1. Year 11 DP Chemistry Rob Slider
Periodicity (AHL)
Year 11 DP Chemistry
Rob Slider

2. First row d-block
The d-block elements are in the middle of the Periodic Table and include the transition metals. Starting in period 4 after the 4s fills, the 3d subshell begins to fill with electrons
A transition metal (TM) is an element that has at least one ion with a partially filled d-subshell. Not all d-block elements are TM. Zn is not considered to be TM (more later...)

3. Properties of TM
Due to the partially filled d-subshell, TM have unique properties including:
Multiple oxidation states
Complex ion formation with ligands
Formation of coloured compounds
Catalytic properties
Magnetic properties depending on oxidation states and coordination numbers

4. Electronic configurationsSc
[Ar] 3d14s2 Ti
[Ar] 3d24s2 V
[Ar] 3d34s2 Cr
[Ar] 3d54s1 Mn
[Ar] 3d54s2 Fe
[Ar] 3d64s2 Co
[Ar] 3d74s2 Ni
[Ar] 3d84s2 Cu
[Ar] 3d104s1 Zn
[Ar] 3d104s2
As we have seen previously, the configurations of the first row d-block mostly fill the 3d subshell in order.
The exceptions come from Cr and Cu where we see more stable configurations from the half-filled and filled 3d subshell.
This is possible because the 4s and 3d subshells are so similar in energy

5. Sc and Zn (not TM?)
Sc forms Sc3+ which has the stable configuration of Ar
Sc3+ has no 3d electrons, therefore it is not considered to be a TM
Note: the discovery of new compounds has resulted in a change in accepted understanding for Sc. A +2 oxidation state can also exist, so Sc is now considered to be TM.
Zn has a configuration of [Ar]3d104s2,
The Zn2+ ion ([Ar] 3d10), therefore is not a typical TM ion

6. Variable oxidation states+2
All transition metals can form the oxidation state of +2 due to the loss of the two s-electrons. In the first row, the 4s.
This is because the 4s fills first, but when ions are being formed, the 4s electrons are also lost first.
Examples:
To write the electronic structure for Co2+:
Co [Ar] 3d74s2
Co2+ [Ar] 3d7
The 2+ ion is formed by the loss of the two 4s electrons
To write the electronic structure for V3+:
V [Ar] 3d34s2
V3+ [Ar] 3d2
The 4s electrons are lost first, then one of the 3d electrons

7. Variable oxidation states
Due to the similar energy levels of the 4s and 3d, other oxidation states in addition to +2 are also possible.
On the left, all of the electrons from the 4s and 3d can be lost forming ions such as Sc3+ and Ti4+. This represents the largest possible OS, so high OS are stable.
On the right, the nucleus has a stronger pull on the outer electrons due to a greater positive charge. This means that +2 is the most stable as there is a greater energy difference between the 3d and 4s(Co, Ni, Cu)
Cu also forms +1 due to the formation of the stable [Ar]3d10

8. In the middle
V, Cr, Mn
It requires too much energy to remove all of the electrons from these elements as the number of valence electrons and high nuclear charge increases. However, they do have stable high oxidation states.
What often occurs is the formation of stable oxyanions, such as VO3-, vanadate(V). Some important ones to note:
Oxidation state
Vanadium
Chromium
Manganese +7
MnO4- permanganate +6
CrO42- chromate Cr2O72- dichromate +5
VO3- +4
MnO2 +3
Cr3+
Note there are other oxidation states possible for these elements. For example Mn can also form +3 (Mn2O3), +6(MnO4 2-). However, the IB course does not specifically require knowledge of these states. Simple ions with high OS’s such as V5+ & Mn7+ are not formed

9. Summary of oxidation statesSc Ti V Cr Mn Fe Co Ni Cu Zn +1 +2 +3 +4 +5 +6 +7
Boxed states are the important ones to know about

