ARRANGEMENT OF ELECTRONS IN

ATOMS

Chapter 4

Visible Light

We are all familiar with light but what is “visible” is just a very, very small portion of the electromagnetic spectrum

What colors make up the rainbow?

Red, Orange, Yellow, Green, Blue, Indigo, Violet (ROYGBIV)

The E-M Spectrum

Gamma Rays(Very Harmful / Cancerous)

X-rays(Cancerous in large doses; small doses medical scanning)

Ultraviolet ( not as harmful - sunburn; black lights)

Infrared(Heat, communication)

Microwaves(Cooking, communications)

Radio Rays(TV, Radio, other communications)

The Development of a New Atomic Model

Wavelength () - length of one complete wave

Frequency () - # of waves that pass a point during a certain time period hertz (Hz) = 1/s

Amplitude (A) - distance from the origin to the trough or crest

Waves

Agreater

amplitude(intensity)

greater frequency

(color)

crest

origin

trough

A

The Electromagnetic Spectrum

AM radio

Short waveradio

Television channels

FM radio

RadarMicrowave

Radio Waves Gamma Rays

X- Raysinfrared

Increasing photon energy

Increasing frequency

Decreasing wavelength

Red Orange Yellow Green Blue Indigo Violet

UV Rays

R O Y G B I V

Visible

Light

Electromagnetic Spectrum

Frequency & wavelength are inversely proportional

c = c: speed of light (3.00 108 m/s): wavelength (m, nm, etc.): frequency (Hz)

Electromagnetic Spectrum

GIVEN: = ? = 434 nm = 4.34 10-7 mc = 3.00 108 m/s

WORK: = c

= 3.00 108 m/s 4.34 10-7 m

= 6.91 1014 Hz

EX: Find the frequency of a photon with a wavelength of 434 nm.

So why is the electromagnetic spectrum so important to chemistry?

• Why is the steel emitting light when it is heated?

• We take it for granted that when things get hot they turn red then orange and finally white; but that isn’t good enough any more

Black Body Radiation Colors

Black Body Radiation Colors

So why is the electromagnetic spectrum so important to chemistry?

Incandescence is heat made visible – the process of turning heat energy into light energy.

Our usage of terms like "red hot," "white hot," and so on, is part of the color sequence black, red, orange, yellow, white, and bluish white, seen as an object is heated to successively higher temperatures.

So why is the electromagnetic spectrum so important to chemistry?

The light produced consists of photons emitted when atoms and molecules release part of their thermal vibration energy.

For increasing temperatures, the sequence of radiated colors is: black, red, orange, yellow-white, bluish-white.

Heat and Light

Planck (1900)

Observed - emission of light from hot objects

Concluded - energy is emitted in small, specific amounts (quanta)

Quantum - minimum amount of energy change

Energy and Light

E: energy (J, joules)h: Planck’s constant (6.6262 10-34 J·s): frequency (Hz)

E = h

The energy of a photon is proportional to its frequency.

Energy and Light

GIVEN:E = ? = 4.57 1014 Hzh = 6.6262 10-34 J·s

WORK:E = h

E = (6.6262 10-34 J·s)(4.57 1014 Hz)

E = 3.03 10-19 J

EX: Find the energy of a red photon with a frequency of 4.57 1014 Hz.

Niels Bohr and the Bohr model of the atomBohr hypothesized that instead of haphazardly orbiting the nucleus, electrons had clearly defined orbits – very similar to the planetary orbits circling our sun

His model is (cleverly) named the Planetary Model

Niels Bohr

Bohr Model (1913)

Bohr’s Proof

Bohr said this: If you assume that the electrons have clearly defined orbits that are congruent to the energy levels…

Bohr’s Proof

… then when an electron gets “excited” it jumps to a higher energy level. When it “relaxes” it emits a certain wavelength of light.

• Bohr showed the energy of an electron in an atom is quantized, which means it has a particular numerical value, not some arbitrary number.

