The Bohr and Quantum Models

Starts with the Study of…..Light

Light is….Made up of electromagnetic

radiation.Waves of electric and magnetic

fields at right angles to each other.

Parts of a wave

l Wavelength

Wavelength = the distance between any point on a wave and the corresponding point on the next wave (crest to crest, trough to trough); m, nm, pm

l lambda

Parts of a wave

Frequency ( )n = number of cycles in one second; hertz, hz1 hertz = 1 cycle/second - n nu

Frequency = n

Kinds of EM waves

There are many different l and n (p. 92) Radio waves, microwaves, x rays and

gamma rays are all examples. Visible light is all our eyes can detect.

GammaRays

Radiowaves

Electromagnetic Spectrum

Energy is Quantized - PlanckQuantum – a packet of energy;

smallest quantity of energy that can be emitted or absorbed

Einstein is next

Light is particulatephoton – a quantum of

electromagnetic radiation

Which is it?

Is energy a wave like light, or a particle?Yes !

Concept is called the Wave -Particle duality (Dual Nature of Light)

What about the other way, is matter a wave? Yes !

Spectrum

The range of frequencies present in light.

White light has a continuous spectrum.–All the colors are possible.–A rainbow.

Electromagnetic Spectrum

Hydrogen spectrum

Line spectrum – series of separated lines of different colors representing photons whose wavelengths are characteristic of one element

410 nm

434 nm

486 nm

656 nm

What this means

Only certain energies are allowed for the hydrogen atom.(Recall relationship b/w energy, wavelength and frequency)

Energy in the atom is quantized. Can calculate Energy based on line

spectra Only the energy we can “see”

Quantum Theory: Colors emitted can be

used to identify elements (absorption of energy and color emitted is a fingerprint of an element)

Kind of like wearing your team colors.

            

Team Oxygen

            

Team Carbon

Let’s Play!!

White light - dispersionFlame Tests

Niels Bohr

The Bohr Ring AtomQuantized energy – energy levelsGround state – ion or atom at its

lowest energy stateExcited state – atom or molecule at

any other state besides ground state–Putting energy (photon) into the atom

moves the electron away from the nucleus.

Bohr

From ground state (lowest energy level) to excited state.

When it returns to ground state (or lower energy level) it gives off light of a certain energy (photon.)

Energy, frequency, wavelength – color of light

The Bohr Ring Atom

n = 3n = 4

n = 2n = 1

Bohr Planetary Model

quantum model of the hydrogen atom

atom was like a solar system; protons in dense nucleus

electrons orbit (specific path electron travels)

ground and excited state

The Bohr Model

Only works for hydrogen atoms. Electrons don’t orbit. Energy quantization correct, but not circling

like planets Introduced concept of energy levels

The Quantum Model

Atomic Models: Old version = Bohr’s Also known as the planetary

atomic model Describes electron paths as

perfect orbits with definite diameters

Good for a visual

New version = Quantum Theory

Most accepted Diagrams electrons of a

atom based on probability of location at any one time

Basic Concepts of the Quantum Model Atoms and molecules can exist only in

certain energy states Atoms absorb or emit radiation as they

change energies Orbitals NOT orbits Energy states are described by sets of

numbers called quantum numbers MATHEMATICAL model

Heisenberg’s Uncertainty Principle We can’t know how the electron is

moving or how it gets from one energy level to another.

Heisenberg Uncertainty Principle:It is impossible to know the exact position AND momentum of a particle simultaneously. (Marco Polo)

Electron Configuration•The distribution of electrons

within their atoms orbitals•Distributed amongst energy levels, sublevels, and orbitals

•Based on a set of stated principles

AUFBAU PRINCIPLE

Electrons are added one at a time to the lowest energy orbitals

“building up” principleReference sheet

Electron Configurations

H 1s1

Energy Level

sublevel

# of electrons in sublevel

How to Determine Electron Configurations and Orbital Diagrams

1. Determine total # of electrons

2. Use Aufbau Principle to “fill”

3. Make sure all electrons of the atom are accounted for!

5s

4p

3d

4s

3p

3s

2p

2s

1s

En

ergy18 electrons

Argon (18 electrons)

5s

4p

3d

4s

3p

3s

2p

2s

1s

En

ergy18 electrons

Argon (18 electrons)

1s22s22p63s23p6

Orbital Diagram – box, circle or line for each orbital in a given energy level, grouped by sublevel, with an arrow indicating electron AND its spin

HUND’S RULE

Hund’s Rule (BUS RULE)- The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons in the orbital.

