Reading Notes for Chapter 5


These are Dr. Bodwin's reading notes for Chapter 5 of "Chemistry 2e" from OpenStax. I am using a local .pdf copy that was downloaded in May 2020.

Chapter Summary:

Chemistry doesn't just happen for random reasons, there must be a reason for a reaction to occur. The most fundamental reason that chemiucal reactions occur is because the universe is lazy and wants to reach the lowest energy possible. {OK, that's an oversimplification...} In order for a chemical reaction to occur, energy must change by moving from one place to another. One of the most common forms of energy that we observe in a chemical reaction is heat energy, so we will start our energy discussion by looking at how heat moves around in a chemical reaction, thermochemistry. Before we get into too much detail, let me emphasize: in many ways, "heat" is just another reactant or product in a chemical reaction. All the rules and tools we have used so far to explore chemical reaction still apply to heat. There are a couple new details, but don't make thermochemistry a whole new subdivision in your brain... it's really just a specific type of stoichiometry problem. Use stoichiometry tools to make thermochemistry work for you.

Energy:

Potential energy is stored energy. It is the energy of position. A rock on the top of a hill has potential energy because of its position. The rock has energy because it required energy to get it to the top of the hill.
Kinetic energy is energy of motion. A rock rolling down a hill has kinetic energy because it is moving (it also has potential energy until it reaches the bottom of the hill...). The rock has energy because it required energy to stop it.
Chemistry isn't really about rocks on hills, so how does this relate to chemistry?
In chemistry, potential energy is stored in bonds (or other interactions), and kinetic energy is still the energy of motion but now it's atomic-level motion.
OK, let's take a little break here to update something that we "know" from a previous chapter. One of the fun things about science is that we know things, but as we learn more, we figure out ways to "know" them better. Are you ready? "Matter cannot be created or destroyed" is wrong. In the context of a normal chemical reaction, it's right enough that we can use it very effectively to balance chemical equations, but as an interesting dude named Einstein told us, energy and matter are related via that handy little "E=mc2" equation. So matter can be "destroyed" if it is converted to energy... which kind of makes matter a type of energy that's stored. Matter is just a form of potential energy. Woah. So the more correct idea is that "energy cannot be created or destroyed" or, stated another way, "the energy of the universe is constant". So we are allowed to interconvert between potential and kinetic energy, but the total energy must remain in balance.
Whenever we're studying energy (and many other things...), it's useful to have a model to work with. A model doesn't have to be complex... in fact, it's usually best to use the simplest possible model to make it easier to understand what is being studied. If you want to study gravity, drop a hammer, don't drop a squirrel. When we are studying thermochemistry, it's usually useful to use a model in which we define some portion of the universe as the thing we're interested in studying (the system) and define the rest of the universe as the other stuff (the surroundings).
If heat is liberated by the system, it must be absorbed by the surroundings. (Remember, we can move energy around and change its form, but it can't just disappear.) When heat flows out of the system, we describe the process as exothermic.
If heat is absorbed by the system, it must be coming from the surroundings. When heat flows in to the system, we describe the process as endothermic.
One thing that can sometimes trip people up is remembering that in most situations you are the surroundings. If a reaction feels warm to you, the system (the reaction) is liberating heat to the surroundings (you), so the reaction is exothermic.

Energy Units:

The SI unit of "heat" is the joule, J. For many chemical processes, the amount of heat moving around is quite large, so we often use kilojoules, kJ.
Calories pop up all the time, so it doesn't hurt to be used to them... but a "calorie" and a "Calorie" are different things... and one of them is listed on the nutritional label of your favorite beverage... It's actually a good example of the importance of capital vs lower case letters. A "Calorie" (sometimes called a "dietary calorie" or "large calorie" in an attempt to be clear) is equal to 1000 "calories".
Originally, calories (like many SI units) were defined by the behavior of water. It's a fascinating adventure to look at all of the water-based definitions. To standardize, we now define 1 calorie as 4.184 joules. So 1 Calorie is 4184 joules or 4.184 kJ.

