1 Chapter 11 Electrons in Atoms Killarney School

1

Chapter 11Electrons in Atoms

Killarney School

2

Section 11.1Models of the Atom

OBJECTIVES: Summarize the development of

atomic theory.

3

Section 11.1Models of the Atom

OBJECTIVES: Explain the significance of

quantized energies of electrons as they relate to the quantum mechanical model of the atom.

4

Greek Idea Democritus and

Leucippus Matter is made up

of solid indivisible particles

John Dalton - one type of atom for each element

5

J. J. Thomson’s Model

Discovered electrons Atoms were made of

positive stuff Negative electron

floating around “Plum-Pudding”

model

6

Ernest Rutherford’s Model Discovered dense

positive piece at the center of the atom- nucleus

Electrons would surround it

Mostly empty space

“Nuclear model”

7

Niels Bohr’s Model He had a question: Why don’t the

electrons fall into the nucleus? Move like planets around the sun. In circular orbits at different levels. Amounts of energy separate one

level from another. “Planetary model”

8

Bohr’s planetary model Energy level of an electron analogous to the rungs of a ladder electron cannot exist between energy

levels, just like you can’t stand between rungs on ladder

Quantum of energy required to move to the next highest level

9

The Quantum Mechanical Model

Energy is quantized. It comes in chunks. A quanta is the amount of energy needed to

move from one energy level to another. Since the energy of an atom is never “in

between” there must be a quantum leap in energy.

Erwin Schrodinger derived an equation that described the energy and position of the electrons in an atom

10

Things that are very small behave differently from things big enough to see.

The quantum mechanical model is a mathematical solution

It is not like anything you can see.


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Has energy levels for electrons.

Orbits are not circular. It can only tell us the

probability of finding an electron a certain distance from the nucleus.


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The atom is found inside a blurry “electron cloud”

A area where there is a chance of finding an electron.

Draw a line at 90 % Think of fan blades


13

Atomic Orbitals Principal Quantum Number (n) = the

energy level of the electron.(1-7) Sublevels- like theater seats arranged

in sections(labeled s,p,d,f) Within each sublevel, the complex math

of Schrodinger’s equation describes several shapes.

These are called atomic orbitals - regions where there is a high probability of finding an electron.

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For any atom there is only one 1s orbital. The "1" represents the fact that the orbital is in theenergy level closest to thenucleus. The "s" refers to the shape of the orbital. S orbitals are spherically symmetric around the nucleus.

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1 s orbital

2 s and 3 s orbitals just larger

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2p and 3p orbitalsAt the first energy level, the only orbital available to electrons is the 1s orbital, but at the second and third levels, as well as the 2s and 3s orbitals, there are also the 2p and 3p orbitals.Unlike an s orbital a p orbital points in a particular direction. At any one energy level it is possible to have three absolutely equivalent p orbitals pointing mutually at right angles to each other. These are arbitrarily given the symbols px, py and pz

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2 px orbital “dumb bell shape”

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19

20

3 d orbital

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Videos of Atomic Orbitals

http://www.humboldt.edu/~chem_dpt/resources/C109_AOSup_1-6.htm

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s2 Principle energy level 1

Principle energy level 2






s2 p6

s2 p6

s2 p6 d10

s2 p6 d10

s2 p6 d10

f14

s2

f14

Increasing energy

p6

d10

23

24

s

p

d

f

# of shapes

Max electrons

Starts at energy level

1 2 1

3 6 2

5 10 3

7 14 4

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Orbital Shapes S - spherical P - dumb bell

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By Energy Level First Energy Level only s sublevel only 2 electrons 1s2

Second Energy Level

s and p sublevels are available

2 e- in s, 6 in p 2s22p6

8 total electrons

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By Energy Level Third energy level s, p, and d

sublevels 2 e- in s, 6 in p,

and 10 in d 3s23p63d10

18 total electrons

Fourth energy level s,p,d, and f

sublevels 2 e- in s, 6 in p, 10

in d, and 14 in f 4s24p64d104f14

32 total electrons

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By Energy Level Any more than

the fourth and not all the orbitals will fill up.

You simply run out of electrons

The orbitals do not fill up in a neat order.

The energy levels overlap

Lowest energy fill first.

