Reading Notes for Chapter 10

These are Dr. Bodwin's reading notes for Chapter 10 of "Chemistry 2e" from OpenStax. I am using a local .pdf copy that was downloaded in May 2020.

Chapter Summary:

This chapter is formally on liquids and solids, but there are a LOT of aspects of liquids and solids that make more sense or are clearer when we also think about gases and their behavior. If you haven't looked over gases recently (or would just like to brush up), check out Chapter 9.

States of Matter:

Before getting too deep into gases, let's look at the bigger picture.
In Gen Chem I, we often engage in what I call "sorting exercises". We define a couple bins (usually no more than 3 or 4) and figure out how to sort the world into those bins. "States of Matter" is one of those sorting activities... look back at Chapter 1. We made our "bins" (solid, liquid, gas, plasma) and tried to classify the world around us into those bins.
"Sorting exercises" are designed to answer "what" questions. "What is this?" - "It's a solid."
Let's move on to "how" and "why" questions.
Matter at the microscopic scale experiences two competing forces - kinetic energy and intermolecular forces. Kinetic energy (E_kin) is the energy of motion, and all matter that "has temperature" is in motion. That motion may only be vibrations of the molecules, but it's still motion. Intermolecular forces (IMFs) are the forces that attract moleclues to each other.
States of matter are the balance between E_kin and IMFs - if a molecule is moving "slowly" (low E_kin) and has "strong" IMFs, then the particles are likely to stick together and be solid; if a molecule is moving extremely fast and has very weak IMFs, then the particles will fly past each other before they stick together and be a gas.

Intermolecular Forces (IMFs):

There are a number of different intermolecular forces with varying strength. In general, their strength from strongest to weakest is in the following order.

Ion-Ion - These are sometimes called "Coulombic Forces", and are (usually) very strong compared to other IMFs. This is why most ionic compounds are solids; the IMF is strong enough to overcome a lot of kinetic energy.
Dipole-Dipole - Polar molecules have a region that is slightly more positive and a regiuon that is slightly more negative. These are also coulombic forces, but because we're dealing with partial positive and negative charges, dipole-dipole interactions are (usually) not as strong as ion-ion interactions.
Hydrogen bonds - Most books discuss hydrogen bonds as if they are a unique IMF, but a hydrogen bond is just a special case of a dipole-dipole interaction. When hydrogen is bound to a significantly more electronegative element (N, O, F especially), the bond is polar with a partial positive charge on H. Because hydrogen is so small, that "partial" positive charge is actually quite concentrated, so the dipole-dipole interaction can be very strong. Because this is a special case, it's been given its own name "hydrogen bonds", but don't forget that this is just a special case of dipole-dipole interactions.
Induced dipole-Induced dipole - These have a number of different names: London Forces, Dispersion Forces, London Dispersion Forces. These are all the same thing. An induced dipole is a dynamic condition in which a non-polar electron cloud distorts to appear polar for a moment. This is easier to do with a large and relatively "squishy" electron cloud, so larger atoms tend to have stronger London forces. For smaller atoms (like carbon), the individual London forces are relatively small, but if there are a lot of carbon atoms connected together in a large molecule, the accumulated London forces can be quite strong. You can lift up a car with thread as long as you have a lot of individual threads. Everything has London forces.

Properties of Liquids:

Liquids have an even more delicate balance between IMFs and Ekin... a small change in Ekin (temperature) can tilt the balance to solid or gas. Because of this, liquids have some unique properties.
Cohesive forces vs Adhesive forces - Many properties of liquids can be explained as a balance between these two forces, and can be simply expressed as "are the particles of the liquid more attracted to each other (cohesive forces) or to the surroundings (adhesive forces)".
Viscosity - viscosity is a liquid's resistance to flow and is a good measure of the cohesive forces in the substance.
Surface tension - In the "middle" of a sample of liquid, all the particles are are surrounded by other similar particles, but at the surface, the particles are interacting with both other liquid particles and particles of the "other" substance. Usually when we think about surface tension, that "other" substance is air.
Capillary Action - When a thin tube is inserted into a liquid, the liquid tend to "climb" up the inside of the tube via capillary action. In some cases, the liquid is pushed down the tube instead of climbing up. Why? Look at the relative strength of the liquid-liquid IMFs (cohesive forces) and the liquid-surface IMFs (adhesive forces). Why doesn't a liquid just keep climbing up the tube until it squirts out the top? There's another important force at work here,,, gravity. The liquid climbs as high as the adhesive forces are able to uvercome the gravitational forces working on the column of liquid.

