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Chapter 7

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

Quantum Theory and the

Electronic Structure of Atoms

2

Properties of Waves

Wavelength (l) is the distance between identical points on

successive waves.

Amplitude is the vertical distance from the midline of a

wave to the peak or trough.

Frequency (n) is the number of waves that pass through a

particular point in 1 second (Hz = 1 cycle/s).

The speed (u) of the wave = l x n

3

Maxwell (1873), proposed that visible light consists of

electromagnetic waves.

Electromagnetic

radiation is the emission

and transmission of energy

in the form of

electromagnetic waves.

Speed of light (c) in vacuum = 3.00 x 108 m/s

All electromagnetic radiation

l x n = c

4

5

l x n = c

l = c/n

l = 3.00 x 108 m/s / 6.0 x 104 Hz

l = 5.0 x 103 m

A photon has a frequency of 6.0 x 104 Hz. Convert this

frequency into wavelength (nm). Does this frequency fall in

the visible region?

l = 5.0 x 1012 nm

l

n

6

Mystery #1, “Heated Solids Problem”

Solved by Planck in 1900

Energy (light) is emitted or

absorbed in discrete units

(quantum).

E = h x n

Planck’s constant (h)

h = 6.63 x 10-34 J•s

When solids are heated, they emit electromagnetic radiation

over a wide range of wavelengths.

Radiant energy emitted by an object at a certain temperature

depends on its wavelength.

7

Light has both:

1. wave nature

2. particle nature

hn = KE + W

Mystery #2, “Photoelectric Effect”

Solved by Einstein in 1905

Photon is a “particle” of light

KE = hn - W

hn

KE e-

where W is the work function and

depends how strongly electrons

are held in the metal

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E = h x n

E = 6.63 x 10-34 (J•s) x 3.00 x 10 8 (m/s) / 0.154 x 10-9 (m)

E = 1.29 x 10 -15 J

E = h x c / l

When copper is bombarded with high-energy electrons, X rays

are emitted. Calculate the energy (in joules) associated with

the photons if the wavelength of the X rays is 0.154 nm.

9

Line Emission Spectrum of Hydrogen Atoms

10

11

1. e- can only have specific

(quantized) energy

values

2. light is emitted as e-

moves from one energy

level to a lower energy

level

Bohr’s Model of

the Atom (1913)

En = -RH ( )1

n2

n (principal quantum number) = 1,2,3,…

RH (Rydberg constant) = 2.18 x 10-18J

12

E = hn

E = hn

13

Ephoton = DE = Ef - Ei

Ef = -RH ( )1

n2f

Ei = -RH ( )1

n2i

i f

DE = RH( )1

n2

1

n2

nf = 1

ni = 2

nf = 1

ni = 3

nf = 2

ni = 3

14

15

Ephoton = 2.18 x 10-18 J x (1/25 - 1/9)

Ephoton = DE = -1.55 x 10-19 J

l = 6.63 x 10-34 (J•s) x 3.00 x 108 (m/s)/1.55 x 10-19J

l = 1280 nm

Calculate the wavelength (in nm) of a photon emitted

by a hydrogen atom when its electron drops from the

n = 5 state to the n = 3 state.

Ephoton = h x c / l

l = h x c / Ephoton

i f

DE = RH( )1

n2

1

n2Ephoton =

16

De Broglie (1924) reasoned

that e- is both particle and

wave.

Why is e- energy quantized?

u = velocity of e-

m = mass of e-

2pr = nl l = h

mu

17

l = h/mu

l = 6.63 x 10-34 / (2.5 x 10-3 x 15.6)

l = 1.7 x 10-32 m = 1.7 x 10-23 nm

What is the de Broglie wavelength (in nm) associated

with a 2.5 g Ping-Pong ball traveling at 15.6 m/s?

m in kgh in J•s u in (m/s)

Electronic

Structure

of Atoms

The Wave Nature of Matter

• Louis de Broglie positedأثبت that if light

can have material properties, matter

should exhibit wave properties.

