6 Water, water, Everywhere
6.1 Water: Properties and Interactions
6.2 Acid/Base Theory
6.5 pH and pOH
6.10 Weak Acids and Bases
6.11 The Water Around Us
6.42 Learning Outcomes
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Time to think about water and just how important it is to pretty much everything here on earth. I need it, you need it, plants need it, animals, the whole earth needs it to "operate" correctly. I'll spare you more drama on it in that respect. I'm really after the whole chemistry of aqueous solutions thing since this is a chemistry class.
Pure Water? Is there really such a thing as pure water? Well yes... but it is really really difficult to just crank out ultra-pure water. It certainly doesn't come from nature pure. Not even rain falling from the sky is pure water. So when I say pure water I'm referring to a rather special reference standard that we often approach, but rarely truly achieve. Truly pure water is nothing but H2O and that's it. But there is more to this than meets the eye...
Water is a very polar molecule with a good partial positive charge (δ+) on the hydrogen atoms and a partial negative charge (δ–) on the oxygen atoms. This in and of itself is why H-bonding is such a big deal and why one of the smallest molecules in the universe has a relatively high melting point and boiling point. Water is only 18 g/mol for its molar mass and is a liquid at room temperature with a boiling point of 100 °C. This same polarity makes water an excellent solvent for hundreds of polar compounds and salts.
Salt dissolves in water by dissociating into its separate cations and anions:
NaCl(s) → Na+(aq) + Cl–(aq)
Let's dissect that equation: Solid NaCl (salt crystals like you sprinkle on food), when put into water, will completely break up into separate aqueous ions. The (aq) in the above equation is the aqueous phase of matter. It really isn't a true independent phase like solid, liquid, or gas are for pure substances - but it is a very very important homogeneous mixture (a solution) in all of science and therefore gets its own (aq) designation.
The aqueous phase is really a species (cation, anion, or neutral species) that is completely surrounded by molecular H2Os . Below is an illustration of both cations and anions being pulled out of the salt crystal to form aqueous ions. The process is called dissolution. This is just a 2-dimensional representation. Remember, it is really occurring in 3-D.
This is a hydrated cation or Na+ in the case of NaCl dissolving. Notice how the oxygens (red atom) are oriented towards the cation. This is due to the large attraction of the positive charge on the cation and the partial negative charge on the oxygen. Hydrated ions like this are best depicted in equations as ion(aq) or in the case of sodium ion, Na+(aq). It is a lot easier to stick on the (aq) than it is to count waters of hydration and explicitly putting them in the formula like this: [Na(OH2)6]+ oof! no thanks, I'll stick with Na+(aq), thank you.
This is a hydrated anion or Cl– in the case of NaCl dissolving. Notice how the hydrogens (gray atoms) are oriented towards the anion. This is due to the large attraction of the negative charge on the anion to the partial positive charge on the hydrogens.
Also know that I've only shown the water molecules that are participating in the dissolution. There would actually be way more molecules present - so much so that our illustration would be totally buried in them.
We make a really big deal about H-bonding in certain sets of molecules. It is like the strongest possible dipole-dipole interaction there is. So much so, it got its own fancy name of H-bonding. Still, even the best H-bonding interaction is small compared to actual ion-dipole interactions. A full blown charge of +1 or -1 is going to have far more pull than any partial charge of δ+ or δ–. So once an ion is in aqueous solution, it typically stays that way until something far stronger comes along to disrupt it.
Check out the Wikipedia link for "aqueous solution" in the nav-menu (margin). There is a lovely illustration of a solvated (hydrated) sodium ion on that page.
Remember, back in section 4.7 I pointed out the details of three important interactions between molecules. These intermolecular forces (IMFs) are present pretty much everywhere. Dipole-dipole interaction is the main interaction occurring between polar molecules, while dispersion forces are the primary ones for most large molecules, and is the only force for non-polar molecules. H-bonding is a special case of dipole-dipole that happens to have relatively strong interactions compared to other IMFs and is the major IMF in water. So that's my little recap of that chapter 4 section.
Because water is such a major molecule here on earth, we often have terminology just for it. The attractive/repulsive forces within water are a major player in the properties of aqueous solutions and any water-like species. For this reason, we often categorize forces of attraction with water as being either hydrophobic (water-fearing) or hydrophilic (water-loving). These are water-specific terms but are really just more IMFs when you really look deeper. But sometimes it is far easier to talk about aqueous interactions as being either dominated by hydrophobic interactions or hydrophilic interactions. I'll explain each below.
The general rule-of-thumb for any hydrophilic species is that is tends to have very favorable (attractive) interactions with water. Simply put, the species is most likely a polar species. Water loves polar compounds. The partial postive (δ+) and partial negative (δ-) regions of water will line up with a considerable amount of attractions with polar species. This certainly increases the solubility of any polar compound or species. As long as there is nothing even more attractive around, most polar substances will dissolve in water. The water molecules will completely hydrate each and every molecule of the species and pull the species into the aqueous phase. I showed this in detail with NaCl in section 6.1 through the process of dissolution. Even though I showed full blown cations and anions with NaCl, the same does occur with any polar molecule. Maybe not as strong as NaCl, but the interaction is very significant.
