Chpt 5MA12.pptx

CHAPTER 5

Although three subatomic particles

had been discovered by the early

1900s, the quest to understand the

atom and its structure had just

begun.

Rutherford proposed

that all of an atom’s

positive charge and

Although his nuclear model was a major scientific development, it lacked detail about how electrons occupy the space surrounding the nucleus. ? ? ? ? ? ? ? ? ? ? ?

Because of this, scientists in the

early twentieth century found

Nor did it address the question of

why the negatively charged

electrons are not pulled into the

atom’s positively charged nucleus.

-Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Energy Levels in Atoms

• It explained only a few simple

properties of atoms.

• It could not explain the chemical

properties of elements.

For example, Rutherford’s

model could not explain why an object such as the iron

scroll shown here first glows dull red, then yellow, and then white when heated to higher and higher temperatures.

Chemists found Rutherford’s

nuclear model lacking because it

did not begin to account for the

differences in chemical behavior

among the various elements.

In the early 1900s,

scientists began to

unravel the puzzle of

chemical behavior.

They had observed that certain

elements emitted visible light when

heated in a flame.

Analysis of the emitted light

revealed that an element’s chemical

behavior is related to the arrangement

of the electrons in its atoms.

ELECTRONS!

When

excited

electrons

drop to lower

energy levels

they release

light!

Behavior of Light

 In order for you to better understand this relationship and the nature of atomic

structure, it will be helpful for you to first understand the nature of light.

 Light is two different things at once:

1. Particles (packets) called photons

2. Waves of energy

 Light actually behaves as both a wave

Electromagnetic Radiation

(Light)

The light that we can see with our eyes is visible light.

Electromagnetic radiation – A series of

electromagnetic waves that travel in a vacuum at a

speed of light. There are many other types of radiation including:

 Radio waves

 Microwave

 Infrared waves

 Visible

 Ultra Violet (UV)

 X-rays

 Gamma rays

 Red Martians Invented Very Useful X-ray Guns.

ROY G. BIV

 ROY G. BIV is an acronym that helps us

remember the order of the visible light spectrum

 Red, orange, yellow, green, blue, indigo, violet.  As you approach violet, the frequency (and

Light and Quantized Energy

 By the year 1900, there was enough experimental evidence to convince

scientists that light consisted of waves.  What are waves?

 Mechanical waves – require a medium to travel

(air, water, or rope).

Light and Quantized Energy

 Wave Properties

 Wavelength (λ or lambda) – distance between

the crests

 1nm = 1 × 10-9 m

 Amplitude – height from origin to crest,

Common wavelength units for electromagnetic radiation

Picometer pm 10-12 _{Gamma ray}

Ångstrom Å 10-10 _X-ray

Nanometer nm 10-9 _X-ray

Micrometer mm 10-6 Infrared

Millimeter mm 10-3 _Infrared

Centimeter cm 10-2 Microwave

Meter m 100 _Radio

Unit Symbol Wavelength, (m) Type of Radiation

Light and Quantized Energy

 Wave Properties

 Frequency ( or nu) – the number of wave

cycles that pass a given point per unit of time.

 The SI unit of cycles per second is called

hertz.

 1 hertz (Hz) = 1 wave per second

 1MHz = 1  106 Hz

 562 Hz = 562 waves/second = 562 = 562s-1

Light and Quantized Energy

 Wave Nature of light

 All electromagnetic light moves at the

speed of 3.00 × 108 m/s and is

represented by the symbol, c.

 The speed of light is the product of the

wavelength (λ) and frequency ().

c

=

l

8 m/s)

 = frequency (Hz = 1/s = s-1)

Light and Quantized Energy

•

Although the speed of electromagnetic

waves is constant, the frequency and the

wavelength may vary.

•

_{As you can see from the equation,}

Light and Quantized Energy

 What is the wavelength of a

microwave having a frequency of 3.44

x 109 Hz?

