Chapter 5

Chapter 5

Quantum Theory and Electron Configurations

It’s all about color…

• In terms of atomic models, so far:Dalton (1803) = Tiny, solid particle

Thomson (1897) = “Plum Pudding” model – Electrons

stuck on the outside of a big positive charge

Rutherford (1911) = Positively-charged nucleus with

electrons moving around it

• Rutherford’s model of the atom not quite rightCould not explain chemical properties of elements

Could not explain color changes when metal is heated

Bohr Model of the Atom

• Niels Bohr’s model of the atom

Electron found only on specific, circular paths around nucleus

Each orbit has fixed energy level

• Hypothesis: When electrons are excited (added

energy), jump into higher energy levels. When they moved back into lower energy levels - gave off light.

• Electrons do not exist between levels (think of rungs on a ladder)

Electrons absorb and emit only certain quanta (amounts) of energy

Quantum of energy = fixed amount of energy

required to move from one energy level to another energy level

Chapter 5

Bohr’s Model

Nucleus

Electron

Orbit

Energy Levels

Bohr’s Planetary Model of the Atom

• Electrons must have enough energy to keep moving around the nucleus

• Electrons orbit nucleus in defined energy levels, just like planets orbit the sun

• Each energy level assigned a principal quantum number n.

• Lowest energy level called ground state (n=1)

• Higher energy levels (n=2, 3, 4...) excited states

• Model worked OK for hydrogen but not so good for other elements

Nucleus

n = 1 n = 2

Chapter 5

Bohr’s ModelIn

crea

sin

g e

ner

gy

Nucleus

First

Second

Third

Fourth

Fifth• Further away from the

nucleus means more

energy.

• There is no “in

between” energy

• Energy Levels

Nucleus

Lowest energy level = ground state

Higher energy levels = excited states

Energy

Level 1

Energy

Level 2Energy

Level 3

Electron starts on lowest energy level (ground state)

Add energy to electron –moves to excited state

Energy levels are not evenly spaced

Nucleus

Electron returns to lower state –emits/gives off quantum of energy

Energy

Level 1

Energy

Level 2

Energy

Electron starts on lowest energy level (ground state)

Lowest energy level = ground state

Higher energy levels = excited states

Add energy to electron –moves to excited state

Energy levels are not evenly spaced

Energy

Level 3

Chapter 5

Bohr used this

theory to explain

the lines in the

atomic emission

spectra for

hydrogen

Chapter 5410 nm

434 nm

486 nm

656 nm

Each of these lines

corresponds to different

energy changes

Chapter 5

Chem I - Mon, 9/22/15

Do Now

Get to work on the PEN worksheet

from last class

Homework

MEAL paragraph if not finished

Agenda

History

Intro to quantum

Electron Config

Quantum-Mechanical Model of the

Atom

• Since the Bohr model had a very limited use, a new and very different model of the atom exists

• The Quantum Mechanical Model

(1926) contains:Quantum energy levels

Dual wave/particle nature of electrons

Electron clouds

• In the new model, don’t know exactlywhere electrons are - only know probabilities of where they could be

• Heisenberg Uncertainty Principle =

impossible to know both the velocity (or momentum) and position of an electron at the same time

Heisenberg Uncertainty Principle

Quantum-Mechanical Model of the

Atom

• Orbital = region around nucleus where an electron with a given energy level will probably (90%) be found

• Four kinds of orbitals

s - spherical in shape, lowest orbital for every energy level

p - dumbbell shaped, second orbital

d - complex “flower” shape, third orbital

f - very complex shape, highest orbital

s-orbitals

• All s-orbitals are spherical.

• As n increases, the s-orbitals get larger.

p- orbitals

• Three p-orbitals: px, py, and pz Lie along the x-, y- and z- axes of a Cartesian

system.

Dumbbell shaped, gets larger as n increases

d and f - orbitals

• There are five d and seven f-orbitals.

Quantum Mechanical Model

• Principle Energy Levels (n)

Labeled from 1-7

First energy level is n=1

Contains sublevels (s, p, d and f)

• Each energy level contains the number of sublevels equal to its value

for n

– If n=3, there are three sublevels

Chapter 5

Quantum Mechanical Model

• In each sublevel there are atomic orbitals

• Atomic orbitals – describe a space where an electron is

likely to be found

Type of

subshell

Shape of

orbitals

Number of

orbitals

Orbital

‘names’

s Spherical 1 s

p Dumbbell 3 px, py, pz

d Cloverleaf

(and one

donut)

5

f Multi-lobed 7

Quantum Mechanical Model

• Each orbital can contain two electrons.

• Since negative-negative repel, these electrons occupy

the orbital with opposite spins.