Phase Changes:

It is important to remember that the temperature of a sample is a measure of the average kinetic energy of the particles in the sample. There are some particles with higher kinetic energy and some with lower. Keep those extremes in mind!
Vapor pressure is an excellent example of the extremes. If a particle with relatively high kinetic energy is located at the surface of a liquid, it may escape to the gas phase. Likewise, if an especially slow gas particle is near the surface fo a liquid, it might be attracted to the surface and be trapped in the liquid. When the rate of liquid particles escaping is equal to the rate of gase particles condensing, the system has reached a dynamic equilibrium state and the observed pressure is constant. That is vapor pressure, and it is a nice way to measure the intermolecular forces in a liquid - the stronger the IMFs, the lower the vapor pressure.
Categorize all the phase changes as either endothermic or exothermic... do they absorb heat or do they release heat.This is important because we sometimes misunderstand due to our experience. An example of this mix-up: we typically think of endothermic reactions as "feeling cold". When water freezes to form ice, it feels cold, but "freezing" is an exothermic process.
Endothermic (need to add heat) - melting, subliming, boiling, vaporization, fusion
Exothermic (liberates heat, needs to cool) - condensation, deposition, freezing, solidification

Math Alert!!!

There's a little mathematical manipulation that is worth pointing out on page 543-544. This specific case is with the Clausius-Clapeyron Equation, but it's some "tricks" we'll use a few times so take a look at it here.
First, review your logarithm math. Common log (base 10), natural log (base e), all the other math bits... logarithms are useful tools because they are one of the few common tools that allow us to get a variable out of the exponent. Useful.
The second thing here is a clue that can help you with some of these manipulations and derivations. 90% of the time, if a chemist starts rearranging a mathematical expression, they're trying to make it into a line. Lines are easy to work with. Any time you you're trying to make sense of a big ugly expression that only has 2 variables and no exponents, see if you wrestle it into the equation of a line.
The third thing we see here is taking a linear expression under two different sets of conditions and making a comparative expression. In the case of the Clausius-Clapeyron, we can use 2 temperature and pressue conditions to find {delta}H, and we don't have to know the value of A! That's a great advantage, depending upon what you need to do.
These are all "tricks" that you'll see multiple times in Gen Chem 2 (and beyond), so looking at them carefully now can be a huge help when you see the same thing happening again later.

Enthalpies of Phase Change:

You can always write a phase change as a chemical equation:
X(s) --> X(l)
Just like every other chemical equation, the equation above has an associated {delta}H. You can do the same "Heat Stoichiometry" process with these phase changes as you would with any other chemical equation. Don't make this a whole new problem, use what you already know!

Heating & Cooling Curves:

Look at Figure 10.29 in your textbook. This heating curve contains an incredible amount of information, and because of that, people often are overwhelmed when they are learning to use them. The thing to remember about heating & cooling curves is that they are just a collection of individual steps that you already know how to do. In Figure 10.29, the horizontal (blue) segments are phase changes, and the heat/energy associated with those phase changes can be found using an enthalpy of phase change. The sloped (red) sections are heat capacity/specific heat processes and can be treated by using specific heat. Take it one small step at a time and these heating & cooling curves are a little easier to use.

Phase Diagrams:

Heating & cooling curves describe a process that is happening at a constant pressure. Change the pressure, and you have to change the curve. That's OK for most "real" experiments because we are often doing an experiment in a lab over a short time period, so the pressure doesn't change appreciably during the experiment.
If we want a more complete picture, we can look at a phase diagram. A phase diagram is the result of taking a bunch of heating & cooling curves at different pressues and stacking them up. Just like with heating & cooling curves, take them one step at a time and they're not too bad.

Solids:

A lot of the information on solids is pretty concise and described well in your textbook. For this class, the info in Table 10.4 is a good summary of the important points.
If you like to think visually and find geometry to be exciting, the sections on crystals and crystallography are great, but they are a bit beyond what we need for this specific course.

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