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Section 11.2Electron Arrangement in Atoms

OBJECTIVES: Apply the aufbau principle, the

Pauli exclusion principle, and Hund’s rule in writing the electron configurations of elements.

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Section 11.2Electron Arrangement in Atoms

OBJECTIVES: Explain why the electron

configurations for some elements differ from those assigned using the aufbau principle.

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Incr

easi

ng e

nerg

y

1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

Aufbau diagram - page 251

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Electron Configurations The way electrons are arranged in

atoms. Aufbau principle- electrons enter the

lowest energy first. This causes difficulties because of the

overlap of orbitals of different energies. Pauli Exclusion Principle- at most 2

electrons per orbital - different spins

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Electron Configuration Hund’s Rule- When electrons

occupy orbitals of equal energy they don’t pair up until they have to.

Let’s determine the electron configuration for Phosphorus

Need to account for 15 electrons

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The first two electrons go into the 1s orbital

Notice the opposite spins

only 13 more to go...

Incr

easi

ng e

nerg

y

1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

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The next electrons go into the 2s orbital

only 11 more...Incr

easi

ng e

nerg

y

1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

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• The next electrons go into the 2p orbital

• only 5 more...Incr

easi

ng e

nerg

y

1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

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• The next electrons go into the 3s orbital

• only 3 more...Incr

easi

ng e

nerg

y

1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

38

Incr

easi

ng e

nerg

y

1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

• The last three electrons go into the 3p orbitals.

• They each go into separate shapes

• 3 unpaired electrons

• = 1s22s22p63s23p3

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The easy way to remember

1s2s 2p3s 3p 3d4s 4p 4d 4f

5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f

• 1s2

• 2 electrons

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Fill from the bottom up following the arrows

1s2s 2p3s 3p 3d4s 4p 4d 4f

5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f

• 1s2 2s2

• 4 electrons

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1s2s 2p3s 3p 3d4s 4p 4d 4f

5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f

• 1s2 2s2 2p6 3s2

• 12 electrons

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1s2s 2p3s 3p 3d4s 4p 4d 4f

5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f

• 1s2 2s2 2p6 3s2

3p6 4s2

• 20 electrons

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1s2s 2p3s 3p 3d4s 4p 4d 4f

5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f

• 1s2 2s2 2p6 3s2

3p6 4s2 3d10 4p6

5s2

• 38 electrons

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1s2s 2p3s 3p 3d4s 4p 4d 4f

5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f

• 1s2 2s2 2p6 3s2

3p6 4s2 3d10 4p6

5s2 4d10 5p6 6s2

• 56 electrons

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1s2s 2p3s 3p 3d4s 4p 4d 4f

5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f

• 1s2 2s2 2p6 3s2

3p6 4s2 3d10 4p6

5s2 4d10 5p6 6s2

4f14 5d10 6p6 7s2

• 88 electrons

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1s2s 2p3s 3p 3d4s 4p 4d 4f

5s 5p 5d 5f6s 6p 6d 6f7s 7p 7d 7f

• 1s2 2s2 2p6 3s2

3p6 4s2 3d10 4p6

5s2 4d10 5p6 6s2

4f14 5d10 6p6 7s2

5f14 6d10 7p6 • 108 electrons

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Exceptional Electron Configurations

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Orbitals fill in order Lowest energy to higher energy. Adding electrons can change the

energy of the orbital. Half filled orbitals have a lower

energy. Makes them more stable. Changes the filling order

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Write these electron configurations

Titanium - 22 electrons 1s22s22p63s23p64s23d2

Vanadium - 23 electrons 1s22s22p63s23p64s23d3

Chromium - 24 electrons 1s22s22p63s23p64s23d4 expected But this is wrong!!

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Chromium is actually: 1s22s22p63s23p64s13d5

Why? This gives us two half filled orbitals. Slightly lower in energy. The same principal applies to

copper.

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Copper’s electron configuration

Copper has 29 electrons so we expect: 1s22s22p63s23p64s23d9

But the actual configuration is: 1s22s22p63s23p64s13d10

This gives one filled orbital and one half filled orbital.

Remember these exceptions: d4, d9

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Section 11.3Physics and the Quantum

Mechanical Model OBJECTIVES:

Calculate the wavelength, frequency, or energy of light, given two of these values.

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Section 11.3Physics and the Quantum

Mechanical Model OBJECTIVES:

Explain the origin of the atomic emission spectrum of an element.

