Transition Metals

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

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

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

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

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

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 ↿ ⇂ ↿ ⇂ ↿ ⇂ ↿ ⇂ ↿ ⇂ ↿

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 ↿ ⇂ ↿ ⇂ ↿ ⇂ ↿ ⇂ ↿ ⇂ ↿ ⇂ ↿ ⇂ ↿ ↿

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

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

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

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

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

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

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

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

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

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