Whenever you hear the word “salt”, what comes to your mind? Most probably the table salt. The white powder which we add to our food to make them taste better. After all who doesn’t like perfectly salted fish n chips! But in chemistry, salts have a much deeper meaning.
In this blog, I am going to discuss in-depth “salts”. After completely reading this blog, you will have a crystal clear idea about the following concepts:
Just properly go through this blog, get your concepts clear and show the nerd in your class that it is you who is the real boss!
In chemistry, salt is just a nickname for ionic compounds. If you are flabbergasted by the word ionic compound and have no clue about it at all, in that case, you better have a look at this article about ionic bonding.
Most ionic compounds or salts are soluble in water because water contains polar molecules. In simple terms, a polar molecule is a molecule with one end slightly positively charged and the other end slightly negatively charged.
So when you place an ionic compound in water, the positive end of the water molecules pulls on the negative ion while the negative part of the water molecule pulls the positive ion of the ionic compound. In this way, the water molecules break apart the ionic lattice and dissolve it. Obviously, there is more to it than that, but for now, just assume this simple story is what happens. You can also go through this article where I explain in detail with diagrams that how water dissolves an ionic compound.
Before you dive into this information goldmine about salts, I think you better have a quick look at this article about acid, there are some important reactions I have mentioned there. This will make sure that you make full use of this article.
To determine whether salt is soluble in water or not, we need to keep in mind the following facts:
**PLEASE KEEP IN MIND THAT THESE FACTS ARE HIGHLY IMPORTANT. JUST BY KNOWING THESE FACTS, YOU CAN EASILY STATE WHETHER A SALT IS SOLUBLE OR INSOLUBLE**
| Salt Name | Solubility |
| NaCl | soluble( all sodium salts are soluble) |
| PbSO4 | insoluble (lead(II)Sulfate is insoluble as we learned) |
| CaCO3 | insoluble( we learned that all carbonate salts are insoluble except (Na2CO3, K2CO3, and (NH4)2CO3, so this salt is insoluble) |
| K2SO4 | soluble (all potassium salts are soluble) |
| AgCl | insoluble (we learned that silver chloride is among one of the insoluble chloride salts) |
In order to make life easier, chemists have divided salts into three types:
Dividing salts into three types makes it easier to conceptualize and prepare them.
Soluble salts – To make soluble salts, simply react an acid with access metal (make sure it is a metal of moderate reactivity, should lie between magnesium and iron in the reactivity series), metal oxide, or metal carbonate
Insoluble salts– Make insoluble salts by the precipitation method (mix two solutions, each solution contains respective ion, when the ions meet when solutions are mixed, they form the insoluble salt)
Potassium, sodium, and ammonium salts – by titration method( by the neutralization reaction between an acid and an alkali)
To better understand the relationship between Acid, Base and Salt I would recommend you read this book “Chemistry: Acids, Bases and Salts (2nd Ed.)” from Amazon.
Here is the summary of which steps to follow while making a salt.
Acid+ metal→ salt+ hydrogen gas
Please make sure to use metals between magnesium and iron in the reactivity series, because it could be dangerous to use a more reactive metal.
Example 1: Mg +2 HCl → MgCl2 + H2
Acid+ metal hydroxide/metal oxide → salt + water
Example 2: CuO + H2SO4 → CuSO4 + H2O
Acid + metal carbonate→ salt+ water+ carbon dioxide gas
Example 2: 2HCl + CaCO3 → CaCl2 + H2O + CO2
From the things we have learned so far, we can easily state that magnesium sulfate is a soluble salt. Magnesium is a moderately reactive metal, so we can directly react to magnesium with an acid to make the salt.
To get the sulfate(SO42-) ions we use sulfuric acid( H2SO4).
We react excess magnesium with dilute sulfuric acid to make sure that all the acid is used up and no H+ ions are left behind.
