1

Atomic Structure

2

Discovery and Properties of Electrons Humphrey Davy (early 1800’s) -

passed electricity through compounds compounds decomposed into elements

compounds are held together by electrical forces

Michael Faraday - (1832-1833) - amount of reaction that occurs during electrolysis is proportional to current passed through compounds

3

Discovery of the Electron J.J. Thomson proved the existence

of the electron by showing that the beam in the Crookes tube experiments could be deflected when passed between two plates containing opposite charges.

4

Discovery of the Electron Electrons carry the - charge since

they come from the negative electrode and go to the positive electrode.

5

e/m = 1.75882 x 108 coulomb/g

Discovery of the Electron J.J. Thomson determined the e/m

ratio by measuring the degree of deflection of cathode rays.

6

Millikan Oil-Drop Experiment

Determined the charge electron.

-1.60218 x 10-

19C/e

7

Electron Mass Millikan’s determination of the

charge of an electron allowed for the determination of the mass of an electron.

electrongxm

CxxCx

gm

/1010940.9

1060218.11075882.1

1

28

198

8

Discovery of Proton Rutherford shot -particles at thin gold

(Au) foil to see if they would be deflected.

Some -particles were deflected back. This could happen only if a highly

concentrated + charge was deflecting the positively charged -particle.

9

Rutherford Model of the Atom Atoms consist of

very small, very dense nuclei surrounded by clouds of electrons at relatively great distances from the nuclei.

10

Atomic Theory All nuclei contain

protons. Protons have a

positive charge. analyzed evidence

from -particle scattering

recognized existence of massive neutral particles - neutrons

James Chadwick-1932

11

Mass Number & Isotopes H.G. J. Moseley (1912-1914) -

recognized that atomic number is the defining difference between elements new understanding of Mendeleev’s

periodic law

12

Atomic Theory All atoms,

except for hydrogen, also contain neutrons.

Neutrons do not have a charge.

13

Atom CompositionThe atom is mostly empty space

protons and neutrons in the nucleus.

the number of electrons is equal to the number of protons.

electrons in space around the nucleus.

extremely small. One teaspoon of water has 3 times as many atoms

as the Atlantic Ocean has teaspoons of water.

Structure of the Atom

Composed of: protons neutrons electrons

protons found in

nucleus relative

charge of +1 relative mass

of 1.0073 amu

Structure of the Atom

Composed of: protons neutrons electrons

neutrons found in

nucleus neutral charge relative mass

of 1.0087 amu

Structure of the Atom

Composed of: protons neutrons electrons

electrons found in

electron cloud relative charge

of -1 relative mass

of 0.00055 amu

17

How Large is an Atom?

Scanning Tunneling Microscopic images of carbon atoms in graphite.

18

10B

11B

IsotopesTwo or more forms of atoms of the same

element with different masses.Atoms contain the same number of protons

but different numbers of neutrons.Boron-10 (10B)

has 5 p and 5 nBoron-11 (11B)

has 5 p and 6 n

19

Isotopes of HydrogenSymbol Nuclide Protons Neutrons Electrons

H 11H 1 0 1

D 21H 1 1 1

T 31H 1 2 1

20

661212

number of protonsnumber of protons

661212CCCarbon-12,Carbon-12,

Mass Number Mass Number: total number of

protons + neutrons (nucleons) in an atom.Mass Number = # protons + # neutrons

mass number = 6 p + 6 n = 12 amu

C

21

# Neutrons = Mass Number - # Protons C12

6 C613 C6

14

6 protons 6 protons 6 protons 12 - 6 = 6 13 - 6 = 7 14 - 6 = 8 6 neutrons 7 neutron

8 neutrons

Isotopes of Carbon

22

Example Consider a neutral atom of the

element phosphorus: Atoms of this element have how

many protons in their nucleus? How many electrons does a neutral

atom of phosphorus have? How many neutrons does an atom of

phosphorus have in its nucleus?

P3115

15

15

16

23

Example Consider a neutral atom of the

element Calcium: Atoms of this element have how

many protons in their nucleus? How many electrons does a neutral

atom of calcium have? How many neutrons does an atom of

calcium have in its nucleus?

20

20

20

Ca4020

24

Example Consider a neutral atom of the

element Calcium: Atoms of this element have how

many protons in their nucleus? How many electrons does a neutral

atom of calcium have? How many neutrons does an atom of

calcium have in its nucleus?