Heat Capacity and Specific Heat:

Heat capacity and specific heat are both the amount of energy required to change the temperature of a substance. The difference is what is meant by "a substance".
Heat capacity is my absolute favorite example of a "pay attention to the units" quantity. Because the concept of heat capacity is used by so many different fields, there are dozens of different unit combinations. It's not just joules vs calories vs Calories, but Celsius vs Fahrenheit vs Kelvin vs Rankine temperature scales, and grams vs moles vs pounds... it's an absolute mess! So rather than memorizing a formula that will probably be wrong, pay attention to the units you are given and use them appropriately.
Formally, "heat capacity" is the heat needed to change the temperature of a specific object. My coffee cup has a specific heat. So does yours. They're different. And they're different for every coffee cup on the planet. For some measurements, this is fine, but it can be a hassle. "Heat capacity" is an extensive property because it depends upon the amount of material that is present.
"Specific Heat Capacity" is a long phrase, so it's shortened to just "specific heat". This is the heat needed to change the temperature of a given amount of a substance. One gram of a specific formulation of ceramic requires the same amount of heat to change its temperature regardless of the shape or size of the coffee cup.
Remember that when we are talking about heat capacity or specific heat, we are talking about a change in temperature. Since 1 degree Celsius and 1 Kelvin are exactly the same size, a change of 5 degrees Celsius is also a change of 5 Kelvins. This is the thing than trips up students most often when talking about heat capacity and specific heat.
There's a nice little table of specific heats in your textbook (Table 5.1)... think about some of the items you use every day and how they respond to heat transfer.

Calorimetry:

There's a lot of specific detail on calorimeters in your textbook. Calorimetry is a lovely application of thermochemistry, but the level of detail is probably a little much for our course.

Enthalpy:

Enthalpy is "heat of reaction".
The concept of state functions is an important one. Essentially, for state functions, it doesn't matter how you get from point A to point B, all that matters is where you started and where you finish. From Moorhead, MN, to Seattle, WA, the change in elevation is negative; Seattle, WA, is at a lower elevation than Moorhead, MN. That's a state function. If you walked from Moorhead, MN, to Seattle, WA, your effort would not be all "downhill"! The effort you put into walking from Moorhead, MN, to Seattle, WA, is not a state function.

Reaction Coordinate Diagrams:

Reaction Coodinate Diagrams are visualization tools that help analyse thermodynamic change. Since this is a graphical representation and that's a little harder to code in to a webpage, here's a .pdf of a powerpoint: Reaction Coordinate Diagrams

Hess' Law:

Hess' Law is really just a direct application of the concept of a state function. It's also something that can be seen in the last slide of the Reaction Coordinate Diagram info posted above. Essentially, it doesn't matter if we get from A to D via a single reaction or by a series of steps, the enthalpy change for the overall process is the same.

Enthalpies of Reaction from Enthalpies of Formation:

We can look up standard enthalpies of formation for every reactant and product in a chemical reaction. The tabulated values are enthalpies of formation specifically.
For every reactant in a chemical reaction, the substance is being comsumed... that's the opposite of "formation", so the magnitude of {delta}H is correct, but it is in the wrong direction. Since it's the opposite direction, the sign is wrong, so change the sign
For every product in a chemical reaction, the substance is being formed, to the enthalphy of formation is the correct magnitude and the correct sign.
The equation of page 268 is how this is usually described, but I always try to avoid subtraction and division in mathematical expressions when possible because if you do those operations in the wrong order, you get the wrong answer. If we instead change the sign by thinking about a substance's role in the reaction we are studying, then we can just add up all the results. Addition can happen in any order and we always get the same answer.


Return to ChemBits General Chemistry Index.

All information on this page is produced by Jeffrey Bodwin, Copper Sun Creations, or curated from the attributed source.
Creative Commons License
This work is licensed under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License.