1

Electrons in Atoms & Periodic Table

2

Quantum Theory or Wave Theory:

a description of the

electron “configuration.”

Where Are the Electrons?

3

Quantum Theory

One of the greatest

achievements of mankind.

4

Arrangement of e- in Atoms

chemical reactivity

bonding between atoms

Periodic Table

many physical properties

Determines:

5

Atomic Models: History

Each atomic model was

eventually replaced because 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

positive charge

8

Rutherford: 1911

Dense 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

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- transition between orbits.

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

Bohr Theory: Failings

• Why do e- only have

certain orbit energies?

• Only explains the

hydrogen atom exactly.

23

Quantum Mechanics

(Wave Theory)

1926: E. Schrodinger

Can not determine exact location

of an electron! Wow!

Currently accepted theory

24

Quantum Mechanics

e- in “atomic orbitals”

Can only determine

the probability of

locating an electron.

electron cloud

25

Models

Dalton Thomson

+

Rutherford

Bohr

++

Quantum

26

Atomic Orbital

Each atomic orbital

can hold 2 e-- -

A region in space around the

nucleus with high probability

of finding an electron.

Analogy: student in a desk

Wave Model

27

Each e- is arranged in an atom

according to its energy.

+1st energy level

(lowest shell)

2nd energy level

3rd energy level

(higher shell)

Overview

28

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

_embedded&v=8ROHpZ0A70I#t=4

Bohr Quantum

29

Regents e- Notation

Regents Periodic Table gives

the number of electrons in each

energy level or shell.

1st shell - 2nd shell - 3rd shell…

[C] = 2 - 4

30

Regents e- “Configuration”

What is the e- configuration for:

1. sodium 2. argon

3. calcium 4. copper

7. lead 6. radium (Ra)

What is the maximum electrons in:

Shell 1? Shell 2? Shell 3?

Regents e- “Configuration”

31

What happens to the electrons in

each shell going from

Ca to Sc? Zn to Ga?

32

Noble Gases

At end of each row in Periodic

Table are the noble or inert gases

with 8 e- in the highest shell.

Stable (not reactive) elements

33

1 1

2 2

3 3

4 3 4

5 4 5

6 5 6

7 6 7

4

5

Periodic Table by Shell

Inner transition elements

Transition elements

Valence Electrons

34

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?

Valence Electrons

35

Note all elements in a Group have

same number of valence electrons.

Group 1: Li, Na, K, Rb, Cs, Fr

Group 16: O, S, Se, Te, Po

Excited State

36

(e- have absorbed energy to

move to a higher energy level)

Remember “excited state”?

What element is 2-8-2-1?

What element is 2-7-6-1?

37

Flame Test (Lab)

Adding energy can cause e-

to “jump” from ground

state (as written in Regents

table) to “excited state”.

When e- fall back, they

emit light.

38

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

39

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 ?

40

Emission Spectrum

Flame test

Neon signs

Fireworks

Fireplace colors

41

Bonding

•Electron configurations are

the key to bonding.

•Atoms become ions to achieve

Noble gas electron configuration.

42

F atom: [F] = 2-7

= [Ne]

Na atom: [Na] = 2-8-1

= [Ne]

Atoms vs. Ions

F- ion: [F-] = 2-8

Na+ ion: [Na+] = 2-8

What does F need to have the e-

configuration of a Noble gas?

43

Practice

Write the e- configuration for

Fe & Cu.

Based on e- configuration,

predict the charge of:

Mg ion S ion

Write two e- configurations for

excited states of calcium.

44

Periodic Relationships

45Early chemists describe the first element.

46

Tabulation of Elements

•Tabulated by chem. &

physical properties

•Arranged by mass

•Predicted missing elements

and properties

Mendeleev (1869)

47

Modern Periodic Table

Now ordered by atomic number,

not mass.

Element 101 (Md)

Argon vs. potassium problem.

48

Periodic Table

Key to understanding chemical

and physical properties

Each group has same electron

configuration for outer shell.

Most important tool in

chemistry

49

Regents Periodic Table

Symbol

Atomic number & atomic mass

Electron configuration

“Charges”

Elements arranged by

atomic no. (#p+)

50

Groups 1, 2 & 13-17

Last digit of group number gives

the number of valence electrons.

