Chapter 3: Electrons in atoms

Learning outcomes:

Energy levels and shapes of orbitals

Electronic configurations

Ionisation energy, trends across a period

The quantum mechanical model

How are electrons arranged?Electrons are not evenly spread but exist in layers called shells. The arrangement of electrons in the shells is called the electron structure or electronic configuration..

3rd shell n=3

2nd shell n=2

1st shell n=1 Max:2

n2

Each shell has a maximum number of electrons that it can hold. Electrons fill the shells nearest the nucleus first.

3rd shell holdsa maximum of18 electrons

2nd shell holdsa maximum of

8 electrons

1st shell holdsa maximum of

2 electrons

Time for a break and practice

Please make check-up 1 on page

Simplified electron configuration

2, 5

2, 8

2, 8, 7

2, 8

7N2nd shell holdsa maximum of

8 electrons

1st shell holdsa maximum of

2 electrons

17Cl

11Na+

8O2-

For “ion” the number of proton is NOT the same as electrons

Electronic configuration This model assumes the electrons

have the same location and energy

Until it was discovered that electrons have different locations and energy

Li Li+ + e- (1st ionisation energy: E1)

Li+ Li2+ + e- (2nd ionisation energy: E2)

Li2+ Li3+ + e- (3rd ionisation energy: E3)

E1 ≠ E2 ≠ E3

Ionisation energy

Li Li+ + e- (1st ionisation energy: E1)

Li+ Li2+ + e- (2nd ionisation energy: E2)

Li2+ Li3+ + e- (3rd ionisation energy: E3)

Li 2,1 Which electron will be easiest to remove?

E1 < E2 < E3

ΔHi1 < ΔHi2 < ΔHi3

Table 3.2 in the book on p. 35 For every element, the successive ionisation energy increases;

for every next electron it is more difficult to remove We can in theory continue removing electrons until only the

nucleus is left We call this sequence the “successive ionisation energy” Sometimes we find a big gap/jump in ionisation energy

Example: sodium The first ionisation energy is quite low,

it is likely quite far from the nucleus

The 2nd to the 9th ionisation energy are in a gradual successive increase indicating these electrons are in the same shell

The 10th and 11th electrons have high ionisation energies compared to the rest, they must be near the nucleus.

The jump between the 9th and 10th

suggests a change in shell

Factors affecting the first ionization energy

Nuclear charge (number of protons) the bigger nuclear charge, the higher 1st ionization energy. Atomic radius (distance effect) the bigger atomic radius, the lower 1st ionization energy. Shielding effect (number of shells) the bigger Shielding effect, the lower 1st ionization energy

The first ionization energies of the first 20 elements in the periodic table is shown below:

Variation of first ionisation energy with atomic number for the first twenty elements

0

500

1000

1500

2000

2500

0 5 10 15 20

atomic number

firs

t io

nis

atio

n e

ner

gy

(kJ

per

mo

le)

Worked example

The model of the atom

A model is what fits logic, experimental observations and mathematical calculations

17Cl 2, 8, 7 3

2nd shell, witha maximum of

8 electrons

Symbol Simple electronic configuration Number of shells (last number is Group)

(=period)

1st shell, witha maximum of

2 electrons

3nd shell, witha maximum of18 electrons

6C

10Ne

11Na

Symbol Simple electronic configuration Number of shells (last number is Group)

(=period)

2, 4 2

2, 8 2

2, 8, 1 3

19K 4

Where in the atom is the electron?

