CHAPTER 5CHAPTER 5ELECTRONS IN ATOMSELECTRONS IN ATOMS

P. 126P. 126

ERNEST RUTHERFORD’S MODEL

Discovered dense + nucleus

•e-s move like planets around sun

•Mostly empty space

Didn’t explain chemical properties of elements

THE BOHR MODEL

Neils Bohr (1885-1962)

•Danish physicist

•Student of Rutherford “W

hy d

on’t

e-s

fall

into

nu

cleu

s”?Why don’t Why don’t

electrons electrons fall into fall into nucleus?nucleus?

Niels Bohr

THE BOHR MODEL

Niels Bohr

I pictured the electrons found in specific circular paths around the nucleus, and can jump from one level to another.

Furthermore, each level has a fixed amount of energy different from other levels

BOHR’S MODEL

fixed energy e- have called Energy levels

Like rungs of laddere- can’t exist btwn energy levels

energy levels not evenly spaced

• High levels closer (less energy needed to jump)

Bohr’s model of the atom 5:17

THE QUANTUM MECHANICAL MODEL• e-’s don’t move like big objects

• Rutherford & Bohr model• Energy - “quantized” (in chunks)

• exact energy needed to move e- 1 energy level called a quantum

• energy never “in btwn” • quantum leap in energy must exist

• Erwin Schrodinger (1926) mathematically described energy & position of e- in atom

SchrodingerSchrodinger

Quantum Leap TV intro

THE QUANTUM MECHANICAL MODEL

energy levels for e-

Orbits not circular Based on probability of

finding e- certain distance from nucleus electron cloud

ATOMIC ORBITALS• Principal Quantum # (n) - energy

level of e- (1, 2, 3 etc.)• atomic orbitals - regions of space regions of space w/w/ high high

probability probability of finding e- (not a true “orbit”)(not a true “orbit”)• within each energy level• Sublevels like rooms in a hotel• s, p, d, and f• Different shapes

Max # of e- that fit in energy level is: 2n2

How many eHow many e-- in level 2? in level 2? level 3?level 3?

s and p orbitals 1:20

d orbitals 3:40

atomic orbitals review (14:28)

ATOMIC ORBITALS

sspherical

pdumbell

dclover leaf

fcomplicated

# of orbitals (regions of

space)

Maximum electrons

First possible

energy level

1 2 1st

3 6 2nd

5 10 3rd

7 14 4th

Summary of Principal Energy Levels, Sublevels, and Summary of Principal Energy Levels, Sublevels, and OrbitalsOrbitals

Principal Principal energy energy levellevel

Number Number of of

sublevelssublevels

Type of Type of sublevelsublevel

Max # of Max # of electronselectrons

Electron Electron configurationconfiguration

n = 1n = 1 11 1s (1 orbital)1s (1 orbital) 22 1s1s22

n = 2n = 2 22 2s (1 orbital2s (1 orbital2p (3 orbitals)2p (3 orbitals)

88 2s2s22

2p2p66

n = 3n = 3 33 3s (1 orbital)3s (1 orbital)3p (3 orbitals)3p (3 orbitals)3d (5 orbitals)3d (5 orbitals)

1818 3s3s22

3p3p66

3d3d1010

n = 4n = 4 44 4s (1 orbital)4s (1 orbital)4p (3 orbitals)4p (3 orbitals)4d (5 orbitals) 4d (5 orbitals) 4f (7 orbitals)4f (7 orbitals)

3232 4s4s22

4p4p66

4d4d1010

4f4f1414

ORDER OF ELECTRON SUBSHELL FILLING:

NOT “IN ORDER”

1s2

2s2 2p6

3p6

4p6

5p6

6p6

7p6

3s2

4s2

5s2

6s2

7s2

3d10

4d10

5d10

6d10

4f14

5f14

1s2 2s2 2p6 3p63s2 4s2 4p65s23d10 5p6 6s24d10 6p6 7s25d104f14 7p66d105f14

energy levels overlap

Lowest energy fill first

Increa

sing e

nerg

y

ELECTRON CONFIGURATION

1s1

Principal energy level

row #

1-7

7 rows sublevel

s, p, d, or f

4 sublevels

group #

# valence e-

s: 1 or 2

p: 1-6

d: 1-10

f: 1-14

Total e- = Atomic #

1

2

3

4

5

6

7

6

7

per

iod

# =

# e

- en

erg

y le

vels 1A

2A

3B 4B5B 6B 7B

8B

1B 2B

3A 4A 5A6A 7A8A

group # = # valence e-

d

f

3d

4d

5d6d

4f

5f

SUBLEVELS D AND F ARE “SPECIAL”

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 133

Aufbau - German for “building up”

SECTION 5.2 ELECTRON SECTION 5.2 ELECTRON ARRANGEMENT IN ATOMS P. 133ARRANGEMENT IN ATOMS P. 133

ELECTRON CONFIGURATIONS…….3 rules explain how e-’s fill their

orbitals:

1) Aufbau principle – e-’s enter lowest energy level first.

