We have already discussed why water is the best solvent all-around in a previous blog article; but what about the things that get dissolved in water? There are sugar and salts and various other stuff that get dissolved in water, but salts have the best solubility in water. Now, why should that be?

In chemistry labs you may have noticed that various chemical salts get dissolved in water so easily; you just have to shake the test tube with the salt and water to make an original solution . But at home, you may have seen that you have to work the sugar with a spoon a lot; even in hot water. Normal table salt gets dissolved much faster, doesn’t it?

There is a very good reason for that. But we have to delve deep into the matter of solubility to understand it.

The definition of solubility tells us that the solubility of a given substance is max how much of it can make a clear, transparent solution with a given solvent at a given temperature. Taking water as the solvent (it is, in most cases), we can see that inorganic salts are more soluble than anything else. The reason for that lies in the manner of how things get dissolved. 

There are two ways something can be dissolved in water, which depends upon the kind of compound it is. As you may know, there are two kinds of chemical bonds that form compounds: (a) Covalent bonds and (b) ionic bonds. Most inorganic salts are formed by ionic bonds and most other stuff are covalent compounds.

Water itself is a covalent compound. That means, each hydrogen atom in H2O shares an electron with the oxygen atom, which shares two of its own with the hydrogens. Similar stuff happens with other covalent compounds and they form generally very stable molecules. 

When a covalent solid compound like sugar touches water, there is no chemical interaction. There are a bunch of sugar molecules huddled together at one side, and touching them are a bunch of water molecules held much loosely together. The water molecules simply seep inside the gaps between the sugar molecules. The process is entirely physical and there are no chemical reactions or interactions going on here. This phenomenon can be quite slow, considering how well the covalent compound is held together.

What happens with ionic compounds ‒ is much more dramatic!

To begin with, there is no such thing as a molecule of an ionic compound. Surprised? Don’t be. The very nature of the ionic bond is that there is no actual bond between two ions ‒ they are just strongly attracted to each other. But unlike covalent compounds, there is no actual binding contract here. 

For this reason, ionic compounds form lattices and crystals much better than other compounds. Take common table salt for example ‒ Sodium chloride. Each Na+ ion had sacrificed an electron to become positively charged; each Cl- ion stole an electron from somewhere and got negatively charged. Opposites attract; and so all Na cations and Cl anions in the vicinity rush together and form a kind of structure where every sodium atom is surrounded by six chlorines and vice-versa. There is no single “NaCl” floating around as we may imagine. They are a joint family or nothing.

This is a pretty strong structure as solid things go. There are few things in the world that may entice a sodium or a chloride ion to leave the group and do something, unless the thing itself is much more charged than them. Like some strong reagent.

We’ll let you in on a little secret. Water is such a charged thing.

How can that be, you must be asking. Wasn’t water supposed to be a covalent compound? How can it be charged? Water molecules exist, don’t they?

Well, let us explain. There is one special thing about water, that makes it such a curious chemical. The oxygen atom in every dihydrogen oxide (water) molecule is a greedy person and attracts all the electrons around it towards itself; even those in the covalent bonds with the hydrogen atoms are pulled closer to the oxygen atom.

Naturally, with all the electrons crowding together around the oxygen atom, that side of the molecule becomes somewhat negatively charged. Whereas the hydrogen sides get positively charged in contrast. This makes water a ‘polarized’ compound ‒ one of the very few.

This is the very property of water that makes it so efficient at pulling apart most ionic compounds. Let’s consider our example again. When the sodium chloride structure comes in contact with polarized water, a silent mayhem begins.

The positive side of the water molecules latches onto the negatively charged chloride ions, and the negative oxide side sticks to the positive sodium ions. This by itself isn’t enough to break the attraction bond between Na and Cl. 

But there are so many water molecules crowding around each of those ions! Like a colony of ants killing a spider, lots of water molecules crowd around each atom, each contributing its own little force. At one point, the combined strength of the water molecules becomes greater than the force between the Na and Cl ions. Like ants pulling apart a dead cockroach, they are mercilessly torn apart.

Each anion and cation of a salt falling into water float around, packed within a bunch of water molecules latched onto them by small electrical charges. Effectively separated, the salt gets quickly spread all around the water body, dissolving better than anything else.

Almost all chemical inorganic salts available in the salt analysis practical of class 10-12 behave like this; that’s why these salts are so easily soluble in water. Of course, there are a few exceptions. Here is a table of common salt solubility for your convenience.

