Chapter 4.1 Properties of Light

- Wave-Particle Nature of Light
- Electrons and light have a dual wave-particle

nature. - Electromagnetic Radiation (EMR)
- Form of energy that exhibits wavelike behavior

and travels at the speed of light. - Speed of Light (C) 3 x 10 8 m/s

Components of a Wave

- Wavelength (?) lambda
- Units any unit of length (m)
- Distance between corresponding points of a wave.
- Crest to Crest or Trough to Trough

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Components of a Wave

- Frequency (?) nu
- Units Hertz (Hz) or 1/s
- How often a wavelength passes a given point in

time.

Components of a Wave

- Amplitude
- Height of the wavelength.
- Measured from the origin to crest or origin to

trough. - Brightness of light.

Wavelength vs. Frequency

- C ? ?
- Inversely proportional.
- As wavelength increases, frequency decreases.

- Spectrums
- Range of wavelengths for a series of waves.
- Electromagnetic Spectrum
- Consist of all electromagnetic radiation.
- Continuous Spectrum

- Spectrum where all wavelengths within a given

range are together. - Examples Visible Light, X-Rays, U.V. Light, etc

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EMR Spectrum

- 7 Parts
- Longest wavelength to Shortest
- Radio
- Microwaves
- Infrared
- Visible Light
- U.V. Light
- X-Rays
- Gamma-Rays

- Problems
- What is the wavelength of EMR that has a

frequency of 7.50 x 10 12Hz?

- Problems
- 1. Determine the frequency of light with a

wavelength of 4.257 x 10-7 cm. - 2. What is the wavelength of U.V light that has a

frequency of 4.50 x 10 16 Hz? - 3. What is the wavelength and color of light,

that has a frequency of 6.00 x 10 14 KHz?

Chapter 4.2 Quantum Theory

- Photoelectric Effect
- Emission of electrons by certain metals when

sufficient light shines on them. - Photoelectric effect

Chapter 4.2 Quantum Theory

- Photoelectric Effect

Chapter 4.2 Quantum Theory

- Photoelectric Effect

Chapter 4.2 Quantum Theory

- Quantum
- Finite quantity of energy that can be gained or

lost by an atom. - Plancks Equation
- h 6.63 x 10 34 Js
- E quantum of energy
- Photon
- An individual quantum of light, caused by

electrons losing quanta of energy.

E h ?

Chapter 4.2 Quantum Theory

- Visible Light Emissions
- As electrons gain quanta of energy they release

it in the form of photons. - Energy States of an Atom
- Ground State- an atoms lowest energy level.
- Excited State- an atoms highest energy level.
- , is produced when

electrons drop from the excited to the ground

states.

Line Spectrum

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Chapter 4.2 Quantum Theory

Chapter 4.2 Quantum Theory

- Problems
- What is the energy of U.V. light with a frequency

of 4.50 x 10 16 Hz? - E h ?, h 6.63 x 10 34 Js

Chapter 4.2 Quantum Theory

- Problems
- Determine the energy of light that has a

wavelength of 450nm.

- Equations

- Problems
- 1. What is the energy of a photon of green light

with a frequency of 5.80 x 1014 1/s? - 2. What is the energy, in joules, of a quantum of

radiant energy whose wavelength is 6.82 x 10 6

cm? - 3. Determine the wavelength of a photon that has

3.11 x 10 19 J of energy. - 4. Determine the frequency, in MHz, of a photon

that has wavelength of 1.36 x 10 10 nm.

Summary

- restricted the amount of energy

that an object emits or absorbs as a quantum. (E

h? ?) - used Plancks theory and

explained the photoelectric effect. - light travels as tiny particles,

photons.

