Modern Atomic Theory

“I can safely say that nobody understands

quantum mechanics”

Physicist Richard Feynman

1965 Nobel Prize winner for contributions to quantum mechanics

History of the atomic model

Dalton’s atomic theory =

Plum Pudding Model =

We left off with Rutherford:

atoms are indivisible spheres

atoms consist of (+) and (-) evenly mixed

• The atom is mostly empty space• Atoms have small, dense, & positively

charged nucleus at center

CRT experiment

Gold foil experiment

Problems with Rutherford Model…

• Where are the electrons?

• Why don’t they fall into the nucleus?

• What about the bright lines on an

emission spectrum of elements?

Models of the atom

Model =

• a theory or explanation for a phenomenon • Cannot prove a model (only disprove!)• Change as experiments dictate

• didn’t explain the arrangement of e-

So we need a new model! The new model came from experiments involving LIGHT!

Quantum Theory:

E (of light and e-) is quantized = composed of discrete bundles called quanta or photons

What was the major shortcoming with Rutherford’s Nuclear model?

Light

The study of light led to the development of the quantum mechanical model.

Light is a kind of electromagnetic radiation.

Electromagnetic radiation includes many kinds of waves

All move at 3.00 x 108 m/s (“c”)

Radio

waves

Micro

waves

Infrared

.

Ultra-

violet

X-

Rays

Gamma

Rays

Low

energy

High

energy

Low

FrequencyHigh

FrequencyLong

WavelengthShort

WavelengthVisible Light

Light is a form of Electromagnetic Radiation

Types of Electromagnetic

Radiation

What is the relationship between wavelength and frequency?

•Draw 3 different size waves on your paper

•Label them A, B & C

•Measure dist. between tops of each wave

•Count the number of “tops”

•Rank A, B, & C by:

•shortest to longest wavelength

•lowest to highest number of waves

•Are frequency and wavelength directly or

inversely related?

•Light and electromagnetic energy has

properties of both:

•Waves

•Particles photons with discrete E

•Electrons also have both particle and

wave properties

Spectra Analysis1) Continuous spectrum = • contains all colors from red to violet

•When passed through prism, see streak of color, not bands or lines

2) Bright line (Atomic/Emission) spectrum =

•contains only certain discrete l (wavelength)

•When passed through prism, see series of lines

•excited atoms emit discrete spectra

Line spectra of elements

What is the significance of a line spectrum?

•atoms can only give off certain E light

•e- can only possess certain amounts of E

•Gives clues to how e- are arranged

H

Hg

Ne

Bohr Model of the Atom:

•First quantum model of atom (e- has a

discrete quantity of E)

• e- arranged in concentric shells around

nucleus (= orbits)

•Move like Planets around the sun

•Electrons are treated like particles

+

n = 1n = 2

n = 3

Excited state=Higher E state

+

n = 1n = 2

n = 3

photon

of light

Ground state

photon

of light

+

n = 1n = 2

n = 3

Ground state=Lowest E state

Bohr Model of the Atom:

1) e- jumps from inner to an outer shell (unstable state)

2) then e- falls from outer to inner shell (more stable)

3) E released as photon of light

When Energy is absorbed:

This emitted E can be used to identify

an element.

Bohr’s Model

Energy Levels

Further away

from the nucleus

means more

energy.

There is no “in

between” energyIncr

easi

ng e

ner

gy

Nucleus

First

Second

Third

Fourth

Fifth

Why did the Bohr Model need to be

replaced?

•Bohr model only explained atomic spectrum

for hydrogen

•Needed a revised model that successfully

explained spectral patterns for multi-

electron elements

+

Bohr’s Model

Based on observation of light (lines)

Move like planets around the sun.

In circular orbits at different levels.

Energy different at each level

Electrons treated as particles

•Mathematical model based on e-

behaving as waves instead of particles

•Electrons are not in orbits but in Orbitals = region with high e- probability = Electron Cloud

•Explain emission lines for all elements(not just hydrogen)!

Schrodinger’s (Wave Mechanical) model of the atom:

The Quantum or “Wave”

Mechanical Model The atom is found

inside a blurry “electron

cloud”

A area where there is a

chance of finding an

electron.

Draw the edge of the

atom at 90 %

Principle Quantum Number

(aka Principle Energy Level)

tells the energy level of the electron.