= 8.72 × 10

m

l

c





l



c

1 -9

8

s

10

3.44

m/s

10

3.00







Light and Quantized Energy

 A helium-neon laser emits light with a

wavelength of 633 nm. What is the frequency of this light? Remember: 1nm=1x10-9m.

c = 3.00  108 m/s

λ = 633 nm

l

c







_l

c

m

10

6.33

s

m

10

3.00













4.74



10

Hz

Light and Quantized Energy

 Particle Nature of Light

 While considering light as

a wave does explain much of its everyday behavior, it fails to adequately

describe important aspects of light’s

interactions with matter.

 Glowing substances

 Photoelectric effect

Light and Quantized Energy



Particle Nature of Light



The wave model of light cannot

explain why heated objects emit

only certain frequencies of light at

a given temperature, or why some

metals emit electrons when

Light and Quantized Energy

 Particle Nature of Light

 In 1900, the German physicist

Max Planck began searching for an explanation as he studied

the light emitted from heated objects.

 matter can gain or lose energy

only in small, specific amounts called quanta.

 Quantum – minimum amount

of energy that can be gained or lost by an atom.

1858–1947

Continuous vs. Quantized

Energy

E

ne

rg

y

A B

Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 330

Light and Quantized Energy



Particle Nature of Light



Planck found that the energy of a

quantum of energy (photon) is

directly proportional to the

frequency.

E = h



h = Planck’s constant (6.626 x10-34 Js)

Light and Quantized Energy

 What is the energy of a photon from

the violet portion of the rainbow if it

has a frequency of 7.23 x 1014 Hz?

E = (6.626 × 10

J·s)(7.23 × 10

s

)

E = 4.79 × 10

J



h

Seeing the Light

 To calculate the energy of

electromagnetic radiation:

E

=

hc

l

E = energy of the electromagnetic radiation

h = 6.626 x 10-34 J . s

c = 3.00 x 108 m/s (speed of light)

Photon Energy Problem 2

What is the energy of blue light that has a wavelength of 450.0 nm? 1 nm = 1  10-9 m

= 4.500  10-7 m

E

=

hc

l

E = (6.626x10-34 Js)(3.00x108

m/s)

4.500x10-7 m

= 4.42  10-19 J

Light and Quantized Energy

 Particle Nature of Light

 Photoelectric effect – electrons, called

photoelectrons, will be emitted from a

Light and Quantized Energy

 Einstein's explanation treated light like

Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

The Quantum Concept and

Photons

No electrons are

ejected because the frequency of the light is below the threshold frequency.

If the light is at or above the threshold frequency, electrons are ejected.

If the frequency is

increased, the ejected electrons will travel faster.

Quantum Theory

•

Einstein

– Concluded - light has properties of both

waves and particles

“

wave-particle duality

”

– Photon - particle of light that carries a

quantum of energy

Light and Quantized Energy

 Particle Nature of Light

 Atomic emission spectrum –

Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

A prism separates light into the colors it contains. White light produces a rainbow of colors.

Light and Atomic Emission

Spectra

Light bulb

Slit Prism

Screen

Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Light from a helium lamp produces discrete lines.

Light and Atomic Emission

Spectra

Slit Prism

Screen

Helium lamp

• _{The energy absorbed by an electron for it to}

move from its current energy level to a higher energy level is identical to the energy of the light emitted by the electron as it drops back to its original energy level.

• _{The wavelengths of the spectral lines are}

characteristic of the element, and they make up the atomic emission spectrum of the element.

• _{No two elements have the same emission}

spectrum.

Light and Atomic Emission

Spectra

Electromagnetic Radiation

Light as a wave

Light as a stream of energy (packets of photons)

Continuous and Line

Spectra

4000 Ao 5000 6000 7000

light

Na

H

Ca

Hg

400 450 500 550 600 650 700 750 nm Visible

spectrum

Quantum Theory and the

Atom

 Energy levels – electrons orbit in circles around the nucleus

at fixed energy amounts (quantized).

 Ground State – an atoms electrons are at the lowest energy levels

 The higher the energy level the farther it is from the nucleus.