In the table above, the cations are arranged horizontally and the anions are arranged vertically. Thereby, you can combine the ions column-by-row and see clearly which salts will be soluble in water, which salts are insoluble in water (but they can be if encouraged by adding a little conc. HCl), and which are slightly soluble in water (meaning you will have to heat the water). 

As you can see, the board is mostly green across lighter molecules. This lets you predict which types of salts will be more soluble in water and which will be less. 

You can follow these rules of thumb to remember which salts will be soluble in water:

• All nitrates are soluble.
• All sodium, potassium and ammonium salts are soluble, except carbonates.
• If it’s a silver or lead salt, chances are it won’t be soluble.

These rules of thumb will let you make your way easier through the chemistry laboratory. Labkafe supplies most of those salts (soluble and insoluble) as part of the lab consumable package; but you can also buy them separately. 

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Last week we discussed the Salt Analysis experiment in chemistry lab class 12 practicals. While talking about performing various tests on the original solution or water extract of the salts, we realized that not everybody can be savvy about how to make the original solutions of salts, or, as they are also called, stock solutions. Hence, we get down to making water extracts of salts for school chem lab experiments.

You’re in luck with the salt ion identification reaction experiment because you have the best solute and solvents on hand.  Water is a universal solvent and salts are very soluble in water ‒ both for good reason. And so you shouldn’t have any trouble dissolving salts in water. However, as you well know, there are orders and methods to anything and everything; and so we will give you the correct instructions on how to prepare the original solution at the beginning of the salt analysis experiment below.

What is a Solution, Anyway?

Students often wonder about this in their studies of chemistry. Let’s talk about what exactly is a solution in the matter of chemistry labs, before we start talking about making them.

Put very simply, a solution is a mixture of a few things done completely evenly. Meaning, the solute (the thing in the smaller amount) should spread completely evenly throughout the solvent (the thing in the bigger amount) in the mixture. This is called a “homogeneous” state of the mix, which is a pretty stable state ‒ thing 1 and thing 2 won’t get separated unless you take some great pains.

To clarify, just mixing something and trying to make it even does not make it a solution. For example, we can take muddy water ‒ looks pretty even, doesn’t it? But just keep it still for a long time and you’ll see the mud precipitate down and the water becoming clear. But dilute some sugar in water and it will stay sweet forever.

If you want to be totally technical about it, we can say that in an actual solution, the solute (like salt) and the solvent (like water) mix so evenly and closely together that the molecules of the solute get in between the molecules of the solvent. If the solute and solvent keep together in their own groups, then no matter how evenly they mix ‒ it’s still not a proper solution. It would be what we call a colloid. 

We will get into more details about solutions, solvents, and solutes in other articles. 

How to Prepare an Original Solution (OS)

Ingredients and equipment required:

1. Given salt, dry, in a good amount
2. Distilled water, about 100 ml
3. Hydrochloric Acid, dilute and concentrated
4. A few  Test tubes
5. Beaker (small) with pouring lip
6. Bunsen burner (or spirit lamp works too)
7. Support apparatus (tongs, test tube clamps, tripod, wire gauge, etc.)

Objective: To prepare a clear and transparent solution of the salt under investigation. This is called an Original Solution (OS) or Initial Solution, Water Extract, or Salt Solution.

First Phase: Try with Water

At first, chances are that you don’t know what salt you have. Unless you have some visibly identifiable salts in your hand (like Copper Sulphate), you are somewhat in the dark. So, at first, we will not try to dilute the whole of the salt at once but try with just a little bit of it.

• Take just a small pinch of the given salt and put it in a clean, fresh test tube. 
• Drown it in about an inch of distilled water in the tube.
• Put a cork or stopper on the test tube to close the mouth well.
• Shake the tube strongly. Try both side-to-side and up-down movements. 

Observe the liquid in the test tube now by holding it at eye level. Is the stuff clear and transparent, or is it murky or opaque? Meaning, can you see things through the liquid in the tube? If yes, then you have a solution of the given salt in class 10/12 practical lab. If the solution is not clear, or the salt remains alone in the test tube, then you have to try something else.

• Light the spirit lamp or the bunsen burner. 
• Remove the cork or stopper and heat the liquid in the test tube gently.
• Is the liquid getting clearer? If not, try shaking it sidewise gently (be careful it doesn’t spill out).