Planck

Einstein

Compton

Bohrs Model

- The Line Spectra demonstrates that the energy

levels of an electron in an atom are quantized - Similar to the rungs of a ladder, nothing exist

in between. - (For Hydrogen (1 p 1 e- )
- 1st Energy Level n 1
- 2nd so on n 2,3,4,5,6, 8
- Only electrons dropping from a Higher Level to a

Lower one emit EMR - A Number of Possibilities for electron drops

Hydrogens Line Spectrum

- Several Series of lines are observed
- Electron Drops to the n 1 Level
- Lyman Series (U.V. Range)
- Electron Drops to the n 2 Level
- Balmer Series (Visible Range)
- Electron Drops to the n 3 Level
- Paschen Series (Infrared Range)

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- The Lines become more closely spaced as the

levels increase - Theres an upper limit to how high the electron

can jump. - The Bohr model explained spectral lines but not

how atoms bonded. - Ultimately Displaced

1924 Louis de Brogile

- French Graduate Student (asked an important

question) - If light behaves as waves particles, can

particles of matter behave as waves? - Derived an Equation

- Predicts that all matter exhibits wavelike

motions.

h Planks Con. m mass v - velocity

- Large Objects
- Small Wavelengths
- 200 g Baseball _at_ 30 m/s - ? 10-32 cm
- Undetectable
- Small Objects Large Wavelengths
- 9.11 x 10-28 g _at_ 30 m/s - ? 10-3 cm
- Very Detectable w/ proper instruments
- New Ballgame Classical Mechanics vs. Quantum

Mechanics - New method for describing the motions of

subatomic particles

Heisenbergs Uncertainty Principle

- It is impossible to know exactly both the

velocity the position of a particle at the same

time. - Accuracy of V ? then Position ?

Classical Vs Quantum

- Classical adequately describes the motions of

bodies much larger than the atoms of which they

are composed. It appears that such a body loses

energy in any amount - Quantum describes the motions of subatomic

particles and atoms as waves. These particles

gain or lose energy in packages called quanta.

Quantum Mechanical Model

- Modern description of the electrons derived from

the mathematical solution to the Schrodinger

equation. - Erwin Schrodinger - used wave mechanics to show

the electrons about the nucleus emit vibration

frequencies that were constant. - Quantum Numbers - specify the properties of

atomic orbitals and their electrons. - distance from the nucleus.
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Chapter 4.4 Quantum Numbers

- Principal Quantum Number (n)
- Main energy level surrounding the nucleus.
- Size of each orbital.
- Primary distance from the nucleus.
- Has values of n 1 to 7, 1 is the closest 7 is

the farthest from the nucleus.

Chapter 4.4 Quantum Numbers

- Orbital Quantum Number (l)
- Shape of the orbitals.
- Referred to as subshells.

Chapter 4.4 Quantum Numbers

p orbital

s orbital

d orbital

f orbital

Chapter 4.4 Quantum Numbers

- Magnetic Quantum Number (m)
- Orientation of an orbital about the nucleus.
- l s
- m 0
- l p
- m -1, 0, 1
- l d
- m -2, -1, 0, 1, 2
- l f
- m -3, -2, -1, 0, 1, 2, 3

Chapter 4.4 Quantum Numbers

- s orbital, 1 orientation.

Chapter 4.4 Quantum Numbers

- p orbital, 3 orientations.

px orbital

py orbital

pz orbital

pxyz orbital

Chapter 4.4 Quantum Numbers

- d orbital, 5 orientations.

Chapter 4.4 Quantum Numbers

- f orbital, 7 orientations.

Chapter 4.4 Quantum Numbers

- Spin Quantum Number(1/2 , -1/2)
- Indicates two possible states on an electron in

an orbital.

Chapter 4.4 Quantum Numbers

- Magnetism
- Caused by the motion of electrons about the

nuclei of atoms. - Diamagnetism substance is weakly repelled by a

magnetic force. - Paramagnetism substance is weakly attracted by

a magnetic force. - Ferromagnetism Strong attraction by a magnetic

force.

