible Light

Radar

Microwaves

Infrared

Radio/TV Ultraviolet

X-Rays

Gamma Rays

Low

Long

High

Short

Red Orange Yellow Green

Blue Indigo Violet

Energy

Energy

Low High

The Electromagnetic Spectrum

http://www.centennialofflight.gov/essay/Dictionary/ELECTROSPECTRUM/DI159G1.htm

Speed of Light

D. (c) 3.0 x 108 m/s or 186,000 miles/sec. The relationship between wavelength and frequency can be shown with the

following equations:

or this is an indirect relationship.

If λ then ν .

cλ= νcν =λ

Quantum Theory

A. Planck’s Hypothesis: (Max Planck 1900)

1. An object absorbs or emits light in

little packet or bundle called a quantum (quanta –plural).

2. Energies are quantized. (Think steps not a ramp.)

3. Equation relating energy (E) to frequency (ν, nu) (h= Planck’s Constant)

E = h This is a direct relationship,

as ν , E

e-

e- Xe-

e-

Light (Electromagnetic Radiation)

1. A quantum of light is called a photon.2. Light travels through space in waves.3. Light acts like a particle when it interacts

with matter.4. This shows the Dual Nature of Light.

3. The AtomA. Atomic Emission Line Spectra: contains

only certain colors or wavelengths ( ) of light.

1. Every element has its own line spectrum (fingerprint).

Double Slit Experiment

Double Slit Experiment Video

This Experiment Proves the Dual Nature of Light – Photons of Light

travel through spaces in waves, but act like particles when they interact

with matter.

Continuous Spectrum – White Light

Line Spectrum – Excited Elements

Line Emission Spectra of Selected Elements

Gas Discharge Tubes

• Electricity is added to the gas which causes the electrons to jump to a higher or excited state. They immediately fall back to the ground state and give off particular wavelengths of light. We see a blending of wavelengths without the spectroscopes.

White Light gives off a Continuous Spectrum

a blending of every possible wavelength

Spectroscope• Uses a diffraction grating to diffract the

light into particular wavelengths of light.

A Line Spectrum results from excited elements - as electrons of an element gain

energy and rise to an excited state they then fall back to their ground state in the same

pattern producing the same energy drop each time which we see as individual wavelengths

of light.

Atomic Spectra

Hydrogen

Lithium

Mercury

Helium

Although Bohr’s atomic model explained the line spectra of hydrogen, it failed for heavier elements.

B. Bohr Model of Hydrogen: (1911)

1. Bohr said the energy of an electron was quantized (only certain orbits that represented different amounts

of energy.)

2. Bohr labeled each energy level with a quantum number (n).

3. n=1 lowest level or ground state.

4. When electrons absorb energy they jump to a higher (excited) state. n=2 n=3 n=4 n=5 n=6 n=7

5. Radiation (light) is emitted when an electron falls back from

a higher level to a lower level.

Electrons release energy as they fall back to a lower

energy level

Excited Atoms Emit Photons of Light:

Electrons absorb energy to rise to a higher or excited state and emit energy in the form of a photon of light as they fall back to their

ground states.

Path of an excited electron as it “falls” back to the Ground State

• When electrons gain energy, they jump to a higher energy level (excited state).

• Electrons are not stable at the excited state and will immediately fall back to a lower level or ground state.

• As they fall, they emit electromagnetic radiation.

• Depending on how far they fall determines the type of radiation (light) released.

electrons fall to n = 1 and give off ultraviolet light.

electrons fall to n = 2 and give off visible light.

electrons fall to n = 3 and give off infrared light.

Hydrogen Atom

Ultraviolet Light

Visible Light

Infrared Light

Werner Heisenberg: (1927)

1. Heisenberg’s Uncertainty Principle: states the

position and momentum (speed, direction & mass) of

a moving object cannot simultaneously be measured and known exactly.

2. You cannot predict future locations of particles.

3. He found a problem with the Bohr Atom - no way to observe or measure the orbit of an electron.

4. Quantum Mechanics

A. Quantum Mechanical Model of the atom combines previous ideas and treats the electron like a wave that has quantized

energy.

Impossible to state the exact position or momentum of an electron, but you can state a probability of where the electron is located.

B. Electron Density

Where the density ofan electron cloud ishigh there is a highprobability that iswhere the electron islocated. If the

electrondensity is low then there is a lowprobability.

C. Atomic Orbitals - region around the nucleus where an electron with a given energy is likely to be found (not the same as Bohr’s orbits!)