2) Pauli Exclusion Principle - 2 e-’s max/orbital (hotel room) - different spins

PAULI EXCLUSION PRINCIPLENo 2 electrons in an atom can have the same four quantum numbers.

Wolfgang Pauli

To show different direction of spin, a pair in the same orbital is written as:

QUANTUM NUMBERS

Each e- has unique set of 4 quantum #’s describing it

1) Principal quantum #2) Angular momentum quantum #3) Magnetic quantum #4) Spin quantum #

ELECTRON CONFIGURATIONS

3) Hund’s Rule- When e-’s occupy orbitals of same energy, they won’t pair up until they must

write e- configuration for Phosphorus

all 15 e-’s must be accounted for

The first 2 e-’s go into the 1s orbital

Notice opposite direction of spinsIn

crea

sing

ene

rgy

1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

The next e-’s go in 2s orbital

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 next e-’s go in 2p orbital

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 next e-’s go in 3s orbital

Incr

easi

ng e

nerg

y

1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

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 3 e-’s go in 3p orbitals

They each go into separate shapes (Hund’s)

• 3 unpaired e-’s

= 1s22s22p63s23p3 Orbital notation

An internet program about electron configurations is:

Electron Configurations

I electron config (song) 3:24

FILLING ORBITALS

Lowest higher energy

Adding e-’s changes energy of orbital

•Full orbitals best situation•half filled orbitals next best

• more stable• Changes filling order

WRITE THE ELECTRON CONFIGURATIONS FOR THESE ELEMENTS:Titanium - 22 electrons

1s2

2s2

2p6

3s2

3p6

4s2

3d2

Vanadium - 23 electrons

1s2

2s2

2p6

3s2

3p6

4s2

3d3

Chromium - 24 electrons

1s2

2s2

2p6

3s2

3p6

4s2

3d4 (expected)

But this is not what happens!!

CHROMIUM IS ACTUALLY:

1s22s22p63s23p64s13d5 Why?

2 half filled orbitals

•Half full slightly lower in energy

•Same applies to copper

COPPER’S E- CONFIGURATION

• Copper has 29 e-s so expect: 1s22s22p63s23p63d94s2

• actual configuration is:

1s22s22p63s23p63d104s1

• 1 more full orbital & 1 half filled

• Exceptions

• d4 • d9

IRREGULAR CONFIGURATIONS OF CHROMIUM AND COPPER

Chromium steals a 4s e- to make its 3d sublevel HALF FULL

Copper steals a 4s electron to FILL its 3d sublevel

SECTION 5.3SECTION 5.3 PHYSICS AND THE QUANTUM PHYSICS AND THE QUANTUM MECHANICAL MODELMECHANICAL MODEL

P. 138P. 138

• Study of light led to quantum mechanical model

• Light is electromagnetic radiation

• EM radiation: gamma rays, x-rays, radio waves, microwaves

• Speed of light = 2.998 x 108 m/s

• “c” - celeritas (Latin for speed)

• All EM radiation travels same in vacuum

Light

- Page 139

“R O Y G B I V”

Frequency Increases

Wavelength Longer

PARTS OF A WAVE

Wavelength

Amplitude

Crest

Trough

Equation:

c =

c = is a constant (2.998 x 108 m/s)

(nu) = frequency, in units of hertz (hz or sec-1) (lambda) = wavelength, in meters

ELECTROMAGNETIC RADIATION PROPAGATES THROUGH SPACE AS A WAVE MOVING AT THE SPEED OF LIGHT.

WAVELENGTH AND FREQUENCY

• inversely related

• one gets bigger, other smaller

• Different frequencies = different colors

• wide range of frequencies (spectrum)

- Page 140

Use Equation: c =

Radiowaves

Microwaves

Infrared .

