Electronic Structure of Electronic Structure of AtomsAtoms(i.e., Quantum Mechanics)(i.e., Quantum Mechanics)

Brown, LeMay Ch 6AP ChemistryMonta Vista High School

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What does light have to do with the atomic model?

Scientists knew the nature of light but knew little about the nature of matter. To understand the nature of matter, scientists studied the changes caused in light by interaction of matter. From these studies, scientist tried to extrapolate information about the nature of matter.

6.1: Light is a Wave6.1: Light is a WaveElectromagnetic spectrum:

A form of radiant energy (can travel without matter)

Both electrical and magnetic (properties areperpendicular to each other)

http://imagine.gsfc.nasa.gov/Videos/general/spectrum.mov

Speed of Light: c = 3.0 x 108 m/s (in a vacuum)

http://www.astronomynotes.com/light/s3.htm

Wavelength (): distance between wave peaks (determines “color” of light), measured in nm, m etc.

Frequency (): # cycles/sec (measured in Hz- Hertz, hz= cycles/s or 1/s)

3c =

6.2: Light is a Particle (Quantum 6.2: Light is a Particle (Quantum Theory)Theory)

Blackbody radiation:* Blackbody: object that absorbs all EM radiation that strikes it; it

can radiate all possible wavelengths of EM; below 700 K, very little visible EM is produced; above 700 K visible E is produced starting at red, orange, yellow, and white before ending up at blue as the temperature increases

◦ discovery that light intensity (energy emitted per unit of time) is proportional to T4; hotter = shorter wavelengths

“Red hot” < “white hot” < “blue hot” Interactive Link

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ch

h E

ch h E

Max Planck(1858-1947)

• Planck’s Theory: (explained blackbody radiation by quantization of energy transfer)

Blackbody radiation can be explained if energy can be released or absorbed in packets of a standard size he called quanta (singular: quantum).

where Planck’s constant (h) = 6.63 x 10-34 J-s Animation Link

The Photoelectric Effect The Photoelectric Effect Spontaneous emission of e- from metal struck by light; first explained by Einstein in 1905A quantum strikes a metal atom and the energy is absorbed by an e-. If the energy is sufficient, e- will leave its orbital, causing a current to flow throughout the metal. To explain photoelectric effect, quantization of light was put forth by Einstein. Animation

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Albert Einstein(1879-1955)

6.3: Bohr’s Model of the H Atom (and 6.3: Bohr’s Model of the H Atom (and only H!)only H!)Applied quantization of energy transfer to the atomic model

Studied atomic spectrum of H to come up with atomic model.

Atomic emission spectra:Most sources produce light that contains many

wavelengths at once. AnimationHowever, light emitted from pure substances may

contain only a few specific wavelengths of light called a line spectrum (as opposed to a continuous spectrum). AnimationAtomic emission spectra are inverses of atomic

absorptionspectra.

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Atomic Emission Spectra of C Atomic Emission Spectra of C and Hand H

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Hydrogen: contains 1 red, 1 green, 1 blue and 1 violet.

Carbon: Contains many more emission lines as compared to H. Why?

Niels Bohr theorized that e-:◦ Travel in certain “orbits” around the nucleus,

or, are only stable at certain distances from the nucleus

◦ If not, e- should emit energy, slow down, and crash into the nucleus.

Allowed orbital energies are defined by:

principal quantum number (n) = 1, 2, 3, 4, …Rydberg’s constant (RH) = 2.178 x 10-18 J

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2

18

2H

n n

10178.2

n

RE

Niels Bohr(1888-1962)

Johannes Rydberg(1854-1919)

Think, Pair, Share ActivityThink, Pair, Share Activity• With your elbow partner, describe

Electromagnetic radiation, blackbody radiation, Plank’s theory and Photoelectric effect. Address each of the above in the following terms:

1. What is it?2. Why was it important?3. What existing theory or concept, it

approved/disapproved.9

As n approaches ∞, the e- is essentially removed from the atom, and E∞ = 0.

• ground state: lowest energy level in which an e- is stable• excited state: any energy level higher than an e-’s ground

state

Incr

easi

ng E

nerg

y, E

Prin

cipa

l Qua

ntum

Num

ber,

n

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3

2

1

E5

E4

E3

E2

E1

ni = initial orbital of e-nf = final orbital of e- in its transitionMovie on e transition

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2

i2

f

Hn

1

n

1RE

2

i2

f

H

n

1

n

1

h

R

h

E

2

f2

i

H

n

1

n

1

h

R

h

E

Figure 1: Line series are transitions from one level to another.

SeriesTransition down to (emitted)

or up from (absorbed)…Type of EMR

Lyman 1 UV

Balmer 2 Visible

Paschen 3 IR

Brackett 4 Far IR

5432

1

n

Theodore Lyman

(1874 - 1954)

JohannBalmer

(1825 – 1898)

FriedrichPaschen

(1865 - 1947)

FrederickBrackett

(1896 – 1988)?

6.4: Matter is a Wave6.4: Matter is a WavePlanck said: E = h c /

Einstein said: E = m c2

Louis DeBroglie said (1924): h c / m c2

h / m cTherefore:

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m = h / c Particles (with mass) have an associated wavelength

h / mcWaves (with a wavelength) have an associated mass and velocity

Louisde Broglie

(1892 - 1987)

Neils Bohr Model: Partner Neils Bohr Model: Partner ActivityActivityOn a sheet of paper, take turns with your

partner drawing Bohr’s model of atom.Draw the following in context of Bohr’s Model:1.nucleus2.energy levels (1,2,3,4)3.an electron in energy level 24. Show an electron transition from energy level 2 to 35. Write formula for calculating this energy change and calculate energy.6. Give each other high fives!!