10. Summary of 1st d-block OS
Higher OS’s become less stable relative to lower ones on moving from left to right across the series(stronger + nuclear force)
On the right, +2 is more common; +3 state highly oxidising.
E.g. Co3+ is a strong oxidising agent (gain e- easily), Ni3+ & Cu3+ do not exist in aqueous solution.
On the left, +2 state highly reducing.
e.g. V2+(aq) , Cr2+(aq) are strong reducing agents (lose e- easily)
Compounds containing TM’s in high OS’s tend to be oxyanions and oxidising agents e.g. MnO4-

11. Complex ion formation
Complex ions have a metal ion at the centre with a number of other molecules or ions surrounding it.
The bonds are coordinate bonds where a lone pair on a molecule/ion is donated to a low energy, unfilled metal orbital such as a d-orbital. These molecules/ions are called ligands.
These are some common ligands found in complex ion formation. Notice they all have lone pairs of electrons
Due to coordinate bonds, these compounds are also referred to as coordinate complexes. In addition to the ligands above, CN- is also another important one.
Ligands are neutral molecules or anions that contain a non-bonding pair of electrons

12. Complex ions Coordination number:
The most common complex ions contain 4 or 6 ligands. These are known as 4-coordinated and 6-coordinated. 2 is also possible.
The charge on the complex is the sum of the metal and the ligands. The Cr is 2+ and the water is neutral leading to a 2+ complex ion charge.
Water forms hexahydrated complex ions (6-coordinated) with most transition metals.
Example: [Fe(H2O)6]3+
Many complex ions form coloured solutions
You try:
Fe(III) + CN- (6-coord)
Cu (II) + Cl- (4-coord)
Ag+ + NH3 (2-coord)

13. Common coordination numbers
There are some common coordination numbers for metal ions depending on:
Size of the metal ion
Charge of the metal ion
Size of the ligand
Charge of the ligand
Ion
Tendency to form and examples
Cu2+
Mostly 6-coordinated complexes (e.g. with water) due to its small size
Will also form 4-coordinated complexes with larger ligands such as Cl-
Ag+
Due to its larger size and single charge, mostly forms 2-coordinated complexes
Ni2+
Similar size to copper, so has same tendency to form 6-coordinated complexes (e.g. with ammonia)
Will form 4-coordinate bonds with ions that have a high relative amount of negative charge such as CN-

14. Complex ion geometry 2-coord complexes form linear geometries
4-coord complexes form tetrahedral or square planar geometries
Source of atom diagrams:
6-coord complexes tend to form octahedral geometries

15. Mono/polydentate ligands
Monodentate Ligands
Polydentate Ligands (Chelates)
These form one coordinate bond with the central atom
Below is a central atom with 6 water molecule ligands each bonded at one site only.
These form two or more coordinate bonds with the central atom
Ethylene diamine forms 2 bonds with a central metal ion (see below)
Mono/polydentate ligands

16. Notice the EDTA ligand can form 6 coordinate bonds with 2 N atoms and 4 O atoms that have bonding lone pairs.
This makes EDTA a very useful substance that has medical benefits (e.g. lead poisoning), water softening capabilities as well as a preservative. (See Oxford text for details)
For more on Chelates, see: “Chemistry The Central Science” page at

17. Complex compounds
Complex ions can be anions or cations and will bond with oppositely charged ions to make salts. Notice how [Cu(NH3)4]2+ is formed:
(CuCl4)2- is an anion that can form a compound with K+ to form [K2(CuCl4)]
Some complexes do not form ions because the charges on the metal and the ligand cancel each other. Example: [Pt(NH3)2(Cl)2]
Would you expect these two compounds to be soluble in water?
This complex ion can then bond with Cl- to form [Cu(NH3)4]Cl2
Yes, they are soluble in water.

