Reading Notes for Chapter 13


These are Dr. Bodwin's reading notes for Chapter 13 of "Chemistry 2e" from OpenStax. I am using a local .pdf copy that was downloaded in May 2020.

Chapter Summary:

Up to this point, we've discussed chemical reactions as if they either "go" or "don't go". In reality, all reactions occur, we just have to determine how far they go. The term equilibrium is used commonly enough that we all probably have some sense of what it means, but how does it apply to chemical reactions? By building upon collision theory and our discussions of kinetics, we can begin to understand what chemical equilibrium is and how to interpret it.

Equilibrium & Equilbrium Constants:

We've actually talked about equilibrium before. Recall the discussion of vapor pressure; when the rate at which liquid molecules escape to the gas phase is equal to the rate at which gas molecules condense into the liquid phase, we observe a constant pressure that we called vapor pressure. This is an equilibrium!
There are many ways to think about equilibrium. Consider the simple example described here.
The most important thing to remember about equilibrium is that it is dynamic; the forward reaction is always happening, and the reverse reaction is always happening.
Because equilibrium is dynamic, it doesn't matter if we start with all reactants, or if we start with all products, or if we start with some mixture of reactants and products, the equilibrium that we react (and the equilibrium constant) will be the same in all of these cases.

Equilibrium Constants - Key Points:

Based upon the balanced chemical equation.
Always positive.
Have no units.
Pure solids and pure liquids do not appear in K
NOTE: For now, "pure solids and pure liquids do not appear in K" works for us, so feel free to lock that into your brain as a "rule". Are you curious why it's true? Formally speaking, equilibrium constants have activities, not concentrations. You get an activity by dividing a concentration in units of "M" by "1 M", so we can just use the molar concentrations directly. That's also why equilibrium constants have no units. What about pure solids and liquids? By definition, their activity is "1", so although in the most formal sense they do appear in the equilibrium constant, they are just multiplying and/or dividing by 1 raised to some power. One raised to any power is... 1. And number multiplied or divided by 1 is... unchanged. So for the level that we are studying equilibrium, we can just ignore pure solids and liquids in constructing our equilibrium constants.
K>1 is product-favored.
K<1 (but always greater than zero...) is reactant-favored.
K values can vary greatly and are usually best expressed using scientific notation.
Think about scale when evlautating "differences". Because K values can be so large, differences become, well, different. For example, an equilibrium constant of 238.8 is not really all that different from an equilibrium constant of 295.3; it is different from an equilbrium constant of 4.286E5!
Most of the equilibria we will discuss are in solution, so we will use Kc. If the reaction is in the gas phase, we could use KP and use partial rpessures instead of concentrations.
We will look at an entire alphabet soup of "different" K values. Although each different subscript describes a different type of reaction, these K values are all conceptually the same, they all mean the same thing, and they all follow the same rules. Don't make them all into completely new things!

Reaction Quotient:

Mathematically, the reaction quotient, Q, has the same form as the equilibrium constant for a reaction.
By comparing Q to K, you can determine if the reaction is at equilibrium, or if the reaction has to shift toward products to reach equilibrium, or if the reaction has to shift toward reactants to reach equilibrium.

Manipulating Equilibrium & Equilibrium Constants:

Changing the reaction we're looking at changes the equilibrium constant, but often in mathematically calculcable ways.
The heart of equilibrium is that all reactions are reversible (to some extent). If a reaction is reversed, reactants become products and products become reactants. This "flips" the numerator and denominator of the equilibrium constant ratio, so the value of K is also inverted.
Coefficients in balanced chemical equations appear as exponents in the equilibrium constant expression, so multiplying all the coefficients in a reaction by a constant raises the K value to that power.

LeChatelier's Principle:

Equilibrium is a balancing act. The concentrations of reactants and products are balanced.
If a system at equilibrium experiences a "stress", the equilibrium will shift to relieve that "stress" if possible.
In this case, "stress" can be:
LeChatelier's Principle is the chemistry version of a concept that shows up in a LOT of places. For the biology/ecology crowd, the concept of "carrying capacity" for predators and prey is just a different version of LeChatelier's Principle. This concept comes up in a number of ways in economics and finance; if you experience the stress of losing your job, you will (hopefully!) shift your spending habits to re-establish an equilibrium where your income and expenses are balanced.

Catalysts and Equilibrium:

All of chemistry can be lumped into either thermodynamics or kinetics.
Thermodynamics compares the energy of reactants to the energy of products. Enthalpy is a thermodynamic concept. Whether a reaction is endothermic or exothermic is not really a factor in the rate of the reaction.
Kinetics explores the activation energy of a reaction, and activation energy determines rate.
Equilibrium is a thermodynamic concept. Catalysts affect the kinetics. Addition of a catalyst to a reaction does not affect the equilibrium of the reaction, it just affects how quickly the reaction reaches equilibrium.

Equilibrium Calculations:

The calculations required to answer equilibrium questions are not especially difficult BUT organizing the information can be a challenge. The easiest way to keep things organized  is by using a table. Your textbook (and many sources) refer to these as "ICE Tables". Personally, I don't care for that term, so I am unlikely to use it, but if it helps you keep things straight, then go for it. I will always use a table of initial concentrations, changes in concentration, and final concentrations to understand equilibrium.
Your textbook has some good examples of an initial/change/equilibrium concentration table starting with Example 13.7. Use tables. Even if you think you "got this", use a table. Even when you've done 284 sample problems, use a table. The 30 seconds it will take you to write out a table and use it to solve your problem is WAY less than the 20 minutes it will take you to find the error without a table if you get the problem wrong.
When we're doing equilibrium problems, we have to be able to use some "super powers" to stop and start time at will. When we mix up a bunch of chemical to start a reaction, we have to think about mixing everything together but not allowing the reaction to start. That's how we get the initial concentrations. Then, we press the magic button and "allow" the reaction to occur. In reality, the reaction starts as soon as the first bits of reactants are mixed together, but to make it easier for us to understand, we get to use this time-stopping super power to help.
Equilibrium problems will require you to exercsie your algebra muscles to rearrange equations. You may also end up having to use tools like the quadratic formula to find some answers. This is why setting up your problem clearly is so important - it helps me (and you!) find where things start going wrong in a problem.

Simplifying Assumptions:

Some equilibrium calculations can get a little "thick", but we can often make some assumptions that will simplify our calculations. Some important notes about assumptions:
  1. Always clearly state the assumption you are making.
  2. Always circle back and check that you assumptions are valid.
Specifically for equilibrium calculations, the assumptions we make are actually related to the concept of significant figures (everyone's favorite concept!!).
Assumptions are nice because they simplify the math, and we can always make an assumption as long as we check that it is valid. If the assumption ends up not being valid, we haven't really lost much time and effort, and we can crank through the "ugly" math that we were hoping to avoid.


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.