Quantum

Quantum Theory and the

Atom

 Building on Planck’s and Einstein’s

concepts of quantized energy (quantized means that only certain values are

allowed), Bohr proposed that the

hydrogen atom has only certain allowable energy states.

Bohr’s Experiment

•

When an atom gains energy, it is

said to be in an

excited state

.

•

Although a hydrogen atom contains

only a single electron, it is capable

of having many different excited

states.

Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

The rungs on this ladder are

somewhat

like the energy levels in Bohr’s

model

of the atom.

stand between the rungs. Similarly, electrons can

only be at

specific energy levels, NOT between levels.

The Bohr Model

e

-e- Ground state

Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

The rungs on this ladder are

somewhat like the energy

levels in Bohr’s model of the

atom.

The Bohr Model

• The energy levels in

atoms are unequally spaced, like the

rungs in this unusual ladder. The higher

An Excited Lithium Atom

Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 326

Photon of red light emitted

Li atom in

lower energy state Excited Li atom

E

n

e

rg

Color = Energy of Photons

Bohr model of the atom

Bohr was able to use his model hydrogen to:

• _{Account for the observed spectral lines.} • Calculate the radius for hydrogen atoms.

His model did not account for:

• _{Atoms other than hydrogen.} • Why energy was quantized.

Quantum Theory and the Atom

 In 1924, a young French

graduate student in physics named Louis de Broglie

proposed an idea that

eventually accounted for the fixed energy levels of Bohr’s model.

 If waves could be treated like a particle, could particles be treated like waves?

 Electron behavior might be explained if we treat

electrons, a particle, as a wave.

 De Broglie knew that if

an electron has wavelike motion and is restricted to circular orbits of fixed radius, the electron is allowed only certain possible wavelengths, frequencies, and

energies. In other words, it would be quantized just like observed.

=

=

Quantum Theory and the Atom

 Developing his idea, de Broglie derived an equation for the wavelength (λ) of a particle of mass (m) moving at velocity (ν).

Does it work?

Experiments show that the smaller the particle, the more it acts like a wave!

v

h

m



Quantum Theory and the

Atom

 Step by step, scientists such as

Rutherford, Bohr, and de Broglie had been unraveling the mysteries of the atom.

• However, a conclusion

reached by the German

theoretical physicist Werner Heisenberg a contemporary of de Broglie, proved to

have profound implications for atomic models.

Quantum Theory and the

Atom

 Heisenberg Uncertainty Principle

 You can’t precisely know both the position

and velocity of a particle at the same time.

No, but I know where I’m at! Do you

know how fast you

were going?

Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

• _{To locate an electron, you might strike it with a}

photon.

• The electron has such a small mass that striking

it with a photon affects its motion in a way that cannot be predicted accurately.

• _{The very act of measuring the position of the}

electron changes its velocity, making its velocity uncertain.

Quantum Mechanics

Before

collision: A photon strikes an electron during an attempt to observe the electron’s position. After

collision: The impact changes the electron’s velocity,

Quantum Theory and the

Atom

 In 1926, Austrian

physicist Erwin

Schrödinger furthered the wave-particle theory

proposed by de Broglie.

•

Schrödinger derived an equation

that described the behavior of the

electron in a hydrogen atom as a

wave.

Quantum Theory and the

Atom

 Remarkably, unlike Bohr’s model,

Schrödinger’s new model for the hydrogen atom seemed to apply

equally well to atoms of all elements!

 The modern description of the electrons in atoms, the quantum mechanical

model, came from the mathematical

Quantum Theory

and the Atom

 In the quantum mechanical model, electron

position and energy are described using energy levels, energy sublevels, orbital shapes, and spin.

