1

2

A Brief HistoryA Brief History

3

Lucretius On the nature of the Universe (55 BC)

• The light and heat of the sun; these are composed of minute atoms which, when they are shoved off, lose no time in shooting right across the interspace of air in the direction imparted by the shove.

http://www.gap-system.org/~history/HistTopics/Light_1.html

4

Electromagnetic RadiationElectromagnetic Radiation

5

6

Properties of Light

• Travels in straight lines. Casts shadows

• Carries energy. Heats objects it strikes.

• Travels fast. No delay when turning on a light switch.

• Comes in different colors.

7

Christiaan Huygens1678

http://www.launc.tased.edu.au/online/sciences/physics/diffrac.html

8

Huygens’ Principle• Throw a rock in a quiet pool, and

waves appear along the surface of the water.

• Huygens proposed that the wavefronts of light waves spreading out from a point source can be regarded as the overlapped crests of tiny secondary waves.

• Wavefronts are made up of tinier wavefronts—this idea is called Huygens’ principle.

9

Huygens’ PrincipleEvery point of a wavefront may be considered the source of secondary wavelets that spread out in all directions with a speed equal to the speed of propagation of the waves.

10

Huygens’ Principle• Plane waves can be generated in

water by successively dipping a horizontally held straightedge into the surface

• As the width of the opening is narrowed, less of the incident wave is transmitted.

• The spreading of waves into the shadow region becomes more pronounced.

11

Characteristics of a WaveCharacteristics of a WaveCharacteristics of a WaveCharacteristics of a Wave

12

Wavelength (λ)

13

wavelength(measured from peak to peak)

wavelength(measured from trough to trough)

Light has the properties of a wave.

14

Frequency (ν)

15

Frequency is the number of wavelengths that pass a particular point per second.

16

Diffraction

17

Diffraction

18

Thomas Young

19

Constructive Interference

20

Destructive Interference

The interference of water is a common sight. In some places, crests overlap crests; in other places, crests overlap troughs of other waves.

Interference patterns of overlapping waves from two vibrating sources.

Thomas Young’s original drawing of a two-source interference pattern.The dark circles represent wave crests; the white spaces between crestsrepresent troughs. Constructive interference occurs where crests overlapcrests or troughs overlap troughs. Letters C, D, E, and F mark regions of destructive interference.

24

Interference Patterns

26

Young’s Experiment 1802

http://www.acoustics.salford.ac.uk/feschools/waves/diffract3.htm

27

Young’s Experiment 1802

28

29

Max Planck

30

• In 1901, Planck developed the quantum theory of light.

• Light is not continuous.

• Light is a stream of distinct units of energy called quanta.

• The amount of energy on one of these units depends on the frequency.

• Light is absorbed or emitted in multiples of whole quanta.

• Quanta is like an atom of light called a photon

ROYGBIV

31

“I think it is safe to say that

no one understands quantum mechanics. Do not keep saying to yourself, if you can possibly avoid it, ‘But how can it be like that?’ because you will go ‘down the drain’ into a blind alley from which nobody has yet escaped. Nobody knows how it can be like that.”

 

Richard Feynman

1918-1988

The Manhattan Project

32

33

The Electromagnetic Spectrum

34

visible light is part of the electromagnetic

spectrum

X-rays are part of the electromagnetic

spectrum

Infrared light is part of the

electromagnetic spectrum

35

36

The Bohr AtomThe Bohr Atom

37

Niels Bohr, a Danish physicist, in 1912-1913 carried out research

on the hydrogen atom

38

A young man working with Rutherford in 1911 was Niels Bohr. They created a halfway successful model of the atom called the Rutherford-Bohr model. They imagined a nucleus at the center with electrons orbiting around it. Twenty-five years later, Bohr described his collaboration with Rutherford...

If, twenty-five years ago, I had the good fortune to give a modest contribution to this development, it was, above all, thanks to the hospitality I then, as a young man, enjoyed in the famous laboratories of England. In particular, I think with grateful emotion of the unique friendliness and straightforwardness with which Rutherford, in the midst of his unceasing creative activity, was always prepared to listen to any student behind whose youthful inexperience he perceived a serious interest.

39

• At high temperatures or voltages, elements in the gaseous state emit light of different colors.

• When the light is passed through a prism or diffraction grating, a line spectrum results.

40

Sir Isaac Newton’s crucial experiment, 1672. he refracts white light with a prism, resolving it into its component colors: red, orange, yellow, green, blue and violet.

41

He Ne Ar Kr Xe

42

Flame test colours

Ba Li Sr Na Cu K

43

Diffraction grating• Composed of a large number of close, equally

spaced slits for analyzing light source

• Produced by spectrometers that disperse white light into colors

44

The lenses of “rainbow” glasses are diffraction gratings. Looking through them causes light to separate into its color components. The spectral patterns seen from light sources such as fireworks and neon signs emit a distinct number of discontinuous colors and do not show all the colors of the rainbow. The spectral patterns from these light sources are the atomic spectra of elements heated in the light sources.

