Chemistry (9701) | General physical and chemical properties of the first row of transition elements, titanium to copper

A-Level Chemistry (9701): Transition Elements (Titanium to Copper)

Hello future Chemists! This chapter takes us into the fascinating world of the Transition Elements, specifically the first row: Titanium (Ti) through Copper (Cu). These elements are incredibly important in industry and biology, largely because they behave very differently from the Group 1 and 2 metals you studied previously.

Don't worry if the details seem numerous—we'll break down the key characteristics (like colour and catalysis) into simple, understandable steps. Ready to dive in?

1. Defining the Transition Elements (LO 28.1)

What is a Transition Element?

A transition element is officially defined as a d-block element which forms one or more stable ions with incomplete d orbitals.

(Wait, what about Zinc? Zinc (Zn) is in the d-block, but its ion, Zn²⁺, has a full d-subshell (3d¹⁰). Therefore, zinc is usually not defined as a true transition element according to the syllabus definition, although it is still often grouped with them physically.)

Electron Configuration: The 3d and 4s Sub-shells

The first row of transition metals involves filling the 3d and 4s sub-shells.

• For atoms, the 4s sub-shell fills before the 3d sub-shell (e.g., K: [Ar] 4s¹, Ca: [Ar] 4s²).
• However, when these metals form ions, the electrons are always removed from the 4s orbital first because the 3d orbital is lower in energy in the resulting ion.
  Example: Fe (atom) is [Ar] 3d⁶ 4s². Fe²⁺ (ion) is [Ar] 3d⁶ 4s⁰.

Key Takeaway: Transition elements are defined by their ability to form ions with incomplete d orbitals, which leads to all their unique properties.

2. The Four Unique Characteristics (LO 28.1.3)

Transition elements display four key properties that distinguish them from typical Group 1 or 2 metals:

1. Variable Oxidation States
2. Ability to act as Catalysts
3. Formation of Complex Ions
4. Formation of Coloured Compounds

2.1 Variable Oxidation States (LO 28.1.4)

Unlike Group 1 metals (which are always +1) or Group 2 metals (always +2), transition metals can exist in many stable oxidation states (O.S.).

The Explanation: The 3d and 4s sub-shells have very similar energies. This means that a relatively small amount of energy is required to remove not only the 4s electrons but also varying numbers of the 3d electrons, leading to different stable ions (e.g., Fe²⁺ and Fe³⁺).

• Did you know? Manganese (Mn) has the largest range, from +2 to +7!
• Common examples: Iron (Fe²⁺, Fe³⁺), Copper (Cu⁺, Cu²⁺), Vanadium (V²⁺, V³⁺, V⁴⁺, V⁵⁺).

2.2 Catalytic Behaviour (LO 28.1.5)

Transition metals and their compounds are excellent catalysts, either homogeneous (same phase as reactants) or heterogeneous (different phase).

The Explanation: The catalytic power comes from two factors:

1. Variable Oxidation States: They can easily change their O.S. (e.g., Fe²⁺ to Fe³⁺), allowing them to transfer electrons between reactants efficiently, providing a pathway with lower activation energy (\(E_a\)).
2. Vacant d Orbitals: They have accessible vacant d orbitals which can form dative bonds with reactant molecules (ligands) during the reaction mechanism. This is particularly important for heterogeneous catalysis.

Real-World Analogy: A catalyst is like a shortcut on a map. Transition metals are the best shortcuts because they are flexible (variable O.S.) and have many 'parking spots' (vacant d orbitals) where reactants can temporarily bond.

• Heterogeneous Example: Iron (Fe) used in the Haber Process.
• Homogeneous Example: Fe²⁺ or Fe³⁺ used to speed up the reaction between iodide (I⁻) and peroxodisulfate (\(S_2O_8^{2-}\)). (LO 26.2.3b)

Key Takeaway: Variable oxidation states and the availability of 3d orbitals are the root causes of both catalysis and variable O.S.

3. Complex Ions and Ligands (LO 28.2)

3.1 Definitions

A complex is a molecule or ion formed by a central metal atom or ion surrounded by one or more ligands.

A ligand is a species (ion or molecule) that contains a lone pair of electrons that forms a dative covalent bond (coordinate bond) to the central metal atom/ion.

Why form Complexes? Transition metals have energetically accessible vacant d orbitals (LO 28.1.6) which can accept the lone pair of electrons from the ligand, forming the dative covalent bond.