• He demonstrated that the relationship

between mass and wavelength was

l =h

mv

Electronic

Structure

of Atoms

The Uncertainty Principleمبدأ عدم التأكد لهايزنبيرج

• Heisenberg showed that the more precisely

the momentum of a particle is known, the less

precisely is its position known:

• In many cases, our uncertainty of the

whereabouts of an electron is greater than the

size of the atom itself!

(Dx) (Dmv) h

4p

Electronic

Structure

of Atoms

• What does the uncertainty principle mean?

• The uncertainty principle also called the Heisenberg Uncertainty

Principle, or Indeterminacy Principle, articulated (1927) by the

German physicist Werner Heisenberg, that the position and the

velocity of an object cannot both be measured exactly, at the same

time, even in theory.

• What did the uncertainty principle state?

• The Heisenberg uncertainty principle states that it is impossible to

know simultaneously the exact position and momentum of a particle.

That is, the more exactly the position is determined, the less known the

momentum, and vice versa.

Electronic

Structure

of Atoms

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Electronic

Structure

of Atoms

Schrodinger’s Wave Equation

22

2 2

8

dh EVm dx

p− + =

Equation for the

probability of a single

electron being found

along a single axis (x-axis)Erwin Schrodinger

(this slide from: J. Hushen’s presentation on Atomic Structure at

http://teachers.greenville.k12.sc.us/sites/jhushen/Pages/AP%20Chemistry.aspx)

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Schrodinger Wave Equation

In 1926 Schrodinger wrote an equation that

described both the particle and wave nature of the e-

Wave function () describes:

1. energy of e- with a given

2. probability of finding e- in a volume of space

Schrodinger’s equation can only be solved exactly

for the hydrogen atom. Must approximate its

solution for multi-electron systems.

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Schrodinger Wave Equation

is a function of four numbers called

quantum numbers (n, l, ml, ms)

principal quantum number n

n = 1, 2, 3, 4, ….

n=1 n=2 n=3

distance of e- from the nucleus

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Where 90% of the

e- density is found

for the 1s orbital

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quantum numbers: (n, l, ml, ms)

angular momentum quantum number l

for a given value of n, l = 0, 1, 2, 3, … n-1

n = 1, l = 0

n = 2, l = 0 or 1

n = 3, l = 0, 1, or 2

Shape of the “volume” of space that the e- occupies

l = 0 s orbital

l = 1 p orbital

l = 2 d orbital

l = 3 f orbital

Schrodinger Wave Equation

27

l = 0 (s orbitals)

l = 1 (p orbitals)

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l = 2 (d orbitals)

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quantum numbers: (n, l, ml, ms)

magnetic quantum number ml

for a given value of l

ml = -l, …., 0, …. +l

orientation of the orbital in space

if l = 1 (p orbital), ml = -1, 0, or 1

if l = 2 (d orbital), ml = -2, -1, 0, 1, or 2

Schrodinger Wave Equation

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ml = -1, 0, or 1 3 orientations is space

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ml = -2, -1, 0, 1, or 2 5 orientations is space

32

(n, l, ml, ms)

spin quantum number ms

ms = +½ or -½

Schrodinger Wave Equation

ms = -½ms = +½

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Existence (and energy) of electron in atom is described

by its unique wave function .

Pauli exclusion principle - no two electrons in an atom

can have the same four quantum numbers.

Schrodinger Wave Equation

quantum numbers: (n, l, ml, ms)

Each seat is uniquely identified (E, R12, S8)

Each seat can hold only one individual at a

time

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35

Schrodinger Wave Equation

quantum numbers: (n, l, ml, ms)

Shell – electrons with the same value of n

Subshell – electrons with the same values of n and l

Orbital – electrons with the same values of n, l, and ml

How many electrons can an orbital hold?

If n, l, and ml are fixed, then ms = ½ or - ½

= (n, l, ml, ½) or = (n, l, ml, -½)

An orbital can hold 2 electrons

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How many 2p orbitals are there in an atom?