20

18

20

Ca2+4020

25

Atomic Weight One amu is exactly 1/12 of the mass

of a 12C atom. 12C is a specific isotope of carbon.

1 g = 6.022 x 1023 amu

26

...2211 AWxfractionAWxfractionAW

Atomic Weight The weighted average of the

masses of its constituent isotopes.

Atomic Weight ∑ fractional

abundance )[( isotopicmass )](x=

27

Increased Deflection

The natural relative abundances fordifferent isotopes can be determinedfrom the mass spectrum.

Neon

Mass Spectrometer

28

Carbon is listed as have an atomic weight of 12.01 amu in the periodic table, based on the weighted average of all carbon isotopes...

isotope abundance mass amu neutrons

12C 98.89% 12.0000 6

13C 1.11% 13.0034 7

14C < .01% n/a 8

29

The atomic weight of carbon is 12.011 amu, computed as follows...

Atomic weight of C = the sum of

(%abundance of isotope) x (its mass) for all stable isotopes. So...

as percentages (98.89%)(12 amu) + (1.11%)(13.0034 amu) =

oras fractions

(0.9889)(12 amu) + (0.111)(13.0034 amu) =

12.011 amu

30

Mass Spectrometer

Neon

Atomic weight of Ne = (90.48%)19.9924 amu +(0.27%)20.9938 amu + (9.25%)21.9914 amu

= (0.9048)*19.9924 amu +(0.0027)*20.9938 amu + (0.0925)*21.9914 amu = 20.1797 amu

31

Mg-24 78.99% 23.98504Mg-25 10.00% 24.98584Mg-26 11.01% 25.98259

24.30 amu

Atomic Weight

What is the average atomic weight of Mg?

AW = 0.7899(23.98504) + 0.1000(24.98584) +

0.1101(25.98259)

32

Atomic Weight The atomic weight of Ga is 69.72;

Ga-69 = 68.9257; Ga-71 = 70.9249

What is the abundance of each isotope?

33

Atomic WeightLet x = abundance Ga-691-x = abundance of Ga-71

x(68.9257) + (1-x)(70.9249) = 69.72

68.9257x + 70.9249 – 70.9249x = 69.72

1.9992x = 1.20

x = 0.600

Ga-69 = 60.0% and Ga-71 = 40.0%

34

Electromagnetic Spectrum

Visible light makes up only a small part of the electromagnetic spectrum.

35

Electromagnetic Spectrum Electromagnetic radiation has a

dual behavior. It has properties of a particle called

a photon and as a wave traveling at the speed of light. Characterized by a wavelength and

frequency.

36

Electromagnetic Radiation

Wavelength- The distance between two

corresponding points on a wave.

37

Electromagnetic Radiation

Frequency- The number of wave crests passing

a given point per unit time.

38

Electromagnetic Radiation

c = c = 3.00 x 108 m/s

39

Electromagnetic RadiationNote that long wavelength small frequencyShort wavelength high frequency

increasing wavelength

increasing frequency

QuickTime™ and aGraphics decompressor

are needed to see this picture.

40

Electromagnetic Radiation Given = 7.31 x 1014s-1, calculate

= 8

14 1

3.00 10 /7.31 10

x m sx s

m

nmxx

9101nm = 4.10 x 10-7 m

= 4.10 x 10-7 m

= 410 nm

41

Planck’s Equation

E = energy of 1 photonh = Planck’s constant, 6.626 x 10-34 J-s= frequency, s-1

= wavelength, mc = speed of light, 3.00 x 108 m/s

hchE

42

Quantization of Energy

E = h Light with large (small ) has a

small E.Light with a short (large ) has

a large E.

43

Quantization of Energy

E = h

44

Planck’s Equation Calculate the energy of a photon

with a wavelength of 4.10 x 10-7m.

mx

smxsJx7

834

1010.4

/1000.31063.6

E = 4.85 x 10-19 J

hc

E

45

Quantum Theory Allowed for the interpretation of

spectra of atoms, ions, and molecules.

Neils Bohr proposed the fundamental hypothesis of the quantum theory.

46

Atomic Line Spectra and Niels Bohr

Bohr’s greatest contribution to science was in building a simple model of the atom. It was based on an understanding of the SHARP LINE SPECTRA of excited atoms.Niels BohrNiels Bohr

(1885-(1885-1962)1962)

47

Line Spectra of Excited Atoms Excited atoms emit light of only

certain wavelengths The wavelengths of emitted light

depend on the element.

