Many inorganic metallic compounds take the form of ions when dissolved in water (i.e. an aqueous solution). The laws that govern their reactions are similar to how compounds react, but with one important emphasis on the role of water in the reactions. In fact, their balanced equations involve water in a generalised stoichiometry form that can be written as:
- M represents any metal element
- The subscript ‘n’ represents the balanced number of water in the ion
- The superscript ‘z’ with positive charge represents the number of ionic charges.
Reactions of ions in aqueous solutions are actually half of the reactions of ionic compounds dissolved in water. They usually form complexes. These reactions and the complexes that they form are pervasive in nature. They’re also very useful in industries, medicine, and other everyday applications.
In this post:
What Are Ionic Aqueous Solutions?
An aqueous solution can either be an ionic (electrolyte) or a nonionic (nonelectrolyte) solution. Our focus in this article, however, is ionic solutions. As the name implies, an aqueous ionic solution consists of water as the solvent and at least one ionic compound as the solute. The dissolved solute dissociates as positive and negative ions in water. The solute can be an acid, a base, or salt.
Ionic aqueous solutions have properties and chemical behaviours that are different from nonionic solutions. For example, ionic solutions can conduct electricity. The ions can also produce electricity when the solutions are arranged into battery cells with two different metals immersed in them, acting as anode and cathode terminals.
When two different ionic solutions are mixed together, chemical reactions may occur, especially if the solutions are an acid solution and a base solution. Acid-base reactions are classified as double-substitution reactions. This type of reaction is relatively straightforward and easy to balance as the elements of the compounds exchange partners.
On the other hand, a special type of reaction that forms complexes is a focus of study in coordination chemistry. This reaction typically involves metal cations with attached ligands, with one of the most common types of ligands being water. These reactions form hydrated metal ions as hexaaqua complexes in an aqueous solution.
What Are Complexes?
You’ll hear complexes called coordination compounds. These are molecules that have a metal as the centre atom. The centre atom has ligands – ions that bind to this centre atom – that are attached to it. The ligands can be individual atoms, ions, or molecules, such as in the case of hydrated hexaaqua complexes.
When metal and non-metal ions react with each other, they either form charged (ionic) or neutral complex products. If it’s the latter, the complex product is usually a solid precipitate.
Complexes are the main focus of a branch of chemistry known as coordination chemistry. The study was pioneered by Alfred Werner, which led to his discovery of the structure of coordination compounds.
Werner studied and experimented with various compounds of cobalt (III) chloride and ammonia. He tried liberating the ammonia from the compounds by adding hydrochloric acid. However, he noticed that the ammonia component could not be removed completely by using HCl alone.
Werner formed the hypothesis that the ammonia component of the compounds must have a very strong bond with the central cobalt ion. When he added silver nitrate to the compounds, he observed that the amount of silver chloride precipitates that were formed had direct correlation with the number of ammonia molecules attached to the cobalt (III) chloride.
Thus, he was able to figure out the reaction patterns and generalised structures of the metal ions with ligands. For example, metal aquo complexes have octahedral shapes, as soon in the illustration below:
If the central metal cation belongs to the main element groups, their complexes are colourless. However, if it belongs to the transition elements, they typically have coloured precipitates.
Here are some examples of the hexaaqua complexes with their corresponding colours.
- [Fe(H2O)6]2+ – pale green
- [Co(H2O)6]2+ – pink
- [Cr(H2O)6]2+ – blue
- [Mn(H2O)6]2+ – pale pink
- [Fe(H2O)6]3+ – yellow
- [Ni(H2O)6]2+ – green
- [Cu(H2O)6]2+ – blue
What Are the Types of Ion Reactions in Aqueous Solutions?
The reactions of ions in aqueous solutions vary depending on the reactants, the coordination number, the concentration of the solution, and the oxidation state of the central metal. While some ions will form solid complex precipitates, others may form either neutral products or ionic compounds.
- Lewis acid-base reactions
The most familiar type of ionic reaction is the Lewis acid-base reaction, wherein one reactant will donate a proton (hydrogen) while another reactant will receive the proton. Typically, this is a neutralisation reaction that forms salt. The metal cations act as Lewis acids when forming bonds with ligands. In the case of aqueous solutions, water may act as the proton acceptor.
Here is an example of a Lewis acid-base reaction between two ammonia molecules and a silver ion, producing a silver-ammonia ion complex:
- Hydrated metal cations
Hexaaqua complexes are the most common type of metal cations in aqueous solutions. As the name suggests, they have six water molecules as ligands. As previously mentioned, transition metal complexes have bright colours. Main group metal complexes and compounds are colourless. Metal cations dissolve in water undergo hydrolysis in dynamic equilibrium. The generalised equation can be written as:
[M(H2O)n]z+ + H2O [M(H2O)n-1(OH)](z-1)+ + H3O+
- Metal complex ions
Metal complex ions can react in several ways, and can undergo two main types of reactions: deprotonation and ligand exchange reactions. Here are some examples:
- Deprotonation by water: [Fe(H2O)6]2+(aq) + H2O(l) ⇌ [Fe(H2O)5(OH)]+(aq) + H3O+(aq)
- Deprotonation by hydroxide ions: [Cu(H2O)6]2+(aq) + 2OH–(aq) ⇌ [Cu(H2O)4(OH)2](s) + 2H2O(l)
- Deprotonation by ammonia: [Cr(H2O)6]3+(aq) + 3NH3(aq) ⇌ [Cr(H2O)3(OH)3](s) + 3NH4+(aq)
- Ligand exchange: [Cu(H2O)6]2+(aq) + 2NH3(aq)⇌ [Cu(H2O)4(OH)2](s) + 2NH4+(aq)
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