1

Modern Atomic Theory

and the Periodic TableChapter 10

Hein and Arena

Eugene Passer

Chemistry Department

Bronx Community College

© John Wiley and Sons, Inc.

Version 1.1

the exam would have to be given earlier

2

A Brief Review

3

Composition of the Atom

1. The atom has a dense Nucleus containing

Protons( p) and Neutrons( n).

2. n have no charge. p have a relative charge

of +1.

3. The atomic number of an element is the

same as the number of p in the Nucleus.

4. Electrons( e-) are located outside of the

Nucleus. e- have a relative charge of -1.

4

Electromagnetic Radiation

5

Energy can travel through space as

electromagnetic radiation.Examples

6

• light from the sun

• x-rays

• microwaves

• radio waves

• television waves

• radiant heat

All show wavelike

behavior.

Each travels at

the same speed,

c, in a vacuum.

c = 3.00 x108 m/s

7

Characteristics of a Wave

8

Wavelength (λ)

9

Wavelength (nm)

(measured from

peak to peak)

Wavelength (nm)

(measured from

trough to trough)

10.1

Light has the properties of a wave.

10

Frequency (ν)

11

Frequency is the number of wavelengths

that pass a particular point per second

(i.e. s-1 or Hz).

10.1

12

Speed (v)

13

Speed is how fast a wave moves through

space (i.e. a non-vacuum environment).

10.1

14

• Light also exhibits the properties of a

particle. Light particles are called

photons, which have discrete energies

(i.e. E = hn, where h is Plank’s Constant).

• The wave model and the particle

model are both used to explain the

properties of light.

15

The Electromagnetic Spectrum

16

10.2

visible light is part of

the electromagnetic

spectrum

Each color in the visible region has a very specific wavelength (l);

red light has a l =~650 nm while blue light has a l =~ 450 nm.

17

• At high temperatures, or voltages,

gaseous atoms are excited; excited

atoms in the gaseous state emit light of

different colors upon relaxation.

• When the light is passed through a

prism or diffraction grating a line

spectrum results.

18

Line spectrum of hydrogen. Each line (i.e. color)

corresponds to the wavelength of the energy

emitted when the electron of a hydrogen atom,

which has absorbed energy falls back to a lower

principal energy level.

These colored lines

indicate that light is

being emitted only at

certain wavelengths.

Each element has its own

unique set of spectral emission

lines that distinguish it from

other elements.

10.3

19

The Bohr Atom

20

Electrons revolve

around the nucleus in

orbits that are located

at fixed distances from

the nucleus.

10.4

An electron has a

discrete energy when it

occupies an orbit.

21

When an electron falls

from a higher energy level

to a lower energy level a

quantum of energy in the

form of light is emitted by

the atom.

10.4

The color of the light

emitted corresponds to

one of the lines of the

hydrogen spectrum.

22

Different lines of the

hydrogen spectrum

correspond to different

electron energy level

shifts.

10.4

23

Light is not emitted

continuously. It is

emitted in discrete

packets called quanta.

10.4

24

Bohr’s H atom

with quantized

energy levels;

e- energies:

E3 > E2 > E1.

E2 E3E1

10.4

25

• Bohr’s calculations succeeded very well

in correlating the experimentally

observed spectral lines with electron

energy levels for the hydrogen atom.

• Bohr’s methods did not succeed for

heavier atoms (i.e. atoms with several e-).

• More theoretical work on atomic structure

was needed.

26

Erwin Schröedinger created a mathematical

model that showed electrons as waves; this lead

to the duality of matter (i.e. e-:wave or particle).– Schröedinger’s work led to a new branch

of physics called wave or quantum

mechanics.

– Using Schröedinger’s equation, the

probability of finding an electron in a

given region of space around the nucleus

can be determined.

– The actual location of an electron within

an atom cannot be determined.

27

• Based on wave mechanics it is clear that

electrons are not revolving around the

nucleus in orbits.

• Instead of being located in orbits, the

electrons are located in orbitals.

• An orbital is a region around the nucleus

where there is a high probability (~90%)

of finding an electron.

28

Energy Levels

of Electrons

29

According to Bohr the energies of

electrons in an atom are quantized.

The wave-mechanical model of the atom

also predicts discrete principal energy

levels for e- of an atom.

30

The first four

principal energy

levels of the

hydrogen atom.

As n increases, the

energy of the electron

increases.

10.7

Each level is

assigned a principal

quantum number n.

31

Each principal energy level

is subdivided into sublevels.10.7, 10.8

32

Within sublevels the electrons are found in

orbitals.

An s orbital is spherical in

shape.

The spherical surface

encloses a space where

there is a 90% probability

that the electron may be

found.

10.10

33

An electron can spin in one of

two possible directions

represented by ↑ (spin up) or

↓ (spin down).

Two electrons that occupy the

same atomic orbital must

have opposite spins.

This is known as the Pauli

Exclusion Principal.10.10

An atomic orbital can hold a maximum of two

electrons.

34

Each p atomic orbital has two lobes.

Each p atomic orbital can hold a maximum of

two electrons.

A p sublevel can hold a maximum of 6

electrons. 10.10

A p sublevel is made up of three p orbitals.

35

The three p orbitals share

a common center, the

nucleus.

10.10

pxpy

pz

The three p orbitals point in

different directions along x,

y and z in 3-space.

36

The five d atomic orbitals all point in different

directions.

Each d atomic orbital can hold a maximum of two

electrons.

A d sublevel can hold a maximum of 10 electrons.10.11

A d sublevel is made up of five d orbitals.