If the liquid gets clear soon, you have a small amount of the original solution. But if it remains murky even till the water starts bubbling, then your salt is not soluble in water naturally and you will have to trick the water into being more, let’s say, persuasive.

Second Phase: Try with HCl

Water, by itself, is a very good solvent. Most salts tend to get dissolved in water. But you know what dissolves things better than water? Acid. So, if you’ve got a particularly nasty salt unwilling to dissolve in simple water, you have to spike the water with hydrochloric acid.

• Take the test tube from the previous step and let it cool. Or, discard the liquid and start fresh with another pinch of salt and distilled water.
• When cool, put one drop of diluted HCl in the test tube. 
• Close it up with a cork or stopper and shake well like before.

Does the salt get dissolved now? If yes, you’ll see a clear, transparent solution, that is OS quality. If it’s still murky, heat the tube like the previous step. Be very careful here since hot acid is not fun at all. Almost all school laboratory salts will get dissolved in this step.

If, for some reason, the solution still remains murky, add a drop of concentrated HCl in the tube (after cooling off) and repeat the previous steps. It is sure to form a clear, transparent solution now. If not, contact your lab instructor, teacher, or lab assistant ‒ you may have been given a wrong chemical. 

Note: You can get an idea about the salt’s solubility from this process. When you have to write about the salt’s character, remember this process:

• If you had managed to get the solution with cool water only, the salt is “water soluble”.
• If you had to heat the water to get the solution, then the salt is “partially soluble in water”.
• If you had to add acid to the water to get the solution, then the salt is “insoluble in water”.
• If you didn’t get a solution even after all the process, then the salt is probably organic.

Third Phase: Prepare the Original Salt Solution

Assuming you had managed to dilute the given salt using the process above (without accidents), then you know how to dilute the salt. Now you are going to do it in volume so you have a decent amount of the Original Solution in hand to use in various identification tests.

Take a 10ml beaker (with pouring lip), and about 3-4 grams of the given salt. Make sure you have some dry salt put aside to do the dry experiments like the flame test, charcoal block test, and other dry tests.

• If you had had a solution in cool liquid, just mix them together in the beaker (add the HCl if required), stirring with a glass rod.
• If you had to heat the test tube to get a solution, set up the burner or lamp and the apparatus to heat the beaker (tripod, wire net, etc.). Then heat it up, stirring with the rod until you get a clear solution. Add the acid if required in the previous steps.

Note: an Original Solution or Initial Solution of the given salt will have a similar color (maybe less intense) as the given salt. If the salt was white, the solution should look like clear water.

Congratulations, now you have the original salt solution (or water extract) and you can move forward to the next step in the salt analysis process. Labkafe sells all kinds of chemical salts, bases, acids, and reagents as part of lab consumables for your chemistry and bioscience experiments. They are available as per-board lab packages or individually. You can order lab consumables online from us, or  get a quote for the full or partial packages.

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Your school laboratory has a lot of chemical compounds to work with. Many of them are solid salts used for various reasons. If you are to be a good chemist or at least want to score some fame in the chemistry lab, then you have to be able to identify various compounds by the color of salts.

The thing is, many of the school laboratory salts are not much to look at ‒ just white powders. But a few important ones have very distinct colors and/or textures. It would be quite humiliating not being able to name them at a glance! Fear not, Labkafe Learning Center is here with a guide to identify common salts by color.

A word of warning, though. This is in no way a perfect idea of identifying a chemical compound ‒ for that, you need tests like the salt analysis . 

Why Do Salts Have Color?

A very good question! Why does anything have color? We could write a whole book on it (actually there are multiple books written on the subject of atomic spectroscopy, the subject of studying colors of things). But we don’t have that much space here, nor do you have the time. So, here’s the absurdly simplified version of how salts get their color.

How do we see colors? They are nothing but different wavelengths of light. When some light falls on any object, it absorbs some of it and reflects the rest. That is how we see an object. That part that got reflected is what we see, those particular wavelengths that the substance didn’t eat up. And we call that the color of that object. For example, a blue salt eats up all other colors but reflects only the blue waves of light, so we think the salt is blue.

Now the question comes, why does a substance absorb only some parts of light? The answer lies deep in the movement of electrons of the ions that make up a compound. Simply put, electrons are not very calm people and they keep hopping around in different orbits. Have you ever hopped around? Then you know it requires a great deal of energy. Electrons, too, require energy to jump from one orbit to another. And light is a ready-made energy source available to them most of the time, so they eat it up in volumes!