 The probability of finding an electron within a

space surrounding the nucleus can be

represented as a fuzzy cloud like region. The cloud is more dense where the probability of finding the electron is high.

nucleus

electron cloud

Models of the Atom

Dalton’s model (1803)

Thomson’s plum-pudding model (1897)

Rutherford’s model (1909)

Bohr’s model (1913)

Quantum mechanical model (present)

Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd _{Edition, 1990, page 125}

Greek model (400 B.C.) + -e e e + + ₊ + + + + + e e e e e e e

Quantum Theory and the Atom

 Principal Energy Level (n)

 Describes distance from the nucleus and

general energy.

 n = 1, 2, 3, 4, ….

 The higher the energy level the greater the

average distance from the nucleus.

 Each energy level contains sublevels

 The number of sublevels on a level is equal to

the energy level (n).

 1st energy level has 1 sublevel (1s)

 2nd energy level has 2 sublevels (2s, 2p)

 3rd energy level has 3 sublevels (3s, 3p, 3d)

Quantum Theory and the Atom

 Each sublevel contains orbitals.

 orbital: a three-dimensional region around the

nucleus in which an electron moves and is found 90% of the time.

 Each orbital can hold up to two electrons.

 The total number of orbitals on a level = n2.  Each sublevel has a different shape of orbital

on the level.

 These shapes are represented by the symbols s, p, d,

Quantum Theory and the Atom

 s Orbitals

 Each level has one s shaped (spherical)

sublevel

 Only 1 orientation on sublevel

 An s sublevel can hold 2 electrons

Quantum Theory and the Atom

 p Orbitals

 2nd energy level and above have a p

sublevel

 3 orientations on each sublevel

 p Sublevel can hold up to 6 electrons

2p_x 2p

Quantum Theory and the Atom

 d orbitals

 3rd energy level and above have a d

sublevel

 5 orientations on each sublevel

 d sublevel can hold up to 10 electrons

) (x2 y2

Quantum Theory and the Atom

 f Orbitals

 4th energy level and above have a f

sublevel

 7 orientations on each sublevel

 f sublevel can hold up to 14 electrons

) 3 (x2 y2 x

Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Atomic Orbitals

Summary of Principal Energy Levels and Sublevels

Princip al energy level Numbe r of sublev els

Type of sublevel

Maximu m number of electro ns

n = 1 1 1s (1 orbital) 2

n = 2 2 2s (1 orbital), 2p (3 orbitals) 8

n = 3 3 3_orbitals)s (1 orbital), 3p (3 orbitals), 3d (5 18

n = 4 4 4orbitals),s (1 orbital), 4p (3 orbitals), 4d (5 4f (7 orbitals)

32

Electron Configurations General

Rules

Rule 1:

Aufbau Principle

– Electrons occupy

the orbitals of the lowest energy

orbitals first.

– “Lazy Tenant Rule”

Aufbau Chart

 1s

 2s 2p

 3s 3p 3d

 4s 4p 4d 4f

 5s 5p 5d 5f

 6s 6p 6d 6f

 7s 7p 7d 7f

1s

2s

2p

3s

3p

3d

4s

4p

4d

4f

5s

5p

5d

5f

6s

6p

6d

6f

s

1

2

3

4

5

6

7

p

2

3

4

5

6

d

3

4

5

6

f

4

5

1.0079 _{2A 3A 4A 5A 6A 7A}

2 He Helium 4.0026 3 Li Lithium 6.941 4 Be Beryllium 9.0122 5 B Boron 10.81 6 C Carbon 12.011 7 N Nitrogen 14.007 8 O Oxygen 15.999 9 F Fluorine 18.998 10 Ne Neon 20.179 11 Na Sodium 22.990 12 Mg Magnesium