Excitation of Hydrogen Atoms

Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 328

Return to Ground State

Bohr’s Proof

n=1

n=2

n=6n=5n=4

n=3

n=7

1= 1.097373 x 107 m-1

λ nr2 ne

2

1 1-

Lyman Series (uv)

Balmer Series (vis and uv)

Paschen Series (ir)

Emission Spectrum of an Element

1 nm = 1 x 10-9 m = “a billionth of a meter”

410 nm 434 nm 486 nm 656 nm

1 nm = 1 x 10-9 m = “a billionth of a meter”

Continuous and Line Spectra

Hydrogen to Steel

If Hydrogen emits 4 distinct wavelengths of light when its one electron is excited what can we extrapolate to that of steel which is made mostly of iron?

http://jersey.uoregon.edu/vlab/elements/Elements.html

Flame Emission Spectra

Photographs of flame tests of burning wooden splints soaked in different salts.Photographs of flame tests of burning wooden splints soaked in different salts.

Include link to web page

http://www.unit5.org/christjs/flame%20tests.htm

methane gas wooden splint strontium ioncopper ionsodium ion calcium ion

Fireworks

Composition of Fireworks

Gunpowder Sulfur, charcoal, potassium nitrate (saltpeter)

Salts (to give color) Red = lithium Green = copper

Good News Bad News

Good News

Bohr’s Model works and moves us along in the development of the Atomic Theory

End of this little unit

Bad (Frustrating) News

Lots of Math

Everything I taught you only works for Hydrogen and therefore is completely wrong and obsolete.

Check for Understanding

c= λν E=hν c=3.0 x108 m/s h=6.626 x 10-34 J s

What is the frequency of a radar photon with an energy of 7.2 x 10-24 J?

What is the frequency of light having a wavelength of 6.20x10-7m?

Models of the Atom

Dalton’s model (1803)

Thomson’s plum-pudding model (1897)

Rutherford’s model (1909)

Bohr’s model (1913)

Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 125

Greek model(400 B.C.)

+--

--

-e

e

e

+

+ +

+

++

++

e

ee

e

e

ee

Your Current View of the Atom

electrons

nucleus

Again… so why is it so important to chemistry?

Einstein (1905) Observed - photoelectric effect

Again… so why is it so important to chemistry?

Einstein (1905) Concluded - light has properties of both waves

and particles

“wave-particle duality”

Photon - particle of light that carries a quantum of energy

Quantum Mechanical Model

Modern atomic theory describes the electronic structure of the atom as the probability of finding electrons within certain regions of space (orbitals).

Modern View

The atom is mostly empty space

Two regions Nucleus

protons and neutrons Electron cloud

region where you might find an electron

Also called the electron cloud model

Modern View of Atom

e-e- Ground state

Excited state

Electrons can only be atspecific energy levels,NOT between levels.

Models of the Atom

Dalton’s model (1803)

Thomson’s plum-pudding model (1897)

Rutherford’s model (1909)

Bohr’s model (1913)

Charge-cloud model (present)

Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 125

Greek model(400 B.C.)

+--

--

-e

e

e

+

+ +

+

++

++

e

ee

e

e

ee

C. Johannesson

Electrons as Waves

Louis de Broglie (1924) Applied wave-particle theory to e-

e- exhibit wave properties

QUANTIZED WAVELENGTHS

C. Johannesson

Quantum Mechanics

Heisenberg Uncertainty Principle Impossible to know both the velocity and position of

an electron at the same time

C. Johannesson

Quantum Mechanics

σ3/2 Zπ

11s 0

eΨ a

Schrödinger Wave Equation (1926) finite # of solutions quantized energy levels

defines probability of finding an e-

Quantum Theory

quantum theory- Describes mathematically the wave

properties of electrons and other small particles

orbital- a region of an atom in which there is a high probability of finding electrons

Today’s atomic model predicts quantized, or particular energy levels for electrons.

does not describe the exact path or location electrons take or can be found around the nucleus

concerned with the probability, or likelihood, of finding an electron in a certain position

Two electrons can occupy each orbital, also called an electron cloud.