48

High EShort High

Low ELong Low

Line Spectra of Excited Atoms Visible lines in H atom spectrum

are called the BALMER series.

49

Bohr Model of Atom An atom has a number of definite

and discrete energy levels in which an electron can exist.

Increasing radius of orbit increases the energy.

Electrons can move from one energy level to another.

Electron moves in circular orbit.

50

The Bohr Model and Quantized Energy

++

ee--

ee--

Excited State

Electron

GroundState Energy

GroundState Energy

High EnergyHigh EnergyOrbitOrbit

Low EnergyLow EnergyOrbitOrbit

High EnergyHigh EnergyOrbitOrbit

Low EnergyLow EnergyOrbitOrbit

Energies are “quantized” in

other words the Energies are

limited to discrete values.

Energies are “quantized” in

other words the Energies are

limited to discrete values.

51

Atomic Spectra Atomic spectra tells us about the

structure of the atom.

52

Atomic Spectra

Continuous spectra from sun contain all wavelengths.

Line spectra have discrete lines from atoms.

53

Atomic Spectra

Emission The process

where an electron moves from a higher to lower energy state, resulting with the loss of energy.

54

Atomic Spectra

Absorption The process

where an electron moves from a lower to higher energy state, resulting with the gain of energy.

55

Spectral LinesSecond Electronic Excited States

First Electronic Excited States

Ground Electronic Excited States

ENERGY

56

Hydrogen Atom Only certain lines are found in the

atomic spectrum of hydrogen. Bohr: electron can only have

certain energy values. Balmer: expressed relationship

mathematically.

57

Rydberg Equation

RH = 2.18 x 10-18 J

ni = initial value of n

nf = final value of n

E Rn nH

i f

1 1

2 2

58

Rydberg Equation What wavelength of light is

associated with a transition from n = 4 to n = 2?

E x J x J

218 10

14

12

409 10182 2

19. .

Jx

smxsJx19

834

1009.4

/1000.31063.6

= 4.86 x 10-7 m = 486 nm

E

hc

59

Atomic Line Spectra and Niels Bohr

Bohr’s theory was a great accomplishment.

Rec’d Nobel Prize, 1922Problems with theory — theory only successful for H. introduced quantum idea

artificially. So, we go on to QUANTUM or

WAVE MECHANICSNiels BohrNiels Bohr

(1885-(1885-1962)1962)

60

Quantum or Wave Mechanics

de Broglie (1924) proposed that all moving objects have wave properties.

= h/mvL. de BroglieL. de Broglie(1892-1987)(1892-1987)

61

Quantum or Wave Mechanics

Baseball (115 g) at 100 mph

= 1.3 x 10-32 cm

electron with velocity =

1.9 x 108 cm/sec = 0.388 nm

Experimental proof of waveproperties of electrons

62

Wave Nature of Electron Calculate the wavelength of an

electron traveling at 1.243 x 107m/s.

hm

x J skg m s

J

x kg x m s

x m

663 101

911 10 1243 10

587 10

342 2

31 7

11

./

. . /

.

Which is similar to the spacing between atoms in crystals.

63

Quantum Mechanical Picture Werner Heisenberg - 1927

Uncertainty Principle It is impossible to determine

simultaneously both the position & momentum of an electron. electron microscopes use this

phenomenon

64

Quantum Mechanical Picture

Werner Heisenberg may have slept here!

65

Quantum Mechanical Picture

devices for detecting motion of electron disturbs its position

like measuring position of a car with a wrecking ball

66

Quantum Mechanical Picture Schrödinger proposed an equation that

contains both wave and particle terms. Solving the equation leads to wave

functions. The wave function gives the shape of the

electronic orbital. The square of the wave function, gives the

probability of finding the electron, that is, gives the electron density for the atom. WAVE FUNCTION –WAVE FUNCTION –

67

Orbitals and Quantum Numbers If we solve the Schrödinger equation,

we get wave functions and energies for the wave functions.

We call wave functions orbitals. Schrödinger’s equation requires 3

quantum numbers: Principal Quantum Number, n. This is the same as Bohr’s n. As n becomes

larger, the atom becomes larger and the electron is further from the nucleus.

68

Quantum Theory Atoms have discrete energy states. Atoms have definite energy. Changes in energy results in

absorption or emission of energy. Allowed energy states are

described by quantum numbers.