24.305 _{3B 4B 5B 6B 7B}

8B

1B 2B 13 Al Aluminum 26.982 14 Si Silicon 28.086 15 P Phosphorus 30.974 16 S Sulfur 32.06 17 Cl Chlorine 35.453 18 Ar Argon 39.948 19 K Potassium 39.098 20 Ca Calcium 40.08 21 Sc Scandium 44.956 22 Ti Titanium 47.90 23 V Vanadium 50.941 24 Cr Chromium 51.996 25 Mn Manganese 54.938 26 Fe Iron 55.847 27 Co Cobalt 58.933 28 Ni Nickel 58.71 29 Cu Copper 63.546 30 Zn Zinc 65.38 31 Ga Gallium 69.72 32 Ge Germanium 72.59 33 As Arsenic 74.922 34 Se Selenium 78.96 35 Br Bromine 79.904 36 Kr Krypton 83.80 37 Rb Rubidium 85.468 38 Sr Strontium 87.62 39 Y Yttrium 88.906 40 Zr Zirconium 91.22 41 Nb Niobium 92.906 42 Mo Molybdenum 95.94 43 Technetium (97) 44 Ru Ruthenium 101.07 45 Rh Rhodium 102.91 46 Pd Palladium 106.4 47 Ag Silver 107.87 48 Cd Cadmium 112.41 49 In Indium 114.82 50 Sn Tin 118.69 51 Sb Antimony 121.75 52 Te Tellurium 127.60 53 I Iodine 126.90 54 Xe Xenon 131.30 55 Cs Cesium 132.91 56 Ba Barium 137.33 71 Lu Lutetium 174.97 72 Hf Hafnium 178.49 73 Ta Tantalum 180.95 74 W Tungsten 183.85 75 Re Rhenium 186.21 76 Os Osmium 190.2 77 Ir Iridium 192.22 78 Pt Platinum 195.09 79 Au Gold 196.97 80 Hg Mercury 200.59 81 Tl Thallium 204.37 82 Pb Lead 207.2 83 Bi Bismuth 208.98 84 Po Polonium (209) 85 At Astatine (210) 86 Rn Radon (222) 87 Fr Francium (223) 88 Ra Radium (226) 103 Lawrencium (260) 104 Rutherforium (261) 105 Dubnium (262) 106 Seaborgium (263) 107 Bohrium (262) 108 Hassium (265) 109 Meitnerium (266) 110 (269) 111 (272) 112 (277) 11 Na Sodium 22.990 57 La Lanthanum 138.91 58 Ce Cerium 140.12 59 Pr Praseodymium 140.91 60 Nd Neodymium 144.24 61 Pm Promethium (145) 62 Sm Samarium 150.4 63 Eu Europium 151.96 64 Gd Gadolinium 157.25 65 Tb Terbium 158.93 66 Dy Dysprosium 162.50 67 Ho Holmium 164.93 68 Er Erbium 167.26 69 Tm Thulium 168.93 70 Yb Ytterbium 173.04 89 Ac Actinium (227) 90 Th Thorium 232.04 91 Pa Protactinium 231.04 92 U Uranium 238.03 93 Neptunium 237.05 94 Plutonium (244) 95 Americium (243) 96 Curium (247) 97 Berkelium (247) 98 Californium (251) 99 Einsteinium (254) 100 Fermium (257) 101 Menelevium (258) 102 Nobelium (259) Atomic Number Element Name Average Atomic Mass

Element Symbol

*Outlined symbols ( ) are not found in nature.

Electron Configurations

2.

Pauli exclusion principle – an

atomic orbtial may describe at

most two electrons, each with

opposite spin direction.

 No 2 electrons in an atom can have the

same four quantum numbers (level, sublevel shape, orientation, and spin).

Carbon ↑↓ ↑↓ ↑ ↑ __

1s 2s 2p Same level,

sublevel, and spin,

but different orientation Same level,

General Rules

•

_{Pauli Exclusion Principle}

– _{Each orbital can hold TWO electrons with}

opposite spins.

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

RIGHT

WRONG

3. Hund’s Rule

–Electrons occupy orbitals of

the same energy in a way that makes the

number of electrons with the same spin

direction as large as possible.

When filing a sublevel with multiple orbitals

(p, d, or f), each orbital must have one

electron before any orbital has a second

electron.

– _{Within a sublevel, place one electron per orbital}

before pairing them.