Quantum Numbers

Four Quantum Numbers: Specify the “address” or “seat” of each

electron in an atom

UPPER LEVEL

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Quantum Numbers

1. Principal Quantum NumberPrincipal Quantum Number ( nn )

Energy level (ladder rungs)

Size of the orbital

Positive integer

1s

2s

3s

Quantum Numbers

1. Principal Quantum NumberPrincipal Quantum Number > number, further away from the

nucleus 1- right next to the nucleus 3- further away from nucleus

> number, higher the energy level n = 2 greater energy level than n = 1 these electrons have more energy

than electrons in the n = 1 level

1s

2s

3s

Quantum Numbers

2. Angular Momentum Quantum #Angular Momentum Quantum # ( ll ) Energy sublevel (orbital) Shape of the orbital Often represented by letters than numbers

s p d fCourtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Quantum Numbers

px pz py

x

y

z

x

y

z

x

y

z

d-orbitals

Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 336

Quantum Numbers

Orbitals combine to form a spherical shape.

2s

2pz2py

2px

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Quantum Numbers

3. Magnetic Quantum NumberMagnetic Quantum Number ( mmll ) Orientation of orbital

Specifies the exact orbital within each sublevel

Shapes of s, p, and d-Orbitals

Quantum Numbers

4. Spin Quantum NumberSpin Quantum Number ( ms ) Electron spin +½ or -½

An orbital can hold 2 electrons that spin in opposite directions.

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Maximum Capacities of Subshells and Principal Shells

n 1 2 3 4 ...n

l 0 0 1 0 1 2 0 1 2 3

SubshellSubshelldesignationdesignation s s p s p d s p d f

Orbitals inOrbitals insubshell subshell 1 1 3 1 3 5 1 3 5 7

SubshellSubshellcapacity capacity 2 2 6 2 6 10 2 6 10 14

Principal shellPrincipal shellcapacity capacity 2 8 18 32 ...2n2

Hill, Petrucci, General Chemistry An Integrated Approach1999, page 320

Filling Rules for Electron Orbitals

Aufbau Principle: Electrons are added one at a time to the lowest energy orbitals available until all the electrons of the atom have been accounted for.

Pauli Exclusion Principle: An orbital can hold a maximum of two electrons.To occupy the same orbital, two electrons must spin in opposite directions.

Hund’s Rule: Electrons occupy equal-energy orbitals so that a maximum number of unpaired electrons results.

*Aufbau is German for “building up”

Diagonal Rule

1s

2s

3s

4s

5s

6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

6d

4f

5f

General Rules

Pauli Exclusion PrinciplePauli Exclusion Principle Each orbital can hold TWO electrons with opposite

spins.

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

RIGHTWRONG

General Rules

Hund’s RuleHund’s Rule Within a sublevel, place one electron per

orbital before pairing them.

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Orbital Diagrams and Electron Configurations

Orbital diagrams Show how electrons are distributed within sublevels Electrons represented by an arrow Orbital is represented by a box Direction of spin represented by direction of arrow

Electron configuration Abbreviated form of orbital diagram

Orbital Diagrams and Electron Configurations

H 1 e-

Orbital diagram

1s

Electron configuration1s1

↑

O

8e-

Orbital Diagrams and Electron Configurations

Orbital Diagram

• Electron Configuration

1s1s22 2s2s22 2p2p44

1s 2s2p

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Ne

10e-

Orbital Diagrams and Electron Configurations

Orbital Diagram

• Electron Configuration

1s1s22 2s2s22 2p2p66

1s 2s 2p

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

neon's electron configuration (1s22s22p6)

Noble Gas Configuration

[Ne] 3s1

third energy level

one electron in the s orbital

orbital shape

Na = [1s22s22p6] 3s1 electron configuration

AA

BB

CC

DD

• Shorthand Configuration

S 16e-

Valence ElectronsValence ElectronsCore ElectronsCore Electrons

S 16e- [Ne] 3s2 3p4

1s2 2s2 2p6 3s2 3p4

Electron Configuration

Longhand Configuration

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

S32.066

16