Is there any difference between getting dissolved and getting dissociated?

In the context of electrolytes only, is there any difference between electrolyte getting dissolved and electrolyte getting dissociated?


According to my professor, they are different. He says a substance may dissolve completely but it may not dissociate completely and a substance may not dissolve completely but amount that dissolves can dissociate completely and, thus, substances which dissolve (either wholly or partially), amount that gets dissolved if dissociate completely are called strong electrolytes and amount of substances which gets dissolved if dissociates sparingly are called weak electrolytes (although it may dissolve completely).


How is it possible? I mean, how can solubility be different from dissociating? 


If something dissociates it means it’s dissolving and in the context of electrolytes if something dissolves in water it has to dissociate into ions and even Arrhenius’ theory of electrolytic dissociation sys that substance which are water soluble, dissociates into ions and get dispersed in water. 


And if someone says they are different to them, I ask the following two questions: 


1. If solubility and dissociation are different then what would be the criteria for getting dissolved and what will be criteria for getting dissociated (for me the criteria for getting dissolved is dissociation).
2. how would you arrange the following in increasing order of 1.solubility 
2.dissociation


NaCl, MgCl2, FeCl2, ZnCl2




Here are four examples of substances dissolving:


$$\ce{C6H12O6(s) <=> C6H12O6(aq)}\tag{1}$$


$$\ce{CH3COOH(l)<=> CH3COOH(aq) <=> CH3COO-(aq) + H+(aq)}\tag{2}$$


$$\ce{NaCl(s) <=> Na+(aq) + Cl-(aq)}\tag{3}$$


$$\ce{SO3(g) <=> SO3(aq);\ \ SO3(aq) + H2O(l) <=> HSO4-(aq) + H+(aq)}\tag{4}$$


Because NaCl is an ionic solid, the only way it can dissolve (3) is to also dissociate. Glucose is a molecular solid, so it only dissolves (1). Glacial acetic acid is a molecular liquid, and it dissolves in water and then partially dissociates (3), whereas sulfur trioxide dissolves in water and then reacts with it (4). 



> 
> $\ce{NaCl, MgCl2, FeCl2, ZnCl2}$
> 
> 
> 


Some transition metals cations hang on to their counter anions, or pick up some more in solution, together with water as a ligand.


For example for zinc cations:



> 
> The highest chloro complex, [ZnCl4]2-, is tetrahedral with a Zn-Cl bond length of 2.294(4) Å. The trichloro complex, [ZnCl3]-, which coordinates one water molecule, is pyramidal with the Cl-Zn-Cl angle 111°. The Zn-Cl and the Zn-H2O bonds are 2.282(4) and 1.9 Å, respectively. The two lower complexes, [ZnCl2] and [ZnCl]+ , cannot be separated by Raman spectra. The average Zn-Cl distance in these complexes is 2.24 Å, and the average Zn-H2O distance is 1.9 Å. In [Zn(H2O)6]2+ the Zn-H2O distance is 2.15 Å.
> 
> 
> 


Source: <https://www.researchgate.net/publication/273506504_The_Structure_of_Zinc_Chloride_Complexes_in_Aqueous_Solution>


So for these metal cations, the story is complicated and there is a complex mixtures of species, with concentrations dependent on anion concentrations, pH and temperature. A classic example is cobalt chloride, where you can follow the equilibrium easily because of color changes, see e.g. <https://www.chemedx.org/blog/multi-colored-equilibrium-experiment>


 4

This is a simple logical and semantic problem, which is not a problem at all. Your professor is right and wrong- both at the same time. He is creating a classification which does not exist and which is meaningless.


Look at the word origin of electrolyte:



> 
> Etymology from *OED*: < electro- comb. form + ancient Greek λυτός that may be
>  dissolved, soluble (see -lyte comb. form). Compare electrolysis n.,
>  electrolyse v.
> 
> 
> and the portion -lyte "< ancient Greek -λυτος, combining form (in e.g.
>  εὐδιάλυτος : see eudialyte n.) of λυτός that may be dissolved, soluble
>  < λύειν to loose (see lysis n.) + -τός, suffix forming verbal
>  adjectives. Compare French -lyte, German -lyt.
> 
> 
> A substance which gives rise to ions when dissolved (typically in
>  water) or fused; a liquid or gel which contains ions and can be
>  decomposed by electrolysis; spec. the conducting fluid used in a
>  battery.
> 
> 
> 


Does this solve the semantics problem?


By definition, an electrolyte is soluble in a given solvent. It is also understood among electrochemists that an electrolyte is added to a solvent to impart conductivity to that solvent. I will never call sugar an electrolyte because nobody would use sugar to increase conductivity of water. You should ask: Does a given substance produce ions in a solution? If yes, it can be called as an electrolyte. 


Look at this logic: grapes is a fruit but not all fruits are grapes.


Electrolytes are soluble in the solvent, but not all soluble substances in that solvent behave as electrolytes. Anything which produces ions in that solvent, can be labelled as an electrolyte.


*Q-1 If solubility and dissociation are different then what would be the criteria for getting dissolved and what will be criteria for getting dissociated (for me the criteria for getting dissolved is dissociation).*


The measure of dissociation is conductivity. How much electrical resistance is offered by that dissolved (and dissociated) substance in that solvent


Q-2 how would you arrange the following in increasing order of 1.solubility 2.dissociation


NaCl, MgCl2, FeCl2, ZnCl2


There is **no** correlation between solubility and dissociation. If I want to study their dissociation I will measure the conductivity of their solutions systematically.


