Atomicstructureoutcomes1-6notes.ppt

(1)

(2)

Section 11.3

Physics and the Quantum

Mechanical Model



OBJECTIVES:



Explain the origin of the atomic

(3)

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

(4)

Parts of a wave

Wavelength

Amplitude

Origin

Crest

(5)

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

(6)

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 =



(7)

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,

(8)

Radio waves Micro waves Infrared . Ultra-violet X-Rays Gamma Rays Low energy High energy Low

Frequency High _Frequency

Long

Wavelength Short _Wavelength

(9)

Prism

 White light is

made up of all the colors of the

visible spectrum.

 Passing it through

(10)

If the light is not white

 By heating a gas

with electricity we can get it to give off colors.

 Passing this light

(11)

Atomic Spectrum

 Each element

gives off its own characteristic

colors.

 Can be used to

identify the atom.

 How we know

(12)

• These are called discontinuous

spectra, or line spectra

• unique to each element.

• These are

emission spectra

(13)

Light is a Particle



Energy is quantized.



Light is energy



Light must be quantized



These smallest pieces of light are

called photons.



Photoelectric effect?

(14)

Energy and frequency



E = h x







E is the energy of the photon







is the frequency



h is Planck’s constant

(15)

The Math in Chapter 11



2 equations so

far:



c =





E = h



(16)

Examples



What is the wavelength of blue light

with a frequency of 8.3 x 10

15

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

(17)

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

(18)

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

(19)

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

(20)

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

(21)

Section 11.1

Models of the Atom



OBJECTIVES:

(22)

Section 11.1

Models of the Atom



OBJECTIVES:



Explain the significance of

quantized energies of electrons

as they relate to the quantum

(23)

Greek Idea

 Democritus and

Leucippus

 Matter is made up

of solid indivisible particles

 John Dalton - one

(24)

J. J. Thomson’s Model

 Discovered electrons

 Atoms were made of

positive stuff

 Negative electron

floating around

 “Plum-Pudding”

(25)

Ernest Rutherford’s Model

 Discovered dense

positive piece at the center of the atom- nucleus

 Electrons would

surround it

 Mostly empty

space

(26)

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

(27)

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

(28)

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

(29)

 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.

(30)

 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.

(31)

 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

(32)

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 -

(33)

For any atom there is only one

1s orbital. The "1" represents the

fact that the orbital is in the

energy level closest to the

(34)

(35)

2p and 3p orbitals

At 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.

(36)

(37)

(38)

(39)

(40)

Videos of Atomic Orbitals

(41)

s2 Principle energy level 1 Principle energy level 2 Principle energy level 3 Principle energy level 4 Principle energy level 5

Principle energy level 6 Principle energy level 7

s2 p6 s2 p6

s2 p6 d10

(42)

(43)

s

p

d

# of

shapes

Max

electrons

Starts at

energy level

1 2 1

3 6 2

(44)

Orbital Shapes

(45)

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

(46)

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

(47)

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

(48)

Section 11.2

Electron Arrangement in Atoms



OBJECTIVES:



Apply the aufbau principle, the

Pauli exclusion principle, and

(49)

Section 11.2

Electron Arrangement in Atoms



OBJECTIVES:



Explain why the electron

(50)

(51)

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

(52)

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

(53)

 The first two electrons

go into the 1s orbital

 Notice the opposite

spins

only 13 more to go...

(54)

 The next electrons

go into the 2s orbital

 only 11 more...

(55)

• The next electrons go

into the 2p orbital

• only 5 more...

(56)

• The next electrons go into the 3s orbital

• only 3 more...

(57)

In

cr

ea

si

ng

e

ne

rg

y

2s 3s 4s 5s 6s 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

(58)

The easy way to remember

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d 6f

7s 7p 7d 7f

• 1s

2

(59)

Fill from the bottom up

following the arrows

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d 6f

7s 7p 7d 7f

• 1s

2

2s

2

(60)

Fill from the bottom up

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d 6f

7s 7p 7d 7f

• 1s

2

2s

2

2p

6

3s

2

(61)

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d 6f

7s 7p 7d 7f

• 1s

2

2s

2

2p

6

3s

2

3p

6

4s

2

(62)

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d 6f

7s 7p 7d 7f

• 1s

2

2s

2

2p

6

3s

2

3p

6

4s

2

3d

10

4p

6

5s

2

(63)

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d 6f

7s 7p 7d 7f

• 1s

2

2s

2

2p

6

3s

2

3p

6

4s

2

3d

10

4p

6

5s

2

4d

10

5p

6

6s

2

(64)

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d 6f

7s 7p 7d 7f

• 1s

2

2s

2

2p

6

3s

2

(65)

Fill from the bottom up

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d 6f

7s 7p 7d 7f

• 1s

2

2s

2

2p

6

3s

2

3p

6

4s

2

3d

10

4p

6

5s

2

4d

10

5p

6

6s

2

4f

14

5d

10

6p

6

7s

2

5f

14

6d

10

7p

6

(66)

(67)

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.

(68)

Write these electron

configurations

 Titanium - 22 electrons

 1s22s22p63s23p64s23d2

 Vanadium - 23 electrons

 1s22s22p63s23p64s23d3

 Chromium - 24 electrons

(69)

Chromium is actually:



1s

2

2s

2

2p

6

3s

2

3p

6

4s

1

3d

5



Why?



This gives us two half filled orbitals.



Slightly lower in energy.



The same principal applies to

(70)

Copper’s electron

configuration



Copper has 29 electrons so we

expect: 1s

2

2s

2

2p

6

3s

2

3p

6

4s

2

3d

9



But the actual configuration is:



1s

2

2s

2

2p

6

3s

2

3p

6

4s

1

3d

10



This gives one filled orbital and one

half filled orbital.

(71)

Explanation of atomic spectra



When we write electron

configurations, we are writing the

lowest energy.



The energy level, and where the

(72)

Changing the energy

(73)

Changing the energy

 Heat or electricity or light can move the

(74)

Changing the energy

 As the electron falls back to ground

(75)

 May fall down in steps

 Each with a different energy

(76)

{

(77)

 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.

(78)

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.

(79)

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

(80)

Moving Electron

Photon

Before

Electron Changes velocity

Photon changes wavelength

(81)

References

This power point presentation

was created by Terry Sproat

dt.edu

mb bell