There are 7 Principle Energy Levels

this corresponds to the 7 periods on periodic table!

Atomic Orbitals

Within each energy level the complex math of Schrodinger’s equation describes several shapes.

These are called atomic orbitals

regions where there is a high probability (90%) of finding an electron.

S orbitalsStarts at the first energy level (period)

1 shape – sphere

There is one s orbital (sphere) per energy level

They overlap each other

Each sphere can hold 2 electrons

P orbitalsStart at the second energy level

3 different shapes

Each shape can hold 2 electrons

6 electrons total per energy level!

P Orbitals

D orbitals Start at the third energy level

5 different shapes

Each shape can hold 2 electrons

10 electrons total per energy level!

F orbitalsStart at the fourth energy level

Have seven different shapes

2 electrons per shape

14 electrons total per energy level!

F orbitals

Summary

s

p

d

f

# of shapes

(orbitals)Max # of

electrons

1 2 1

3 6 2

5 10 3

7 14 4

Sublevel

Starts at energy level

3-D Orbitals - Scandium

http://www.youtube.com/watch?v=sYRX9O3a-bM

Principal Quantum Number, n

Indicates main energy levels

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

Each main energy level has sub-levels

Is the mostly the same as the Period on

the Periodic table except for transitional

(period -1) and inner transitional (period

-2)

Orbital Quantum Number, ℓ(Angular Momentum Quantum Number)

Indicates shape of orbital sublevels

ℓ = n-1

ℓ sublevel

0 s

1 p

2 d

3 f

4 g

11s

value of energy level

sublevel

no. ofelectrons

spdf NOTATION

for H, atomic number = 1spdf Notation

Orbital Box Notation

Arrows show electron spin(+½ or -½)

ORBITAL BOX NOTATIONfor He, atomic number = 2

1s

21 s

2 ways to write electron configurations

Electron ConfigurationsFirst Energy Level

only s sublevel (1 s orbital)

only 2 electrons

1s2

Second Energy Level

s and p sublevels (s and p orbitals are available)

2 in s, 6 in p

2s22p6

8 total electrons

Third energy level

s, p, and d orbitals

2 in s, 6 in p, and 10 in d

3s23p63d10

18 total electrons

Fourth energy level

s,p,d, and f orbitals

2 in s, 6 in p, 10 in d, and 14 in f

4s24p64d104f14

32 total electrons

Electron Configurations

The way electrons are arranged in atoms.

Aufbau principle- electrons enter the lowest energy first.

This causes difficulties because of the overlap of orbitals of different energies.

Pauli Exclusion Principle- at most 2 electrons per orbital - different spins (+1/2 & -1/2)

Periodic Table with Orbitals

Electron Configuration

Hund’s Rule- When electrons occupy orbitals

of equal energy they don’t pair up until they

have to .

Fill one orbital at a time until all filled with

one then add the second to make a pair

Half filled orbitals are more stable than

partly filled filled orbitals but not as stable

as completely filled orbitals

Orbitals fill in order

Lowest energy to higher energy.

Adding electrons can change the energy

of the orbital.

Half filled orbitals have a lower energy.

Makes them more stable.

Changes the filling order

Orbital

Diagram for

Hydrogen

Orbital

Diagram for

Helium

Orbital

Diagram for

Lithium

Orbital

Diagram for

Beryllium

Orbital

Diagram for

Boron

Orbital

Diagram for

Carbon

Orbital

Diagram for

Nitrogen

Orbital

Diagram for

Fluorine

Standard Notation

of Fluorine

Sublevels

Number of electrons

in the sub level 2,2,5

1s2 2s2 2p5

Write these electron

configurations

Titanium - 22 electrons

1s22s22p63s23p64s23d2

Vanadium - 23 electrons

1s22s22p63s23p64s23d3

Chromium - 24 electrons

1s22s22p63s23p64s23d4 is expected

But this is wrong!!

Chromium is actually

1s22s22p63s23p64s13d5

Why?

This gives us two half filled orbitals.

Slightly lower in energy.

The same principal applies to copper.

Copper’s electron

configuration

Copper has 29 electrons so we expect

1s22s22p63s23p64s23d9

But the actual configuration is

1s22s22p63s23p64s13d10

This gives one filled orbital and one half filled orbital.

Remember these exceptions