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IBM – Almaden:IBM – Almaden:

“Stadium “Stadium Corral”Corral”

This image shows a ring of 76 iron atoms on a copper (111) surface. Electrons on this surface form a two-dimensional electron gas and scatter from the iron

atoms but are confined by boundary or "corral." The wave pattern in the interior is due to the density distribution of the trapped electrons. Their

energies and spatial distribution can be quite accurately calculated by solving the classic problem of a quantum mechanical particle in a hard-walled box.

Quantum corrals provide us with a unique opportunity to study and visualize the quantum behavior of electrons within small confining structures.

Heisenberg’s Uncertainty Heisenberg’s Uncertainty Principle (1927)Principle (1927)

It is impossible to determine the exact position and exact momentum (p) of an electron.

p = m vTo determine the position of an e-, you

have to detect how light reflects off it.But light means photons, which means

energy. When photons strike an e-, they may change its motion (its momentum).

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WernerHeisenberg

(1901 – 1976)

Electron density distribution Electron density distribution in H atomin H atom

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6.5: Quantum Mechanics & Atomic 6.5: Quantum Mechanics & Atomic OrbitalsOrbitalsSchrödinger’s wave function:Relates probability () of

predicting position of e- to its energy.

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dt

dihU

dx

d

m

hE

2

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2

Where: U = potential energyx = position t = timem = mass i =√(-1)

http://daugerresearch.com/orbitals/index.shtml

ErwinSchrödinger

(1887 – 1961)

Probability plots of 1s, 2s, Probability plots of 1s, 2s, and 3s orbitalsand 3s orbitals

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6.6: Representations of Orbitals6.6: Representations of Orbitalswww.orbitals.com; animation 1, www.orbitals.com; animation 1, Draft of a letter from Bohr to Draft of a letter from Bohr to Heisenberg (never sent)Heisenberg (never sent)

s orbital

p orbitals

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d orbitals

f orbitals: very complicated

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1. Aufbau principle: e- enter orbitals of lowest energy first

Incr

eas i

ng E

nerg

y

1s

2s

3s

4s

5s

6s

7s

3d

4d

5d

6d

4f x 7

5f x 7

2p

3p

4p

5p

6p

7p

• Relative stability & average distance of e- from nucleus

6.7: Filling Order of Orbitals

Animation for filling of Animation for filling of OrbitalsOrbitals

Use the “diagonal rule” (some exceptions do

occur).

Sub-level maxima: s = 2 e-

p = 6 e-d = 10 e-f = 14 e-…

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1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d

7s 7p

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2. Pauli exclusion principle (1925): no two e- can have the same four quantum numbers; e- in same orbital have opposite spins (up and down)

3. Hund’s rule: e- are added singly to each equivalent (degenerate) orbital before pairing

Ex: Phosphorus (15 e-) has unpaired e- inthe valence (outer) shell.1s 2s 2p 3s 3p

WolfgangPauli

(1900 – 1958)

FriedrichHund

(1896 - 1997)

6.9: Periodic Table & Electronic 6.9: Periodic Table & Electronic ConfigurationsConfigurations

s block p blockd blockf block

s1 s2

p1 p2 p3 p4 p5 p6

d2 d3 d5 d5 d6 d7 d8 d10d10

f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14

s2

1s2s3s4s5s6s7s

2p3p4p5p6p7p4f

5f

3d4d5d6d

3d4d5d6d

d1

Notable Exceptions:Cr & Mo: [Ar] 4s1 3d5 not [Ar] 4s2 3d4

Cu, Ag, & Au: [Ar] 4s13d10 not [Ar] 4s23d9

Electronic ConfigurationsElectronic Configurations

Element Standard ConfigurationNoble Gas Shorthand

Nitrogen

Scandium

Gallium

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[He] 2s22p3

[Ar] 4s23d1

[Ar] 4s23d104p1

1s22s22p3

1s22s22p63s23p64s23d1

1s22s22p63s23p64s23d104p1

Element Standard ConfigurationNoble Gas Shorthand

Lanthanum

Cerium

Praseodymium

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[Xe] 6s25d1

[Xe] 6s25d14f1`

1s2 2s22p6 3s23p6 4s23d104p6

5s24d105p6 6s25d1

1s2 2s22p6 3s23p6 4s23d104p6

5s24d105p6 6s25d14f1

[Xe] 6s24f31s2 2s22p6 3s23p6 4s23d104p6

5s24d105p6 6s24f3

Electron Configuration for Electron Configuration for IonsIons

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Cr +

Cr3+

Valence Electrons: Only s and p e are valence electrons. The maximum number of valence e that an atom can have is 8. WHY? Write the electron configurations for the following ions:

Ground State Electron Config. V. Excited State Electron Configuration

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Ways to Represent Electron Configuration

1.Expanded Electron Configuration2.Condensed Electron Configurations3.Orbital Notation4.Electron Dot Structure

Write the above four electron configurations for Zinc, Zinc ion and Cu ion.

ParamagneticDiamagnetic

Why are some ions colored and some aren’t?

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Electron Configuration and Para- and Diamagnetism demo + activity