ELECTRONS IN ATOMS

LIMITATIONS OF RUTHERFORD’S ATOMIC MODEL

•Did not explain the chemical properties of atoms

– For example, it could not explain why metals or compounds of

metals give off characteristic colors when heated in a flame

– It could not explain why objects might change color when

heated to higher temperatures

T H E H Y D R O G E N AT O M

A N D B O H R

THE BOHR

MODEL

• In 1913 Niels Bohr

developed a new atomic

model

• Experiment: Tested

Hydrogen atoms

• Conclusion: Bohr

proposed electrons

orbit around the

nucleus in fixed energies

ENERGY LEVELS

• Fixed energy levels of an electron are similar to rungs of a ladder

– Electrons cannot exist between energy levels

– To move from one energy level to the next, an atom must gain or lose the correct amount of energy

– A quantum of energy is the amount of energy required to move an electron from one energy level to the next

• The amount of energy an electron

gains or loses is not always the same

– Higher energy levels are closer together

– It takes less energy to move from one

rung to the next near the top of the

ladder

– The higher the energy level occupied by

an electron, the less energy it takes the

electron to move from that energy level

to the next higher energy level

THE HYDROGEN ATOM

• Excited State (based off electron placement around nucleus)

– Bohr proposed that an electron moves into an orbit or higher

energy level further from the nucleus when an atom absorbs

energy

•Ground State

– Electron returns here (home) after being excited

THE HYDROGEN SPECTRUM

T H E Q U A N T U M M E C H A N I C A L

M O D E L A N D S C H R O D I N G E R

THE QUANTUM MECHANICAL MODEL• Starting Point: Both Bohr and Rutherford’s model of the atom described the

path of a moving electron

• Experiment: Austrian physicist Erwin Schrodinger used calculations and results

to devise and solve a mathematical equation describing the behavior of the

electron in a hydrogen atom

• Conclusion: There is a probability that describes how likely it is to find an

electron in a particular location around the nucleus of an atom

– Location is described as an ‘electron cloud’ that is dense and the probability of finding an

electron there is high

– For each energy level, the Schrodinger equation also leads to a mathematical expression

called an atomic orbital

• Model: Devised from the mathematical solutions to the Schrodinger equation

which is the modern description of the electrons

T H E Q UA N T U M M E C H A N I C A L

M O D E L D E T E R M I N E S T H E

A L L OW E D E N E R G I E S A N

E L E C T R O N C A N H AV E A N D

H OW L I K E LY I T I S TO F I N D

T H E E L E C T R O N I N VA R I O U S

L O C AT I O N S A R O U N D T H E

N U C L E U S

The propeller blade has the

same probability of being

anywhere in the blurry

region, but you cannot tell

its location at any instant.

The electron cloud of an

atom can be compared to a

spinning airplane propeller.

ATOMIC ORBITALS (CLOUDS)

• An atomic orbital is often

thought of as a region of

space in which there is a

high probability of finding

an electron.

– The probability of finding

an electron within the

‘electron cloud’ is 90%

Like the Bohr model, the quantum mechanical model

restricts the energy of electrons to certain values. Unlike the

Bohr model, the quantum mechanical model does not

specify an exact path the electron takes around the nucleus.

O R B I TA LSS , P, D, A N D F

S - O R B I TA L

• S-orbitals are

spherically shaped

• Smaller atoms have

fewer electrons and

take up less space

• Larger atoms have

more electrons and

take up more space

p-orbitals are

“dumbell” shaped.

z-axis

p-orbitals are

“dumbell” shaped.

x-axis

p-orbitals are

“dumbell” shaped.

y-axis

p-orbitals together

x, y, & z axes.

ATOMIC ORBITALS

•Different atomic orbitals are denoted by letters.

• Four of the five d orbitals have the same shape, but

different orientations in space.

ATOMIC ORBITALS

•Describe the probability of finding an electron at various

locations around the nucleus

– s orbitals: groups 1 and 2 on the periodic table

– p orbitals: groups 13-18

– d orbitals: groups 3-12

– f orbitals: lanthanide and actinide series

ORBITAL SUBLEVELS

• Each energy sublevel (subshells) correspond to one or

more orbitals of different shapes.

– The orbitals describe where an electron is likely to be found.

Summary of Principal Energy Levels and Sublevels

Principal energy

level

Number of

sublevelsType of sublevel

Maximum

number of

electrons

n = 1 1 1s (1 orbital) 2

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

n = 3 33s (1 orbital), 3p (3 orbitals),

3d (5 orbitals)18

n = 4 44s (1 orbital), 4p (3 orbitals),

4d (5 orbitals), 4f (7 orbitals)32

SUBLEVELS

• The Principal Quantum Number, n, always equals the number of

sublevels within that principal energy level

• The number of orbitals in a principal energy level is equal to n2

• A maximum of two electrons can occupy an orbital

• Therefore, the maximum number of electrons that can occupy a

principal energy level is given by the formula 2n2

VA L E N C E E L E C T R O N S

VA L E N C E

E L E C T R O N S

• The electrons in the outermost,

furthest from the nucleus, electron

shell are called valence electrons

• The number of valence electrons in

orbitals s and p (not transition

metals) is the same as the group

number

• The number of electron shells with

electrons in them is the same as the

period number

NOBLE GAS STABILITY

• Noble gases are usually unreactive

• This is because they have a full valence shell

• For two atoms to bond, they must gain, lose, or share electrons

– Metals tend to lose electrons

– Non-metals tend to gain electrons

ELECTRON CONFIGURATION

BLOCK TABLE

• The periodic table shows the different blocks located on the periodic table

• It also shows the electron configuration order

– 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d

E L E C T R O N P L A C E M E N TA N D T H E R U L E S T H AT F O L L OW

AUFBAU PRINCIPLE• Electrons are placed in the lowest energy levels first

PAULI EXCLUSION PRINCIPLE• Only 2 electrons can be held in an orbital, different than an

