Electron Configuration of the Elements

Hydrogen Emission Spectrum When hydrogen gas (H 2 ) is placed in a CRT and a high voltage electrical current passed through it, the tube glows a violet colour. Johann Balmer 

When viewed through a spectroscope (prism), we observe four discrete lines— and NOT a continuous spectrum:

When an electron in the ground state of H absorbs energy, it gets promoted into a higher energy level. The electron is unstable in this higher energy level.

When the electron falls back to the ground state, energy is given off. This explains the “bands” of light emitted from a hydrogen discharge tube.

Here’s another way to look at it:

Hydrogen Emission Spectrum Electrons can only exist in certain energy levels (n) n = 1, n = 2, n = 3, n = 4, etc Energy levels in atom are quantized. This means that only certain E levels are allowed.

Each E level has one or more sublevels called orbitals An orbital is a region of space where there is a high probability of finding an electron

Each orbital can hold a maximum of two electrons. Electrons in an orbital will have opposite spin, designated ↑ (clockwise spin) or ↓ (counterclockwise spin).

high “probability” ? Heisenberg Uncertainty Principle We cannot simultaneously know the position and the momentum of an electron

Back to orbitals... For n = 1 there is only one sublevel, called an s orbital. Since this orbital is in the first energy level, it is called a 1s orbital. s orbitals are spherical.

For n = 2 there are two sublevels: 2s orbital (one of these) 2p orbital (three of these) a p orbital looks like this A set of three p orbitals looks like this We refer to the individual p orbitals as p x, p y, p z.

Let’s put these orbitals together...

For n = 3 (the third energy level) there are three sublevels: 3s orbital (one of these) 3p orbital (three of these) 3d orbital (five of these—see next slide) NB. Each orbital holds a maximum of 2 electrons

The d-orbitals

A funky look at d-orbitals

Your “bottom line” with d-orbitals: There are five of them in each set. eg. there are five 3d orbitals; five 4d orbitals, etc 2 electrons in each, for a maximum of 10 electrons How many columns are in the Transition Metal block (d-block) in the periodic table? 10 columns in the transition metals (5x2).

For n = 4 (the fourth energy level) there are four sublevels: 4s orbital (one of these) 4p orbital (three of these) 4d orbital (five of these) 4f orbital (seven of these—see next slide)

f-orbitals

How do electrons fill orbitals? Aufbau Principle aka “Building-up” Principle Electrons occupy orbitals beginning from the lowest energy orbital (i.e. the orbital closest to the nucleus) Start by filling 1s orbital How many electrons per orbital? Each orbital can hold a maximum of two electrons—of opposite spin, don’t forget

Here is the order in which orbitals are filled...

Note the peculiarity... 3s is followed by 3p, which is followed by 4s, which is followed by 3d. There are others... (help is on the way)

How do the electrons of 7 N fill the orbitals? 1s 2 2s 2 2p 3 Overall for 7 N: 1s 2 2s 2 2p 3

Hund’s Rule More stable than...

Hund’s Rule When filling p, d, f orbitals, pair electrons only when necessary

Aufbau Principle Mnemonic Device

Let’s write some electron configurations... 1 H 1s 1 2 He 1s 2 3 Li 1s 2 2s 1 4 Be 1s 2 2s 2 5 B 1s 2 2s 2 2p 1 6 C 1s 2 2s 2 2p 2 ↓ ↓ 10 Ne 1s 2 2s 2 2p 6

11 Na 11 Na 1s 2 2s 2 2p 6 3s 1 12 Mg 1s 2 2s 2 2p 6 3s 2 13 Al 1s 2 2s 2 2p 6 3s 2 3p 1 ↓ ↓ 18 Ar 1s 2 2s 2 2p 6 3s 2 3p 6 19 K 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 20 Ca 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 21 Sc 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1

In the Periodic Table, what is the connection between the outermost electron configuration and family (column)? Alkali metals end in s 1 Alkali earth metals end in s 2 Halogens end in p 5 Noble Gases end in p 6

Groups (families) in PT

Putting it all together... To write the electron configuration of any element, use the periodic table (play “Battleship”) and the Aufbau Principle mnemonic device.

Write the complete electron configuration for 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 1 ____________________________________ [Kr] preceding noble gas only 49 In [Kr] 5s 2 4d 10 5p 1

Practice Write the electron configuration for each of the following 15 P 15 P 1s 2 2s 2 2p 6 3s 2 3p 3 33 As [use noble gas core abbreviated form] 33 As [Ar] 4s 2 3d 10 4p 3

more practice Al 13 Al 1s 2 2s 2 2p 6 3s 2 3p 1 26 Fe 26 Fe 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 44 Ru [ ] 44 Ru [Kr] 5s 2 4d 6 52 Te [ ] 52 Te [Kr] 5s 2 4d 10 5p 4

Exceptional Electron Configurations 1. Write the expected electron configuration of 24 Cr [Ar]4s 2 3d 4 Actual electron configuration is [Ar]4s 1 3d 5 Special stability associated with half-filled p, d, f orbitals

Now write the electron configuration for 42 Mo 42 Mo [Kr] 5s 1 4d 5 Notice any similarity with Cr?

2. Write the expected electron configuration of 29 Cu: [Ar]4s 2 3d 9 Actual electron configuration is [Ar]4s 1 3d 10 In this way Cu has completely filled 3 rd energy level (Copper is a very stable metal)

Now write the electron configuration for silver ( 47 Ag) and gold ( 79 Au). Use the noble gas core abbreviated forms. 47 Ag [Kr] 5s 1 4d Au [Xe] 6s 1 4f 14 5d 10