1. Orbitals have characteristic shapes, sizes, &

energies.2. Orbitals do not describe how the electron

moves.3. The drawing of an orbital represents the

3-dimentional surface within which the electron is found 90% of the time.

4. Sublevels can have 4 different shapes

s – orbital spherical

1s, 2s & 3s orbitals Superimposed on one another

Electron-Cloud Models

p-orbital – dumbbell shaped

p-orbital - dumbbell shaped

d-orbital - double dumbbell or fan blades

s,p and d orbitals

For a more complete representation and presentation of atomic orbitals go to http://winter.group.shef.ac.uk/orbitron/

x

y

z

x

y

z

x

y

z

x

y

z

x

y

z

x

y

z

x

y

z

x

y

z

x

y

z

s orbital p orbitals

d orbitals

Models of d-orbitals

f-orbital – more

complex!

Quantum Numbers - Finding an address for each

electron:1) “state” Principle Quantum Number (n) or

the energy level ranges from n=1 to n=7

2) “city” Sublevel shape either s, p, d,or f.

3) “street” Orbital The orientation in space ex) x,y,z axis

4) “house” Spin The cw or ccw motion of electrons.

1s, 2s,2p and 3s orbitals superimposed on each other

• Model of s and p Orbitals Together

2. The number of Sublevels in an energy level

equals the Principle Quantum Number (n).

3. Orbital: There are always an odd number of orbitals.

s-sublevel has 1 orbitalp-sublevel has 3 orbitalsd-sublevel has 5 orbitalsf -sublevel has 7 orbitals

• Orbitals in higher principle levels get larger.• A max of 2 electrons fit in each orbital.

Electron Spin

a. Two electrons in each orbital have opposite

spins. (clockwise and counterclockwise)

b. The opposite spin is written: or ___

c. Pauli Exclusion Principle:

1. Each orbital can only hold 2 electrons.

2. The electrons must have opposite spins.

s-sublevel = max 2 electronsp-sublevel = max 6 electronsd-sublevel = max 10 electronsf-sublevel = max 14 electrons

incorrect: ↑↑↑ incorrect: ↑↑ correct: ↑↓

d. Hund’s Rule:• Electrons will

spread out among the orbitals before they pair up.

incorrect ↑↓ ↑ __

correct ↑ ↑

↑

E. Electron Configurations:

1. Shows the distribution of electrons among the orbitals. Describes where the electrons are found & what energy they possess.

2. The Aufbau Principle: electrons are added up one at a time to the lowest energy orbital.

Aufbau Diagram/ Diagonal Rule

Electron Configuration Examples:

Ex) electron configuration for Na:

1s2 2s2 2p6 3s1

Ex) orbital filling box diagram for Na:

yx z

1s 2s 2p 3s

_

3. Electron Dot Diagrams:

Write the symbol for the element.

Place dots around the symbol to represent the

valence s & ps & p electrons only.

Do NOT include d & f orbitals in diagram.

p orbital electrons s orbital electrons

Electron Configuration

Aufbau Diagram Order

Aufbau Diagram

Visualizing The Electron~

Model of Bohr Atom - Electron Movement

Line Emission Spectra of Selected Elements

The Aufbau Principle(Diagonal Rule)

1s2

2s2 2p6

3s2 3p6 3d10

4s2 4p6 4d10 4f14

5s2 5p6 5d10 5f14

6s2 6p6 6d10 6f14

7s2 7p6 7d10 7f14

Incr

easi

ng e

nerg

y

1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

Pauli Exclusion Principle: No more than 2 e- are put in each orbital and they must have opposite spin.Hund’s Rule: electrons spread out among equal energy orbitals in a sublevel (like charges repel)Aufbau Principle: Electrons fill lowest energy levels first (n=1)

Electron Blocks on the

Periodic Table

Electron Configuration Orbital Box Diagram Electron-dot Diagram

yx z

1s 2s 2p

y yx z x

1s 2s 2p 3s 3p

z

168O

3517 Cl 1s22s22p63s23p5

12752Te 1s22s22p63s23p64s23d104p65s24

d105p4

1s22s22p4

y y y yx z x z x z x z

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p

What does the Tellurium electron-dot resemble???

The Copper Atom

Mark your Periodic Tables

1 2 13 14 15 16 17 18

Scanning Tunneling Microscope

In 1926, Erwin Schrodinger derived an equation that described the energy and

position of the electrons in an atom

22

2 2

8dh EV

m dx

Equation for the probabilityprobability of a single electron being found along a single axis (x-axis)

Erwin SchrodingerErwin Schrodinger