1

Electrons in Atoms

It’s Unreal !! Check your

intuition at the door.

Quantum Theory or Wave Theory

description of the electronic

structure in atoms

2

Quantum Theory

1900~ 1930

One of the greatest

achievements of mankind.

3

Quantum Theory

wave-particle duality

probability of electron location

“orbital” shapes

Unlike anything in our

macroscopic world.

4

e- Arrangement in Atoms

chemical reactivity

bonding between atoms

Periodic Table

many physical properties

Determines:

5

Atomic Models: History

Each atomic model was eventually

replaced in light of new

experimental evidence.

1 2 3

6

Dalton: 1803

Concept of the atom as

smallest unit of an element.

Indivisible particle

7

Thomson: 1897

Atom has parts!!

Discovered the e-

“Plum pudding” model

electron

+charge

8

Rutherford: 1911

Nucleus with

positive charge

Au foil experiment

Nuclear model+

Most of atom is empty space

9

Nuclear Model: Problem

What keeps the electrons

and nucleus apart?+

10

Neils Bohr: 1913

e- held in “orbits”

Motion of e- keeps them from

“falling” into nucleus

Similar to planets around sun

11

Bohr: “Planetary” Model

e- move in circular

orbits around nucleus,

and each orbit has a

certain energy.

12

Bohr: “Planetary” Model

+

E1

E2

E3

“Quantized”

energy levels

13

“Bright Line Spectrum” of Hydrogen

Stair Analogy: H spectrum due to e- “transitions”.

14

E1

E2

E3

E4

E5

ener

gy

Stairs are quantized.

Not a ramp

e- in Ground State

15

E1

E2

E3

E4

E5

ener

gy

Ground state is lowest

energy of the e-.

e- in Excited State

16

E1

E2

E3

E4

E5

ener

gy

e- absorbs energy to move

to a higher energy level.

e- in Excited State

17

E1

E2

E3

E4

E5

ener

gy

e- Returning to Ground

18

E1

E2

E3

E4

E5

ener

gy photon

Elight=Eexcited-Eground

e- gives off energy as light

19

Elight=Eexcited-Eground

The energy of the light is the

difference between the higher and

lower energy level of the electron.

Each energy of light corresponds to a

unique color of light.

e- Returning to Ground

20

E1

E2

E3

E4

E5

ener

gy

lower energy

photon

21

Bohr: Hydrogen Emission Spectrum

+

E3

E2

E1

e- absorbs energy

(heat, elec.)

e- falls to lower E

and gives off

energy as light

Elight=E3-E1

22

Emission Spectrum

Flame test

Neon signs

Fireworks

Fireplace colors

23

Bohr Theory: Failings

• Why do e- only have

certain orbital

energies?

• Only explains the

hydrogen atom exactly.

24

Quantum Mechanics

Or Wave Model

1926: E. Schrodinger

e- location (atomic orbital) described

by a probability function:

Y

25

Schrodinger Equation

d2Ydx2

+d2Ydy2

d2Ydz2

+ +

8p2m

h2(E – U) Y = 0

Many solutions.

26

Quantum Mechanical Modelor Wave Model

Electrons are in “atomic orbitals”

Can only determine the

probability of locating

an electron

e- “cloud”

27

Models

Dalton Thomson

+

Rutherford

Bohr

++

Quantum

+

28

Atomic Orbital

Each atomic orbital

can hold 2 e- maximum- -

A region in space around the

nucleus with high probability

of finding an electron.

Analogy: student in a desk

29

e- orbital (location) determined by:

1. Principal quantum number

-- “shell” or

-- “principle energy level”

Analogy- floor number

2. Sublevel

-- “subshell”

Analogy: apartment number

Only certain combinations are allowed.

30

Principal Quantum Number (n) or Principal Energy Level

n = 1, 2, 3, 4, 5 …

(integer values)

Gives overall energy of an e- and

its distance from nucleus.