Lamps of a chandelier seen through diffraction-grating glasses.

46

47

Line spectrum of hydrogen. Each line 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.

Incandescent bulb has a continuous spectrum

Hydrogen

Sodium

Mercury

49

The Bohr Atom

50

Electrons revolve around the nucleus in orbits that are located at fixed distances from the nucleus.

An electron has a discrete energy when it occupies an orbit.

51

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.

The color of the light emitted corresponds to one of the lines of the hydrogen spectrum.

52

Different lines of the hydrogen spectrum correspond to different electron energy level shifts.

53

Light is not emitted continuously. It is emitted in discrete packets called quanta.

54

An electron can have one of several possible energies depending on its orbit.

E2 E3E1

55

• 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.

• More theoretical work on atomic structure was needed.

56

Baseball Waves

57

58

• In 1924 Louis De Broglie suggested that all objects have wave properties.– De Broglie showed that

the wavelength of ordinary-sized objects, such as a baseball, are too small to be observed.

– For objects the size of an electron the wavelength can be detected.

59

60

In this view, the electron is thought of not as a particle located at some point in the atom but as if its mass and charge were spread out into a standing wave surrounding the atomic nucleus with an integral number of wavelengths fitting evenly into the circumferences of the orbits.

(a) An orbiting electron forms a standing wave only when the circumference of its orbit is equal to a whole-number multiple of the wavelength.

(b) When the wave does not close in on itself in phase, it undergoes destructive interference. Hence, orbits exist only where waves close in on themselves in phase.

The circumference of the innermost orbit, according to this picture, is equal to one wavelength.

• The second has a circumference of two electron wavelengths.

• The third orbit has a circumference of three electron wavelengths, and so forth.

• The electron orbits in an atom have discrete radii because the circumferences of the orbits are whole-number multiples of the electron wavelength.

• This results in a discrete energy state for each orbit.

• This model explains why electrons don’t spiral closer and closer to the nucleus, causing atoms to shrink to the size of the tiny nucleus.

• If each electron orbit is described by a standing wave, the circumference of the smallest orbit can be no smaller than one wavelength.– No fraction of a wavelength is possible in a circular (or

elliptical) standing wave.

• As long as an electron carries the momentum necessary for wave behavior, atoms don’t shrink in on themselves.

65

For a fixed circumference, only some wavelengths are self-reinforcing.

Wire loop at rest.

Tacoma Narrows Bridge Collapse "Gallopin' Gertie"

Nov. 7, 1940

• http://www.youtube.com/watch?v=j-zczJXSxnw

67

Davisson & Germer 1927

68

The diffraction on the left was made by a beam of x-rays passing through thin aluminum foil. The diffraction on the right was made by a beam of

electrons passing through the same foil.

69

Baseball Waves

70

71

• In 1926 Erwin Schrödinger created a mathematical model that showed electrons as waves.

72

– Schröedinger’s work led to a new branch of physics called wave or quantum mechanics.

– Using Schröedinger’s wave mechanics, the probability of finding an electron in a certain region around the atom can be determined.

– The actual location of an electron within an atom cannot be determined.

73

• 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 of finding an electron.

74

Energy LevelsEnergy Levelsof Electronsof Electrons

75

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 within the atom

76

The first four principal energy levels of the hydrogen atom.

As n increases, the energy of the electron increases.

Each level is assigned a principal quantum number n.

77

Each principal energy level is subdivided into sublevels.

78

Within sublevels the electrons are found in orbitals.

An s orbital is spherical in shape (s from sharp).

The spherical surface encloses a space where there is a 90% probability that the electron may be found.

79

An electron can spin in one of two possible directions represented by ↑ or ↓.

The two electrons that occupy an atomic orbital must have opposite spins.

This is known as the Pauli Exclusion Principal.

An atomic orbital can hold a maximum of two electrons.

80

Each p orbital has two lobes.

Each p orbital can hold a maximum of two electrons.

A p sublevel can hold a maximum of 6 electrons.

A p sublevel is made up of three orbitals.

81

The three p orbitals share a common center.

pxpy

pz

The three p orbitals point in different directions.

82

The five d orbitals all point in different directions.

Each d orbital can hold a maximum of two electrons.

A d sublevel can hold a maximum of 10 electrons.

A d sublevel is made up of five orbitals.

