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:
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.
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.
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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