18. How does the metal attract so many ligands??
You may be wondering why a metal ion will attract more ligands than it has charges. +2 should attract -2 and +3 should attract -3, right??
Let’s look at an example:
Fe(H2O)6 3+
Fe: 1s22s22p63s23p63d64s2
Fe3+: 1s22s22p63s23p63d5
In Fe3+, the 4s is now empty and there are 5 unpaired e-. You might expect 5 ligands, but the ion uses six orbitals from the 4s, 4p and 4d to accept lone pairs from six water molecules. It hybridises six new orbitals all with the same energy.
Note: the 3d e- are not used in bonding as the water molecules form coordinate bonds using the empty 4s, 4p and 2 of the 4d orbitals that have been hybridised.
Why not 4 or 8? Six is the maximum number of water molecules it is possible to fit around an iron ion (and most other metal ions). By making the maximum number of bonds, it releases most energy and so becomes most energetically stable.

19. Isomers (cis,trans)
Metal complexes sometimes have more than one type of ligand attached. This leads to possible isomerism with complexes having different ligand arrangements. These are called stereoisomers.
cis
When ligands are adjacent to each other they are said to be cis-
trans
When ligands are opposite to each other they are said to be trans-
Molecule diagram source:
cis-[CoCl2(NH3)4]+
trans-[CoCl2(NH3)4]+

20. Optical isomers
Some isomers are mirror images of one another. Therefore, they cannot be superimposed on one another. These two mirror image compounds are known as optical isomers or enantiomers.
Easiest way to visualise this is to look at the human hands which are mirror images of each other. They are called optical isomers because they rotate the plane of polarised light. Chemically, they behave the same unless they are interacting with other optically active molecules. The complexes above show bidentate ligands (attached at 2 locations on the metal).

21. Perceived colours Try this complementary colour interactive
When a solution absorbs some of the wavelengths of light, what we see is a mixture of the remaining wavelengths.
The solution below is [Cu(NH3)4]2+ which absorbs red-orange leaving us to perceive the complementary colours blue-green
White light
Blue-green appearing light
Complementary colours are on opposite sides of the colour wheel
Try this complementary colour interactive

22. Coloured complexes
Many d-complexes are coloured. These characteristic colours are specific to individual ions and depend upon:
Metal oxidation state
Ligands attached
Coordination number/geometry
Source of diagrams:
Same metal/different OS
Same metal/different ligand
Same metal/different coord

23. Why coloured? d,d transitions
The d-orbitals shown above, have various arrangements around the x, y and z axes. When a 6-coord complex is formed with a d-block element, the ligands will approach along the axes of an octahedral, to minimise repulsions of bonding e-.
Source of diagrams:
The approach of the ligands raises the energy level of the d-orbitals, but the orbitals that lie on the axes (4,5 above) will experience more repulsion and thus will be a slightly higher energy level (than 1,2,3).
This means the d orbitals are split.

24. d,d transitions
Movement between the d-orbitals by the e- represents an energy change, ΔE.
Remembering ΔE=hv, a transition between d-orbitals represents a specific frequency/wavelength that is specific to a complex. The value of ΔE determines the colour(s) absorbed.
Considering the four d-block elements above, only 2 and 3 have possible transitions. They are coloured due to the excitation of e- to higher d-orbitals. This transition absorbs specific frequencies and we perceive the remaining frequencies.
Perceived colour example: Copper (II) complexes are often blue-green due to the red frequencies being absorbed
Why are Sc3+ and Zn2+ colourless?

25. Colour depends on metal ion
The lone pairs of the ligands interact more strongly with the d-orbitals of ions with higher nuclear charge. A stronger interaction means ligands get closer and the closer they get, the greater the splitting of the d-orbitals.
For example:
[Mn(H20)6]2+ – Mn2+ (absorbs green region/appears pink)
[Fe(H20)6]3+ – Fe3+ (absorbs blue region/appears yellow-brown)
Mn(II) and Fe(III) have the same electron configuration, but a different charge. This shows that the higher charge has a greater interaction with the ligands indicated by a larger energy transition between the split d-orbitals. This is seen in the absorption of the higher energy colour blue.
Question:
Which of the 2 compounds above has a higher energy transition between d-orbitals?