Examples: oxygen Group 16(6 VE)

sodium Group 1

(1 VE)

Representative Elements

51

Representative Elements

Group 1: alkali metals

Group 2: alkaline earth metals

Group 17: halogens

Some groups have special names

52

Noble Gases

Last element in each Period

8 VE, except He

very stable

(non-reactive)

53

Transition Elements

e.g. Iron

Regents: 2-8-14-2

Compounds with these

elements have colored solutions.

e- filling

54

Inner Transition Elements

At bottom of

Periodic Table

for convenience.

55

Trends in Atomic Size

Atomic size is

measured by radius. Table ‘S’

For chlorine:

radius = 100. pm

What is its radius

in meters?R

56

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

OK

(model)

57

Atomic Radius

Down a Group: size

increases due to adding

electrons to higher energy

levels (shells) further from

the nucleus.

58

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 2-8-5 2-8-8

p+ =11 12 13 14 15 16 17 18

59

Atomic Size: Across a Period

+5 +6

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

smaller

60

Atomic Radius

smallerRow: greater

nuclear charge

larg

er Column: e- in

higher shell

61

Atomic Radius

Try It:

Arrange these atoms in order of

increasing size.

N, O, P, S

O < N < S < P

62

Ionization Energy (I)

Chemical properties determined

by valence electrons.First ionization energy (I): energy

(kJ/mol) to remove an e-

from an atom.

If I high, e- held tightly.

63

I + X(g) X+(g) + e-

I is endothermic (need to put

energy in to pull off an e-)

Ionization Energy

ionization

energy

64Atomic Number

I1

Atomic Number

I1

Ionization Energy: Table ‘S’

I1 across Period

I1 down Group

65

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?

66

I Predicts Ionic Charges

Element I1(kJ/mol)

I2(kJ/mol)

I3(kJ/mol)

Na 496 4565 6912

Mg 738 1450 7732

Na atom 2-8-1 lose 1 e-

Na+ ion 2-8

Mg atom 2-8-2 lose 2 e-

Mg+2 ion 2-8

67

Ionization Energy

Which has smaller I and why?

O or S

Ge or Br

68

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

69

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

70

Ionic Radii

cations anions

(model)

71

Ionic Radii

Place in order of increasing size.

Fe, Fe2+ and Fe3+

72

Try It !!!

1. Use e- configuration to

predict the charge of Ca ion.

2. Is this ion larger or smaller

than its atom?

73

Electronegativity

The tendency of an atom to

attract bonding electrons.

HH

OWater: which atom

“wins the battle” for

the bonding e-?

74

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)

75

Electronegativity

Why don’t the Noble gases

have electronegativity values?

76

Electronegativity

Example: Water

HH

O

2.1 2.1

3.4

HH

O

d+ d+

d-

Water is a “polar” molecule.

slightly

77

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.

(Same trend as ionization energy.)

78

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)

79

“Diagonal

Relationships”

Smallest R

Largest I1Largest EN

Least metallic

Largest R

Smallest I1Smallest EN

Most metallic

80

Warm-up

81

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

82

What was Bohr’s explanation

for the emission or bright-line

spectrum of hydrogen?

+

Warm-up

83

What is the name of the region

outside the nucleus where electrons

are most probably found?

Warm-up

84

What is the name of the region outside

the nucleus where electrons are most

probably found?

Write the Regents electron

configuration for arsenic.

What does each of the numbers mean?

How many valence electrons does

arsenic have?

Warm-up

85

•What is 2-8-7?

•What is 2-7-8?

•What is the e- configuration of gold?

•How many valence electrons does

manganese have?

•What is the electron configuration

of the nitride ion?

Warm-up

86

What are the names of

Groups: 1, 2, 17, and 18?

How many valence e- in Co?

What is the trend size:

-down a group?

-across a row?

What is e- config. of Al+3?

Why?

Warm-up

87

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

Which element, P or S, is bigger

(larger radius)? Explain.

Warm-up

88

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

89

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

90

http://www.privatehand.com/flash

/elements.html