Introduction
This series of posts seeks to present the material covered in the first semester of a college level general chemistry course, in an easily digestible steemit blog post format. The series is intended to be read, and experienced in sequential order starting with Post 1. The material will build upon itself, and potential exercises included (problem sets), will pertain to the post they are contained in, or any previous post. Each post will pick up immediately where the previous in the series left off. Please check out the #chemistry-lesson tag for all posts in this series. I hope you find this series to be informative and beneficial toward your understanding of chemistry and science in general.
Immediate Preceding Post
Part 7: Aqueous Chemical Reactions 2
Legend for This Section
As subscripts do not currently work on steemit, to symbolize when a number should be sub-scripted I will be writing it in the text as follows: /x/.
Gasses
I am sure you all remember from the start of this series where we talked about the three states of matter Solid, Liquid and Gas. To this point we have really only been dealing with molecules and compounds in the solid or liquid phase, we have been leaving out a whole state of matter! Gasses are a state of matter that can expand to take on the total volume of whatever container they are stored in. Gasses have very low density and are extremely compressible which means a large amount of them can be stored in a relatively small container with very high pressure. We live every day with gasses all around us, including one extremely important gaseous mixture that we don’t often think about: Air. Air is composed of the gaseous states: it is 78.09% Nitrogen, 20.8% Oxygen, 0.93% Argon and 0.04% Carbon Dioxide, and a variety of other gasses in trace amounts (Source).
Gasses Exert Pressure
When you place a gas into a container, the molecules of that gas are whizzing around at a certain velocity in there and smacking into the walls of the container, the force they exert by coming in contact with the containers walls is called pressure. To understand where this comes from lets digress into physics for a bit:
(Information Source)
Finally this information, along with knowing the area of the surface of the container the gas is coming in contact with allows for us to determine the pressure it is exerting on the vessel. So you can also see from these equations that the pressure of a gas is related to the velocity that the gas molecule is moving right? If we change the velocity we will also change the pressure, the same goes for the area of the container, if the container gets smaller but the same amount of force is being applied, then pressure would increase.
Gas pressure is typically referred to with units in terms of the pressure exerted by the Earth’s atmosphere called atm. While the pressure from the atmosphere varies based upon your height on the planet (the lower you are the more atmosphere is on top of you and so the more pressure it exerts, if you are on a mountain there is less atmosphere above and around you and the pressure is lower). The standard was set is that 1 atm is equal to the amount of pressure that supports a tube that is 76 centimeters long completely full of mercury.
Some Gas Laws lead up to the Ideal Gas Equation
Most of the laws dealing with gasses all revolve around aspects of the math behind the calculation of pressure we described above! The first law called Boyles Law (as it was studied by a man named Robert Boyle) states that the pressure exerted by a gas is inversely proportional to the volume of the container it is in. This relates to the math we discussed earlier as decreasing the volume of a container makes the area of the walls the gas strikes smaller. Smaller area means larger pressure, assuming the amount of gas and thus the amount of force exerted, remains the same.
Boyles law is written like this:
This law states that for a consistent temperature and amount of gas, changing the volume will change the pressure it exerts.
The next gas law is called Charles’s Law. Charles’s law deals with the relationship between temperature and a gasses volume. The rate that molecules move is proportional to their temperature, at higher temperatures molecules move faster than they do at lower temperatures. If you think about the equations we wrote out above, we know that velocity is related to the acceleration of the atoms, so increasing the temperature on a gas increases the velocity that the molecules move. Thinking this way it is easy to understand that temperature and the properties of a gas are very tightly related.
Charles saw that the volume and pressure of a gas were directly proportional to its temperature so we know that for any change in temperature a relative change in the volume of the gas would be required.
Charles’s law is written like this:
The final law we (Avogadro’s Law) will discuss is from our old friend Avogadro (who if you will recall the number of atoms in a mole was named after). Avogadro noticed that the volume a gas wanted to occupy was directly related to how many gas molecules were there. He also noticed that the pressure generated by a gas was related to the amount of gas molecules present, but that all gasses had a similar effect regardless of their molecular composition. This allows us to write chemical equations describing the reactions of gasses and know what effect that reaction is going to have on the pressure of the system.
Consider the following reaction:
Here two NO/2/ molecules react together and form 1 N/2/O/4/ molecule, this results in a decrease in the number of gas molecules present in the reaction vessel, and would thus decrease the pressure as the reaction proceeds. For every two moles of NO/2/ that were present to start, we end up with only one mole of N/2/O/4. Less molecules to bang against the walls of the reaction vessel, means less force, and that translates to less pressure!
(Information Source 1 and Information Source 2)
Finally we can put all of these laws together into one big Ideal Gas Law:
This law states that Pressure (P) x Volume (V) is equal to number of gas moles (n) x a gas proportionality constant (R) x Temperature (T). Now you may be wondering what is that gas constant R? It is a value which allows for us to relate together all of the other properties together.
Or as we will write R = 0.08206 (L•atm)/(K•mol), the values in here allow us to relate Volume (liters) with Pressure (atmospheres) with Temperature (in kelvin, the absolute temperature scale) and also number of moles (mol).
Let us end this section with a problem that we can use the ideal gas law to solve: Calculate the pressure of 10 g of H/2/ (g) when placed into a 2 L flask at 25 °C.
We can do this using the ideal gas equation, but first we must convert our information into the appropriate forms for being in the equation (we know those forms based on the units in the gas constant!) We need Volume in Liters, which we have. Moles of substance which we must calculate, and temperature in kelvin which we must also calculate. We want to pull out pressure, which will be in atm (as it should be) so let’s begin:
That 2 liter reaction vessel has 60 times more pressure exerted on its walls then the atmosphere exerts, that’s quite a bit!
We will continue in the next blog with more examples of gas reaction calculations, how to calculate pressures for individual gaseous components in a mixture, and some theory behind the kinetics of how the gas molecules move.
Future Posts
Subsequent posts will cover: Electronic Configuration of Atoms, Chemical Bonding, and Molecular Geometry, and more.
Reference Figure: Periodic Table
Other References
Constants and Conversions List
Source for Additional Constants
Some Common Ions
An Open Source Chemistry Text Book
Additional Sourcing Information
General Chemistry: Principles, Patterns, and Applications
General Chemistry: The Essential Concepts
General Chemistry
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