Ultra-violet

X-Rays

GammaRays

Low Frequency

High Frequency

Long Wavelength

Short WavelengthVisible Light

Low Energy

High Energy

Long =Low Frequency

=Low ENERGY

Short =High Frequency

=High ENERGY

ATOMIC SPECTRA

White light all colors of visible spectrum

• prismprism separates it

according to λ

IF THE LIGHT IS NOT WHITE

heating gas with electricity will emit colors

• this light thru prism is different

ATOMIC SPECTRUM

elements emit own characteristic colors

• composition of stars determined thru spectral analysis

• atomic emission spectrum

• Unique to each element, like fingerprints!

• ID’s elements

LIGHT IS A PARTICLE?Energy is quantizedLight is energy…..light must be quantized

photons photons smallest pieces of lightPhotoelectric effect –

• Matter emits e- when it absorbs energy• Albert Einstein Nobel Prize in chem

Energy & frequency: directly related

Planck-Einstein Equation: E = hEE = Energy, in units of Joules (kg·m = Energy, in units of Joules (kg·m22/s/s22)) (Joule…metric unit of energy)(Joule…metric unit of energy)

hh = Planck’s constant (6.626 x 10 = Planck’s constant (6.626 x 10-34-34 J·s) J·s)(reflecting sizes of energy quanta)(reflecting sizes of energy quanta)

= frequency, units of hertz (hz, sec= frequency, units of hertz (hz, sec-1-1))

ENERGY (E ) OF ELECTROMAGNETIC RADIATION DIRECTLY PROPORTIONAL TO FREQUENCY () OF RADIATION.

THE MATH IN CHAPTER 5There are 2 equations:

1) c = 2) E = h Put these on your 3 x 5

notecard!

EXAMPLES1) What is the wavelength of

blue light with a frequency of 8.3 x 1015 hz?

2) What is the frequency of red light with a wavelength of 4.2 x 10-5 m?

3) What is the energy of a photon of each of the above?

EXPLANATION OF ATOMIC SPECTRA

electron configurations written in lowest energy.

energy level, and where electron

starts from, called it’s ground ground statestate - lowest energy level.

CHANGING THE ENERGYLet’s look at a hydrogen atom, with only one electron, and in the first energy level.

Changing the energyHeat, electricity, or light can move e-’ up to different energy levels. The electron is now said to be ““excitedexcited””

Changing the energyAs electron falls back to ground state,

it gives energy back as lightlight

Experiment #6, page 49-

may fall down in specific steps

Each step has different energy

Changing the energy

{{{

Lyman series (UV) Balmer series

(visible)

Paschen series

(infrared)

further they fall, more energy released = higher frequency

orbitals also have different energies inside energy levels

All electrons can move around.

Ultraviolet Visible Infrared

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

combine E=mc2 , c=, E = 1/2 mv2 and E = hthen we can get:

= h/mv (from Louis de Broglie)Calculates wavelength of a particle.

called de Broglie’s equation • He said particles exhibit properties of waves

WAVE-PARTICLE DUALITYJ.J. Thomson won the Nobel prize for describing the electron as a particle.

His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron.

The electron is a particle!

The electron is an energy

wave!

CONFUSED? YOU’VE GOT COMPANY!

“No familiar conceptions can be woven around the electron;

something unknown is doing we don’t know what.”

Physicist Sir Arthur Eddington

The Nature of the Physical World

1934

THE PHYSICS OF THE VERY SMALL

Quantum mechanics explains how very small particles behave

•Quantum mechanics is an explanation for subatomic particles and atoms as waves

Classical mechanics describes the motions of bodies much larger than atoms

HEISENBERG UNCERTAINTY PRINCIPLEimpossible to know exact location and velocity of particle

better we know one, less we know other

Measuring changes properties.

True in quantum mechanics, but not classical mechanics

HEISENBERG UNCERTAINTY PRINCIPLE

You can find out where the electron is, but not where it is going.

OR…

You can find out where the electron is going, but not where it is!

“One cannot simultaneously determine both the position and momentum of an electron.”

Werner Heisenberg

IT IS MORE OBVIOUS WITH THE VERY SMALL OBJECTS

To measure where e-, we use light

But light energy (photon) moves e- due to small mass

And hitting e- changes frequency of light

Moving Electron

Photon

Before

Electron velocity changes

Photon wavelengthchanges

After

Fig. 5.16, p. 145