31

Energy Sublevels(subshell)

Gives shape of the e- cloud

s, p, d, f

32

Energy Sublevels

s has 1 orbital

p has 3 orbitals

d has 5 orbitals

f has 7 orbitals

Each orbital can hold 2 e-

How many e-

can fit on each

sublevel?

33

Sublevel Shapes

sspherical

“90%”

px py pz

dumb bell

p

34

Allowed Combinations

n Sublevels # Orbitals #e

1 1s 1 2

2 2s 2p 4 8

3 3s 3p 3d 9 18

4 4s 4p 4d 4f 16 32

(model)

35

e- Configurations: 3 Rules

3. Hund’s Rule: maximize the

number of “parallel spin” e-

when filling a sublevel

1. Aufbau: arrange e- by

lowest energy level first

2. Pauli Exclusion Principle:

only 2 e- per orbital

36

1. Aufbau: Arrange e- by

lowest energy level first

lower than 3d !!

Increasing Energy

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s…

37

1s hold 2e-

3s hold 2e-

4s hold 2e-

3d hold 10e-

3p hold 6e-

2p hold 6e-

2s hold 2e-

38

Orbital “Box” Diagram

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

energy

3d 4p 5s

energy continued

39

2. Pauli Exclusion Principle: only 2 e- per orbital

1s orbital

Paired electrons: two e- in an

orbital have opposite “spins”

Electrons have spin!

40

3. Hund’s Rule

• Maximize parallel (same) spins

when filling a sublevel

• In a sublevel, put one e- in each

orbital before pairing

Example: 4 e- in a ‘p’ sublevel

Overview

41

https://www.youtube.com/watch?feature=player_emb

edded&v=8ROHpZ0A70I#t=4

42

Hydrogen: 1 electron

1s 2s 2p 3s

Recall for neutral atom #e- is same

as # p+ (atomic number).

43

Electron “Configuration”

1s

Hydrogen

1s1

n sublevel

# of e-

s

44

Regents Table Notation

Regents Periodic Table gives only

the number of electrons in each

principle energy level, ‘n’.

Hydrogen

1s1

1st level - 2nd level - 3rd level…

1Regents

45

Helium: 2 electrons

1s 2s 2p 3s

configuration:

1s2 2

Regents

46

Lithium: 3 electrons

1s 2s 2p 3s

configuration:

1s22s1 2-1

Regents

47

Boron: 5 electrons

1s 2s 2p 3s

configuration:

1s22s22p1 2-3

Regents

48

Nitrogen: 7 electrons

1s 2s 2p 3s

configuration:

1s22s22p3 2-5

Regents

49

Fluorine: 9 electrons

1s 2s 2p 3s

configuration:

1s22s22p5 2-7

Regents

50

Neon: 10 electrons

1s 2s 2p 3s

configuration:

1s22s22p6 2-8

Regents

51

Sodium: 11 electrons

1s 2s 2p 3s

configuration:

1s22s22p63s1 2-8-1

Regents

52

Shorthand for Sodium

1s22s22p63s1

[Ne] 3s1

Use the preceding [noble gas]

[ ] = “electron configuration of ”

53

Other 3rd Period Elements

Al 1s22s22p63s23p1 2-8-3

Ar 1s22s22p63s23p6 2-8-8

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

energy

54

Recall: 4s Out-of-Order

Energy

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s…

lower than 3d !!

So the next element, K, is

2-8-8-1 rather than 2-8-9

55

Other Elements

K [Ar]4s1 2-8-8-1

Sc [Ar]4s23d1 2-8-9-2

Cr [Ar]4s13d5 2-8-13-1weird

56

Noble Gases

At end of each row in Periodic Table

are the noble or inert gases with filled

ns and np orbitals.

Stable (not reactive) elements

57

1s 1s

2s 2p

3s 3p

4s 3d 4p

5s 4d 5p

6s 5d 6p

7s 6d

4f

5f

Periodic Table by Subshell

Inner transition elements filling

‘f’ orbitals

Transition elements

filling ‘d’ orbitals

58

Let’s Predict One

Using the Periodic Table, predict

the full, shorthand, and Regents e-

configuration for zinc.

59

You Try It !!

