Reading Notes for Chapter 7

These are Dr. Bodwin's reading notes for Chapter 7 of "Introduction to Chemistry". I am using a local .pdf copy that was downloaded in August 2020.

Chapter Summary:

Chemical reactions involve transformation of atoms, ions, and molecules, but they also involve the transfer of energy in various forms.

Energy definitions and units:

Potential energy is energy that is "stored". A rock on the top of a hill has potential energy because it can roll down the hill; energy has been "stored" in the rock by bringing it to the top of the hill. Potential energy can also be stored in chemical bonds.
Kinetic energy is energy of motion. When the rock is rolling down the hill, it has kinetic energy. In chemical systems, kinetic energy is the energy of molecular motion, including the vibration of molecules in solids.
Energy does work, energy and work result in heat.
Joules - SI unit of energy, J.
calories - water-based unit of energy. 1 calorie will raise the temperature of 1 gram of water 1 degree Celsius. 1 cal = 4.184 J
Calories - these are dietary calories. 1 Calorie = 1000 calories = 1 kcal.
If you're looking up information about energy and heat, you will find many different units used in different sources. It's important to be very careful and observant to keep those energy units stright. This is especially true for "calories" vs. "Calories"... always read the fine print and think about context when looking at energy units!
Temperature - temperature is a way for use to measure energy and heat. More on temperature below...

Energy Conversions:

Energy conversions are (mostly) just like any other unit conversion... find an appropriate conversion factor, make it into the correct fraction, multiply.
Temperature - converting temperature is not quite as easy as converting joules to calories.
Formulas for temperature conversion - there are online calculators for temperature conversion, and formulas for the conversions, but they are easy to mix up if you just try to blindly apply the formulas. Think about the formulas when you apply them! Here's a couple examples of temperature conversions that don't really use the formulas, but work through the conversions: Temperature Conversions (slides)
Change in Temperature - This one gets people in trouble all the time. When you're working with energy, heat, and temperature, it is important to keep track of whether you are talking about the temperature of a substance, or the change in temperature of the substance. If I grab a 90ºF piece of steel, it's not a big deal, but if I grab a piece of steel that has increased it's temperature by 90ºF from room temperature, I'm probably going to regret it! Again, read carefully!

Heat, Temperature, and Heat Transfer:

Temperature is a measure of the kinetic energy (molecular motions...) of a sample.
When discussing heat, remember that "heat" is a thing, "cold" is not. What we describe as "cold" is just something with less heat than the surrounding, or something that heat is flowing into. When you pour a hot cup of coffee and set it on your kitchen counter, it "gets cold" because it is losing heat, not because it is "gaining cold".
Specific Heat is the amount of heat associated with changing the temperature of a substance

CAUTION: Different sources are sometimes a little "loose" with how they define what your textbook calls "specific heat". If you're web-surfing, you will likely run into a variety of similar sounding terms like "specific heat capacity" or "heat capacity" or others. You are also likely to stumble upon some sources that use grams, others that use moles, others that use pounds as well as some that use ºC, some that use ºF, and others that use Kelvin temperature units. The best thing you can do is ALWAYS make sure you write down the units of the numbers you find, and make sure you use measurements that match the units.

Whenever we start talking about heat transfer (or energy transfer in general), it becomes important to pay special attention to some mathematical formality and to use your intuition and observations to guide your solutions to problems. Here are a couple important points that should help:

Whenever we talk about a "delta" function to describe a change in conditions (such as ΔT), the strict mathematical treatment is "final condition minus initial condition". If you are using a very careful mathematical approach, the order here is critical.
It is often helpful to (in your mind) separate the amount of heat (the magnitude of change), from the direction of heat flow (the sign of the result). For many problems, we can work with magnitudes and add the sign later.
The sign and direction of heat transfer can often be indicated by words in a problem. This is often done to try and be very clear about the process that is being described. "You have 2 metal cubes, A and B. In the experiment, 25.83J of heat is absorbed by Cube A from Cube B." In this example, "absorbed" tells us the direction of heat flow.
At the end of the problem (and at the beginning and the middle as well!), think about the system and make sure your answer makes sense. If I'm saying that metal Cube A is absorbing heat, then its final temperature better be higher than its initial temperature. If Cube A is at 31.65ºC and Cube B is at 197.23ºC, when I put them together and let the temperature equilibrate, the final temeprature better be higher than 31.65ºC and lower than 197.23ºC. You have been observing heat flow for your entire life outside of science classes, make sure you use your experience to inform your answers!

Specific Heat problems
To simplify problems, we often assume that a system is completely isolated - one substance liberates heat, another substance absorbs heat, and we don't include loss of heat to the surroundings. That's a very handy assumption when we are trying to understand heat and heat transfer, but when you're in the real world, remember that the there is always some tranfer of heat to or from the surroundings. There are many things we can do to minimize that transfer, but it's always there...

Phase Changes:

When a solid, liquid, or gas is heated or cooled, that's a heat capacity. But when the material gets to a temperature where the phase changes, it does thorough an energy process where the substance changes phase but the temperature doesn't change. When water boils, its temperature is 100ºC until the whole sample boils. Energy is continually being added to the system, but it is being used to accomplish the phase change, not to change the temperature.
When a substance is heated through a temperature range that includes a phase change, break the process down into the different steps to determine the energy of the whole process. For example, if a sample of water at room temperature is cooled down to -15ºC, there are 3 steps to the process: 1)liquid water cooling down (heat capacity); liquid water freezing (phase change); 3) solid water cooling down (heat capacity). Don't try to make it into a single problem!

Bond Energy:

Chemical potential energy is stored in bonds.
Chemical bonds require energy to break them.
"Stable" molecules are made of bonds that require more energy to break them.

Enthalpy:

Your book doesn't really use the term "enthalpy", but it's an important term to know.
The heat associated with a chemical reaction is called the enthalpy. Enthalpy is the relationship between the bond energies of the reactants and the bond energies of the products in a chemical reaction.
"Enthalpy" is often called "heat of reaction".
If a chemical reaction has reactants with a higher bond energy than the bond energy of the products, heat is released by the reaction. This is an exothermic process.
If a chemical reaction has reactants with a lower bond energy than the bond energy of the products, heat must be added (absorbed from the surroundings) as the reaction occurs. This is an endothermic process.
Look at Figures 7.3 & 7.4. These are very helpful ways to visualize the enthalpy of a reaction.

Applying Enthalpy:

Most of the biochemical systems that have evolved in nature are really just ways to store and retrieve energy. Basic nutrition considers carbohydrates, proteins, and fats... the Calories associated with each of these molecules is related to the relative stability of the chemical bonds that hold them together.
Fats have a higher Calorie density (9Cal/g) than carbohydrates & protein (4Cal/g). Carbs & protein have more diverse functions in nature than fats... fats evolved (almost exclusively) to store energy, so they should be the most efficient at that function.
The real "energy currency" in biology is the ATP-ADP cycle (Figure 7.5).

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