In this article, we learn about complex ions in chemistry and their ligands, including the various types of complexes, the importance of coordination number, and the chemistry of ligands including ligand substitution reactions.
In inorganic chemistry, a “complex” describes a structure involving a central metal ion with coordinate covalent bonds to one or more ligands. Chemists define “ligand” as a molecule capable of donating a pair of electrons to form a coordinate covalent bond. Since ligands donate pairs of electrons, they are classified as Lewis bases. Conversely, central metal ions count as Lewis acids because they accept pairs of electrons. Often, chemical complexes have an ionic charge and thus form complex ions. Complex ions are also known as coordination complexes.
The properties of a complex ion depend on the identity of the central atom or ion and the ligands it is bonded to. For example, the color of a complex ion can depend on the type of ligands it contains and the way they are arranged around the central atom or ion. The stability of a complex ion can also depend on the nature of the bonds between the central atom or ion and the ligands, as well as the interaction between the ligands themselves.
Complex ions are important in many fields of chemistry, including inorganic chemistry, biochemistry, and analytical chemistry. They are also used in a variety of applications, such as in the synthesis of drugs and other compounds, in the separation and purification of molecules, and in the detection of specific ions or molecules.
Overall, complex ions are an important part of coordination chemistry, and they play a key role in many chemical processes and applications.
In chemistry, a ligand is a molecule or ion that binds to a central atom or ion to form a coordination complex. Ligands are often used in coordination chemistry to form coordination compounds, which are molecules or ions that consist of a central atom or ion bonded to one or more ligands.
Ligands are sometimes classified based on at how many spots they attach to the central atom. A mondentate ligand only attaches at one spot. A bidentate ligand will attach at two spots. And a polydentate ligand attaches through three or more atoms. A common example of a bidentate ligand is ethylenediamine. The more spots a ligand is binding through, the more strongly it is attached. This increase in attachment strength is called the “chelate effect”.
Ligands are typically Lewis bases, which means that they have at least one pair of non-bonded electrons that they can use to form a bond with the central atom or ion. This allows the ligand to coordinate (or bind) to the central atom or ion, creating a coordination complex.
Some common ligands include water, ammonia, and other simple molecules that can easily donate a pair of electrons to form a bond with the central atom or ion. Ligands can also be more complex molecules or ions, such as organic compounds or metal ions. Other common ligands include carbon monoxide, acetylacetonate, and ethylenediamine, as well as metal ions, such as chloride, cyanide, and nitrate.
In addition to forming bonds with the central atom or ion, ligands can also interact with each other through various types of chemical bonds, such as hydrogen bonds or electrostatic interactions. These interactions can affect the properties and behavior of the coordination complex, such as its stability and reactivity.
Overall, ligands play a key role in coordination chemistry, allowing the formation of complex molecules and ions with a wide range of properties and applications.
One common class of coordination complex ions involves six ligands of water molecules. Chemists use the term “hexaaqua ions” to describe such complexes. Cobalt (II), copper (II), iron (II), iron (III), and many other transition metals can serve as the central ion in hexaaqua complexes. Importantly, because water molecules have a neutral charge, the charge of the entire complex ion remains the same as that of the central ion.
These complexes tend to form spontaneously when metal cations dissolve into an aqueous solution. Further, the coordinate bonds allow the complex to absorb light at certain frequencies, resulting in the color of certain metal ion solutions.
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Additionally, when a hexaaqua ion reacts with water, the complex releases a proton, making it an Arrhenius acid. This is why dissolving metal cations into a solution consequently lowers its pH.
Chemists define homoleptic complexes as complexes with only one species of ligand. Hexaaqua ion complexes count as homoleptic complexes since they only involve water. Other homoleptic complexes include those with only ammonia ligands, such as [Co(NH3)6] 2+ , or only carbon monoxide ligands, such as [W(CO)6], for instance.
Conversely, chemists define heteroleptic complexes as those with at least two distinct species of ligands. Examples of heteroleptic complexes include [Fe(NH3)4Cl2] + , [Co(NH3)2(H2O)4] 2+ , [Al(H2O)2(OH)2Cl2] – , and [Pt(NH3)2Cl].
When predicting and characterizing complex ions, the ion’s coordination number is important to keep in mind. Chemists define “coordination number” (CN) or “ligancy” as the maximum number of ligands that can bind to a central metal ion. Physically, the coordination number indicates the number of sites to which a ligand can bind. An ion’s coordination number varies from 2-8, with 4 and 6 being the most common. No rule or trend reliably predicts a metal ion’s coordination number, but factors such as ligand sterics and metal atomic radius appear to contribute somewhat.
Importantly, the coordination number of complex ions determines its molecular geometry, following similar trends to VSEPR theory.
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Many different molecules can serve as ligands in a coordination complex ion. One way to classify ligands involves their charge:
Another way to classify ligands involves the number of ligand sites to which they can bind on a metal ion. Chemists call this property the “denticity” of a ligand, from the Greek term for tooth, because it determines the maximum number of “bites” a ligand can take out of an ion.
Many of the simplest ligands count as “monodentate”, meaning they only bind to one site on the metal ion. Examples include: F – , Cl – , H2O, NH3, CO, NO + , and most other relatively small ligands.
Chemists use the term “polydentate” or “chelate” for ligands capable of binding to multiple ligand sites. Often, these ligands tend to have rather large organic structures. Metal ions tend to have a higher affinity to polydentate in second and later bindings compared to the first binding. This “polydentate effect” is explained by entropy; multiple bindings of a polydentate consequently displace other ligands, which results in an increase of molecules in the system.
Examples of polydentate ligands include:
Another important type of ligand goes by the name “bridging ligand”, which can bind to multiple metal ions. Bridging ligands then allow the formation of complex ions with multiple central metals. Examples include: CO, OH – , N3 – , and NH2 – .
To synthesize different metal ion coordination complexes, chemists perform ligand exchange reactions, which replace the ligands bound to the metal ion. Most complex ions tend to be synthesized from the homoleptic aqua ion since those complexes form easily from dissolving a soluble metal cation salt in water.
In this example, this hexaaquacobalt(II) complex forms the basis from which virtually all cobalt (II) complexes can be synthesized.
Importantly, the formation of any metal ion complex is reversible. Thus, Le Chatelier’s Principle governs all ligand exchange reactions. If we want to replace the ligands in our hexaaquacobalt(II) complex with ammonia, we would therefore need to add excess concentrated ammonia.
If instead we only add a little ammonia, instead of performing the ligand exchange reaction, the hexaaqua ion becomes deprotonated. This is due to the acidic properties of hexaaqua ions.
Importantly, these deprotonated hydroxide complexes often form a precipitate at significant concentrations.
Additionally, some ligand replacements change the coordination number of the metal ion. If we place our hexaaquacobalt(II) ion in concentrated hydrochloric acid, we consequently form the tetracobalt complex. The coordination number then changes from 6 to 4, due to the added bulk of chloride ions compared to water.
Finally, some ligand exchanges remain incomplete, even when we have an excess of a new ligand. For instance, hexaaquacopper(II) only forms tetraaminediaquacopper(II) when in excess of ammonia.
