Reading Notes for Chapter 11

These are Dr. Bodwin's reading notes for Chapter 11 of "Chemistry 2e" from OpenStax. I am using a local .pdf copy that was downloaded in May 2020.

Chapter Summary:

Although we can learn a lot from studying pure compounds, the vast majority of the real substances we encounter every day are mixtures. Ofthese mixtures, those that can be described as solutions are the most useful because they have some unique qualities. When thinking about solutions, it is still very helpful to look at the intermolecular forces (IMFs) present in all the components because those IMFs determine most of the solution's behaviour.

Formation of Solutions:

Back in Gen Chem I, you probably learned some "solubility rules". It might not be a bad idea to review those now because we're going to be looking at solutions in much closer detail. Having the big-picture "sorting rules" in mind will help us discuss the reasons they are true, and sometimes not true.

Types of Solutions - just a little aside here... take a moment to sit back, close your eyes, and picture a solution being made. No, really, do it...
OK, do you have that picture? Most of you probably pictured a solid (that probably looked like white crystals or powder) being added to a liquid (that probably looked like water). If you were feeling "sciencey", your picture might have been taking place in an Erlenmeyer flask, maybe even with a magnetic stir bar. That's a great picture of "solution", but remember that a solution is "a homogeneous mixture of 2 or more components"... that definition doesn't say anything specific about states of matter. A solution can be a solid dissolved in a liquid, or a liquid in a liquid, or a gas in a liquid, or a gas in a solid, or a solid in a solid... any combination of states of matter can make a solution under the right conditions.

So again, why do solutions form? Think about those intermolecular forces, and more importantly the balance between the IMFs of the solute and the solvent. A couple cases:

If the IMFs between solute particles are extremely strong compared to the IMFs between solute and solvent, then the (relatively) strong solute-solute IMFs will be hard to break and the solute is unlikely to dissolve.
If the solute-solute IMFs are (relatively) weak compared to the solute-solvent IMFs, then the solute is likely to break apart and dissolve quite readily.
If the solute-solute and solute-solvent IMFs are ofsimilar strength, then the solute is likely to dissolve only slowly or only in a limited amount.

When we're forming a solution, the microscopic processes that are taking place are: we must break solute-solute interactions, we must break solvent-solvent interactions, and we must form solute-solvent interactions.Breaking bonds (or IMFs) requires energy; forming bonds (or IMFs) releases energy. The balance between those energies determines whether or not a solution will form.
When a solution has reached its limit of solubility, it is said to be saturated. If more solute is added to a saturated solution, it will not dissolve.

Gases dissolved in liquids:

This is an important one. It's how fish breathe. It's how YOU breathe! It's also one that most people see first-hand every day. Open a container of carbonated beverage and you see Henry's Law in action!
Keep thinking about IMFs, and some of the things that we looked at with vapor pressure or pure substances... Temperature affects the solubility of a gas in because at higher temperatures, the particles are moving faster, so they are more likely to escape form the solution. Back to the carbonated beverage example: open an ice-cold soda and a room temperature soda... which one "fizzes" more? Open 2 sodas. Put one on the room-temperature countertop and put the other in the fridge. Come back in 45 minutes and take a drink of each... which one is more flat?

Liquids dissolved in liquids:

We don't always thing about these as "dissolving", we use terminology like "mixing". Liquid-liquid systems are described as "miscible" (able to mix) or "immiscible" (unable to mix). Think of water and vinegar (miscible) versus water and oil (immiscible).

Solids dissolved in liquids:

Unlike gases in liquids, increasing the temperature usually increases the solubility. This is not always true (see Figure 11.16 in your textbook), but when in doubt, heat up the mixture to encourage solids to dissolve.
This is also a good place to circle back to "saturated" solutions. Our Gen Chem I sorting rules just plop everything into 2 categories - "soluble" and "insoluble". Sodium chloride is soluble. If I take a spoonful of table salt and add it to a glass of water, it dissolves. What if I take a spoonfu. of water and add it to a glass full of table salt? Even with heat, that's not all likely to dissolve. Gen Chem I "sorting activities" are a good place to start, but they need to be used responsibly.

Concentration Units:

When discussing solutions, it's handy to have a concentration unit that describes the relative amounts of the various components. Molarity (M) is the one we've used most, and will continue to use most, but there are others that become important in different settings. That's why it's so important to include units and be very careful to use the units in any calculations we might do.