1. Predict the full, shorthand, &

Regents e- configuration of:

Ca As

2.How many unpaired electrons are

there in these elements?

3. Write the shorthand notation for:

Zr

Valence Electrons

60

How many “valence electrons” in:

Li Fe Cu

Electrons that are in the highest energy

level are called “valence electrons.”

These are the most important electrons

when atoms bond. Why?

Excited State

61

(e- have absorbed energy to move

to a higher energy level)

Remember “excited state”?

What atom is 1s22s22p33s1?What atom is 2-8-2-1?

[Al] is 2-8-3

normal Al

[Al]* could be 2-8-2-1

excited Al

62

Flame Test for Copper

Cu atom in excited state:

2-8-17-2

Cu atom in ground state:

2-8-18-1

Can return to ground state by

emitting energy as light

63

Flame Test for Copper

Which photon has greater energy:

An e- the falls from E5 to E3 or

An e- the falls from E5 to E2 ?

64

Weirdos !

Some ground state elements on

the Periodic Table have

1 or 2 electrons out of order.

e.g. Cu and Cr

You are not responsible for

determining the standard electron

configurations, just the Regents.

65

Bonding (the octet rule)

Electron configurations are the key

to bonding. Some atoms will

become ions to achieve Noble gas

electron configuration.

Octet rule

66

F-

1s 2s 2p

F

F atom vs. F- ion

= [Ne]extra e-

67

1s 2s 2p 3s

Na

Na+

Na atom vs. Na+ ion

= [Ne]missing e-

Practice

68

Write full electron configuration for:

P, Tc, S-2

69

Periodic Relationships

70Early chemists describe the first element.

71

Tabulation of Elements

•Tabulated by chem. &

physical properties

•Arranged by mass

•Predicted missing

elements and properties

Mendeleev (1869)

72

Modern Periodic Table

Now ordered by atomic

number, not mass.

Element 101 (Md)

Argon vs. potassium problem.

73

Groups 1, 2 & 13-17

Filling the s or p subshells

Last digit of group number gives

the number of valence electrons.

Representative Elements

74

Representative Elements

Group 1: alkali metals

Group 2: alkaline earth metals

Group 17: halogens

Some groups have special names

75

Noble Gases

filled p subshell (8 valence e- except He)

e.g. neon 1s22s22p6

very stable

(non-reactive)

76

filling d subshell

“d block”

Transition Elements

e.g. Iron

Regents: 2-8-14-2

Salts yield colored solutions.

77

Inner Transition Elements

filling the f subshell

“f block”

78

Trends in Atomic Size

Atomic size is

measured by radius.Table ‘S’

For chlorine:

= 100. pm

= ? mR

79

Atomic Radius: Trends?????????

OK

(model)

80

Atomic Radius

Down a Group: size increases

due to adding electrons to

higher energy levels (shells)

further from the nucleus.

81

Atomic Size: Across a Period

Electrons added to same shell

Nuclear charge increases (more p+)

Greater inward pull on the electrons

Atoms get smaller

2-8-1 2-8-3 2-8-6 2-8-8

p+ =11 12 13 14 15 16 17 18

82

Atomic Size: Across a Period

+5 +6

Boron (2-3) vs. Carbon (2-4)

smaller

83

Atomic Radius

smallerRow: greater

nuclear charge

larg

er Column: e- in

higher shell

84

Atomic Radius

Try It:

Arrange these atoms in order of

increasing size.

N, O, P, S

O < N < S < P

85

Ionization Energy (I)

Chemical properties determined

by valence electrons.

Ionization energy: energy (kJ/mol)

to remove an e- from an atom.

If I is high, e- held tightly.

86

1st ionization energy or I1

energy + X(g) X+(g) + e-

I is endothermic (need to put

energy in to pull off an e-)

Ionization Energy

87Atomic Number

I1

Atomic Number

I1

Ionization Energy: Table ‘S’

I1 across Period

I1 down Group

88

Trends in I (due to size)

I1 decreases going down a Group.

The e- are farther from the nucleus.

