1. Transition Metals

2. Energy levels
In the modern atomic model of the atom, electrons are arranged in energy levels (shells)
Each energy level is given a number (1, 2, 3 or 4)
Energy level 1 is closest to the nucleus
The energy of each level increases the further away it is from the nucleus
Level 4
Level 3
Increasing energy
Level 2
Level 1
Nucleus

3. Sub-levels Each energy level contains sub-levels
Each sub-level is given a letter (s, p, d or f)
Each sub-level (s, p, d or f) can hold a maximum number of electrons:
s sub-level can hold 2 electrons
p sub-level can hold 6 electrons
d sub-level can hold 10 electrons
The energy of the sub-levels increases from s to p to d

4. The electrons are found in the orbitals
Each sub-level consists of orbitals
Each orbital can hold a maximum of two electrons and the electrons must have opposite spin
Electrons exist in two spin states, spin up or spin down
Each sub-level contains different number of orbitals
s sub-level has 1 orbital (max 2 electrons)
p sub-level has 3 orbitals (max 6 electrons)
d sub-level has 5 orbitals (max 10 electrons) ↿ ⇂
The electrons are found in the orbitals

5. Each time you have a new energy level you gain a sub-level as well
Arranging electrons
The order in which the energy levels are filled with electrons is:
1s2
2s2 2p6
3s2 3p6 3d10
4s2 4p6 4d10
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 4d10
The number of electrons has to be written as a superscript
Remember s, p, d
Each time you have a new energy level you gain a sub-level as well
4s is before the 3d

6. This is called the electron arrangement or electron configuration
Arranging electrons
In atoms, the atomic (proton) number = number of electrons
Examples:
1) Sodium (11 electrons)
1s2 2s2 2p6 3s1
2) Phosphorus (15 electrons)
1s2 2s2 2p6 3s2 3p3
This is called the electron arrangement or electron configuration

7. Arranging electrons
We can draw energy level diagrams to show electron arrangements
The electrons fill the orbitals in each sub-level as single unpaired electrons, all with the same spin
The electrons are paired up when there are no more empty orbitals in that sub-level
Example: Sodium (11 electrons)
1s2 2s2 2p6 3s1 1s 2s 2p 3s ↿ ⇂ ↿ ⇂ ↿ ⇂ ↿ ⇂ ↿ ⇂ ↿

8. Arranging electrons Example 1: Phosphorus (15 electrons)
1s2 2s2 2p6 3s2 3p3
Example 2: Sulfur (16 electrons)
1s2 2s2 2p6 3s2 3p4 1s 2s 2p 3s 3p ↿ ⇂ ↿ ⇂ ↿ ⇂ ↿ ⇂ ↿ ⇂ ↿ ⇂ ↿ ↿ ↿ 1s 2s 2p 3s 3p ↿ ⇂ ↿ ⇂ ↿ ⇂ ↿ ⇂ ↿ ⇂ ↿ ⇂ ↿ ⇂ ↿ ↿

9. Arranging electrons Example 3: Potassium (19 electrons)
1s2 2s2 2p6 3s2 3p6 4s1
Example 4: Calcium (20 electrons)
1s2 2s2 2p6 3s2 3p6 4s2
Once the 4s sub-level is full (2 electrons), the 3d is filled
Example: 5: Scandium (21 electrons)
1s2 2s2 2p6 3s2 3p6 4s2 3d1
4s is before the 3d

10. There is repulsion between paired electrons
Arranging electrons
Example: 6: Vanadium (23 electrons)
1s2 2s2 2p6 3s2 3p6 4s2 3d3
You would expect chromium (24 electrons) to have an electron configuration of: 1s2 2s2 2p6 3s2 3p6 4s2 3d4
Instead the 4s electron is promoted to the 3d sub-level
This means both the 4s and 3d sub-level are half full and this is more stable
1s2 2s2 2p6 3s2 3p6 4s1 3d5
There is repulsion between paired electrons