– “Empty Bus Seat Rule”

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Electron Configurations

 Orbital Notation

 shows every occupied orbital in every

sublevel with electrons. Arrows or slashes are used to show the electrons and the

direction of spin (↑ or ↓) or (/ or \)

Electron Configurations

 Since:

 s orbitals can hold 2 electrons, draw 1

square.

 p orbitals can hold 6 electrons, draw 3

squares.

 d orbitals can hold 10 electrons, draw 5

squares.

 f orbitals can hold 14 electrons, draw 7

Electron Configurations

 Electron Configuration Notation

 each sublevel with electrons is described

with the number of electrons in the sublevel as a superscript.

Carbon:

O 8e

-

_{Electron Configuration}



_{Orbital Notation}

1s

2

2s

2

2p

4

Notation

1s

2s

2p

O

15.9994

H : 1s1

1s

He : 1s2

1s

Li : 1s2 _2s1

1s 2s

Be : 1s2 2s2

1s 2s

C : 1s2 2s2 2p2

1s 2s 2p

S : 1s2 2s2 2p6 3s2

3p4

1s 2s 2p 3s 3p

Fe : 1s2 2s22p63s23p64s23d6Iron has ___

electrons.

26

Electron Configurations

 Noble Gas or Shorthand Notation

 Electron configuration except the inner

neon's e- configuration (1s22s22p6)

Shorthand Configuration

[Ne] 3

s

third energy level

one electron in the s orbital

orbital shape

[1

s

2

s

2

p

] 3

s

A

B

C

D

1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p2

Xe

Pb: [Xe]6s24f145d106p2

Na

Shorthand Configuration

[Ar] 4s2

Electron configuration Element symbol

[Ar] 4s2 3d3

[Rn] 7s2 5f14 6d4

[He] 2s2 2p5

Exception [Kr] 5s1 4d10

[Kr] 5s2 4d10 5p5

[Kr] 5s2 4d10 5p6

Ca

V

Sg

F

Ag I

Xe

Fe [Ar] 4s23d6

Exceptions to the predicted

electron configurations

Two elements of the first 40 elements have

electron configurations different from what would be normally predicted.

 Predicted:  Cr:  [Ar] 4s23d4    

 Actual:       Cr:  [Ar] 4s13d5    

 Chromium gains stability with a half-full d-sublevel.

 Applies to Cr and Mo.

4s 3d

Exceptions to the predicted

electron configurations

 Predicted: Cu: [Ar] 4s23d9

 Actual: Cu: [Ar] 4s13d10

 Copper gains stability with a full d-sublevel.

 Applies to all atoms in the same column as copper.

4s 3d

Electron Configurations

 Electron Dot Notation

 The element symbol represents the inner level

electrons and dots are used to show the valence(outside) electrons.

 Valence electrons are the electrons that are usually

involved in reactions

 The total number of valence electrons=outer s and p

electrons

 Space out electrons with no more than 2 to a side

Electron Dot Diagrams

1A 2A 3A 4A 5A 6A 7A 8A

= valence electron

s1 _s2 s2p1 s2p2 s2p3 s2p4 s2p5 s2p6

1 2 13 14 15 16 17 18

Stable Electron

Configurations

 All atoms react to achieve noble gas configuration.

 Noble gases have two s and six p electrons.  Eight valence electrons .

 Also called the octet rule.

Electron Configurations for

Cations

 Metals lose electrons to attain noble gas

configuration.

 They make positive ions.

 If we look at electron configuration it

makes sense.

 _{Na [Ne]3s}1 _{- 1 valence electron}

Ca

2 +

Electron Dots For Cations

 Metals will have few valence electrons

 These will come off

 Forming positive ions 40.078

Ca

20

Electron Configurations for

Anions

 Nonmetals gain electrons to attain noble

gas configuration.

 They make negative ions.

 If we look at electron configuration it

makes sense.

 S [Ne]3s23p4 - 6 valence electrons

 S2- [Ne]3s23p6 -noble gas configuration.

Electron Dots For Anions

 Nonmetals will have many valence

electrons.

 They will gain electrons to fill outer

shell.