I1 increases going across a Period.

The e- are closer to the nucleus.

Which corner of Periodic Table has:

-highest I1?

-lowest I1?

89

I Predicts Ionic Charges

Element I1(kJ/mol)

I2(kJ/mol)

I3(kJ/mol)

Na 496 4565 6912

Mg 738 1450 7732

1s 2s 2p 3s

Na

90

Ionization Energy

Which has smaller I1 and why?

O or S

P or Cl

91

Trends in Ionic Size

Cation is smaller than its atom.

(less e- with same # protons)

Na-1e- Na+

160 pm 95 pm Al-3e- Al+3

124 pm 50 pm

92

Trends in Ionic Size

Anion is larger than its atom.

(more e- with same # protons)

Cl-+1e-Cl

100 pm 181 pmF-

+1e-F

60 pm 136 pm

93

Ionic Radii

cations anions

(model)

94

Ionic Radii

Place in order of increasing size.

Fe, Fe2+ and Fe3+

95

Try It !!!

1.Use e- configuration to

predict the charge of Ca ion.

2.Is this ion larger or smaller

than its atom?

96

Electronegativity

The tendency of an atom to

attract bonding electrons.

HH

OWater: which atom

“wins the battle” for

the bonding e-?

97

Electronegativity

Low attraction High attraction

for e- in bond for e- in bond

Least EN Most EN

An arbitrary scale from 0 to 4.

0 4

Fr (0.7) F (4.0)

98

Electronegativity

Why don’t the Noble gases

have electronegativity values?

99

Electronegativity

Example: Water

HH

O

2.2 2.2

3.4

HH

O

d+ d+

d-

Water is a “polar” molecule.

slightly

100

Electronegativity

Group Trend: EN decreases going

down a group. Atoms get larger, so

bonding e- are farther from the nucleus.

Period Trend: EN increases going

across a period. Atoms get smaller, so

bonding e- are closer to the nucleus.

(Similar to ionization energy.)

101

Metallic Character

Metals lose e- to become cations.

Which element is the most metallic?

(smallest ionization energy)

Nonmetals gain e- to become anions.

Which element is the least metallic?

(largest ionization energy)

102

“Diagonal

Relationships”

Smallest R

Largest I1Largest EN

Least metallic

Largest R

Smallest I1Smallest EN

Most metallic

103

Warm-up

104

What did Rutherford’s gold foil

experiment show about the structure of

the atom?

How did Bohr’s model of the atom

differ from the prior model of the atom?

Warm-up

105

What was Bohr’s explanation

for the emission or bright-line

spectrum of hydrogen?

+

Warm-up

106

What two quantum properties

determine the location of an electron

in an atom?electron

neutron

proton

Warm-up

107

n Sublevels No. orbitals No. e-

1

2

3

4

What is the relationship between the value

of n and the number of electrons?

Warm-up

108

• Describe the 3 rules for placing

electrons in orbitals around an atom.

• What is the order of energy sublevels,

listed in increasing energy?•

Warm-up

109

For oxygen & sulfur write:

•box diagram

•electron configuration

•shorthand notation

•Regents configuration

Warm-up

110

Write the shorthand e- config. for:

•iron

•cadmium

What is the max. number of e- that can

be on the 4th principal energy level in

the ground state?

Warm-up

111

How many valence e- in Cr?

What is the atomic size trend:

-down a group?

-across a row?

Which is larger, Si or As?

What is e- config. of Al+3?

Why?

Warm-up

112

Define first ionization energy, I1.

What is the trend in I1 across a row and

down a group? Explain.

Place the following elements in order

of increasing I1: P, Cl, As

Warm-up

113

What is “metallic character”?

How is metallic character related to

ionization energy?

What happens to metallic character

going down Group 15?

Which has greater metallic

character: Fe or Na?

Warm-up

114

Define each term, state the trend, and

explain why:

•Atomic radius across a row

•Ionization Energy down a group

•Electronegativity across a row

•Metallic character down a group

Element Song

115

http://www.privatehand.com/flash

/elements.html