3.2 Types of Ligands (LO 28.2.3)

Ligands are classified by their denticity (how many dative bonds they form to the central ion).

1. Monodentate Ligands: Form one dative bond.
   Examples: \(H_2O\), \(NH_3\), \(Cl^-\), \(CN^-\).
2. Bidentate Ligands: Form two dative bonds.
   Examples: 1,2-diaminoethane (en, \(H_2NCH_2CH_2NH_2\)), and ethanedioate ion (\(C_2O_4^{2-}\)).
3. Polydentate Ligands: Form multiple dative bonds (more than two).
   Example: EDTA⁴⁻ (Ethylenediaminetetraacetate ion). This forms six dative bonds!

Memory Aid: Think of "dentist" – teeth are used for gripping. Denticity refers to how many 'grips' (dative bonds) the ligand has.

3.3 Coordination Number (C.N.) and Geometry (LO 28.2.5, 6)

The coordination number is the total number of dative bonds formed between the ligands and the central metal ion. This determines the complex geometry.

Coordination Number (C.N.)	Geometry (Shape)	Example
2	Linear	[Ag(NH₃)₂]⁺ (often silver, not Ti-Cu)
4	Tetrahedral	[CuCl₄]²⁻, [CoCl₄]²⁻
4	Square Planar	[Pt(NH₃)₂Cl₂] (Cisplatin)
6	Octahedral	[Cu(H₂O)₆]²⁺, [Co(NH₃)₆]³⁺

Predicting Formula and Charge (LO 28.2.6):
The complex charge is simply the sum of the charge on the central metal ion and the charges of all the ligands.
Example: A Cobalt(II) ion (\(Co^{2+}\)) surrounded by four chloride ions (\(Cl^-\)) in a tetrahedral arrangement (C.N. 4).
Metal charge: +2. Ligand charge: 4 x (-1) = -4.
Total Complex Charge: +2 + (-4) = -2. Formula: \([CoCl_4]^{2-}\).

4. Ligand Exchange Reactions (LO 28.2.1, 7)

Ligand exchange occurs when one type of ligand surrounding the central metal ion is replaced by another ligand. This is usually accompanied by a dramatic colour change.

We must know the specific reactions for Copper(II) and Cobalt(II).

Copper(II) Ions (\(Cu^{2+}\))

Aqueous copper(II) ions are normally blue: \([Cu(H_2O)_6]^{2+}(aq)\).

1. Addition of aqueous Ammonia (\(NH_3\)):
   Small amounts of \(NH_3\) first precipitate pale blue copper hydroxide:
   \([Cu(H_2O)_6]^{2+} + 2NH_3 \rightarrow [Cu(H_2O)_4(OH)_2] + 2NH_4^{+}\) (Pale blue ppt.)
   Excess \(NH_3\) causes ligand exchange, dissolving the precipitate to form a deep blue solution:
   \([Cu(H_2O)_4(OH)_2] + 4NH_3 \rightarrow [Cu(NH_3)_4(H_2O)_2]^{2+} + 2OH^- + 2H_2O\) (Deep blue solution)
2. Addition of concentrated Chloride ions (\(Cl^-\)) (e.g., concentrated HCl):
   This is often a change in coordination number (from 6 to 4) and geometry (octahedral to tetrahedral).
   \([Cu(H_2O)_6]^{2+} + 4Cl^- \rightleftharpoons [CuCl_4]^{2-} + 6H_2O\)
   Colour change: Blue solution changes to yellow/green.

Cobalt(II) Ions (\(Co^{2+}\))

Aqueous cobalt(II) ions are normally pink/pale red: \([Co(H_2O)_6]^{2+}(aq)\).

1. Addition of aqueous Ammonia (\(NH_3\)):
   Small amounts form a grey-green precipitate (cobalt hydroxide).
   Excess \(NH_3\) causes ligand exchange to form a complex ion that is often straw-coloured/brown: \([Co(NH_3)_6]^{2+}\) or \([Co(NH_3)_6]^{3+}\) (as it may be air-oxidised).
2. Addition of concentrated Chloride ions (\(Cl^-\)):
   Change from octahedral (C.N. 6) to tetrahedral (C.N. 4).
   \([Co(H_2O)_6]^{2+} + 4Cl^- \rightleftharpoons [CoCl_4]^{2-} + 6H_2O\)
   Colour change: Pink/pale red changes to deep blue. (This reaction is useful as it is reversible with heat/water).

