• Each shell is divided into subshells called s, p, d, f…. • There is one extra subshell for each new shell
10e 18e -
8e-
2eElectrons
1
Recap – Last Lecture
1s 2s
2p
3rd shell
3s
3p
3d
4th shell
4s
4p
4d
4f
5th shell
5s
5p
5d
5f
etc
2
• In hydrogen the energy of the sub-shells of a given n are degenerate (of the same energy). • In all other atoms the sub-shells are of different energies.
• Consequently electrons occupy a 3-D area. • Uncertainty means we have electron ‘clouds’ though often represented by a surface incorporating 90% of the electron density.
p orbital
1st shell 2nd shell
Sub-shell energy
• Electrons as waves.
s orbital
– First shell: 1s – Second shell: 2s and 2p – Third shell: 3s, 3p and 3d – Fourth shell: 4s, 4p, 4d and 4f
d orbital 3
4
1
Sub-shell energy
Electron configuration • There are three rules in determining electron configuration:
• The order in which the sub-shells are filled becomes important with the orbital energy increasing in the order:
• Pauli exclusion principle - no two electrons can have the same four quantum numbers. i.e. maximum of 2 electrons in any one orbital.
• Aufbau principle - fill up low energy orbitals first before high energy ones.
• Follow the arrows, starting at the top, to get the order of energies of the subshells.
• Hund’s rule - orbitals with the same energy (i.e. the same sub-shell) have the maximum number of unpaired electrons. 5
Electron Configuration
Electron Configurations
2 electrons can fit into a s subshell 6 electrons can fit into a p subshell 10 electrons can fit into a d subshell 14 electrons can fit into a f subshell
s: 2 electrons p: 6 electrons d: 10 electrons
• Ground state (no ‘excited’ electrons) – H has 1 electron: 1s1 – He has 2 electrons: 1s2 1s2
(1s now full)
2s1
– Li has 3 electrons: – Be has 4 electrons: 1s2 2s2 – B has 5 electrons: 1s2 2s2 2p1
energy increases
energy increases
• • • •
6
(2s now full)
1s
1s 2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
etc
7
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
etc
8
2
s: 2 electrons p: 6 electrons d: 10 electrons
Electron Configurations • Period 2, 3 & 4
s: 2 electrons p: 6 electrons d: 10 electrons
Electron Configurations • Period 4
– Ne has 10 electrons: 1s2 2s2 2p6 – Na has 11 electrons: 1s2 2s2 2p6 3s1 – Mg has 12 electrons: 1s2 2s2 2p6 3s2 – Ar has 18 electrons: 1s2 2s2 2p6 3s2 3p6
or: [Ne] 3s1 or: [Ne] 3s2
– Ca has 20 electrons: 1s2 2s2 2p6 3s2 3p6 4s2 or: [Ar] 4s2 – Fe has 26 electrons: [Ar] 4s2 3d6 – Zn has 30 electrons: [Ar] 4s2 3d10 – Se has 34 electrons: [Ar] 4s2 3d10 4p4
– K has 19 electrons: 1s2 2s2 2p6 3s2 3p6 4s1 or: [Ar] 4s1 – Ca has 20 electrons: 1s2 2s2 2p6 3s2 3p6 4s2 or: [Ar] 4s2 No-ce! We have filled 4s before 3d 9
Periodic Table • Group 1:
Box representation of orbitals • A box represents an orbital and an arrow represents an electron. • Indicates occupancy of orbitals.
• Group 2: 2s1
– Li: [He] – Na: [Ne] 3s1 – K: [Ar] 4s1
10
2 s2
– Be: [He] – Mg: [Ne] 3s2 – Ca: [Ar] 4s2
3s 2s
core valence electrons electrons
2p 2s
2p
1s
• Elements in the same group have the same valence electron configurations and, as a result, very similar chemical properties 11
1s
C
C: 1s2 2s2 2p2
Na Na: 1s2 2s2 2p6 3s1 12
3
Ions
Ions
• Add or subtract electrons as appropriate.
• In forming a cation, electrons are always removed from the sub-shell with the highest principle quantum number first. • So for period 4 elements this means the 4s electrons are removed first.
Eg Na: 1s2 2s2 2p6 3s1 or [Ne] 3s1 Na+:1s2 2s2 2p6 or [Ne]
Eg K: 1s2 2s2 2p6 3s2 3p6 4s1 or [Ar] 4s1 K+:1s2 2s2 2p6 3s2 3p6 or [Ar]
Applications Mendeleev constructed his periodic table guided by the
• Electrons can be excited to higher energy levels 1 s2
2s2
2p6
– Ground state Na has – One excited state of Na has 1s2 2s2 2p6 3s0 3p1 – Excited state relaxes to ground state emitting yellow light
3s1
sodium line
1s2 2s2 2p6 3s0 3p1 The Sun
1s2
2s2
2p6
14
3s1
Na
chemical properties of the elements. The reason this works is that elements in the same Group have similar valence p shell electron configurations. s Periodic Table has blocks of elements: “s block” (two elements wide) “p block” (6 elements wide) “d block” (10 elements wide) “f block” (14 elements wide)
d
16
emission spectra sodium street light
15
f
4
Learning Outcomes:
Questions to complete for next lecture:
• By the end of this lecture, you should:
1. List the sub-shells present in the third shell and indicate the maximum number of electrons each may contain.
− be able to write out the electron configuration for atoms and ions. − recognise ground state and excited state electron configurations. − Be able to use the box notation to represent orbital occupancy. − Understand the three rules associated with determining electron configurations.
2. Write the electron configuration of the following atoms or ions: a) N b) P3c) Ca d) Mn
3. What feature do all elements in the ‘p’ – block of the Periodic Table share?
− be able to complete the worksheet (if you haven’t already done so…) 17
18
Questions to complete for next lecture: 4. The electron configuration of He is 1s2. Why do you think it is included in Group 18 in the Periodic Table rather than Group 2? 5. Which of the following are impossible electron configurations? Give your reason(s). (a) 1s2 2p1 (b) 1s2 2s3 (c) 1s2 2s2 2p6 3d1 (d) 1s2 2s2 2p7 (e) 1s2 2s0.5 19
