Electron

Configuration

IB Chemistry Power Points

Topic 12

Atomic Structure

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HL Topic 12.1 – Electron Configuration

Ionization energy is defined as the amount of energy required to remove a mole of electrons from a mole of gaseous atoms of a particular element.

E(g)

E+(g)

+ e-

Trends in ionization energy occur in the Periodic Table. Ionization energy decreases down a group and increases across a period.

Trends in ionization energy occur in the Periodic Table. Ionization energy decreases down a group and increases across a period.

WHY?

Effective nuclear charge is the net positive charge felt by an electron in an atom.

The basics of electrostatics imply that each and every electron is attracted to each and every proton in the nucleus and repelled by every other electron. However . . .

Electrons between the valence electron and the nucleus provide a shielding effect – weakening the electrostatic force on the valence electron.

Across a Period:

• shielding remains constant• atomic number increases so effective nuclear charge

increases• ionization energy increases

Down a Group:

• shielding increases AND atomic number increases• effective nuclear charge does not change significantly• valence electrons further from nucleus• so weaker electrostatic force and lower ionization

energy

Lithium (Z=3)

Sodium (Z=11)

Hydrogen (Z=1)+ e-

++

+

e-

8e-2e-

++++

e-

+e-

H+

++

+ e-

2e-

Li+

e-

8e-2e-

++++

Na+

2e-

This diagram shows how 1st ionization energy decreases down a group and that trends in ionization energy also occur across a period. Look at the H, Li, Na, K, Rb, Cs values!

.

Looking at just the trend across the 1st period, what does the graph imply?

The theory is . . . Across a period, number of p+ increases so effective nuclear charge increases.

As a result, the valence electrons are more strongly held, and are closer to the nucleus (radius decreases)

This does not explain the drop in ionization energy (decreased stability) observed between Be and B and between N and O.

NEW IDEA – suborbitals (or subshells)

Within a given energy level (shell), there are different subshells that electrons can occupy that have slightly different energy levels

Old Idea expanded: 2,8,8,2 configuration with energy shells 1,2,3,4 . . . but each shell subdivided

Electron arrangement in atoms can be described by terms called quantum numbers – no two electrons can have the same quantum

number!

1st Term: Shell (n)- principle energy level

n = 1

n = 2

n = 3

lone electronof Hydrogen

2nd Term: subshell - designated by s, p,d,f

1s

n = 2

n = 3

The first energy shell (1) has one subshell (s).

2nd Term: subshell - designated by s, p, d, f - designates the sub-energy level within the shell.- refers to the shape(s) of the volume of space in which the electron can be located.

1s

n = 2

n = 3

The first shell (1) has one subshell (s).

The s subshell has 1 spherical shaped orbital

orbitals are volumes of space where the probability of finding an electron is high

The Electronic Configuration of Hydrogen

1s

Hydrogen has one electron located in the first shell (1). (Aufbau principle)

The first shell has only one subshell (s). The s subshell contains 1 spherical orbital.

1s1

shell

subshell

# of electrons present

energy

1s

Electronic configuration

Orbital Energy Level Diagram

The Electronic Configuration of Helium He: Atomic # of 2, 2 electrons in a neutral He atom

H 1s1

He 1s2

He 1s 1s

the maximum number of electrons in an orbital is TWO

if there are 2 electrons in the same orbital they must have an opposite spin.

This is called Pauli’s Exclusion Principle

1s

Lithium (Li)Li: Z=3 Li has 3 electrons.

2nd shell

1s

The 2nd shell (n= 2) has 2 subshells which are s and p.

The s subshell fills first! (Aufbau Principle)

2s 2p

Li 1s22s1

2s

Li 1s

Electronic configuration

Orbital Energy Level Diagram

Berylium (Be)Be: Z=4 Be has 4 electrons.

Be 1s22s2 2s Be 1s

Electronic configurationOrbital Energy Level Diagram

1s2nd

shell2s 2p

B 1s22s22p1

2p 2s

B 1s

Boron (B) has 5 electrons, the s subshell is full so the 5th electron occupies the first orbital in the p subshell

Subshells so far - designated by s, and p - refers to the shape(s) of the volume in which the electron can be located. - also designates an energy level within the shell. - relative energy: s < p

s subshell: spherical1 orbital

p subshell: pair of lobes, 3 orbitals, each holds 2 electrons

x y z

x

y

z

Carbon (C)C: Z=6 C has 6 electrons.