According to quantum mechanics it is most likely to find the electron for the of the H-atom at 0.0000000000529 meter (52.9 pm) from the nucleus

Shells Principal quantum shells (n=1, n=2

etc.) Remember for each the max number

of electrons is 2n2 (so for n=2, max 8 electrons)

We know from experiments and calculations these 8 electrons have different energies…. so we need a new model of the atom where we can distinguish between electron energy

n = 1

n = 2

Quantum shell Subshells

The quantum mechanical model

Simplified model Realistic model

Principal quantum shell

Number of Sub-shells

Name of the Sub-shell

Max. number of electrons

n = 1 1 1s 2

n = 2 2 2s2p

26

n = 3 3 3s3p3d

2610

Subshells and their shapes

Atomic orbital is a space around the nucleus holding 1 or 2 electrons

n = 1 1s

Where in the atom are the electrons?

energy

2He 2

2

Simple electronic configuration Complicated electronic configuration

1s2

Principle quantum shell

Sub-shellNumber of electrons

n = 1

n = 2

1s (e<2)

2s (e<2)

2p (e<6)

Where in the atom are the electrons?

energy

8O 2, 6

2

24

Simple electronic configuration Complicated electronic configuration

1s22s22p4

n = 1

n = 2

1s (e<2)

2s (e<2)

2p (e<6)

Where in the atom are the electrons?

energy

11Na 2, 8, 1

226

Simple electronic configuration Complicated electronic configuration

1

n = 33d (e<10)

3p (e<6)3s (e<2)

n = 1

n = 2

1s (e<2)

2s (e<2)

2p (e<6)

Where in the atom are the electrons?

energyn = 3

3d (e<10)

3p (e<6)

3s (e<2)

n = 4

4s (e<2)

4d (e<10)

4p (e<6)

4f

Subshells and atomic orbitals

From simple to complicated electron configuration to noble gas electronic configuration notation

2, 6 1s22s22p4 [He] 2s2sp4

2, 8, 7 1s22s22p63s23p5 [Ne] 3s23p5

8O

17Cl

19K

35Br

Element: Simple: Complicated: Noble gas:

Note the following:

Potassium: 1s22s22p63s23p64s1

The 3d subshell: 3d<4p Chromium and copper are exceptions: Cr: [Ar] 4s13d5 rather than [Ar] 4s23d4 and Cu: [Ar] 4s13d10 rather than [Ar] 4s23d9

The blocks of the periodic table

Elements in Group 1 and Group 2 are in the s-block and have their outer electrons in an s subshell.

Elements in Group 3 to 18 have outer electrons in a p subshell. Elements that add electrons to the d subshells are called the d-block

elements.

Use the electronic configuration to find the group

…s2 is in group:

…p1 is in group:

…p3 is in group:

…p6 is in group:

…d3 is in group:

…d7 is in group:

In which period, group and block of the following electron configuration?

period group block

1s22s22p5

1s22s22p63s23p64s23d104p2

1s22s1

1s22s22p63s2

1s22s22p63s23p64s23d5

RULES FOR FILLING ENERGY LEVELS

Aufbau Principle “Electrons enter the lowest energy orbital first”

Pauli’s Exclusion “Sub-Orbitals can hold a max. of 2 electrons provided they have opposite spin”

Hund’s Rule “Orbitals of the same energy remain singly occupied before pairing up.

Examples

N = 1s22s22p3

O = 1s22s22p4

From simple to complicated electron configuration

2, 5 becomes 1s22s22p3

2, 6 becomes 1s22s22p6

2, 8, 7 becomes 1s22s22p63s23p5

2, 8, 8 becomes 1s22s22p63s23p6

7N

8O2-

17Cl

19K+

Ionisation: trend across a period

General increase across period Rapid decrease between last

element of a period and 1st of a new period

Be and B because 2s and 2p N and O

The first ionization energies of the first 20 elements in the periodic table is shown below:

Variation of first ionisation energy with atomic number for the first twenty elements

0

500

1000

1500

2000

2500

0 5 10 15 20

atomic number

firs

t io

nis

atio

n e

ner

gy

(kJ

per

mo

le)

Ionisation: Trend down a group

General trend decrease further away from the nucleus increased shielding despite increased nuclear charge

Li = 519 kJ/molNa = 494 kJ/molK = 418 kJ/molRb = 403 kJ/mol

Worked example