The following reaction below occurs between magnesium and sulfuric acid:
Mg(s) + H2SO4 (aq) → MgSO4 (aq) + H2 (g)
During the reaction, we see the bubbling of an invisible gas. That invisible gas is hydrogen. Excess magnesium is left behind in the solution of magnesium sulfate. The magnesium simply sits in the bottom of the beaker as a deposit.
To obtain crystals of magnesium sulfate we need to follow the crystallization procedure because magnesium sulfate needs water of crystallization to form crystals. If we heat the solution of magnesium sulfate strongly, only powder of anhydrous magnesium sulfate will be left behind.
MgSO4 (aq) + 7H2O (l) → MgSO4.7H2O (s)
We cannot react copper with acid because copper is less reactive than hydrogen. Copper does not react with any acid at all! Maybe a little bit of highly concentrated sulfuric acid, but that’s dangerous! However, we can react copper (II) oxide with an acid to form our desired salt solution Copper (II) Sulfate.
CuO(s) + H2SO4(aq) → CuSO4(aq) +H2O(l)
We use sulfuric acid because we need the sulfate(SO42-) ion from the sulfuric acid.
In a beaker of dilute sulfuric acid, add excess black powder of copper (II) oxide. After the reaction occurs, a blue solution of copper (II) sulfate is formed. Excess black copper (II) oxide powder is left behind as a deposit at the bottom of the beaker.
Firstly filter, to remove the excess copper (II) oxide powder, and yet again follow the crystallization procedure as I described before.
CuO(s) + H2SO4(aq) → CuSO4(aq) +H2O(l)
CuSO4 (aq) + H2O (l) → CuSO4.5H2O(s)
When making sodium, potassium, or ammonium salts, we face an odd problem, which is that all of their compounds/salts are soluble in water.
For example, if you reacted excess sodium carbonate with hydrochloric acid to make sodium chloride, all the acid would have reacted to form the sodium chloride salt solution. But the problem is that the excess remaining sodium carbonate would have also dissolve in the solution, hence contaminating our desired salt solution.
So to solve this problem, we need to do a titration, to make salts of these kinds. Meaning we need to react exactly the required amount of acid with an alkali. So no leftover reactants will remain behind to contaminate our salt solution.
To make sodium sulfate crystals, we can react sodium hydroxide (NaOH) with sulfuric acid (H2SO4)
NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2H2O(l)
Firstly use a pipette to transfer 25 cm3 of sodium hydroxide to a conical flask.
Then add a few drops of methyl orange indicator into the conical flask.
The colors of methyl orange are given below
| Alkaline | yellow |
| Neutral | Orange |
| Acidic | Red |
When the few drops are added to the conical flask it turns yellow, as it contains an alkaline solution of sodium hydroxide.
Use a burette to add dilute sulfuric acid to the sodium hydroxide solution. When the sulfuric acid will neutralize all of the sodium hydroxides, the color will change from yellow to orange. But if too much acid is added the solution will red, and we will have to repeat the whole thing again.
When the indicator color turns orange, note down the volume of acid used.
In a separate beaker mix 25 cm3 of NaOH with the volume of acid used. Because the titration solution is now contaminated with the indicator methyl orange.
The sodium sulfate salt also requires water of crystallization to form crystals, so we must follow the crystallization procedure yet again.
Na2SO4 (aq) + 10H2O (l) → Na2SO4.10H2O(s)
To make sodium chloride crystals, you can react sodium hydroxide with hydrochloric acid, by titration method.
NaOH (aq) + HCl(aq) → NaCl(aq) + H2O(l)
Sodium Chloride does not require water of crystallization to form crystals, so just heat the solution to evaporate all the water, sodium chloride crystals will remain behind.
To make (NH4)2SO4 crystals react aqueous ammonia (NH4OH) with sulfuric acid (H2SO4):
NH4OH (aq) +H2SO4 (aq) → (NH4)2SO4 (aq) + H2O (l)
OR NH3 (aq) + H2SO4 (aq) → (NH4)2SO4 (aq)
Ammonium salts do not have water of crystallization, but still, we must follow the crystallization method. As heating dry ammonium salts will cause them to break up or decompose.
