Coordination Bonds and Complex Ions (6.4.5) | IB DP Chemistry HL 2025 Notes

Coordination Bonds: An Elaborate Perspective

Definition: A coordination bond, often referred to as a dative covalent bond, is a specialised kind of covalent bond. It's characterised by the donation of an electron pair by one atom, the donor, to another atom, the acceptor.
- This is unlike a regular covalent bond where each atom typically contributes one electron to the bond.
Lewis Bases and Lewis Acids: The crux of coordination bond formation lies in the interaction between Lewis bases and Lewis acids.
- Lewis Base: It's an electron-pair donor. They have lone pairs of electrons that can be shared. Examples include water (H₂O) and ammonia (NH₃).
- Lewis Acid: This is an electron-pair acceptor. They have vacant orbitals ready to accept a lone pair. A classic example is boron trifluoride (BF₃).
- The bond formed in the product NH₃BF₃ between nitrogen and boron is a coordination bond. Here, ammonia serves as the Lewis base, donating its lone pair, and BF₃ acts as the Lewis acid, accepting the lone pair.

Diagram showing the formation of a dative covalent bond between ammonia and boron trifluoride forming a Lewis adduct.

The formation of a dative covalent bond between ammonia and BF3 forming a Lewis adduct. Arrows originating from the donor atom (Lewis base) and culminating at the acceptor atom (Lewis acid).

Image courtesy of Ben Mills

Role of Nucleophiles and Electrophiles in Coordination

Nucleophiles as Lewis Bases: In essence, nucleophiles are species that have a keen affinity for positive centres, primarily because they have lone pairs or pi bonds ready to be donated. This inherent characteristic aligns them with Lewis bases.
- Examples: Hydroxide ion (OH-), cyanide ion (CN-), and even molecules like water can act as nucleophiles.
Electrophiles as Lewis Acids: Electrophiles, on the other hand, are "lovers of electrons". They are always on the prowl for areas of high electron density, to which they can attach. This quality makes them perfect candidates for being Lewis acids.
- Examples: Protons (H+), carbocations, and species like BF3.

Drawing and Interpreting Lewis Formulas

Visual Representation: Lewis formulas offer a visual depiction of molecular structures and the interactions within them.
- Coordination Bonds: They are displayed as arrows, originating from the donor atom (Lewis base) and culminating at the acceptor atom (Lewis acid).
- An example is the reaction between ammonia and BF₃. The arrow will extend from the nitrogen of ammonia and terminate at the boron of BF₃.

Comparison: Coordination Bonds vs. Traditional Covalent Bonds

Origin of Electrons: As discussed earlier, the electron origin differentiates the two. Covalent bonds involve shared responsibility, whereas in coordination bonds, the donor atom is the primary contributor.
Bond Strength: The strength of coordination bonds varies. Often, their strength depends upon the Lewis base and acid involved. They are generally weaker than covalent bonds but stronger than Van der Waals interactions.
Bond Length: The bond lengths in coordination complexes tend to be longer than typical covalent bonds. The reason can be attributed to the fact that the electrons in coordination bonds come from one atom, possibly making the bond less “tightly held”.

Diagram showing the difference between Coordination Bonds and Traditional Covalent Bonds

Image courtesy of RMIT University

The Fascinating World of Complex Ions with Transition Element Cations

Peculiarity of Transition Metals: The d-orbitals of transition metals can accommodate extra electron pairs. This makes them exemplary Lewis acids, eagerly waiting for electron pairs from Lewis bases.
Complex Ions: When a transition metal cation reacts with several molecules or ions (known as ligands), a complex ion is formed.
- For instance, the copper(II) ion (Cu²⁺) interacting with six water molecules results in the formation of the complex ion [Cu(H₂O)₆]²⁺.

A diagram showing the chemical structure of hexaaquacopper(II) ion ([Cu(H2O)6]2+)

Copper(II) ion (Cu²⁺) interacting with six water molecules results in the formation of the complex ion [Cu(H₂O)₆]²⁺.

Image courtesy of Ewen

Deducting Charge on Complex Ions: A Methodical Approach

Charge Calculation: A systematic approach is required to deduce the charge on a complex ion.
- Metal Charge: Begin by identifying the charge of the central metal ion.
- Ligand Charge: Calculate the total charge contributed by the ligands.
- Net Charge: Combine the two to get the overall charge on the complex ion.
Using the [Cu(H₂O)₆]²⁺ example: Cu²⁺ has a charge of +2. As water is neutral, the total charge remains +2.