26. Colour depends on metal OS
The oxidation state of the metal ion determines the number of d electrons which influences the amount of electron repulsion between the ligands and the metal ion. This affects the strength of the interactions between the metal ion and the ligands. The greater the repulsion, the greater the energy transition.
For example:
[Fe(H20)6]2+ – Fe2+ (absorbs violet region/appears green-yellow)
[Fe(H20)6]3+ – Fe3+ (absorbs blue region/appears orange-brown)
Fe(II) has one more electron than Fe(III) meaning that there is greater e- repulsion. This leads to a greater difference between the split d-orbitals energy levels which is shown by the higher energy violet colour being absorbed by Fe(II)
Question:
Which of the 2 compounds above has greater e- repulsion and therefore a higher energy transition between d-orbitals? Why?

27. Colour depends on charge density of the ligand
The greater the charge density of the ligands, the larger the split in the d-orbitals. This again has to do with the electron repulsions between ligands and the central metal ion.
For example:
[CoF6]3- – F- (has a lower charge density)
[Co(CN)6]3- – CN- (has a higher charge density)
The F ligand has a lower charge density than CN, so this would lead to the F complex absorbing a smaller wavelength of light and a higher frequency than the CN complex. The CN complex could be purple, orange, red, yellow or green.
Question:
If [CoF6]3- appears blue, estimate how the wavelengths and frequencies would be different for [Co(CN)6]3- . What colour could it be.

28. Hexa-aqua complex colours
This shows the colours of 6-coord aqua complexes of the first row d-block. Cu(I) is an exception which only forms simple colourless compounds. Notice there are no possible transitions for Sc3+ and Zn2+, so they are typically colourless.

29. Complex colours-examples

30. TM as catalysts
Transition metals and their compounds function as catalysts due to:
their ability to change oxidation state
In the metal’s ability to adsorb other substances on to their surface and activate them in the process.
Iron in the Haber Process
The Haber Process combines hydrogen and nitrogen to make ammonia using an iron catalyst.
Source of digram:
Nitrogen and hydrogen molecules are adsorbed on to the metallic iron surface. The hydrogen almost immediately splits into its component atoms by sharing or exchanging electrons with the catalyst surface

31. Catalyst examples V2O5 in the Contact Process
This is the conversion of sulfur dioxide to sulfur trioxide by passing the gaseous reactants over a solid vanadium (V) oxide
MnO2 in the decomposition of hydrogen peroxide
This speeds up the spontaneous decomposition of hydrogen peroxide by manganese (IV) oxide
Enzymatic catalysis
Fe in haemoglobin for carrying oxygen
Co in vitamin B12 to help produce red blood cells
See (Green, p92 for structures)
Nickel in the hydrogenation of C=C bonds
This reaction see the conversion of alkenes to alkanes
Catalytic converters
Pt and Pd are used to convert NOx and CO to harmless gases

32. Magnetic nature of some TM
The magnetism of a TM is determined by the number of unpaired electrons it has in its configuration. In complexes, the coordination and structure are also factors.
Because e- spin, they create a magnetic field and act like tiny magnets. When they are paired, they cancel each other out, but when unpaired, they can be attracted to an external magnetic field. These are known as paramagnetic substances.
Others contain no unpaired electrons and repel magnetic fields. These are known as diamagnetic substances. These are generally weaker than paramagnetic forces.
Ferromagnetic substances form permanent magnets that are the strongest of all types.
Types of magnetic behavior. (a) Diamagnetic; no centers (atoms or ions) with magnetic moments. (b) Simple paramagnetic; centers with magnetic moments are not aligned unless the substance is in a magnetic field. (c) Ferromagnetic; coupled centers aligned in a common direction Source: