Ions Dissolving and Dissociating in Aqueous Solution

Introduction

When a chemical dissolves, its molecules or formula units (depending on the chemical) separate from each other and immerse themself in the chemical which dissolved them. In aqueous solutions, that chemical is water. Some chemicals also dissociate--they are broken into their component ions, if there are any. In this article we will explore dissociativity in relation to a chemical's electrolytic qualities, and show how this relates to other key properties of chemicals.

What is an Electrolyte?

An electrolyte is a substance that, once dissolved, is able to perform electrical work. Place two wires connected to a lightbulb on both ends into your solution. The brightness of the lightbulb gives you the strength of the electrolyte. However, there is a more theoretical way to classify chemicals based on their electrolytic quality. An electrolyte can perform electrical work because the chemical dissociates into ions--some of these are anions which have excess electrons. The anions will attempt to get rid of said electrons to revert to a neutral charge as long as they are not in close proximity to cations (in which case the cations and anions will form the original chemical). These free electrons travel through the wire and power the lightbulb. Chemicals that do not have this ability are called nonelectrolytes. When nonelectrolytes are dissolved in water, the molecules remain intact, seemingly impervious to attractive forces with different ends of the water molecule.

Strong Electrolyte vs. Weak Electrolyte

There are actually two different types of electrolytes. A strong electrolyte completely dissociates into its component ions, leaving none of the original solute. Weak electrolytes, however, only partially dissociate into component ions--some molecules of the original chemical will still float around in the solution. Note that all ionic compounds are strong electrolytes. Molecular, or colvantly-bonded compounds, are almost all either weak electrolytes or nonelectrolytes, except for acids, which are also strong electrolytes. There is no "magic" rule for differentiating between weak electrolytes and nonelectrolytes. 

Example 1: Is \(Cr(OH)_3\) a strong electrolyte, a weak electrolyte, or a nonelectrolyte? If it is an electrolyte, write out the ions it dissociates into.

Solution: Note that Chromium (III) hydroxide is an ionic compound, so it completely dissociates into ions in aqueous solution and is thus a strong electrolyte. The ions formed are the cation and anion that comprise this compound; those are \(Cr^{3+}\) cations and \(OH^{-}\) anions, respectively.

The Intermolecular Forces Explanation of Electrolytes

The idea of why some compounds dissociate in water more easily than others can be explained using the concepts of intermolecular forces. Note that intermolecular forces are sometimes taught well after aqueous solutions are introduced, so you, as the student or teacher, may opt to skip this section for now.

Here is a copy of the Lewis Structure for water, the solvent in all solutions being covered here:

Now, oxygen has a greater electronegativity than hydrogen, which means that in a hydrogen-oxygen bond, the electrons shared between the two atoms are more attracted to the oxygen. Thus the oxygen acquires a slight, partial negative charge (less than the charge of an electron) and the hydrogen acquires a slight partial positive charge. Thus the hydrogen end of the molecule will be more attracted to anions, and the oxygen end of the molecule will be more attracted to cations. Thus you can see why ionic compounds tend to be "ripped apart" in water. The attractive forces between the hydrogen atoms and the ionic compound's anion alongside the attractive forces between the oxygen atoms and the ionic compound's cation exceed the strength of the attractive force that holds the ionic compound together. In nonelectrolytes, the forces holding the atoms of the original solute exceed the strength of any attractive forces between ends of the solute and ends of the water molecule. In weak electrolytes, all of the aforementioned attractive forces are more equal in strength. At this point, the decision as to whether an individual molecule gets ripped apart depends on its proximity to the water molecules.

Example 2: Chromium (III) hydroxide, as discussed in Example 1, is a strong electrolyte. The chromium cations are attracted to the oxygen atoms in the water, and the hydroxide anions are attracted to the hydrogen atoms in the water molecules.

Example 3: Consider the dichromate ion, \(Cr_2O_7^{2-}\), pictured below:

This particular diagram is actually very helpful, because it illustrates that some of the oxygen atoms in this polyatomic anion have negative charges.Thus these oxygen atoms will be attracted to the hydrogen atoms in the water. However, all other parts of the anion have a neutral charge, so the oxygen atoms in the water molecules will not be attracted to or repelled to any atoms in particular. Obviously this is not a strong electrolyte, but keep in mind two things for this particular example:

I. It is not important whether this is a weak electrolyte or a nonelectrolyte. It is more important to be able to analyze the intermolecular forces between atoms and see why it is not a strong electrolyte.

II. As this is an anion, it was likely initially accompanied by a cation as it entered the water solvent. For instance, if this cation was magnesium, then the initial chemical, \(MgCr_2O_7\), is a strong electrolyte because it is an ionic compound. The magnesium and dichromate ions will separate before any intermolecular forces between the dichromate ions and the water molecules can take place.

Dissociability and Electrolytic Properties are one example of taking a big-picture approach to chemistry. While it is important to be able to identify electrolytes and how they are defined, also remember how the use of other topics in chemistry can make this topic much more clear.

Resources

Thanks to WolframAlpha for allowing use of this widget: http://www.wolframalpha.com/widgets/view.jsp?id=689aa5a01c216d8b16ed0250cebdc702