37

n=1 1s

1

n=2 2s

1

2p 2p 2p

3

n = 3 3s

1

3p 3p 3p

3

3d 3d 3d 3d 3d

5

n = 4 4s

1

4p 4p 4p

3

4d 4d 4d 4d 4d

5

4f 4f 4f 4f 4f 4f 4f

7

Distribution of Sublevels

and Orbitals by Principal

Energy Level

Note: The number in RED indicates the number of orbitals in a given sublevel

38

The Hydrogen Atom

• In the ground state

hydrogen’s single

electron lies in the

1s orbital.

• Hydrogen can

absorb energy and

the electron will

move to an excited

state.10.12

39

Electron Configurations

of the First 18 Elements

40

Electron Configuration

Step by Step

41

1. Maximum of

two electrons

per orbital

10.10

42

1 s orbital

2. Lowest energy orbitals occupied first; higher

energy orbitals occupied only after the lower

energy orbitals are filled.

3. Orbital energies: s < p < d < f for a given

value of n.

2 s orbital

10.10

43

4. Each orbital in a sublevel is occupied

by a single electron before a second

electron enters. For example, all three

p orbitals must contain one electron

before a second electron enters a p

orbital.10.10

44

Electron Configuration

Arrangement of

electrons within their

respective sublevels. 2p6

Principal

energy levelType of orbital

Number of

electrons in

sublevel orbitals

45

Orbital Filling

46

• The following diagrams boxes

represent orbitals.

• Electrons are indicated by arrows: ↑ or

↓.

– Each arrow direction represents one of

the two possible electron spin states.

47

Filling the 1s Sublevel

48

↑ 1s2↓

H ↑ 1s1

Hydrogen has 1 electron. It will occupy the orbital of lowest

energy which is the 1s.

He

Helium has two electrons. Both helium electrons occupy the

1s orbital with opposite spins.

49

Filling the 2s Sublevel

50

Li

1s22s2

The 1s orbital is filled. Lithium’s third electron will enter the

2s orbital.

↑↓ ↑

↓

1s 2s

↑

1s22s1

Be ↑↓

The 2s orbital fills upon the addition of beryllium’s third and

fourth electrons.

1s 2s

51

Filling the 2p Sublevel

52

B 1s22s22p1

1s 2s 2p

↑↓ ↑↓ ↑

Boron has the first p electron. The three 2p orbitals have the same

energy. It does not matter which orbital fills first.

C

1s 2s 2p

↑↓ ↑↓

The second p electron of carbon enters a different p orbital than the

first p electron so as to give carbon the lowest possible energy.

1s22s22p2

↑ ↑ ↑N

1s 2s 2p

↑ ↓↓ ↑

The third p electron of nitrogen enters a different p orbital than its

first two p electrons to give nitrogen the lowest possible energy.

1s22s22p3

↑ ↑

53

↑ 1s22s22p4↑ ↑

1s 2s 2p

↑↓ ↑↓O

There are four electrons in the 2p sublevel of oxygen. One of the

2p orbitals is now occupied by a second electron, which has a spin

opposite to that of the first electron already in the orbital.

↑↓ ↑↓ ↑

2p

F

1s 2s

↑↓ ↑↓

There are five electrons in the 2p sublevel of fluorine. Two of the 2p

orbitals are now occupied by a second electron, which has a spin

opposite to that of the first electron already in the orbital.

1s22s22p5

↓

↑↓ ↑↓ ↑↓

2p

Ne

1s 2s

↑↓ ↑↓

There are 6 electrons in the 2p sublevel of neon, which fills the

sublevel.

1s22s22p6

This electronic configuration is referred to as a

“full octet” (ns2np6). All noble gases, except He,

have a full octet! It is this electron configuration

that make noble gases un-reactive.

The valence electrons occupy the valence shells (i.e. ns2np6)

55

Filling the 3s Sublevel

56

Na 1s22s22p63s1

1s 2s 2p 3s

↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑

The 2s and 2p sublevels are filled. The next electron enters the

3s sublevel of sodium.

Mg

1s 2s 2p 3s

↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑

The 3s orbital fills upon the addition of magnesium’s twelfth

electron.

1s22s22p63s2↓

The valence electrons occupy the valence shells (i.e. ns2np6)

57

58

Which atom above has the largest number of unpaired electrons?

P

59

Electron Configurations

and the Periodic Table

60

10.15

Elements in the same group, or family, have

similar chemical properties because the valence

electron (i.e. ns2np6) configuration is the same in

each group, or family.

61

10.15

With the exception of helium which has a filled s

orbital, the nobles gases have filled p orbitals.

62

The electron configuration of any of the

noble gas elements can be represented by

the symbol of the element enclosed in

square brackets; the core.

B 1s22s22p1 [He]2s22p1

Cl 1s22s22p63s23p5 [Ne]3s23p5

Na 1s22s22p63s1 [Ne]3s1

63

The electron configuration of argon is

Ar 1s22s22p63s23p6

Ca 1s22s22p63s23p6 4s2 [Ar]4s2

K 1s22s22p63s23p64s1 [Ar]4s1

The elements after argon are potassium

and calcium. Instead of entering a 3d

orbital, the valence electrons of these

elements enter the 4s orbital.

64

10.16

d orbital filling

d orbital numbers are 1 less

than the period number

Arrangement of electrons

according to sublevel being filled.

65

10.16

f orbital filling

f orbital numbers are 2 less

than the period number

Arrangement of electrons

according to sublevel being filled.

66

10.17

Period number corresponds with the

highest principle energy level occupied by

electrons in that period.

67

10.17

The group numbers for the representative

elements are equal to the total number of

valence electrons in the atoms of the group.

The elements of a group have the same

valence electron configuration except that the

electrons are in different principle energy

levels.

68