5. Stability Constants (\(K_{stab}\)) and Ligand Exchange (LO 28.5)

5.1 Defining \(K_{stab}\)

The stability constant, \(K_{stab}\), is the equilibrium constant for the formation of a complex ion in a solvent from its constituent ions or molecules. It tells us how stable a complex is.

For the general reaction (often simplified for the syllabus):
\(M^{n+}(aq) + xL \rightleftharpoons [ML_x]^{n+}\)

The expression for \(K_{stab}\) is:
$$K_{stab} = \frac{[ML_x]^{n+}}{[M^{n+}][L]^x}$$

• Crucially: The concentration of the solvent, water (\([H_2O]\)), is not included in the \(K_{stab}\) expression, even though water molecules are usually the initial ligands being replaced.
• A large \(K_{stab}\) value means the equilibrium lies far to the right, indicating a very stable complex.

5.2 \(K_{stab}\) and Ligand Exchange

Ligand exchange reactions occur because one complex is more stable than the other.

When an aqueous ion \([M(H_2O)_6]^{n+}\) reacts with a new ligand (L), the reaction proceeds if the complex formed with L has a significantly larger \(K_{stab}\) than the water complex.

For polydentate ligands (like EDTA⁴⁻), the resulting complex is often exceptionally stable due to the chelate effect, leading to very large \(K_{stab}\) values, making these exchanges highly favourable.

Key Takeaway: Ligand exchange is driven by the formation of a more stable complex, reflected by a larger equilibrium constant, \(K_{stab}\).

6. The Origin of Colour in Complexes (LO 28.3)

The most striking feature of transition metals is their ability to form brightly coloured compounds.

6.1 Degenerate d Orbitals and Splitting

In an isolated transition metal atom or ion, the five d orbitals all have the same energy; they are called degenerate orbitals.

When ligands approach the central ion to form a complex, the repulsion between the electrons of the ligands and the electrons in the d orbitals causes the d orbitals to split into two non-degenerate sets of different energy.

• The energy difference between these sets is called \(\Delta E\) (or the crystal field splitting energy).
• Octahedral complexes: Split into two higher energy orbitals and three lower energy orbitals.
• Tetrahedral complexes: Split into three higher energy orbitals and two lower energy orbitals.

Analogy: Imagine five identical people (d-orbitals) standing on a flat floor (degenerate). When a crowd (ligands) rushes towards them, they jostle and climb onto a small staircase (non-degenerate sets) to avoid the crowd. They are no longer all at the same energy level.

6.2 Absorption and Colour

When white light (which contains all colours) shines on the complex ion, an electron from a lower energy d orbital is promoted to a higher energy d orbital.

The energy required for this jump (\(\Delta E\)) corresponds to the energy of a specific frequency of visible light. This specific frequency (colour) is absorbed.

The colour we observe is the combination of the remaining colours in the white light spectrum—this is the complementary colour to the one absorbed.

• Example: If a complex absorbs yellow light, the remaining light (mostly blue/violet) is transmitted, so the complex appears purple/blue.

6.3 Factors Affecting Colour (LO 28.3.4)

The magnitude of \(\Delta E\) determines which frequency (colour) of light is absorbed, and therefore what colour is observed. \(\Delta E\) is affected by:

1. The nature of the ligand: Different ligands cause different amounts of splitting (\(\Delta E\)). Strong field ligands (like \(CN^-\) or \(NH_3\)) cause large splitting, while weak field ligands (like \(Cl^-\)) cause small splitting.
   Demonstration: Ligand exchange changes the colour (e.g., \([Cu(H_2O)_6]^{2+}\) (blue) vs \([CuCl_4]^{2-}\) (yellow/green)).
2. The oxidation state of the metal ion: Higher O.S. usually leads to a larger \(\Delta E\).
3. The coordination number/geometry: Tetrahedral splitting is generally smaller than octahedral splitting.

Key Takeaway: Colour results from the specific energy of light absorbed (\(\Delta E\)) required to promote a d-orbital electron. Different ligands change \(\Delta E\), changing the colour.

7. Stereoisomerism in Complexes (LO 28.4)

Transition metal complexes, particularly those with a Coordination Number of 6 (octahedral) or certain C.N. 4 (square planar), can exhibit stereoisomerism.

7.1 Geometrical Isomerism (Cis/Trans) (LO 28.4.1a)

This occurs when the ligands occupy different spatial positions relative to each other.