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Light The study of light led to the

development of the quantum mechanical model.

Light is a kind of electromagnetic radiation.

Electromagnetic radiation includes many kinds of waves

All move at 3.00 x 108 m/s = c

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Parts of a wave

Wavelength

AmplitudeOrigin

Crest

Trough

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Parts of Wave - p.255 Origin - the base line of the energy. Crest - high point on a wave Trough - Low point on a wave Amplitude - distance from origin to crest Wavelength - distance from crest to

crest Wavelength is abbreviated by the Greek

letter lambda =

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Frequency The number of waves that pass a

given point per second. Units: cycles/sec or hertz (hz or sec-1) Abbreviated by Greek letter nu =

c =

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Frequency and wavelength Are inversely related As one goes up the other goes down. Different frequencies of light are

different colors of light. There is a wide variety of frequencies The whole range is called a spectrum,

Fig. 11.15, page 256

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Radiowaves

Microwaves

Infrared .

Ultra-violet

X-Rays

GammaRays

Low energy

High energy

Low Frequency

High Frequency

Long Wavelength

Short WavelengthVisible Light

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Prism White light is

made up of all the colors of the visible spectrum.

Passing it through a prism separates it.

61

If the light is not white By heating a gas

with electricity we can get it to give off colors.

Passing this light through a prism does something different.

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Atomic Spectrum Each element

gives off its own characteristic colors.

Can be used to identify the atom.

How we know what stars are made of.

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• These are called discontinuous spectra, or line spectra

• unique to each element.

• These are emission spectra

• The light is emitted (given off)

• Sample 11-16/17 p.257/58

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Light is a Particle Energy is quantized. Light is energy Light must be quantized These smallest pieces of light are

called photons. Photoelectric effect? Energy & frequency: directly related.

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Energy and frequency E = h x E is the energy of the photon is the frequency h is Planck’s constant h = 6.6262 x 10 -34 Joules x sec. joule is the metric unit of Energy

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The Math in Chapter 11

2 equations so far:

c = E = h Know these!

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Examples What is the wavelength of blue light

with a frequency of 8.3 x 1015 hz? What is the frequency of red light

with a wavelength of 4.2 x 10-5 m? What is the energy of a photon of

each of the above?

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Explanation of atomic spectra When we write electron

configurations, we are writing the lowest energy.

The energy level, and where the electron starts from, is called it’s ground state- the lowest energy level.

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Changing the energy Let’s look at a hydrogen atom

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Changing the energy Heat or electricity or light can move the

electron up energy levels (“excited”)

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Changing the energy As the electron falls back to ground

state, it gives the energy back as light

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May fall down in steps Each with a different energy

Changing the energy

73

{{{

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Further they fall, more energy, higher frequency.

This is simplified the orbitals also have different energies

inside energy levels All the electrons can move around.

Ultraviolet Visible Infrared

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What is light? Light is a particle - it comes in chunks. Light is a wave- we can measure its

wavelength and it behaves as a wave If we combine E=mc2 , c=, E = 1/2 mv2

and E = h We can get: = h/mv called de Broglie’s equation Calculates the wavelength of a particle.

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Sample problem What is the approximate mass of a

particle having a wavelength of 10-7 meters, and a speed of 1 m/s? Use = h/mv

= 6.6 x 10-27

(Note: 1 J = N x m; 1 N = 1 kg x m/s2

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Matter is a Wave Does not apply to large objects Things bigger than an atom A baseball has a wavelength of

about 10-32 m when moving 30 m/s An electron at the same speed has

a wavelength of 10-3 cm Big enough to measure.

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The physics of the very small Quantum mechanics explains how

the very small behaves. Classic physics is what you get

when you add up the effects of millions of packages.

Quantum mechanics is based on probability

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Heisenberg Uncertainty Principle

-It is impossible to know exactly the location and velocity of a particle.

The better we know one, the less we know the other.

Measuring changes the properties. Instead, analyze interactions with

other particles

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More obvious with the very small

To measure where a electron is, we use light.

But the light moves the electron And hitting the electron changes the

frequency of the light.

81

Moving Electron

Photon

Before

Electron Changes velocity

Photon changes wavelength

After

Fig. 11.26, p. 265

Documents

1 Chapter 11 Electrons in Atoms Killarney School