In conclusion, understanding the nuances of coordination bonds and complex ions enables students to appreciate the myriad ways in which atoms and molecules can interact. This knowledge is crucial for anyone hoping to master inorganic chemistry and its applications.

FAQ

Not all coordination compounds are complex ions, but all complex ions are coordination compounds. A complex ion consists of a metal ion at its centre with a number of other molecules or ions surrounding it. These can be negatively charged, like Cl-, or neutral, like H₂O. When these ions/molecules surround the metal ion through dative covalent bonds, they form a complex ion. However, some coordination compounds are neutral. These don't form ions in solution and hence aren't considered complex ions. To summarise, while all complex ions fall under the broader category of coordination compounds, not every coordination compound is a complex ion.

Transition metals are particularly adept at forming complex ions due to several factors. Firstly, they have vacant d-orbitals that can accept lone pairs of electrons from ligands, facilitating the formation of dative covalent bonds. Secondly, transition metals often exist in multiple oxidation states, which means they can form complex ions of varying charges. Furthermore, their relatively small size and high charge make them polarising, enabling them to distort electron clouds and accept lone pairs more readily. Lastly, many transition metals have the ability to form various geometries around them, allowing for flexibility in the kinds and number of ligands they can accommodate.

Chelating ligands are a special type of ligand that can form multiple bonds to a single metal ion. The term "chelate" originates from the Greek word for "claw", reflecting the way these ligands can "grab" a metal ion at two or more attachment points. Ethylenediamine is a classic example of a chelating ligand, as it has two nitrogen atoms that can both form dative covalent bonds with a metal ion. This ability to form rings with metal ions often gives complexes with chelating ligands increased stability, a phenomenon known as the "chelate effect". In contrast, regular ligands typically bind at only one point on the metal ion.

The coordination number of a metal ion in a complex is determined by counting the number of ligand attachment sites around the metal ion. This doesn't just mean the number of ligands, but the number of places on the metal where these ligands are attached. For instance, the complex [Fe(CN)₆]^3- has a coordination number of 6 because there are six cyanide ligands attached to the iron. Another example would be the ethylenediamine ligand, which can attach at two different points; so, if a metal ion is surrounded by three ethylenediamine ligands, its coordination number would be 6. Understanding coordination numbers is essential for predicting the geometry and properties of complex ions.

Ligand field theory (LFT) is an adaptation of molecular orbital theory that describes the bonding, orbital arrangement, and electron configurations in coordination compounds. LFT provides a method for understanding the electronic structure of transition metal ions in complex ions, especially the splitting of d orbitals. When ligands approach a central metal ion, they can cause the degenerate set of d orbitals to split into different energy levels. This splitting can affect the colour, magnetism, and reactivity of the complex. In essence, ligand field theory offers a deeper insight into the nature of coordination bonds and how the surrounding environment impacts the behaviour of metal ions in complexes.

Practice Questions

In the context of Lewis acid-base theory, explain the nature of the bond formed when ammonia reacts with boron trifluoride. Using Lewis structures, depict this interaction.

When ammonia (NH₃) reacts with boron trifluoride (BF₃), a coordination bond or dative covalent bond is formed. This bond arises from the donation of a lone pair of electrons by the nitrogen atom in ammonia (acting as a Lewis base) to the boron atom in BF₃ (acting as a Lewis acid). As boron trifluoride has an empty p-orbital, it can accept the lone pair from ammonia. The Lewis structure would show an arrow originating from the nitrogen in ammonia pointing towards the boron in BF₃, signifying the donation of the lone pair to form the coordination bond.

Given a complex ion [Fe(H2O)5Cl]2+, deduce the charge on the central metal ion and explain the formation of this complex ion in terms of Lewis acid-base interactions.

To determine the charge on the central metal ion (iron, Fe), consider the charges of the ligands surrounding it. Water molecules (H₂O) are neutral and hence contribute no charge, while the chloride ion (Cl-) contributes a charge of -1. Given that the overall charge of the complex ion is +2, the charge on the iron must be +3 (because +3 and -1 combine to give +2). The formation of this complex ion involves the water molecules and the chloride ion, all acting as Lewis bases, donating electron pairs to the iron (Fe³⁺), which acts as a Lewis acid, thereby forming coordination bonds.

Try All Topic Practice Questions

Written by:

Dr Shubhi Khandelwal

Qualified Dentist and Expert Science Educator

Shubhi is a seasoned educational specialist with a sharp focus on IB, A-level, GCSE, AP, and MCAT sciences. With 6+ years of expertise, she excels in advanced curriculum guidance and creating precise educational resources, ensuring expert instruction and deep student comprehension of complex science concepts.