To make insoluble salts, we will follow the precipitation method.
The precipitate is a fine solid that is formed by a chemical reaction involving liquids or gases.
Suppose you want to make a salt AB, where A+ is the cation and B– is the anion. So we take a solution that contains A+ ions and takes another solution that contains B– ions, then we mix the solutions in a beaker, hence we get the precipitate of the salt AB.
But to get an A+ ion in a solution, we must dissolve a soluble salt in the water which would release the A+ ions, the same thing goes for B+.
From our previous lessons now we know that silver chloride is an insoluble salt, so we must follow the precipitation method to make it.
We must take a solution that will contain Ag+ ions, so we can dissolve silver nitrate (AgNO3) to obtain it. Because all nitrate salts are soluble in water, so they would dissolve to release ions of Ag+
We also must take a solution which will contain Cl– ion, so we can dissolve sodium chloride (NaCl) to obtain it, because all sodium salts are soluble in water, so it would dissolve to release ions Cl–
Overall equation: AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)
Ionic equation: Ag+(aq) + Cl–(aq) → AgCl(s)
Note: The ionic equation is known as a precipitation reaction.
Silver nitrate contains silver ions and nitrate ions in the solution. The positive and negative ions are attracted to each other but the attractions aren’t strong enough to make them stick together. Similarly, the sodium chloride solution contains sodium ions and chloride ions- again, the attractions aren’t strong enough for them to stick together in the water.
When you mix the two solutions, the various ions meet each other. When silver ions meet chloride ions, the attractions are so strong that the ions clump together and form a solid. The sodium and nitrate ions remain in the solution because they aren’t sufficiently attracted to each other.
Take a solution containing Ba2+ ions, for example, Barium nitrate [Ba(NO3)2]
Take solution containing SO42- ions, for example, sodium sulfate [Na2SO4]
Mix the solutions in a beaker and the precipitate of BaSO4 will be formed.
Overall equation: Ba(NO3)2 (aq) + Na2SO4 (aq) → BaSO4(s) +2 NaNO3(aq)
Ionic equation: Ba2+(aq) + SO42-(aq) → BaSO4(s)
From the overall equation, we can see that the precipitate of BaSO4 is in the solution of sodium nitrate,
Firstly we must filter to remove the BaSO4, and then wash it with distilled water to remove any solution
clinging on to it. Then just leave it to dry.
Take a solution containing Pb2+ ions, for example lead(II)nitrate [Pb(NO3)2]
Take a solution containing I– ions, for example, sodium iodide [NaI]
Mix the solutions in a beaker and the precipitate of PbI2 will be formed.
Overall equation: Pb(NO3)2(aq) +2 NaI(aq) → PbI2(s) + 2NaNO3(aq)
Ionic equation: Pb2+(aq) + 2I–(aq) → PbI2(s)
Follow the same procedure to obtain the pure dry sample of lead(II)Iodide.
CaCO3(s) + HCl(aq) → CaCl2(aq) + CO2(g) + H2O(l)
CaCO3(s) + H2SO4(aq) → CaSO4(s) + CO2(g) + H2O(l)
In reality, the calcium carbonate reacts normally with hydrochloric acid, fizzing of carbon dioxide gas occurs. But in the case of sulfuric acid if we use big lumps of calcium carbonate, nothing seems to happen, we may see a few bubbles of carbon dioxide only, but why?
Answer: If we recall, we can see that calcium sulfate is insoluble in water. When the calcium carbonate reacts with the sulfuric acid, the insoluble calcium sulfate forms a layer of insoluble calcium sulfate around the calcium carbonate, which prevents the acid from getting to the calcium carbonate and reacting with it.
Thank you for reading this informational blog. We have covered a ton of important concepts about salts. These will definitely help you in your exams, doesn’t matter what level of chemistry you are studying.
If you have any queries, please feel free to comment below with your questions.
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Before you leave, have a look at the other chemistry blogs I have written. You may find something useful!
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