Chapter 4.4 Quantum Numbers

Principal Energy Level Sublevels Orbitals

N1 1s 1s

N2 2s, 2p 2s(one) 2p(three)

N3 3s, 3p, 3d 3s(one) 3p(three) 3d (five)

N4 4s, 4p, 4d, 4f 4s(one) 4p(three) 4d (five) 4f(seven)

Chapter 4.4 Quantum Numbers

Principal Q.N. Orbitals per Main level (n2) Electrons per main level (2n2)

1 1 2

2 4 8

3 9 18

4 16 32

Chapter 4.4 Quantum Numbers

Orbital Max electrons

s 2

p 6

d 10

f 14

Rules Governing Electron Configurations

- Electron Configuration arrangement of electrons

in the atom - Rules
- Aufbau Rule electron occupies the lowest energy

level that will receive it. - Hunds Rule orbitals of equal energy each

receive one electron (equal spin) before any

receive two. - Paulis Exclusion Principle no two electrons

can have the same set of 4 quantum numbers

(n,l,m,s)

Orbital Notation

- Orbital Notation
- Orbital represented by a line ____ or
- Electron is represented by an ½ Arrow ???
- ½ (?)
- - ½ (?)
- Examples
- H (1 e-1) He (2 e-1) ?

Order of Energy Levels

- 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s

5f 6d - Number - principal quantum number, the main

energy level - Letter orbital quantum number, the shape
- Useful Diagram

1s 2s 3s 4s 5s 6s 7s

2p 3p 4p 5p 6p 7p

3d 4d 5d 6d 7d

4f 5f 6f 7f

Orbital Notation

- Write the orbital notation for the following

elements - Al
- Zn
- P
- Cl Atomic Orbitals

1s 2s 3s 4s 5s 6s 7s

2p 3p 4p 5p 6p 7p

3d 4d 5d 6d 7d

4f 5f 6f 7f

Electron-Configuration Notation

- Eliminates the lines arrows
- Superscripts are used to illustrate the number of

electrons in the sublevel - Same order of sublevels
- Examples
- H (1 e-1) - 1s1
- He (2 e-1) - 1s2
- Li (3 e-1) - 1s22s1

Electron-Configuration Notation

- Write the short-hand electron-configuration for

the following - Br
- Pb
- Cs
- Kr

1s 2s 3s 4s 5s 6s 7s

2p 3p 4p 5p 6p 7p

3d 4d 5d 6d 7d

4f 5f 6f 7f

Exceptions to Aufbau

- All elements prefer a more stable configuration

of electrons. - Fully filled and ½ filled orbitals are more

stable than others. - Elements that are 1 shy of a full or ½ filled d

orbital configuration will have electrons

transfer from the s to the d to reach this stable

state. - Example if you have a 4s2 and 3d4 the actual

configuration should be 4s1 and 3d5.

Practice

- Ba
- Mg
- W
- Ag
- Sb

Noble Gas Configuration

- Shortest method of writing the electron

configurations. - Use the last noble gas to occur prior to the

element that is being configured. - Start at the nS where n equals the period in

which the element being configured can be found. - Example Zr
- Noble gas would be Kr and start configuration at

5s. - Kr 5s24d2

Writing in noble gas configuration

- Al
- Write the noble gas configuration for the

following - V
- Rb
- I
- Hg
- U
- W

Identifying Electrons

- Paired electrons when 2 electrons are within

the same orbital. - Unpaired electrons when a single electron is

within an orbital.

Identifying Electrons

- How many unpaired electrons does the following

elements have? - Na
- O
- B

Electron Dot Notation

- Describes the number of electrons in an atoms

outer electron cloud, or highest energy level, as

dots around the symbol. - Valence electron - refers to the highest energy

level electron within an atom.

Electron Dot Notation

- Draw the electron-configuration and electron-dot

notation for Nitrogen.

Electron Dot Notation

- Write the noble gas-configuration and electron

dot notation for the following. - Al
- As
- I
- Sr