11. Arranging electrons Example: 7: Manganese (25 electrons)
1s2 2s2 2p6 3s2 3p6 4s2 3d5
Example: 8: Nickel (28 electrons)
1s2 2s2 2p6 3s2 3p6 4s2 3d8
You would expect copper (29 electrons) to have an electron configuration of: 1s2 2s2 2p6 3s2 3p6 4s2 3d9
Instead the 4s electron is promoted to the 3d sub-level
The 3d sub-level is full and this is more stable
1s2 2s2 2p6 3s2 3p6 4s1 3d10

12. Electron configuration
Transition Metals
Element
Atomic number/
Electron number
Electron configuration
Scandium 21
1s2 2s2 2p6 3s2 3p6 4s2 3d1
Titanium 22
1s2 2s2 2p6 3s2 3p6 4s2 3d2
Vanadium 23
1s2 2s2 2p6 3s2 3p6 4s2 3d3
Chromium 24
1s2 2s2 2p6 3s2 3p6 4s1 3d5
Manganese 25
1s2 2s2 2p6 3s2 3p6 4s2 3d5
Iron 26
1s2 2s2 2p6 3s2 3p6 4s2 3d6
Cobalt 27
1s2 2s2 2p6 3s2 3p6 4s2 3d7
Nickel 28
1s2 2s2 2p6 3s2 3p6 4s2 3d8
Copper 29
1s2 2s2 2p6 3s2 3p6 4s1 3d10

13. Blocks in the Periodic Table
The Periodic Table is divided into four blocks
p block
s block
d block
f block

14. Transition Metal Ions Transition metals form positive ions (d block)
Electrons are not lost from the 3d sub-shell
Electrons are lost from the 4s sub-shell first then the 3d
Example:
Iron, Fe (26 electrons)
1s2 2s2 2p6 3s2 3p6 4s2 3d6
Iron ion, Fe2+(24 electrons)
1s2 2s2 2p6 3s2 3p6 3d6

15. Electron configuration of atom Electron configuration of ion
Transition Metal Ions
Atom Z
Electron configuration of atom
Ion
Electron configuration of ion Fe 26
1s2 2s2 2p6 3s2 3p6 4s2 3d6
Fe2+
1s2 2s2 2p6 3s2 3p6 3d6 Ti 22
1s2 2s2 2p6 3s2 3p6 4s2 3d2
Ti2+
1s2 2s2 2p6 3s2 3p6 3d2 V 23
1s2 2s2 2p6 3s2 3p6 4s2 3d3
V3+ Cr 24
1s2 2s2 2p6 3s2 3p6 4s1 3d5
Cr3+
1s2 2s2 2p6 3s2 3p6 3d3 Cu 29
1s2 2s2 2p6 3s2 3p6 4s1 3d10
Cu2+
1s2 2s2 2p6 3s2 3p6 3d9 Mn 25
1s2 2s2 2p6 3s2 3p6 4s2 3d5
Mn7+
1s2 2s2 2p6 3s2 3p6
Electrons are lost from the 4s sub-shell first then the 3d

16. Transition Metals
The d block in the periodic table contains the transition metals
Definition: A transition metal is an element whose atom or ion has an incomplete d sub-level
Example 1: Scandium, Sc (21 electrons)
1s2 2s2 2p6 3s2 3p6 4s2 3d1
Example 2: Iron, Fe (26 electrons)
1s2 2s2 2p6 3s2 3p6 4s2 3d6

17. The d sub-level is complete when it has 10 electrons
Transition Metals
Definition: A transition metal is an element whose atom or ion has an incomplete d sub-level
Example 3: Zinc, Zn (30 electrons)
1s2 2s2 2p6 3s2 3p6 4s2 3d10
Zinc(II), Zn2+
1s2 2s2 2p6 3s2 3p6 3d10
Zinc is in the d block but it is not a transition metal
This is because zinc or its ions do not have an incomplete d sub-level
The d sub-level is complete when it has 10 electrons

18. Transition Metals There are four properties of transition metals:
Catalytic activity
Variable oxidation states
The ability to form complexes
The ability to form coloured ions
Transition metals are more stronger, denser and less reactive compared to the metals in groups 1 and 2