• Cis: Identical ligands are next to each other (at 90°).
• Trans: Identical ligands are opposite each other (at 180°).

Square Planar Examples: \([Pt(NH_3)_2Cl_2]\) (C.N. 4).
The cis- isomer (Cisplatin) is a famous anti-cancer drug, while the trans- isomer is biologically inactive. Their polarity differs: the cis- isomer is polar (dipoles don't cancel), while the trans- isomer is non-polar (dipoles cancel). (LO 28.4.2)

Octahedral Examples: \([Co(NH_3)_4(H_2O)_2]^{2+}\) or \([Ni(H_2NCH_2CH_2NH_2)_2(H_2O)_2]^{2+}\).

7.2 Optical Isomerism (LO 28.4.1b)

This occurs in complexes that are non-superimposable mirror images of each other (like your hands). These complexes must lack a plane of symmetry.

It is most common in octahedral complexes involving bidentate ligands (like 'en', 1,2-diaminoethane). The bidentate ligands wrap around the central metal ion, creating a chiral shape.

Octahedral Examples: \([Ni(H_2NCH_2CH_2NH_2)_3]^{2+}\) or \([Ni(H_2NCH_2CH_2NH_2)_2(H_2O)_2]^{2+}\).

Key Takeaway: Know the difference between cis/trans (positional) and optical (non-superimposable mirror image), and which specific shapes/ligands lead to which isomerism.

8. Redox Reactions Involving Transition Elements (LO 28.2.8, 9, 10)

Because transition elements can exist in multiple stable oxidation states, they are crucial components in many redox reactions. We can use Standard Electrode Potentials (\(E^{\ominus}\)) to predict the feasibility of these reactions.

8.1 Feasibility Prediction using \(E^{\ominus}\) (LO 28.2.8)

A redox reaction is feasible (thermodynamically possible) if the cell potential (\(E^{\ominus}_{cell}\)) is positive.

$$E^{\ominus}_{cell} = E^{\ominus}_{reduction} - E^{\ominus}_{oxidation}$$

For a reaction to happen spontaneously:

• The species being reduced must have a more positive \(E^{\ominus}\) value.
• The species being oxidised must have a more negative \(E^{\ominus}\) value.

Simplified rule: A substance on the left side of a half-equation will oxidise a substance on the right side of a half-equation that is higher up (more negative \(E^{\ominus}\)) in the electrochemical series.

8.2 Important Redox Reactions and Calculations (LO 28.2.9)

You need to be able to describe and perform calculations for these specific systems:

1. Permanganate and Oxalate: \((MnO_4^- / C_2O_4^{2-})\)
   In acidic solution, the purple permanganate ion (\(MnO_4^-\)) is a very strong oxidising agent and is reduced from O.S. +7 to +2 (colourless \(Mn^{2+}\)). Oxalate (\(C_2O_4^{2-}\)) is the reducing agent, oxidised to \(CO_2\).
   Equation (Half-cells):
   Oxidation: \(C_2O_4^{2-} \rightarrow 2CO_2 + 2e^-\)
   Reduction: \(MnO_4^- + 8H^+ + 5e^- \rightarrow Mn^{2+} + 4H_2O\)
2. Permanganate and Iron(II): \((MnO_4^- / Fe^{2+})\)
   Permanganate oxidises \(Fe^{2+}\) (O.S. +2, pale green) to \(Fe^{3+}\) (O.S. +3, yellow/brown) in acidic solution. This is a common titration.
   Equation (Half-cells):
   Oxidation: \(Fe^{2+} \rightarrow Fe^{3+} + e^-\)
   Reduction: \(MnO_4^- + 8H^+ + 5e^- \rightarrow Mn^{2+} + 4H_2O\)
3. Copper(II) and Iodide: \((Cu^{2+} / I^-)\)
   Copper(II) ions (\(Cu^{2+}\)) act as the oxidising agent and are reduced to Copper(I) iodide (CuI), which is a white precipitate. Iodide ions (\(I^-\)) are oxidised to iodine (\(I_2\)).
   Equation (Overall):
   \(2Cu^{2+}(aq) + 4I^-(aq) \rightarrow 2CuI(s) + I_2(aq)\)
   (The production of iodine allows this reaction to be followed by titration with thiosulfate.)

Key Takeaway: Transition metals' variable O.S. makes them powerful redox agents. Feasibility is checked using \(E^{\ominus}\) values, where the reaction runs with the most positive overall cell potential.