83Number of Orbitals in a Sublevel

84

f Orbitals

86

Web Sites

• http://web.mit.edu/3.091/www/orbs/

• http://winter.group.shef.ac.uk/orbitron/AOs/1s/index.html

87

Main energy

Level (n)

Maximum Sublevels (n)

Maximum Orbitals

(n2)

Maximum Electrons

(2n2)

1 1: s 12 = 1 2

2 2: s and p 22 = 4 8

3 3: s, p, and d 32 = 9 18

4 4: s, p, d, and f 42 = 16 32

Distribution of Sublevels by Principal Energy Level

Atomic Models

John Dalton’s atom 1803 Thomson’s pudding model 1897 Rutherford’s nuclear model 1911

Bohr’s orbit model 1913Schrödinger electron cloud model 1926

89

Atomic Structure of the Atomic Structure of the First 18 ElementsFirst 18 Elements

90

To determine the electronic structures of atoms the following guidelines are used.

91

1. No more than two electrons can occupy one orbital and then only if they have opposite spin. This is the Pauli Exclusion Principle.

92

Wolfgang Ernst Pauli

93

1 s orbital

2. Electrons occupy the lowest energy orbitals available. They enter a higher energy orbital only after the lower orbitals are filled. This is called the Aufbau Principle.

3. For the atoms beyond hydrogen, orbital energies vary as s<p<d<f for a given value of n.

2 s orbital

94

4. Each orbital on a sublevel is occupied by a single electron before a second electron enters. This is Attila the Hun’s Rule (Hund’s Rule). For example, all three p orbitals must contain one electron before a second electron enters a p orbital.

95

Friedrich Hund1896-1997

96

Attila the Hun

97

Nuclear makeup and electronic structure of each principal energy level of an atom.

number of electronsin each sublevel

number of protons and neutrons in the nucleus

98

Electron Configuration Using Spectroscopic Notation

Arrangement of electrons within their respective sublevels. 2p6

Principal energy level

Type of orbital

Number of electrons in

sublevel orbitals

99

Orbital FillingOrbital FillingOrbital FillingOrbital Filling

100

• In the following diagrams, boxes represent orbitals.

• Electrons are indicated by arrows: ↑ or ↓.– Each arrow direction represents one of

the two possible electron spin states.

105

Filling the 1s Sublevel

106

↑ 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.

107

Abbreviated Electron Configurations

Use the symbol of the nearest preceding noble gas to represent the electron configuration of the core electrons.

Phosphorus: 1s2 2s2 2p6 3s2 3p3

[Ne] 3s2 3p3

Core Electrons

Valence Electrons

108

Filling the 2s Sublevel

109

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

[He]2s1

[He]2s2

110

Filling the 2p Sublevel

111

B 1s22s22p1

1s 2s 2p

↑↓ ↑↓ ↑

Boron has the first p electron. The three 2p orbitals have the same energy (degenerate). It does not matter which orbital fills first.

C1s 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

↑ ↑ ↑N1s 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

↑ ↑

[He]2s22p1

[He]2s22p2

[He]2s22p3

112

↑ 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 that of the first electron already in the orbital.

↑↓ ↑↓ ↑

2p

F1s 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 that of the first electron already in the orbital.

1s22s22p5

↓[He]2s22p4

[He]2s22p5

113

↑↓ ↑↓ ↑↓

2p

Ne1s 2s

↑↓ ↑↓

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

1s22s22p6

114

Filling the 3s Sublevel

115

Na 1s22s22p63s1

1s 2s 2p 3s

↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑

The 2s and 2p sublevels are filled. The next electron enters the 3s sublevel of sodium.

Mg1s 2s 2p 3s

↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑

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

1s22s22p63s2↓

[Ne]3s1

[Ne]3s2

116

117

118

Electron Structures and Electron Structures and the Periodic Tablethe Periodic Table

119

In 1869 Dimitri Mendeleev of Russia and Lothar Meyer of Germany independently published periodic arrangements of theelements based on increasing atomic masses.

Mendeleev’s arrangement is the precursor to the modern periodic table.

120

Dimitri Mendeleev

121

122

Glenn Seaborg1912-1999

123

Horizontal rows are called periods.

Period numbers correspond to the highest occupied energy level.

124

Elements in the A groups are designated

representative elements.

Elements in the B groups are designated transition elements.

Elements with similar properties are organized in groups or families.

Groups are numbered with Roman numerals.

125

Some A groups have family names.

Alkali metals (H not included)

Alkaline earth metals Halogens

Noble gasesChalcogens

126

127

128

For A family elements the valence electron configuration is the same in each column.

The chemical behavior and properties of elements in a family are associated with the electron configuration of its elements.

129

With the exception of helium which has a filled s orbital, the nobles gases have filled p orbitals.

130

d orbital filling d orbital numbers are 1 less

than the period number

Arrangement of electronsaccording to sublevel being filled.

131

f orbital filling f orbital numbers are 2 less

than the period number

Arrangement of electronsaccording to sublevel being filled.

132

Period number corresponds with the highest energy level occupied by

electrons in that period.

133

The group numbers for the representative elements are equal to the total number of

outermost electrons in the atoms of the group.

The elements of a family have the same outermost electron configuration except that the electrons are in different energy levels.

134