Fraction-based Units - Divide the amount of what you're interested in by the total amount. That's a fraction. Multiply by 100, it's a percent. Multiply by a million, it's a part-per-million. By a billion? Part-per-billion. These aren't as mysterious as people try to make them. The biggest challenge is to define what you mean by "amount". If you use moles, it's a "mole fraction". If you use grams, it's a mass fraction. Milliliters? Yep, volume fraction.
Molarity - "M" - This is the one you know and love, moles of solute per liter of solution.
Normality - "N" - This is very closely related to molarity. Not used very often, but still pops up once in a while.
Formality - "F" - An old-fashioned version of molarity used mostly for soluble ionic compounds. Useful because it takes into account the relationships in the balanced chemical formula, but almost never actually used. Most sources just talk about "molarity of nitrate ions in in a calcium nitrate solution" rather than call it "formal concentration"... but be careful... "formal" is a nice, common English word that can be used as a modifier of the word "concentration" in other contexts...
Molality - "m" - This is why units matter and careful reading and writing of units matter. Molality is moles of solute per kilogram of solvent. Numerically, it is often close to molarity, but it is always just a little bit different. Molality is important because mass (kg of solvent) is a measure of the amount of material that is present, while volume is a measure of how much space that matter occupies. Volume changes when temperature changes, mass does not.

There are other concentration units, but as long as you're careful about reading and consistent about writing concentration units, you'll be fine.

Colligative Properties:

Some properties of solutions depend on the number of "pieces" in the solution, not really on the identity of those pieces. These are colligative properties.

Vapor Pressure Depression (VPD)- When a non-volatile solute is added to a pure solvent, some of the solute particles get in the way of solvent particles that are trying to escape from the surface of the solution into the gas phase. If less solvent can escape into the gase phase (but it can still condense equally well into the liquid phase), there will be less solvent particles in the gas phase, therefore the vapor pressure will be lower in the solution that it was in the pure solvent.We use Raoult's Law for vapor pressure depression. NOTE: remember that Raoult's Law uses the mol fraction of the solvent.
Boiling Point Elevation (BPE)- A liquid "boils" when its vapor pressure is equal to the atmospheric pressure. If adding a solute decreases the vapor pressure, then we must have to heat the solution up more to reach the boiling point. The expression for calculating BPE is a very simple equation, keep an eye on the units of the constant to guide you. BPE uses molality, m, as a concentration unit. NOTE: The equation shown on p625 of your textbook calculates the change in the boiling point. Make sure you remember that when answering a question... I've had many people tell me that an aqueous solution boils at 1.74°C because they report the change rather than the actual boiling point of 101.74°C!
Freezing Point Depression (FPD) - When a pure liquid freezes, the particles slow down, align and stick together to form a crystalline solid. When there are solute particles in the way,the solvent particles have to slow down more so the solution has to be brought to a lower temperature. The expression looks just like the BPE expression and works the same way. NOTE: Different sources use different conventions! Some put a negative sign in the equation (because the freezing point is depressed), others put a negative sign on the constant. Your book doesn't use a negastive sign anywhere (which is how I prefer it) and relies on you to know that the freezing point goes down. NOTE2: Because we often work in water (whose freezing point is 0°C), it's easy to get a little complacent, but remember that the equation on page 628 of your textbook is still calculating a change in freezing point.
Osmotic Pressure - This one's a little differen from the other three listed above, but it's still a colligative property. Osmosis is the tendency for solvent concentration to equalize across a semipermeable membrane. Your textbook has some excellent descriptions and U-shaped tube graphics that illustrate osmosis and osmotic pressure. Osmosis is the pump that makes all life work. Your cells function because of osmosis. Plants live because of osmosis. Osmotic Pressure is the pressure that is required to stop osmosis. That's probably easiest to visualize in a U-shaped tube model (Figure 11.24 in your textbook). For a fun at-home experiment, try making a "Naked Egg" by soaking an egg in vinegar for a few days (with occassional scrubbing) to dissolve and remove the shell. The inner membrane of an egg is semipermeable. Soak it in pure deionized water and it swells; soak it in highly salted water and it shrivels (much like the blood cells in Figure 11.27 of your textbook). Science!

Since all of these colligative properties depend upon the number of solute particles and not their identity, ionic solutes provide an interesting extra bit for us. Since ionic solutes dissolve and dissociate to form ions, each mole of ionic solute formula provides more than 1 mole of solute particles. The number of particles an ionic solute forms in solution is given by the van't Hoff factor, i.

Colloids:

Colloids are an interesting intermediate phase that's not quite a solution, but not quite a suspension. That makes them interesting (and challenging!) to study. As you can see from Table 11.4 in your textbook, many colloids are beautiful or delicious (or both) and you encounter them every day.

Amphiphiles:

One convenient rule of thumb for solutions is "like dissolves like" meaning that substances with polar IMFs tend to dissolve in solvents with polar IMFs, while substances that only have dispersion forces tend to dissolve better in solvent that only have dispersion forces. Vinegar (polar) dissolves in water (polar), but oil (non-polar) does not dissolve well in water. How do you get grease off of a plate when you're doing dishes? With a soap! Soaps are amphiphilic, meaning they can interact with both polar and non-polar substances. On a molecular level, this is because they have a polar, ionic end and "greasy" end. They form micelles that trap greasy bits.

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