Transition elements- CIE iGCSE Chemistry Notes - New Syllabus
Transition elements for iGCSE Chemistry Notes
Core Syllabus
- Describe the transition elements as metals that:
(a) have high densities
(b) have high melting points
(c) form coloured compounds
(d) often act as catalysts as elements and in compound
Supplement Syllabus
- Describe transition elements as having ions with variable oxidation numbers, including iron(II) and iron(III)
Properties of transition elements
Properties of transition elements
- Transition elements are metals located in the d-block of the Periodic Table (groups 3–12).
- They have partially filled d-orbitals in one or more of their common oxidation states, which gives rise to characteristic properties.
Physical properties
- High densities: Atoms are closely packed in the metallic lattice and have high atomic masses → large number of protons and neutrons → high density.
- High melting and boiling points: Strong metallic bonding arises from delocalised electrons in d-orbitals → large amounts of energy required to break metallic bonds. Mercury is an exception with low melting point.
- Hardness: Most transition metals are hard due to strong metallic bonds.
Chemical properties
- Formation of coloured compounds: Partially filled d-orbitals allow for d-d electron transitions when light is absorbed → compounds appear coloured. Example: Cu²⁺ (blue), Fe³⁺ (yellow-brown), Cr³⁺ (green).
- Catalytic activity: Transition metals can act as catalysts due to variable oxidation states and ability to adsorb reactant molecules on their surfaces.
- Example: Fe catalyses the Haber process.
- Example: V₂O₅ catalyses the Contact process (conversion of SO₂ to SO₃).
- Multiple oxidation states: Partially filled d-orbitals allow transition metals to form ions in more than one oxidation state → important in redox chemistry.
Explanation of trends
- High density arises from strong metallic bonding and closely packed atoms.
- High melting points are due to strong metallic bonds from delocalised electrons including d-electrons.
- Coloured compounds arise from absorption of visible light causing d-electron transitions.
- Catalytic activity occurs because transition metals can temporarily change oxidation states, providing pathways with lower activation energy for reactions.
Example
Explain why iron has higher density and melting point than sodium.
▶️Answer/Explanation
Iron atoms are heavier and closely packed in a metallic lattice.
Metallic bonding is stronger due to partially filled d-orbitals → high melting point.
Sodium is lighter and has weaker metallic bonding → lower density and melting point.
Example
Explain why Cu²⁺ ions in solution are blue.
▶️Answer/Explanation
Cu²⁺ has partially filled d-orbitals.
Absorption of visible light promotes an electron from a lower-energy d-orbital to a higher-energy d-orbital (d-d transition) → remaining light is transmitted or reflected → solution appears blue.
Example
Give an example of a transition metal acting as a catalyst and explain how it works.
▶️Answer/Explanation
Iron (Fe) in the Haber process provides a surface for N₂ and H₂ molecules to adsorb.
Temporary bonding and electron transfer lower the activation energy → ammonia forms more quickly.
Variable oxidation numbers of transition elements
Variable oxidation numbers of transition elements
- Transition metals can form ions with more than one oxidation state due to the availability of both 4s and 3d electrons (or equivalent d-orbitals in other periods) for bonding.
- This is a key chemical property distinguishing them from main-group metals, which usually have only one common oxidation state.
Examples of variable oxidation states
- Iron: Fe²⁺ (iron(II)) and Fe³⁺ (iron(III))
- Copper: Cu⁺ (copper(I)) and Cu²⁺ (copper(II))
- Chromium: Cr²⁺, Cr³⁺, Cr⁶⁺
- Manganese: Mn²⁺, Mn³⁺, Mn⁴⁺, Mn⁷⁺
Explanation
- Partially filled d-orbitals allow the metal to lose different numbers of electrons → multiple oxidation states.
- The stability of each oxidation state depends on electronic configuration, lattice energy, hydration energy, and ligand effects.
- Iron example: Fe²⁺ has configuration [Ar] 3d⁶, Fe³⁺ has [Ar] 3d⁵ → both are reasonably stable in aqueous solutions.
Importance of variable oxidation states
- Allows transition metals to participate in redox reactions and act as catalysts.
- Responsible for the formation of coloured compounds as different oxidation states absorb different wavelengths of visible light.
- Enables compounds to exhibit magnetic properties and varying chemical reactivity.
Example
Explain why iron forms both Fe²⁺ and Fe³⁺ ions.
▶️Answer/Explanation
Iron has the electron configuration [Ar] 4s² 3d⁶. It can lose two electrons to form Fe²⁺ ([Ar] 3d⁶) or three electrons to form Fe³⁺ ([Ar] 3d⁵). Both ions are stable in aqueous solutions, giving iron variable oxidation states.
Example
State the oxidation states of copper and explain why it forms coloured compounds.
▶️Answer/Explanation
Copper forms Cu⁺ and Cu²⁺ ions. Cu²⁺ has partially filled d-orbitals → d-d electron transitions → absorbs visible light → solutions appear blue (e.g., CuSO₄ solution).
Example
Explain how variable oxidation states allow manganese to act as an oxidising agent.
▶️Answer/Explanation
Manganese can exist in multiple oxidation states (Mn²⁺, Mn³⁺, Mn⁷⁺). Mn⁷⁺ in KMnO₄ is a strong oxidising agent because it can gain electrons and be reduced to lower oxidation states such as Mn²⁺ or Mn⁴⁺.