About why the jumpy electrons would eat only a certain wavelength, there is a good answer too. When you hop from one place to another, you will notice that all the jumps are not the same. Each new place you land pretty much dictates how much energy will it require and in what way you’ll jump and whether you will break your leg or not. The same is true for electrons ‒ when they jump from one orbit to another, the nature of that movement dictates the nature of the light that will get absorbed in the process.

For example, in dry cobalt chloride (CoCl2), the electrons in the cobalt atoms are quite unhappy with their space in the D and F orbitals. So, they try to move to more stable places and these jumps consume most of the yellow, red and green wavelengths of light. Only the blue parts are left and that’s what gets reflected. Thus, we see cobalt chloride as a blue powder.

Funny enough, add a few drops of water to that dry CoCl2 powder and you’ll see it promptly turn purple-pink. That is because now water molecules intervene in the crystalline structure of the salt, giving the Cobalt electrons some new kind of buzz. They are still jumpy, but now they hop to and from different orbitals that consume some blue and most of the yellow and green light. This returns only red and some blue, mixing up and looking pinkish purple. Or purplish pink. Or onion color. Or whatever, I’m not good with color names!

Colors of Some Common Salts

Copper Sulphate

This is probably the most easily recognizable compound in most chemical labs. We know it as a brilliant blue crystalline solid. Note that this is the ‘wet’ version of the compound with five molecules of water attached to each molecule of CuSO4. Otherwise, in its anhydrous form, the ‘dry’ copper sulphate looks like a fine powder of pale blue color. It absorbs the 750 nm wavelength of light the best.

Ferrous Sulphate

Anhydrous ferrous sulphate is clear white and powdery, but it absorbs water to become crystalline ferrous sulphate (FeSO4.7H2O) which looks like light bluish-green salt. Heat it a little bit to see the color disappear. Alchemists of the olden days called it green vitriol or iron vitriol. Note that this substance can create a variety of stable forms with different levels of water absorption.

Cobalt Chloride

We have already discussed Cobalt Chloride above ‒ it is a fine blue when dry but forms pinkish violet crystals with water. It is very much hygroscopic ‒ means it will absorb water when left in the air and you can see the blue turning red slowly. It also has a light sharp odor.

Ferric Chloride

While yellow in color naturally, Ferric Chloride (FeCl3.6H2O) will absorb more water from the air if left alone for a long time and become brownish in color. Which is also the color of its watery solution. But if you heat the salt to dry it out and evaporate the water molecules, it will turn very black absorbing most of the light that falls on it.

Potassium Dichromate

One of the few pure chemical substances with such a bright orange color, Potassium Dichromate (K2Cr2O7) is easily recognized in the school chem lab. Unlike most other chemicals we talk about here, it doesn’t absorb water from the air ‒ but it is extremely toxic! Don’t even touch it with your bare hands. Very aptly colored salt, won’t you agree? Because dangerous things, in this world, are generally bright-colored. Except for flowers, of course.

Copper Chloride

Also known as Cupric Chloride (CuCl2·2H2O), it is a fine greenish blue crystalline salt. When heated, it loses the water and turns yellowish-brown. It gives the same blue-green color in a flame test too. For this exact reason, copper chloride is used in fireworks to make that azure color. Be warned, though ‒ it is not a safe compound and can be toxic to living things. Be careful not to confuse Copper Chloride with Copper Sulphate ‒ those two are very close in color and texture and it is hard to point which is which by just looking unless you put them side-by-side.

Nickel Chloride

It is hard to get the anhydrous yellow form of nickel chloride, but we can easily get our hands on NiCl2·6H2O in our school labs, which is a nice green-colored grainy sticky powder. Did you know that nickel chloride is used to electroplate metal objects? Its cousin, nickel iodide, also has the exact same color but looks more crystalline.

You may often need to handle these salts in your chemistry laboratory. Please do be careful when you use them, especially in their dry (anhydrous) forms. They can be harmful to your bare skin. And for the love of your life, don’t try to directly smell, or, heavens forbid, taste them. Speaking from personal experience, that’s NOT a good idea. 

So, one may naturally ask, where can I get these colorful salts? The answer is Labkafe. We have all of these salts ‒ and many, many more ‒ as part of our  laboratory consumables supply. You can buy them separately or buy them as curriculum-based or experiment-based packages.

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