Quantum Mechanical Model

• The total number of orbitals of an energy level is n2.

For the third principle energy level, n=3, which means there are

9 orbitals

• These orbitals are 3s, 3px, 3p

y, 3p

z and the 5 d orbitals

• Remember, we no longer think of orbitals as concentric

circles, but we can say that n=4 extends farther from the

nucleus than n=1.

Valence Electrons

• Only those electrons in the highest principle energy

level

Electron Configuration and Orbital Notation

• Aufbau Principle – electrons fill lower energy orbitals first,

“bottom-up”

n=1 fills before n=3

• Will an electron fill the 1s or the 2s orbital first?

1s

2s2px 2py 2pzE

ne

rg

y

Electron Configuration &Orbital Notation

• Hund’s Rule – electrons enter same energy orbitals so that

each orbital has one electron before doubling up

Each of the first electrons to enter the equal energy orbitals must

have the same spin

If we have 7 electrons, how will they fill in the below orbitals?

1s

2s2px 2py 2pzE

ne

rg

y

Electron Configuration and Orbital Notation

• Pauli Exclusion Principle – an orbital can contain no more

than 2 electrons. Electrons in the same orbital must have

different spins.

• If we have 8 electrons, how will they be arranged?

1s

2s2px 2py 2pzE

ne

rg

y

Apartment Analogy

• Atom is the building

• Floors are energy levels

• Rooms are orbitals

• Only two people per room

Orbital Diagrams

• Draw each orbital as a box.

• Each electron is represented using an arrow.

Up arrows – clockwise spin

Down arrows – counter-clockwise spin

• Determine the total number of electrons involved.

• Start with the lowest energy level (1s) and start filling in

the boxes according the rules we just learned.

Orbital Diagram

2s

1s

3s

4s

2p

3p

4p

3d

Energ

y

Chapter 5

Incr

easi

ng e

ner

gy

1s

2s

3s

4s

5s

6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

Chapter 5

• The first to electrons go into the

1s orbital

• Notice the opposite spins

• only 13 more

Incr

easi

ng e

ner

gy

1s

2s

3s

4s

5s

6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

Chapter 5

• The next electrons go into the 2s

orbital

• only 11 more

Incr

easi

ng e

ner

gy

1s

2s

3s

4s

5s

6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

Chapter 5

• The next electrons go

into the 2p orbital

• only 5 more

Incr

easi

ng e

ner

gy

1s

2s

3s

4s

5s

6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

Chapter 5

• The next electrons go

into the 3s orbital

• only 3 more

Incr

easi

ng e

ner

gy

1s

2s

3s

4s

5s

6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

Chapter 5

Incr

easi

ng e

ner

gy

1s

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

• 3 unpaired electrons

• 1s22s22p63s23p3

Orbital Diagrams

• Orbital diagrams are used

to show placement of

electrons in orbitals.

• Need to follow three

rules (Aufbau, Pauli,

Hund’s) to complete

diagrams

Li

Be

B

C

N

Ne

Na

Chapter 5

Orbitals and Energy Levels

Principal

Energy Level

Sublevels Orbitals

n = 1 1s 1s (one)

n = 2 , 2s 2p 2s (one) + 2p (three)

, , n = 3 3s 3p 3d 3s (one) + 3p (three) + 3d (five)

n = 4 4s, 4p, 4d, 4f 4s (one) + 4p (three) + 4d (five)

+ 4f (seven)

Chapter 5

Summary

s

p

d

f

shapes Max

electrons

Starts at

energy level

Orbitals and Energy Levels

and so on....

1s

n = 1 2s

2p

n = 2 3s

3p

3d

n = 3

4s

4p

4d

4f

n = 4

Incre

asin

g e

nerg

y

Chapter 5

Electron Configuration

• Let’s determine the electron configuration for

Phosphorus

• Need to account for 15 electrons

Writing Electron Configuration

• Determine the total number of electrons.

• Write the principle energy level number as a coefficient,

the letter for the subshell, and an exponent to represent

the number of electrons in the subshell.

• He: 1s2

The Kernel (Noble Gas) Notation

• Determine the total number of electrons

• Find the previous noble gas and put its symbol in brackets

• Write the configuration from that noble gas forward as

usual

Chapter 5

Writing electron configurations

• Examples

• O 1s2

2s2

2p4

• Ti 1s2

2s2

2p6

3s2

3p6

3d2

4s2

• Br 1s2

2s2

2p6

3s2

3p6

3d10

4s2

4p5

• Core format

• O [He] 2s2

2p4

• Ti [Ar] 3d2

4s2

• Br [Ar] 3d10

4s2

4p5