1. Different sorts of salts
While the salt we eat is the real OG, chemists have expanded the definition of salts to include a whole lot of different ionic compounds. The definition of salt is as follows:
Let’s use this definition to explain why the OG — sodium chloride, NaCl — is defined as a salt. It can be formed when we replace the hydrogen ion of hydrochloric acid, HCl, with a sodium ion, Na+.
The definition holds an important clue to how we can prepare salts in the laboratory: by reacting acids with metals, carbonates, or bases.
- Acids react with reactive metals to form salt and hydrogen gas
- Acids react with carbonates to form salt, water and carbon dioxide gas
- Acids undergo neutralisation reaction with bases (including insoluble basic oxides and soluble alkalis) to form salt and water only
2. Like acid, like salt
|Acid Reactant||Anion in Salt|
|Hydrochloric acid, HCl||Chloride, Cl–|
|Nitric acid, HNO3||Nitrate, NO3–|
|Sulfuric acid, H2SO4||Sulfate, SO42-|
Like father, as the adage goes, like son. The same can be said for salt preparation. You can see the remnant of the acid in the anion of the salt. To put it in a cheem way, the anion of an acid becomes the anion of a salt during a reaction.
2HCl(aq) + CuO(s) ⟶ CuCl2(aq) + H2O(l)
Let’s take the example of the reaction between hydrochloric acid and copper(II) oxide, a basic oxide. It forms a chloride salt, copper(II) chloride.
H2SO4(aq) + 2KOH(aq) ⟶ K2SO4(aq) + H2O(l)
Likewise, reacting sulfuric acid with potassium hydroxide gives a sulfate salt, potassium sulfate.
3. Easy to say, hard to do
Some may think that salt preparation is simple: add something to an acid and voila, a salt! If anyone ever tells you this, do please, take it with a pinch of salt.
Indeed, a single reaction is enough to produce a salt. Yet, it is far from enough to get a pure sample of solid salts. This is because the salt product is formed in a reaction mixture, containing water and any reactants that are added in excess and are not used up completely.
To purify the salt formed, we have to consider separation techniques and drying. Given these considerations, there are two main methods to prepare pure samples of soluble salts from acids.
4. The spam method to prepare soluble salts: addition of excess insoluble solid to acid
This method involves the reaction between an acid and a solid of insoluble metal, metal oxide, or metal carbonate. It takes advantage of the insolubility of the solid. As it does not dissolve in the reaction mixture, any excess can be easily separated by filtration.
Let’s consider how we can prepare copper(II) sulfate, CuSO4, from sulfuric acid, H2SO4. The excess solid to use is insoluble copper(II) oxide, CuO. The copper(II) ions from copper(II) oxide will replace the hydrogen ions in sulfuric acid. The steps are as follows:
- Addition of excess solid to acid: H2SO4(aq) + CuO(s) ⟶ CuSO4(aq) + H2O(l)
- Filtration: discard the residue containing excess copper(II) oxide and keep the filtrate containing the dissolved copper(II) sulfate salt
- Evaporation and crystallisation: evaporate water to form hot saturated copper(II) sulfate solution, which is then allowed to cool to form crystals
- Drying: press the crystals between sheets of filter paper
For Metals, Use Moderately Reactive Ones
Besides using insoluble copper(II) oxide as a reactant, insoluble copper(II) carbonate could also be used. However, we cannot use copper, as it is a non-reactive metal.
In general, only magnesium, zinc, and iron are suitable reactants. Anything more reactive, like sodium and potassium, will go kaboom in acid. They are too dangerous to react! Anything less reactive, like copper and silver, may take a million years and we ain’t got time for that.
5. The precise method to prepare soluble salts: acid-alkali titration
When both reactants and their product are all soluble, separation of any excess reactant becomes nigh impossible. This is when titration becomes necessary. It is a precise method that allows for the exact amount of each reactant to be used. There is no excess reactant, avoiding the need for separation.
We can use it to prepare soluble sodium chloride from dilute hydrochloric acid and sodium hydroxide solution. The exact steps are as follows:
- Titration: titrate twice, the first time with an indicator to determine how much sodium hydroxide is needed to completely react with hydrochloric acid, and the second time without an indicator to prevent the contamination of the sodium chloride salt produced
- Evaporation to dryness: sodium chloride can be evaporated to dryness as it does not decompose at high heat
6. What about preparing insoluble salts?
- Chlorides, bromides, and iodides: silver halides and lead(II) halides only
- Sulfates: barium sulfate, calcium sulfate, and lead(II) sulfate
- Carbonates: all carbonates, except sodium carbonate, potassium carbonate, and ammonium carbonate
While we could theoretically react acid with an excess solid of metal, metal oxide, or metal carbonate to form insoluble salt, it is practically impossible. There would be a very low yield, as the insoluble product would coat the insoluble reactant to prevent further reaction.
7. The way out for insoluble salts: precipitation
Precipitation is easy to conduct, with just a few steps. Let’s use the example of preparing insoluble lead(II) sulfate, PbSO4. To introduce the lead(II) ions and sulfate ions for precipitation, we can mix soluble solutions of lead(II) nitrate and sodium sulfate. The sodium and nitrate ions are so soluble that they will simply remain in the solution, without interfering the precipitation.
- Precipitation: mix lead(II) nitrate solution and sodium sulfate solution
- Filtration: filter and keep the residue that contains the lead(II) sulfate precipitate
- Drying: dry the residue between sheets of filter paper
8. Not all methods are good methods
To decide on the best method to prepare a salt, we have two main considerations: whether the salt product is soluble and whether both reactants are soluble.