69

Quantum Numbers Principal quantum number-n main energy level n = 1, 2, 3, ...

70

Quantum Numbers Subsidiary quantum number-l shape of orbital l = 0, 1, 2, ... n-1

l = 0 s orbital l = 1 p orbital l = 2 d orbital l = 3 f orbital

71

Atomic Orbitals s orbitals

72

Atomic Orbitals p orbitals

73

Atomic Orbitals d orbitals

74

Atomic Orbitals f orbitals

75

Quantum Numbers

Magnetic quantum number-ml

spatial (xyz) orientation of orbital -l, -1, 0, 1, 2 ..., l

l = 0 s orbital ml = 0 l = 1 p orbital ml = -1, 0, 1 l = 2 d orbital ml = -2, -1, 0, 1, 2 l = 3 f orbital ml = -3, -2, -1, 0,

1, 2, 3

76

Atomic Orbitals There are three p-orbitals px, py, and pz.

(The three p-orbitals lie along the x-, y- and z- axes of a Cartesian system. The letters correspond to allowed values of ml of -1, 0, and +1.)

The orbitals are dumbbell shaped. As n increases, the p-orbitals get larger. All p-orbitals have a node at the

nucleus.

77

The p Orbitals

78

The d and f Orbitals There are 5 d- and 7 f-orbitals. Three of the d-orbitals lie in a plane

bisecting the x-, y- and z-axes. Two of the d-orbitals lie in a plane

aligned along the x-, y- and z-axes. Four of the d-orbitals have four lobes

each. One d-orbital has two lobes and a collar.

222 zy-xxzyzxy d ,d ,d ,d ,d

79

The d Orbitals

80

Atomic Orbitals f orbitals

start with n = 4 most complex shaped orbitals 7 per n level, complicated names

l = 3 ml = -3,-2,-1,0,+1,+2, +3 7 values of

ml important effects in lanthanides & actinides

81

Quantum Numbers Spin quantum number-ms

+ ½ or - ½ Wolfgang Pauli - 1925

exclusion principle no two electrons in an atom can have

the same set of 4 quantum numbers.

82

Pauli Exclusion Principle No two electrons in

an atom can have the same set of four quantum numbers.

No two electrons with the same spin in an atom can fill one orbital.

83

Quantum Numbers Write quantum numbers for each

electron in nitrogen.

n l ml ms

1 1 0 0 ½2 1 0 0 -½3 2 0 0 ½4 2 0 0 -½5 2 1 -1 ½6 2 1 0 ½7 2 1 1 ½

84

Quantum Numbers What are the values for n and l for:

n l1s4d3p4f

1

4

3

4

02

1

3

85

Aufbau Principle The electron that distinguishes an

element from the previous element enters the lowest-energy atomic orbital available

86

Energy Levels The order of

filling is determined by the energy of each orbital.

87

IA VIIIA1 2

H He1.008 IIA IIIA IVA VA VIA VIIA 4.0033 4 5 6 7 8 9 10

Li Be B C N O F Ne6.941 9.012 10.81 12.01 14.01 15.99 19 20.1811 12 13 14 15 16 17 18

Na Mg Al Si P S Cl Ar22.99 24.31 IIIB IVB VB VIB VIIB VIIIB IB IIB 26.98 28.09 30.97 32.07 35.45 39.9419 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36

K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr39.1 40.08 44.96 47.88 50.94 52 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.9 83.8

37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

85.47 87.62 88.91 91.22 92.91 95.94 -98 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.355 56 57 72 73 74 75 76 77 78 79 80 81 82 83 85 86

Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn132.9 137.3 138.9 178.5 180.9 183.9 186.2 190.2 192.2 195.1 197 200.6 204.4 207.2 209 (209) (210) (222)87 88 89

Fr Ra Ac(223) 226 227

58 59 60 61 62 63 64 65 66 67 68 69 70 71Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu

140.1 140.9 144.2 (145) 150.4 152 157.3 158.9 162.5 164.9 167.3 168.9 173 17590 91 92 93 94 95 96 97 98 99 100 101 102 103

Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr232 231 238 237 (244) (243) (247) (247) (251) (252) (257) (258) (259) (260)

s electrons p electrons

d electrons

f electrons

88

Order of Filling

89

Examples-Orbital Diagram

1s 2s

H æ

He å

Li å æ

Be å å

90

Electronic Configuration1s 2s 2p Configuration

Li å æ 1s22s1

Be å å 1s22s2

B å å æçç 1s22s22p1

HUND’S RULE. When placing electrons in a set of orbitals having the same energy, we place them singly as long as possible.

