Transition metals are the workhorses of chemistry - they're everywhere...
Chemistry433 views·Updated 23 Aug 2026·4 pagesUnderstanding Transition Metals in Chemistry
Katie Rose@katierose
1 / 4
1
of 4
Electronic Configuration and Properties
Transition metals sit between groups 2 and 3 on the periodic table and are called d-block elements because they have an incomplete d subshell in at least one of their ions. This is what gives them their special properties that other metals just don't have.
Most d orbitals fill up following the normal aufbau principle, but chromium and copper are rebels - they break the rules because having a half-filled or completely filled d subshell is more stable. When these metals form ions, they always lose their 4s electrons first, not the 3d ones.
Oxidation states are basically how many electrons an atom has gained or lost when it's part of a compound. Transition metals are brilliant because they can have multiple oxidation states, which explains why they form compounds with different colours.
Quick Tip: Remember that oxidation numbers in a neutral compound always add up to zero - this is your go-to check when working out oxidation states!
2
of 4
Oxidation States and Redox Chemistry
Working out oxidation numbers is actually straightforward once you know the rules. Elements on their own are always zero, single-atom ions equal their charge, oxygen is usually -2, and hydrogen is usually +1.
Transition metals can exist in different oxidation states within their compounds, and this flexibility makes them incredibly useful. High oxidation states often make compounds that are oxidising agents, while low oxidation states tend to create reducing agents.
The stability of different oxidation states depends on electronic structure, bonding type, and stereochemistry. This is why some oxidation states are more common than others for each metal.
Remember: Oxidation = increase in oxidation number, Reduction = decrease in oxidation number. OIL RIG still works!
3
of 4
Ligands and Complex Formation
Ligands are electron donors that attach to transition metal ions to form coordination compounds or complexes. Think of them as molecules or ions with lone pairs of electrons that can form coordinate bonds with the metal.
Ligands are classified by how many bonds they can form - monodentate (one bond), bidentate (two bonds), up to hexadentate (six bonds). The coordination number tells you how many bonds the central metal ion has to ligands, which determines the shape of the complex.
Naming complexes follows specific rules that might seem complicated at first, but they're logical. You write ligands alphabetically, add prefixes for multiples (di, tri, tetra), and change endings for negative ion ligands to end in 'o'. Common ligands like water become 'aqua' and ammonia becomes 'ammine'.
Pro Tip: The formula goes in square brackets with the overall charge outside - this makes it clear what's part of the complex!
4
of 4
Colour and Catalysis
The amazing colours of transition metal complexes come from d-d transitions. When ligands approach the metal ion, they split the d orbitals into different energy levels. Electrons can jump between these levels when they absorb light energy.
The spectrochemical series ranks ligands by their ability to split d orbitals. Strong field ligands like CN⁻ cause large energy gaps, while weak field ligands like I⁻ cause smaller gaps. You see the complementary colour to whatever's absorbed.
Transition metals make excellent catalysts because of their variable oxidation states and available d orbitals. Heterogeneous catalysts are in a different state to reactants, whilst homogeneous catalysts are in the same state. Both work by providing alternative reaction pathways with lower activation energies.
Key Point: Catalyst poisoning happens when substances bind irreversibly to active sites, so industrial processes need to be carefully controlled!
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Chemistry433 views·Updated 23 Aug 2026·4 pagesUnderstanding Transition Metals in Chemistry
Katie Rose@katierose
Transition metals are the workhorses of chemistry - they're everywhere from the coins in your pocket to the catalysts that make industrial processes possible. These d-block elements have some seriously cool properties that make them stand out from other metals,...
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Electronic Configuration and Properties
Transition metals sit between groups 2 and 3 on the periodic table and are called d-block elements because they have an incomplete d subshell in at least one of their ions. This is what gives them their special properties that other metals just don't have.
Most d orbitals fill up following the normal aufbau principle, but chromium and copper are rebels - they break the rules because having a half-filled or completely filled d subshell is more stable. When these metals form ions, they always lose their 4s electrons first, not the 3d ones.
Oxidation states are basically how many electrons an atom has gained or lost when it's part of a compound. Transition metals are brilliant because they can have multiple oxidation states, which explains why they form compounds with different colours.
Quick Tip: Remember that oxidation numbers in a neutral compound always add up to zero - this is your go-to check when working out oxidation states!
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Oxidation States and Redox Chemistry
Working out oxidation numbers is actually straightforward once you know the rules. Elements on their own are always zero, single-atom ions equal their charge, oxygen is usually -2, and hydrogen is usually +1.
Transition metals can exist in different oxidation states within their compounds, and this flexibility makes them incredibly useful. High oxidation states often make compounds that are oxidising agents, while low oxidation states tend to create reducing agents.
The stability of different oxidation states depends on electronic structure, bonding type, and stereochemistry. This is why some oxidation states are more common than others for each metal.
Remember: Oxidation = increase in oxidation number, Reduction = decrease in oxidation number. OIL RIG still works!
3
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Ligands and Complex Formation
Ligands are electron donors that attach to transition metal ions to form coordination compounds or complexes. Think of them as molecules or ions with lone pairs of electrons that can form coordinate bonds with the metal.
Ligands are classified by how many bonds they can form - monodentate (one bond), bidentate (two bonds), up to hexadentate (six bonds). The coordination number tells you how many bonds the central metal ion has to ligands, which determines the shape of the complex.
Naming complexes follows specific rules that might seem complicated at first, but they're logical. You write ligands alphabetically, add prefixes for multiples (di, tri, tetra), and change endings for negative ion ligands to end in 'o'. Common ligands like water become 'aqua' and ammonia becomes 'ammine'.
Pro Tip: The formula goes in square brackets with the overall charge outside - this makes it clear what's part of the complex!
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Colour and Catalysis
The amazing colours of transition metal complexes come from d-d transitions. When ligands approach the metal ion, they split the d orbitals into different energy levels. Electrons can jump between these levels when they absorb light energy.
The spectrochemical series ranks ligands by their ability to split d orbitals. Strong field ligands like CN⁻ cause large energy gaps, while weak field ligands like I⁻ cause smaller gaps. You see the complementary colour to whatever's absorbed.
Transition metals make excellent catalysts because of their variable oxidation states and available d orbitals. Heterogeneous catalysts are in a different state to reactants, whilst homogeneous catalysts are in the same state. Both work by providing alternative reaction pathways with lower activation energies.
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