•Try to find visual of uncertainty principle.

Electronic Structure of AtomsChapter 6

Niels Bohr tried to explain the location of electrons….

Based his ideas on the relationship between matter and light (electromagnetic radiation)

We need background information on electromagnetic radiation before we can discuss Bohr’s ideas

LIGHTProperties of light: •Part of the electromagnetic radiation

spectrum•Has a constant velocity (symbol is c)

c = 3.00 X 108 m/sec *watch the units!

There are 2 theories that can explain light: wave theory & particle theory.

on AP equation sheet!

WAVE THEORY OF LIGHT BEHAVIOR light has a wave-like nature

wavelength ( l ) - distance per cycle of wave

 frequency ( n ) - number of cycles per

unit of time

Velocity is related to wavelength and frequency

For any wave, velocity = l • n For any electromagnetic wave, c = l • n

3.00x108 m/sec= l • n

l and n are inversely related for electromagnetic waves

on AP equation sheet!

Electromagnetic Radiation Spectrum - the type of radiation is related to its wavelength

-11 -9 -7 -5 -3 -1 1 310 10 10 10 10 10 10 10 (l in

cm) | | | | | | | | | | | | | | | | | | | | | Gamma | X-rays | Ultra |V| Infrared | Micro | TV & radio radiation | | Violet |i| light | waves | signals | | light |s| | |

smaller wavelength larger wavelengthlarger frequency smaller frequencylarger energy smaller energy

VISIBLE LIGHT

PARTICLE THEORY OF LIGHT BEHAVIOR

•Max Plank figured out that light can be released from and absorbed by atoms in “chunks” that have a minimum size.

•Light is made up of particles called photons or quanta- bundles of energy

There is a mathematical equation that ties these 2 theories together

Energy of photon = (Plank’s constant) (frequency)

E = h • nh = Plank’s constant = 6.63 X 10-34 Joule-

second (J-s)

E and n are directly relatedE and l are inversely related

on AP equation sheet!

Niels Bohr Model of the AtomElectrons are found in “orbits” that surround the nucleus. The concept of an orbit means that the electron travels around the nucleus at a fixed distance and at a fixed speed that keeps it on that particular orbit.

Yes, Bohr was wrong about orbits….let’s pretend he meant energy levels(because that’s what they really are!)

Bohr Model (cont.)

When an electron absorbs a certain amount of energy, called a “quanta,” it can jump to an energy level farther from the nucleus.

Bohr ModelWhen electrons absorb energy to jump up a level, they immediately release the same amount of energy to get back to the ground state.

Bohr Model•Sometimes, the energy that is absorbed

or released by excited electrons can be seen as different colors of light!

•The color depends on the wavelength or energy of the light emitted!

Neils Bohr•Examined the light released by atoms of

certain elements•Worked mostly with hydrogen since it has

only 1 electron•Bohr used a spectroscope to analyze the

light released from a vacuum tube - the result is a bright line spectrum more commonly called an EMISSION SPECTRUM (black background with bright colored lines)

Only certain wavelengths or photon energies of light are present in the emission spectrum

An element’s emission spectrum is unique to that element (like a fingerprint!)

EMISSION SPECTRUM

Energy level diagram

A B C D

H emission spectrum

4 x 10-5 λ 7 x 10-5

smaller λ larger λ

A C D

again, 21 possible emission lines

U V Visible I R

violet red

again, also 21 possible absorption linesabsorption lines (arrows) would point UP

larger E smaller E

ENERGY

1

2

3

4

567

Arrow D is the longest arrow in the visible region which means it has the highest energy and corresponds to the violet line in hydrogen’s emission spectrum.Arrow A is the shortest arrow in the visible region which means it has the lowest energy and corresponds to the red line in hydrogen’s emission spectrum.

B

Calculations:•To find the energy associated with each orbit

in the hydrogen atom, Bohr used the equation:

1. Subtract the energy of the orbits between which an electron moves to find the energy released/absorbed.

2. Use E=hν to convert energy to frequency3. Use c=λν to convert frequency to

wavelength4. These correspond to the colors seen in the

emission spectrum!

18

2

2.178 10xE Joulesn

On AP Equation Sheet!

Bohr tried to apply his ideas about electron behavior of hydrogen atoms to atoms of other elements.

Bohr tried to explain the arrangement of elements on the periodic table by showing that elements in a column of the periodic table have similar chemical properties because they have similar electron structures.

does not work for #21 scandium

K Ca

After Bohr….•Loius De Broglie…if light can behave like

matter (photons)…can matter behave like a wave??

YEP!