C å å ææç 1s22s22p2

91

Electronic Configuration1s 2s 2p Configuration

Li å æ 1s22s1

Be å å 1s22s2

B å å æçç 1s22s22p1

C å å ææç 1s22s22p2

N å å æææ 1s22s22p3

O å å åææ 1s22s22p4

F å å ååæ 1s22s22p5

Ne å å ååå 1s22s22p6

92

Electronic Configuration1s 2s 2p 3s 3p Configuration

Na å å å å å æ 1s22s22p63s1

Mg å å å å å å 1s22s22p63s2

Al å å å å å å æçç 1s22s22p63s23p1

Si å å å å å å ææç 1s22s22p63s23p2

P å å å å å å æææ 1s22s22p63s23p3

S å å å å å å åææ 1s22s22p63s23p4

Cl å å å å å å ååæ 1s22s22p63s23p5

Ar å å å å å å ååå 1s22s22p63s23p6

93

Electronic Configuration

4s 3d 4p

Configuration

K [Ar] æ [Ar] 4s1

Ca [Ar] å [Ar] 4s2

Sc [Ar] å æçççç [Ar] 4s23d1

Ti [Ar] å ææççç [Ar] 4s23d2

V [Ar] å æææçç [Ar] 4s23d3

Cr [Ar] æ æææææ [Ar] 4s23d4

94

Electronic Configuration Exceptions

[Ar]4s2 3d4 [Ar]4s13d5

95

Electronic Configuration

4s 3d 4p

Mn [Ar] å æææææ [Ar]

4s23d5

Fe [Ar] å åææææ [Ar] 4s23d6

Co [Ar] å ååæææ [Ar] 4s23d7

Ni [Ar] å åååææ [Ar] 4s23d8

Cu [Ar] æ ååååå [Ar] 4s13d10

Zn [Ar] å ååååå [Ar] 4s23d10

96

Electronic Configuration

4s 3d 4p

Ga [Ar] å ååååå æçç [Ar] 4s23d104p1

Ge [Ar] å ååååå ææç [Ar] 4s23d104p2

As [Ar] å ååååå æææ [Ar] 4s23d104p3

Se [Ar] å ååååå åææ [Ar] 4s23d104p4

Br [Ar] å ååååå ååæ [Ar] 4s23d104p5

Kr [Ar] å ååååå ååå [Ar] 4s23d104p6

97

Electronic ConfigurationSodiumSodium1s1s2 2 2s2s2 2 2p2p6 6 3s3s11 1s1s2 2 2s2s2 2 2p2p66

NaNa ŽNa+ + e

ChlorineChlorine1s1s2 2 2s2s2 2 2p2p6 6 3s3s2 2 3p3p5 5 1s1s2 2 2s2s2 2 2p2p6 6

3s3s2 2 3p3p66

ClCl + e Ž Cl-

98

Electronic Configuration of Inert Gases

The noble gases are chemically stable as individual atoms and have a full complement of outer groups s and p electrons. 2He = 1s2

10Ne = 1s22s22p6

18Ar = 1s22s22p63s23p6

36Kr = 1s22s22p63s23p64s23d104p6

54Xe = 1s22s22p63s23p64s23d104p65s24d105p6

99

Electronic Configuration of Inert Gases Because they have complete

shells, they neither gain nor lose electrons easily; as a result, they do not form compounds readily.

They don’t form diatomic molecules with each other like H2, N2, O2, F2, Cl2, Br2, I2, and At2

100

Dalton’s Model of Atom All matter is composed of tiny,

indivisible particles called atoms Atoms of each element are alike Atoms of different elements have

different masses Atoms of different elements can

join to form compounds

101

Thomson’s Model of Atom Atoms are not solid spheres, they

contain particles Particles are negatively charged

called electrons

102

Rutherford’s Model of Atom Protons are concentrated in a

small area at the center of an atom

103

Bohr’s Model of Atom Electrons have fixed amount of energy,

which keeps the electron moving around nucleus

Area that the electron moves in is called an energy level

Each energy level is further from nucleus

Electrons can move from one level to another, but can’t be between levels

104

Electron Cloud Model of Atom Electrons do not orbit the nucleus Move in changing paths Most of the path falls within a

region called the electron cloud High probability that the electron

exists in electron cloud