 4

[EDIT] I start with the [following definition](https://www.britannica.com/science/dissociation):



> 
> Dissociation, in chemistry, the breaking up of a compound into simpler constituents that are usually capable of recombining under other conditions. In electrolytic, or ionic, dissociation, the addition of a solvent or of energy in the form of heat causes molecules or crystals of the substance to break up into ions (electrically charged particles). Most dissociating substances produce ions by chemical combination with the solvent. The idea of ionic dissociation is used to explain electrical conductivity and many other properties of electrolytic solutions.
> 
> 
> 


I now present an example of two molecules, chlorine oxides ($\ce{Cl\_xO\_y}$) and also chlorine, all of which are gases at room temperature, to state my case that there can be a major difference between dissolve and dissociate creating distinct and important properties in a solvent of water.


For example, $\ce{ClO\_2}$ is a gas that is highly soluble in water (some 8 g/L at 20 °C). However, [per Wikipedia](https://en.wikipedia.org/wiki/Chlorine_dioxide) , to quote:



> 
> It does not hydrolyze when it enters water, and is usually handled as a dissolved gas in solution in water.
> 
> 
> 


Although technically, there are many (including important) 'produce ions by chemical combination of the molecule with' water (see my note below).


In contrast, we have $\ce{Cl\_2O}$, which [per Wikipedia](https://en.wikipedia.org/wiki/Dichlorine_monoxide) is described as:



> 
> At room temperature it exists as a brownish-yellow gas which is soluble in both water and organic solvents. Chemically, it is a member of the chlorine oxide family of compounds, as well as being the anhydride of hypochlorous acid.
> 
> 
> 


As a result of the latter, we have:


$\ce{Cl2O + H2O <=> 2 HOCl}$


And, further aqueous hypochlorous acid can create ions ([hence an electrolyte](https://www.scientificamerican.com/article/what-are-electrolytes/), albeit, very weak):


$\ce{HOCl <=> H+ + OCl-}$


In fact, dichlorine monoxide is very soluble due to hydrolysis in water (some 143 g of $\ce{Cl\_2O}$ per 100 g water), but also can dissolve in organics (like CCl4 and CHCl3).


So in water, $\ce{ClO\_2}$ mostly dissolves and $\ce{Cl\_2O}$ largely dissociates, which seemingly presents significant differences. And, finally, in contrast to chlorine, it is but mildly soluble (1/10 of $\ce{ClO\_2}$) and mildly dissociates, per below, which varies significantly with temperature:


$\ce{Cl2 + H2O <=> H+ + Cl- + HOCl}$


Therefore, it appears that all three gases have distinct (and important) properties connected to solubility and propensity to undergo dissociation. If a chemistry student of mine were to conflate solubility and dissociation, in regard to these compounds, I would feel bad if I contemplated assigning a final grade above C. 


A technical note, with respect to $\ce{ClO\_2}$, I explain the dissociation of chlorine dioxide due to its self-reaction in water as chlorine dioxide is an example of a stable free radical:


$\ce{.ClO2 + .ClO2 <=> Cl2O4}$


$\ce{Cl2O4 + H2O <=> HClO2 + HClO3}$


where both created acids form ions in water. [EDIT] [A supporting source](https://www.epa.gov/sites/production/files/2014-11/documents/chlorinedioxide_final_volume5_2007.pdf) makes the statement:



> 
> In air, chlorine dioxide readily dissociates both thermally and photochemically and may form chlorine, oxygen, hydrogen chloride, HClO3, HClO4.ClO, chlorine peroxide, and/or chlorine trioxide, dependent on temperature and humidity. Chlorine dioxide dissociates in water into chlorite and chloride, and to a lesser extent into chlorate (Budavari et al. 1996). 
> 
> 
> 


I would also note that $\ce{Cl\_2O}$ dissolves before undergoing dissociation, and the dichlorine monoxide reversibly, to an extent, will escape from solution (especially concentrated or on heating) and is responsible for the 'chlorine-smell' of bleach ([per Watts, p. 16](https://books.google.com/books/reader?id=ijnPAAAAMAAJ&dq=hclo%20decomposition&printsec=frontcover&output=reader&pg=GBS.PA16#v=onepage&q=hclo%20decomposition&f=false)) and as also reported by Watts, the ability to concentrate HOCl by distilling off half of the solution (due to the partial release of the $\ce{Cl\_2O}$).


[EDIT] Now, as to the answer to the 2nd question, the answer depends on a substance's ability to dissolve and undergo dissociation, to rank one must measure the electrical current produced by my adopted definition. However, I would argue that the results can vary in the presence of dissolved oxygen for acidic ferrous salt, for example. The explanation is that the ferrous can be converted into ferric per an electrochemical reaction in the presence of H+ (per hydrolysis of FeCl2) and O2 as dissolved or from air, upon exposure over time. As such, the FeCl2 with time and air exposure may seemingly present different electrical properties (due actually to dissociation). Similarly, per this work: [The reaction of oxygen with magnesium chloride](https://www.sciencedirect.com/science/article/abs/pii/004060317785003X), again measurable changes in voltage may occur. Apparently, MgCl2/O2 is NOT exactly inert either, so again a possible departure in ranking with time in the presence of gaseous or dissolved oxygen.


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