Λ = wavelength (m)h= Plank’s constant (J-sec)m= mass (kg)V = velocity (m/sec)

hmv

l on AP equation sheet!

The Uncertainty Principle•Werner Heisenberg….because of the dual

nature of matter both the path and velocity of a particle with the size and mass of the electron can not be determined at the same instant in time - in other words, electrons cannot be in orbits!

•*The uncertainty in the calculation of the position of the electron is larger than the entire atom because the electron is so small!

Breakdown of the Bohr model of the atom1. Does not explain the behavior of atoms

with more than 20 electrons

2. Violates the Heisenberg Uncertainty Principle

SCHRÖDINGER •He developed the idea of wave mechanics -

treating the electron as if it had wave-like properties.

•He mathematically considered the region of space in the atom where an electron of a given energy would MOST PROBABLY BE LOCATED.

•This is a radical departure from earlier work. Schrödinger says to forget exactly where the electron is (Heisenberg says we can’t find out anyway) and concentrate on where the electron MIGHT be.

RESULTS OF THE MODERN THEORY OF ELECTRON BEHAVIOR

•The math involved in wave equations is very complicated and beyond the scope of this course. Fortunately, the results of this theory are easier to understand than the mathematics.

ORBITAL- The region of space in an atom where it

is highly probable that an electron of a given energy will be located

- About 90 % of the solution set to the wave equation

- Max of 2 electrons/orbital

OrbitalOrbital

Energy Level Diagram

Each box = electronEach column = orbitalEach group = sublevelBig # out front = Energy level

Multi-electron atoms:•Atoms fill the lowest energy positions first

•2 methods to know the order:▫Periodic table▫Wavy line diagram

The Periodic Table12

7

34

6

1

5

ns

d(n-1)

p

f(n-2)

Wavy-line Diagram

Principle Quantum Number (n)

•Indicates a specific energy LEVEL (energy levels values found by Bohr)

•Indicates the relative size of the energy level - larger value means larger size

•Indicate the relative energy of the electron – larger value means higher energy.

•Possible values are n = 1, 2, 3, 4, 5, 6, 7, ...

Angular Momentum Quantum Number ( )•Indicates a specific energy SUBLEVEL

•The number of energy sublevels per energy level is equal to the value of the energy level (example: 3 sublevels in the 3rd energy level)

•Possible values are = 0, 1, 2, 3, ..., max. value of n-1

•more commonly = s, p, d, f, ... (alphabetically after f)

Indicates the shape of the orbitals in that sublevel- s orbital has a spherical shape - p orbital has a dumb-bell shape (technically, a pair of ellipsoids)

- d orbitals - we won’t worry about their shape

- f orbitals - we won’t worry about their shape

Magnetic Quantum Number (m )

•Indicates a specific ORBITAL of an energy sublevel

•Indicates the spatial orientation of the orbital on an x, y, z axis system in a magnetic field

•Possible values are m = - to + (integer values)

Spin Quantum Number (ms ) • Indicates a specific ELECTRON of an

orbital• Indicates the electron spin direction -

clockwise or counter-clockwise• Possible values are s = -1/2 or +1/2

PAULI EXCLUSION PRINCIPLE•No two different electrons in the same

atom can have the EXACT same set of quantum numbers.

A Complete Set of Quantum Numbers( n, , m, s )

electronorbital

sublevellevel

Electron Configuration - A representation of an atom’s electron

structure using energy levels and sublevels

- All electrons are shown

Example: for 6C 1s22s22p2

Orbital Notation -A representation of an atom’s electron

structure using arrows for the electrons and lines/boxes for the orbitals

- Hund’s Rule: Within the same sublevel, the electrons go in separate orbitals until they are forced to pair up

Example: for 6C

↑↓ ↑↓ ↑ ↑ __1s 2s 2p

Vocabulary:•Paramagnetic: an atom with one or more

unpaired electrons (electron alone in an orbital)

•Diamagnetic: an atom with only paired electrons (two electrons in each filled orbital)

Excited State Electron Configurations:•We predict ground state electron

configurations.•Sometimes when an atom gains energy an

electron jumps to a higher energy level.

Ground state of iron: [Ar]4s23d6

But if it gains some energy one of the 4s electrons might jump into the 3d sublevel to make the electron configuration:

[Ar]4s13d7

Special Stability•Half filled sublevels have an extra level of

stability compared to other configurations – because of this some ground state configurations are not what we predict!

Examples:•Cr & Mo – s1d5 (not s2d4)•Cu & Ag – s1d10 (not s2d9)

Chapter 6 Book Work:•6(6, 13, 15, 17, 28, 37, 45, 50a, 50b ,53, 55,

69, 85)