C

1s2 2s2 2p2

Details

Valence electrons- the electrons in the outermost energy levels (not d).

Core electrons- the inner electrons.

TRY!

Calcium and Manganese

5s

4p

3d

4s

3p

3s

2p

2s

1s

En

ergy

Calcium (20 electrons)

1s22s22p63s23p64s2

Try Oxygen, Iron, Strontium, Molybdenum, Tungsten!

Orbital Diagrams AND Electronic Configuratiions

Using the Periodic Table!

Noble Gas Configuration!

P

E

R

I

O

D

S

1-7

Blocks and Sublevels

d (n-1)1

2

3

4

5

6

7

4

5

Label your blank periodic table.

Read it “like a book”

1. C

2. Kr

3. Ca

4. Fe

5. Hg

WRITE the Electron Configuration using the periodic table:

Electron Configuration – Noble Gas Configuration Electron Configuration demonstrates a

periodic trend, so you can write shorthand electron configuration using the electron configuration of the noble gases in Group 18 of the periodic table.

Noble gases have stable configurations.

Noble Gas Configuration

When writing shorthand e- config for an element, refer to the noble gas in the energy level (period) just above the element.

Write the symbol of the noble gas in brackets.

Write out the remaining e-config based on the energy filling diagram.

Electron ConfigurationNa = 1s22s22p63s1

Al = 1s22s22p63s23p1

Ne = 1s22s22p6

Shorthand Electron ConfigurationNa = [Ne] 3s1

Al = [Ne] 3s23p1

EX: NaStep 1: Na is in period 3 so refer to the

noble gas in period 2 which is Neon.

Step 2: Write Ne in brackets. [Ne]

Step 3: Now write remaining electrons in standard form. 3s1.

Step 4: Combine. [Ne]3s1

Noble Gas Configuration

Now try:

1. I

2. Kr

3. Na

4. Cu

Nobel Gas Configuration

Electron Configuration with Ions

When we write the electron configuration of a positive ion, we remove one electron for each positive charge:

Na → Na+

1s2 2s2 2p6 3s1 → 1s2 2s2 2p6

When we write the electron configuration of a negative ion, we add one electron for each negative charge:

O → O2-

1s2 2s2 2p4 → 1s2 2s2 2p6

Now try:

1. Ca+2

2. Fe-3

Electron Configuration with Ions

Details

Elements in the same column have the same electron configuration.

Put in columns because of similar properties.

Similar properties because of electron configuration.

Noble gases have filled energy levels.

Exceptions

Ti = [Ar] 4s2 3d2 V = [Ar] 4s2 3d3

Cr = [Ar] 4s1 3d5

Mn = [Ar] 4s2 3d5

Half filled orbitals. Scientists aren’t sure of why it happens same for Cu [Ar] 4s1 3d10

More exceptions

Lanthanum La: [Xe] 6s2 5d1

Cerium Ce: [Xe] 6s2 4f1 5d1

Promethium Pr: [Xe] 6s2 4f3 5d0

Gadolinium Gd: [Xe] 6s2 4f7 5d1

Lutetium Pr: [Xe] 6s2 4f14 5d1

We’ll just pretend that all except Cu and Cr follow the rules.

Quantum Numbers Quantum Numbers represent 4 solutions to

Schroedinger’s Equation Describe the distribution of electrons in

space BASED ON PROBABILITY!!!

1. Principal Quantum Number(n)

Describes main energy level an electron occupies

n=1,2,3,4,5,6,7

2. Angular Momentum Quantum(l)

sublevel designates the shape of the region in space an

electron occupies (orbital). integer values from 0 to n-1 l = 0 is called s Sublevels include s (l=0, p=1, d=2,

f=3) Not all energy levels have all

sublevels

3. Magnetic quantum number(Orbitals ml)

m l

designates the spatial orientation of orbital integer values between - l and + l s: ml = 0

p: ml = -1, 0, +1 (i.e. px, py, pz) s-1, p-3, d-5, f-7

S orbitals

P orbitals

P Orbitals

D orbitals

F orbitals

F orbitals

4. Spin Quantum Number

m s

Electrons are negatively charged Behave as if spinning on axis – repel so… either +1/2 or -1/2

Pauli Exclusion Principle

NO TWO ELECTRONS MAY HAVE THE SAME FOUR QUANTUM NUMBERS

(Spin must be different!)