Formation and Shapes of Complex Ions (OCR A-Level Chemistry A): Revision Notes

Formation and Shapes of Complex Ions

Introduction to complex ions

Transition metals have a unique ability to form complex ions, which is one of their most important chemical properties. When you dissolve certain metal compounds in water, they often produce coloured solutions. For example, when hydrated copper(II) sulfate ($\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$) dissolves in water, it creates a distinctive blue solution containing the complex ion $[\text{Cu}(\text{H}_2\text{O})_4]^{2+}$.

What are complex ions?

A complex ion forms when one or more molecules or negatively charged ions bond to a central metal ion. The molecules or ions that attach themselves to the metal are called ligands. Understanding how these ligands interact with the metal ion is fundamental to grasping complex ion chemistry.

Ligands and coordinate bonds

A ligand is a molecule or ion that donates a pair of electrons to a central metal ion, forming what we call a coordinate bond (also known as a dative covalent bond). This type of bond is special because both electrons in the shared pair come from the same atom - in this case, from the ligand rather than being contributed by both bonding partners.

Coordination number

The coordination number tells you how many coordinate bonds are attached to the central metal ion. This number is crucial because it determines the shape of the complex ion. The most common coordination numbers you'll encounter are 6 and 4, though other values are possible.

Representing complex ions using formulas

When writing the formula of a complex ion, chemists use a specific notation system that provides important information about the structure:

• The entire complex ion is enclosed in square brackets
• The ligands are written inside the brackets, shown in round brackets with a subscript indicating how many of each ligand are present
• The overall charge of the complex is written outside the square brackets as a superscript

Types of ligands

Ligands can be classified based on how many lone pairs of electrons they can donate to the central metal ion. Understanding the difference between monodentate and bidentate ligands is essential for A-Level chemistry.

Monodentate ligands

A monodentate ligand is able to donate just one pair of electrons to a central metal ion. The word "monodentate" comes from "mono" (one) and "dentate" (tooth), suggesting the ligand has one "binding site".

Common monodentate ligands include both neutral molecules and negatively charged ions:

Notice that neutral ligands like water ($\text{H}_2\text{O}$) and ammonia ($\text{NH}_3$) have no charge, while ionic ligands like chloride ($\text{Cl}^-$), cyanide ($\text{CN}^-$), and hydroxide ($\text{OH}^-$) carry a negative charge. This affects the overall charge of the resulting complex ion.

Bidentate ligands

Some ligands can donate two lone pairs of electrons to the central metal ion, forming two coordinate bonds. These are called bidentate ligands (from "bi" meaning two). The most common bidentate ligands you need to know are:

• 1,2-diaminoethane (often abbreviated as "en")
• Ethanedioate ion (also called oxalate ion)

In 1,2-diaminoethane, each nitrogen atom has a lone pair of electrons available to donate, allowing the ligand to form two coordinate bonds to the central metal ion. Similarly, in the ethanedioate (oxalate) ion, each negatively charged oxygen atom can donate a lone pair of electrons to form a coordinate bond.

An example of a complex containing bidentate ligands is $[\text{Co}(\text{NH}_2\text{CH}_2\text{CH}_2\text{NH}_2)_3]^{3+}$:

In this complex:

• The oxidation state of cobalt is +3
• The coordination number is 6 (because there are three bidentate ligands, and each forms two coordinate bonds: $3 \times 2 = 6$)

Shapes of complex ions

The three-dimensional shape of a complex ion depends directly on its coordination number. Understanding these shapes and their associated bond angles is crucial for exam success.

Six-coordinate complexes: octahedral shape

Many complex ions have a coordination number of six, which results in an octahedral shape. In this arrangement, the ligands are positioned at the corners of an octahedron around the central metal ion, with bond angles of approximately 90° between adjacent ligands.

A common example is the manganese complex formed when manganese sulfate ($\text{MnSO}_4$) dissolves in water, producing $[\text{Mn}(\text{H}_2\text{O})_6]^{2+}$:

The octahedral shape is also shown by the cobalt complex $[\text{Co}(\text{H}_2\text{NCH}_2\text{CH}_2\text{NH}_2)_3]^{3+}$ mentioned earlier, which has six coordinate bonds (three bidentate ligands, each forming two bonds), giving it an octahedral geometry with 90° bond angles.

Four-coordinate complexes

Complexes with a coordination number of four can adopt two different shapes: tetrahedral or square planar. The shape that forms depends on the specific metal ion and its electronic configuration.

Tetrahedral complexes

The tetrahedral shape is by far the more common of the two four-coordinate geometries. In this arrangement, the four ligands are positioned at the corners of a tetrahedron around the central metal ion, with bond angles of approximately 109.5°.

Common examples include $[\text{CoCl}_4]^{2-}$ and $[\text{CuCl}_4]^{2-}$:

Square planar complexes

A square planar shape is less common and occurs mainly with transition metals that have eight d-electrons in their highest energy d-sub-shell. Metals that commonly form square planar complexes include platinum(II), palladium(II), and gold(III).

In a square planar complex, the four ligands are arranged at the corners of a square around the central metal ion, all in the same plane. The bond angles between adjacent ligands are 90°, similar to those in an octahedral complex, but unlike the octahedral shape, there are no ligands above or below the plane.

An example is $[\text{Pt}(\text{NH}_3)_4]^{2+}$, where four ammonia ligands are arranged in a square around the platinum ion.

Colours in transition metal chemistry

The characteristic colours of transition metal complexes arise from electronic transitions within the d-orbitals of the metal ion. The colour you observe depends on several factors: the identity of the metal, its oxidation state, and the ligands coordinated to it.

When white light passes through a solution containing a complex ion, certain wavelengths are absorbed while others are transmitted or reflected. The colour we see is complementary to the colour absorbed. For instance:

• If a compound absorbs red light, it appears cyan (blue-green)
• If a compound absorbs blue light, it appears yellow/orange
• If all wavelengths are absorbed, the compound appears black
• If no wavelengths are absorbed, the compound appears colourless or white

Summary of key shapes and bond angles

Coordination Number	Shape	Bond Angles	Example
6	Octahedral	90°	$[\text{Cr}(\text{H}_2\text{O})_6]^{3+}$
4	Tetrahedral	109.5°	$[\text{CoCl}_4]^{2-}$
4	Square planar	90°	$[\text{Pt}(\text{NH}_3)_4]^{2+}$

Exam focus checklist