electron energy shell, and they must have opposite spins

HUND’S RULE• Every orbital within a sublevel gets an electron before any gets paired

ELECTRON CONFIGURATIONS

• The electron configuration of an atom is a shorthand method of

writing the location of electrons by sublevel

• The sublevel is written followed by a superscript with the number

of electrons in the sublevel

– If the 2p sublevel contains 2 electrons, it is written:

2p2

Number of electrons

Energy level

Energy sublevel

ELECTRONS IN SUBLEVELS REVIEW

• s-orbital

– 1 orbital, 2 electrons

• p-orbital

– 3 orbitals, 6 electrons

• d-orbital

– 5 orbitals, 10 electrons

WRITING ELECTRON CONFIGURATIONS• First, determine how many electrons are in the atom

– For example, Iron has 26 electrons

• Arrange the energy sublevels according to increasing energy

– 1s 2s 2p 3s 3p 4s 3d

• Fill each sublevel with electrons until you have used all the

electrons in the atom

– Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d6

• The sum of the superscripts equals the atomic number of iron (26)

ELECTRON CONFIGURATION PRACTICE

• Write a ground state electron configuration of a neutral atom:

• K:

• Ne:

A SHORTCUT!1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f 5g

6s 6p 6d 6f 6g 6h

7s 7p 7d 7f 7g 7h 7i

Do not exist in

normal ground

state atoms

MORE PRACTICE!

• Write a ground state electron configuration of a neutral atom using

the shortcut:

• Cl:

• Rb:

N O B L E G A S C O N F I G U R AT I O N

NOBLE GAS CONFIGURATION

• The Noble Gases are:

– He, Ne, Ar, Kr, Xe, Rn

• Notice that each noble gas finishes a row, or energy level

• Noble gas configurations take advantage of this by condensing what

you have to write

– Example: He: 1s2

– Example: C: 1s2 2s2 2p2

• Noble Gas Configuration for C: [He] 2s2 2p2

MORE EXAMPLES

• The ground state configuration for Arsenic (As) is:

– 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3

• Notice, that the part in purple is the same as Argon’s configuration:

– 1s2 2s2 2p6 3s2 3p6

• The noble gas configuration will state with the gas in the row

before it

– [Ar] 4s2 3d10 4p3

NOBLE GAS CORE ELECTRON CONFIGURATIONS

• Core Electrons: Electrons in [Noble Gas]

• Valence Electrons: Electrons outside of [Noble Gas]

• Recall, the electron configuration for Sodium (Na) is:

– Na: 1s2 2s2 2p6 3s1

• We can abbreviate the electron configuration by indicating the innermost

electrons with the symbol of the preceding noble gas

• The preceding noble gas with an atomic number less that sodium is neon, Ne.

We rewrite the electron configuration:

– Na: [Ne] 3s1

NOBLE GAS CONFIGURATION PRACTICE

• Write the noble gas configuration for the following neutral atoms:

• Cu:

• Sr:

O R B I TA L D I A G R A M S

THE AUFBAU

PRINCIPLE

• Each electron

occupies the lowest

energy orbital

• All orbitals related to

an energy level are of

equal energy

• Example: The three 2p

orbitals are the same

energy level

ORBITAL FILLING DIAGRAM

PAULI EXCLUSION PRINCIPLE• A maximum of two electrons

may occupy a single orbital, but

only if the electrons have

opposite spins

– Spin: Electrons have an associated

‘spin’, either one way or the

other

– These spins are called ‘spin up’

and ‘spin down’

– In the example to the right:

• Box = orbital

• Arrow = electron

HUND’S RULE

• Single electrons with the

same spin must occupy

each equal-energy orbital

before additional

electrons with opposite

spins can occupy the

same orbitals

Example: Nitrogen

1s2 2s2 2p3

1s2 2s2 2p3

1s2 2s2 2p3

NOT

THE ORDER OF THINGS…

•Electrons fill up

the empty

orbitals before

sharing orbitals

EXCEPTIONAL ELECTRON CONFIGURATIONS

• Some actual electron configurations differ from those

assigned using the aufbau principle because half-filled

sublevels are not as stable as filled sublevels, but they are

more stable than other configurations.

–Exceptions to the aufbau principle are due to subtle

electron-electron interactions in orbitals with very

similar energies

ORBITAL DIAGRAM PRACTICE

• Draw the orbital diagram for the following neutral atoms:

• N:

• Al:

MORE PRACTICE!

• Ti:

• Mg:

• As: