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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We cannot imagine life without water. All living things depend upon water in various degrees, and for a good reason. The reason is, water is a universal solvent. What does that mean, and why is water a universal solvent? Today we will set out to explore this.

Water is, of course, the most common chemical compound in the world. Hell, two-thirds of the planet’s surface is covered by it! And there are few things occurring in the natural world that do not get dissolved in water, to some varying degree. And that exactly is why we call water a universal solvent ‒ it dissolves mostly everything.

This is an extremely important matter in the cycle of nature. Water dissolves all the required salts and nutrients and other chemicals required to support life and carries them from one point to another. All life depends upon this phenomenon. For example, plants are built to suck water out of the ground and there is all the food dissolved in that water that the plant needs. Water is the largest part of the blood that circulates through all living things’ bodies, carrying nutrients to every cell and waste material out of the body as urine and sweat.

The definition of a universal solvent is difficult to give. Naturally, you’d think that means a substance that can dilute anything. But as you can further imagine, this is far from practical ‒ every solvent will have at least one (hundreds!) similar chemical with similar properties that will be inert in that solvent, not a solute. So we have to widen our definition of a universal solvent as the substance that can dissolve the most amount of chemicals.

Water is, by and far, the best solvent in the world. It can make solutions out of most compounds, save a few organic materials. The reason for that is that water is one of the best polarized compounds capable of breaking up the ionic bonds present in most compounds.

What does that mean? Let’s discuss.

The proper chemical name of water is Dihydrogen Oxide, and the chemical symbol is H2O. In this compound, there are two hydrogen atoms and one oxygen atom bound by a covalent bond. This means that this compound should be perfectly balanced electrically. The oxygen atom, however, is somewhat greedy and pulls the shared electrons closer to itself. This results in a slightly negative charge on the oxygen’s side and an equal positive charge on the hydrogen atoms’ side.

Now, that is very attractive to chemicals that have bonds based on electrical charge; that is, ionic-bonded compounds. Let’s take the example of Sodium Chloride, the common salt. The copper part of each molecule of NaCl carries a strong positive charge, while the Chloride is the anion with an equal negative charge. Normally, they are connected to each other in this opposites-attract sort of situation and form crystals

When you drop some of that crystallized powder into some water, silent mayhem starts. The water molecules are polarized with one side negative and the other positive, and they normally live together in harmony aligned by that. But as soon as you drop something with a bit more charge ‒ like this NaCl in question ‒ the molecules break their communal hydrogen bond and rush the ionic compound.

The H2O molecules crowd around the cations and anions present in the salt, pulling them apart. All the water molecules near the cation Na+ spin to face the negative oxygen side to face the cation, while the water molecules near the negative Cl- ions latch onto them by the positive hydrogen side. 

The pull into the crowd of the H2O molecules grows stronger as more and more of them attach themselves to the salt ions. The Sodium and the chloride ions are very strongly bonded, but there are just too many water molecules around them hungrily clamping on. Like ants dismembering a dead cockroach, the anion and the cation are pulled apart mercilessly. 

Once the salt is broken down into its basic cation and anion, these clusters of ions with water molecules drift around since they are perfectly balanced now. So the salt becomes homogeneously spread all over the water body.

This ability of H2O to break compounds into ions is what makes it the best solvent all around. Obviously, this power is not the same for all chemicals ‒ there are some chemicals like salt and sugar that just love to get dissolved in water, and there are some carbonates, etc. that have great resistance against it. However, the dissolving power of water gets much more intense when it is a little bit acidic ‒ which is very easy to happen in nature. That is exactly what causes  natural corrosion in structures .

Water is such a good solvent that it can dissolve not only solids but also other liquids (duh) and even gasses. Yes, you heard that right. How do you think fish breathe? They get the oxygen that is dissolved in the water. By just being in contact with air, much of it gets dissolved in water. 

Frustratingly for scientists, this super-solvent capability of water is exactly why it is very hard to get it in completely pure form. Pure H2O is never found ‒ there is almost always something diluted in it, be it simple gases from the air. Truly, given pure enough water and enough time, it can even break down plastic which is supposedly impervious to water! 

This leads us to comment on one thing. Since water cannot, apparently, live alone, the more you purify it the more it will become hungry for salts. This is why it is not a good idea to drink water that is too pure ‒ it will dilute whatever salts and minerals and nutrients it can find in your body and take it away while going out, making you weaker in the process. 

That’s one of the biggest reasons why drinking water purifiers are intentionally made to never completely purify the water even when they can. Sure, they kill off all the germs in water (even that may not be a very good idea) but keep the salts in. Of course,  water purifiers used in laboratories are different ‒ they produce much purer water that is highly dangerous to drink. You can also buy  demineralized water from Labkafe which is pretty much free of any cations and anions.

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