1s2nd

shell2s 2p C 1s22s22px

1py1

2p 2s

C 1s

C 1s22s22p2

The 6th electron occupies an empty p orbital. This illustrates “Hund’s Rule” – electrons do not pair in orbitals until each orbital is occupied with a single electron.

The electron configuration is

But always written as

2p 2s

N 1s 1s22s22p3

2p 2s

O 1s 1s22s22p4

2p 2s

Ne 1s 1s22s22p6

Can we relate the filling of the subshells with the ionization energy data?

Ionization energy trends

Down a group : ionization energy decreases- ENC constant but atoms larger so easier to

ionize

Across a period : ionization energy increases- increasing ENC therefore smaller size (e- closer

to nucleus) so harder to ionize

Explaining the “dips” – support for s and p orbital model

Be to B “dip”- because s shields p and lowers ENC

N to O “dip” - because repulsions between electron pair in first full orbital (experimental evidence supporting Aufbau and Hund)

Electron Configurations and the Periodic Table

So far, we have seen how the subshell model provides and explanation for the patterns in ionization energy we see in the periodic table.

You have also seen how to write electron configurations

Example CALCIUM 1s22s22p63s23p64s2

Principle energy level subshell # of e-

Calcium can also be written shorthand as:

[Ar]4s2

Practice

Use the sheets provided to fill out orbital diagrams and determine the electron configuration for the following elements 

1. Fluorine 2. 56Fe3. Magnesium - 224. 131I5. Potassium – 426. 75Ge7. Zirconium – 908. 41Ca2+

Practice

Use the sheets provided to fill out orbital diagrams and determine the electron configuration for the following elements 

1. Fluorine 1s22s2p5

2. 56Fe 1s22s2p63s23p64s23d6

3. Magnesium – 22 1s22s2p63s2

4. 131I 1s22s2p63s23p63d104s24p64d105s25p5

5. Potassium – 42 1s22s2p63s23p64s1

6. 75Ge 1s22s2p63s23p64s23d104p2

7. Zirconium – 90 1s22s2p63s23p64s23d104p65s24d2

8. 41Ca2+ 1s22s2p63s23p6

The organization of the Periodic table correlates directly to electron structure

Condensed electron configurations – for example the electron configuration of bromine can be written [Ar] 4s23d104p5

Read questions carefully – many IB questions require you to write the FULL electron configuration

Electron configuration of ions:

The exception: TRANSITION METAL IONS

In general, electrons will be removed from orbitals (ionization) in the reverse order that the orbitals were filled. In other words, electrons vacate higher energy orbitals first.

When these ions form, electrons are removed from the valence shell s orbitals before they are removed from valence d orbitals when transition metals are ionized.

For example: Cobalt has the configuration [Ar] 4s2 3d7 OR [Ar] 3d7 4s2

The Co2+ and Co3+ ions have the following electron configurations.

Co2+: [Ar] 3d7 Co3+: [Ar] 3d6

Condensed electron configurations – for example the electron configuration of bromine can be written [Ar] 4s23d104p5

1. Si ___________________________2. S2- ___________________________3. Rb+ ___________________________4. Se ___________________________5. Ar ___________________________6. Nb ___________________________7. Zn2+ ___________________________8. Cd ___________________________9. Sb ___________________________

You are responsible for configurations up to Z = 54 (Xe). The table works well for this with the exception of Cr and Cu

Chromium’s configuration is:

[Ar]4s13d5

Copper’s configuration is:

[Ar]4s13d10

These configurations are energetically more stable than the expected arrangements. KNOW THEM!

1st 737.7

2nd 1450.

7

3rd 7732.

7

35458 31653 25661 21711 18020 13630 10542.5

16998

8

189367

.7

Successive ionizationenergy data supports the electron configuration model

Review: the principles involved

Hund’s Rule: the most stable arrangement of electrons in orbitals of equal energy is where there is the maximum number of unpaired electrons all with the same spin.

Aufbau Principle: electrons will fill the lowest energy orbitals first

Pauli’s Exclusion Principle: A maximum of two electrons can occupy a single orbital. These electrons will have opposite spins.