2p

n=2

l = 1

If l = 1, then ml = -1, 0, or +1

3 orbitals

How many electrons can be placed in the 3d subshell?

3d

n=3

l = 2

If l = 2, then ml = -2, -1, 0, +1, or +2

5 orbitals which can hold a total of 10 e-

37

Energy of orbitals in a single electron atom

Energy only depends on principal quantum number n

En = -RH ( )1

n2

n=1

n=2

n=3

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Energy of orbitals in a multi-electron atom

Energy depends on n and l

n=1 l = 0

n=2 l = 0n=2 l = 1

n=3 l = 0n=3 l = 1

n=3 l = 2

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“Fill up” electrons in lowest energy orbitals: Building – Up Principle

(Aufbau principle)

H 1 electron

H 1s1

He 2 electrons

He 1s2

Li 3 electrons

Li 1s22s1

Be 4 electrons

Be 1s22s2

B 5 electrons

B 1s22s22p1

C 6 electrons

? ?

40

C 6 electrons

The most stable arrangement of electrons in

subshells is the one with the greatest number of

parallel spins (Hund’s rule).

C 1s22s22p2

N 7 electrons

N 1s22s22p3

O 8 electrons

O 1s22s22p4

F 9 electrons

F 1s22s22p5

Ne 10 electrons

Ne 1s22s22p6

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Order of orbitals (filling) in multi-electron atom

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s

42

Electron configuration is how the electrons are

distributed among the various atomic orbitals in an

atom.

1s1

principal quantum

number n

angular momentum

quantum number l

number of electrons

in the orbital or subshell

Orbital diagram

H

1s1

43

What is the electron configuration of Mg?

Mg 12 electrons

1s < 2s < 2p < 3s < 3p < 4s

1s22s22p63s2 2 + 2 + 6 + 2 = 12 electrons

Abbreviated as [Ne]3s2 [Ne] 1s22s22p6

What are the possible quantum numbers for the last

(outermost) electron in Cl?

Cl 17 electrons 1s < 2s < 2p < 3s < 3p < 4s

1s22s22p63s23p5 2 + 2 + 6 + 2 + 5 = 17 electrons

Last electron added to 3p orbital

n = 3 l = 1 ml = -1, 0, or +1 ms = ½ or -½

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Outermost subshell being filled with electrons

45

Some Anomalies

( شذوذ عن القاعدة(

Some irregularities

occur when there

are enough

electrons to half-fill

s and d orbitals on

a given row.

Some Anomalies

For instance, the electron

configuration for

chromium is

[Ar] 4s1 3d5

rather than the expected

[Ar] 4s2 3d4

Copper is [Ar] 4s1 3d10

rather than the expected

[Ar] 4s23d9

Some Anomalies

• This occurs

because the 4s and

3d orbitals are very

close in energy.

• These anomalies

occur in f-block

atoms, as well.

49

Paramagnetic

unpaired electrons

2p

Diamagnetic

all electrons paired

2p

© 2009, Prentice-Hall, Inc.

The shielding Effectتأثير الحجب في الذرات متعددة الإلكترونات

• The shielding effect describes the attraction between an electron and the nucleus in any atom with more than

one electron shell. Shielding effect can be defined as a reduction in the Effective Nuclear Charge on the electron

cloud, due to a difference in the attraction forces of the electrons on the nucleus.

• In hydrogen, or any other atom in group 1A of the periodic table (those with only one valence electron), the

force on the electron is just as large as the electromagnetic attraction from the nucleus. However, when more

electrons are involved, each electron (in the nth-shell) experiences not only the electromagnetic attraction from

the positive nucleus, but also repulsion forces from other electrons in shells from 1 to n. This causes the net

force on electrons in outer shells to be significantly smaller in magnitude; therefore, these electrons are not as

strongly bonded to the nucleus as electrons closer to the nucleus. This phenomenon is often referred to as the

orbital penetration effect. The shielding theory also contributes to the explanation of why valence -shell

electrons are more easily removed from the atom