Non-polar molecules (hydrocarbon chains and the like) really have nothing substantial for water to grab hold of. Water is far more attractive to itself than to non-polar molecules or regions of non-polarity. This generally causes the non-polar molecule to just stick to itself or other non-polar species instead of water. Water will generally "stay outta the way" of these non-polar interactions. This is why it is called a hydrophobic interaction. Typically, this just leads to the general rule that non-polar compounds are insoluble in water. Greases and oils (mainly long chain hydrocarbons) will bead up on water and float on the surface (less dense than water). As they say... "oil and water don't mix" and that is true because the oil is a very hydrophobic substance. There is even a hydrophobic effect which is when these hydrophobic species find each other in aqueous solutions and congregate/aggregate building little islands and globs of non-polarity. This is further discussed in the emulsification section further down this page.
All of this leads up to the chemists rule-of-thumb about solubility in general:
"Likes dissolve likes"
So simple, and so true. If you want to dissolve greases and oils, get yourself a non-polar solvent like carbon tetrachloride or hexane. Many sticky adhesives are hydrophobic which makes cleaning off the sticky residue difficult with any aqueous solution. But using a more non-polar solvent like original version of Goof-Off or Goo Gone will dissolve that sticky stuff pretty quick. But if your sticky stuff is aqueous based - like a sugary mess (maple syrup) - then by all means, use water. Even better, use warm or hot water because the dissolving process for most solids in water is endothermic which means a little heat will help things get pushed more forward (dissolving).
So oil and water don't mix. Right? Yes, oil will completely separate out and float as a separate layer ontop of water. You can mix, stir, shake, whatever... but you will not get them to stay mixed because of the very real set of hydrophilic (water-to-water) and hydrophobic (water-avoiding-oil) forces at play. But we can get the two to mix in a very clever way. Add an emulsifier. An emulsifier is a large molecule with two distint regions on each end. One end is a long-chain hydrocarbon and is known as the hydrophobic "tail". The other end is typically a deprotonated carboxylic acid or a sulfate group. That means those ends are highly polar and even fully charged negative (not just a partial charge). This makes the end very hydrophilic and is known as the "head" group. So the "head" end loves water and will have favorable interactions while the "tail" portion will flee from water and just try to find more hydrophobic tail groups to strengthen those dispersion forces. This is the building block for making micelles. Below is an example using stearate ion (the deprotonated form of stearic acid) showing how the actual line structure is depicted in a schematic drawing with a circle for the head group and a long line for the tail group.
A micelle is a relatively large molecular super-structure that is generally spherical and is a bunch of these emulsifier molecules like stearate making the sphere. The entire surface of the sphere is the head ends of the molecules - all pulled together with a full water layer reinforcing the structure. At the same time, the tail groups are all maximizing their interactions with each other to build the interior of the micelle.
So the overall picture of a micelle is a little (super little as in you can see them) micro-droplet of "oil" (the tails) with an outer shell of water-loving charged heads. Below is a schematic drawing/rendering of a typical micelle (a cross-sectional view). The same schematic of the stearate ion shown above is now used 24 times to create an entire circle. Realize that it actually makes and entire sphere and not just a ring.
If you add enough emulsifier such as stearate, micelles will form and any oil floating on the water can easily be emulsified which means the oil gets broken into very very small oil droplets (smaller than your eye can see!) that are in the "cozy" hydrophobic environment of the center of the micelle. So when you shake/stir/mix up the oil and water, you get a stable emulsion which to the eye looks like a homogeneous solution. It looks as though everything completely mixes and is soluble. In reality, the oil is just broken up into tiny-ass droplets encased by the emulsifier. This is show below.
BTW... if you are in mostly oil and just a little water, the micelle will invert which results in the hydrophilic heads all congregating to the center (avoiding the non-polar oil) and the hydrophobic tails pointing out into the oil. Neato. Have a look...
What is one of the most common emulsified substances of oil and water? Mayonnaise. Oil is emulsified with an egg. The egg provides the molecules that act as emulsifiers. Use lemon juice for your "water" to get the right taste and mix thoroughly. Bam - mayonnaise. Keep in mind, not all emulsifications end up thick and spreadable. But many do. Lotions and creams are examples of cosmetics that are emulsions (they have micelles in them).
There are many other examples of getting hydrophobic substances together with hydrophilic substances. Soaps and detergents are excellent emulsifiers and help you get clothes and yourself clean.
Ok - that's it for now. There is so much more on this but I'm taking a rest. Hydrophobic/hydrophilic interactions are important across many disciplines. Read up on some.
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