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. 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 1: Introduction to the Study of Chemistry

Solving Problems with the Factor-Label Method

In order to determine accurate values, careful use of the numbers we determine (including significant figures), the units they have and careful calculations are necessary. In general chemistry, the method primarily taught to allow for accurate calculations is called the factor-label method. This method requires us to think of things in terms of equivalences. The easiest way to explain this is to show a few examples:

Volume

A volume is simply a length cubed (aka length x length x length). 1 cm³ = (1x10^-2 m)³ = 1x10^-6 m³. It is also important to know that 1 cm³ = 1 mL. Let us now use factor labeling to determine the number of cm³ in a 20 fl oz bottle of coke. For this example we will need the following equivalence 1 fl oz = 29.57 mL.

Example 1: Factor Labeling and Volume Conversions

What we did here was list in order the equivalences we know, we start with one 20 fl oz bottle of coke which we place over 1. We know that for every 29.57 mL we have one fl oz, so we multiply by that equivalence and cancel out the units that we can (fl oz). We also know that 1 mL = 1 cm³ so we set up that final equivalence and cancel out the remaining units that we can (mL). We can now multiply the terms together and write down the number along with the only unit that remains (cm³). If you follow this method for calculations in science, you will never not know what units you should have at the end of your calculation (a problem a large number of undergraduates who take general chemistry find themselves with).

Temperature

Some common temperature conversions are between Celsius (°C), Fahrenheit (°F) and Kelvin (K). Note that Kelvin has no degrees symbol, this is because it is an absolute scale (0 K is absolute zero, it is the point where there is no longer any molecular movement), that little degrees symbol means it is relative to something. For °C it is relative to the boiling and freezing points of water (100 °C is where water boils and 0 °C is where water freezes). For °F it is based on some (bullshit) measurements by Daniel Gabriel Fahrenheit… (0 °F is supposedly the temperature where a solution of brine freezes, brine is a 50/50 salt/water mixture. And 96 °F was his best guess as to the temperature of the human body... which isn’t even the correct temperature, and from this the wondrous Fahrenheit scale was born…).

We can convert between °C and °F fairly simply. Any temperature in °C is equal to a temperature in °F – 32 °F then we multiply that number by (5/9). The inverse is also true, any temperature in °F is equal to a temperature in °C * (9/5) + 32°F. We can also do conversions between K and °C, these are easily convertible by taking the temperature in °C and adding 273. So let is do a similar factor labeling conversion for temperature. How many K is 74 °F?

Example 2: Factor Labeling Temperature Conversion

Here again we used the conversion factors we know, we must go to Celsius first and then convert to kelvin. So we subtract 32 from the temperature in Fahrenheit, then multiply by (5/9). This gives us a temperature in Celsius, which we can add 273 to and arrive at the temperature in Kelvin.

End of Introductory Materials Problem Set

Were you to be enrolled in a general chemistry course there would usually be some problem sets to probe your understanding of the material. I will provide you this same opportunity throughout this series as well. If you choose to do it, the following problem set involves things from Part 1 and any material prior to this text in Part 2. Additionally at the bottom of the document is a set of reference data that you may or may not need to use to solve the problems in the set. The problem sets will become significantly more “chemistry” based as the series progresses, however right now getting used to the basics is very important!

Problem Set 1

Problem Set 1 Answer Key

Introduction to Atomic Theory

Atomic theory, like many other parts of science and mathematics has its roots in ancient Greece. Two Greek philosophers Leucippus and Democritus thought up the concept for the atom in the 5th century BCE. They postulated that all matter was composed of tiny particles which could not be divided any further that they termed atomos (which means uncuttable) 1. In addition to atomos they proposed other aspects of atoms which we now know to be true including the space between atoms containing no matter. Interestingly enough, other prominent Greek philosophers (cough Aristotle) thought the idea was ridiculous.

Atomic theory took a step forward in 1805 due to the work of John Dalton and his theory about the nature of matter. There are a few tenants to this theory 2:

1.) All matter is made of atoms, which are indivisible and cannot be destroyed
2.) All atoms of any given element are identical in their properties and their masses
3.) Compounds are formed by a combination of two or more different kinds of atoms
4.) Chemical reactions are rearrangements of atoms

Atomic Structure

Thanks to Dalton’s work we have our starting official definition of an Atom. An atom is the smallest unit of an element which can partake in a chemical reaction. What Dalton didn’t realize was that atoms are not “indivisible” and are actually themselves composed of particles, that we term subatomic particles. These particles are protons, neutrons and electrons.

Electrons

The first subatomic particle to be discovered was the electron, and it was figured out through the study of radiation. Radiation is defined as the emission of energy as waves or subatomic particles. Radiation was originally studied through use of an Crookes tube (more commonly known as a cathode ray tube).

Figure 1: Crookes tube

Source

How this tube works: First air is removed from the tube (so the particles don’t hit any gas molecules as they travel). Next there are two metal plates which are connected to a power supply. When the power supply is turned on, the negatively charged plate (the cathode) begins to emit an invisible “ray” which is drawn to the positively charged plate (the anode). When the ray strikes the surface of the Crooks tube it glows. The glowing comes from a fluorescent chemical (Zinc Sulfide, ZnS) that the inside of the tube has been painted with. Researchers identified that the particles traveling in the tube were electrons through putting magnets on the outside of the tube. These magnets changed the direction that the ray was traveling. Because the ray was traveling in the direction of the positively charged piece of metal, the particles were known to be negatively charged, which we now call electrons.

Protons

Now researchers knew that atoms contained electrons, yet somehow they are electrically neutral. In order to maintain that neutrality an atom must have an equal number of positive charges for every negative charge. Originally it was thought that this positive charge was spread throughout the entire sphere of the atom, however a cleaver researcher named Ernest Rutheford came up with an experiment using thin gold foil and a source of radiation (α-particles which are actually the nucleus of a Helium atom). In Rutheford’s experiments the α-particles were shot at the gold foil which was sitting inside a ring of a screen which could detect the particles hitting it. Were all of the charges to be evenly distributed in the atoms then the particles should be able to pass straight through the foil and strike a spot directly in line with the radiation source. However that is not what happened! Rutherford observed that a small amount of the α-particles were deflected, and the only way this could happen was if there was a very dense source of charge inside the atom for the α-particles to deflect off of. Rutheford discovered the nucleus, the location in the atom where the positive charge came from to balance the negative charge of the electrons.

Neutrons

Soon researchers figured out that Hydrogen the smallest atom contained only one proton, and Helium contained two protons. However, Rutherfords model had a problem, he knew that there was a dense nucleus with a positive charge but the masses of atoms didn’t match up. By his model, helium should have weighed two times as much as hydrogen, however it weighs 4 times as much! Thus Rutheford hypothesized that there must be another subatomic particle in the nucleus that accounts for the difference in mass. It wasn’t until 1932 when a researcher named James Chadwick performed a similar experiment to Rutherford’s gold bombardment experiment, that the neutron was discovered. Chadwick blasted α-particles at a thin sheet of beryllium, and when he did this a very high energy radiation was emitted, only this radiation was electrically neutral, hence these particles were named neutrons.

We have finally arrived at the more modern depiction of an atom. Here a dense nucleus containing positively charged protons and neutral neutrons, is surrounded by a large cloud of electrons. This model is called the electron cloud model.

Figure 2: Electron Cloud Model of a Helium Atom

Source

What we see in the image above helps put into perspective the size of the nucleus of the Helium in comparison to the over-all electron cloud, the nucleus is over 100,000 times smaller! Also the figure depicts the fact that electrons are more likely to be in areas closer to the nucleus then they are further out (that’s why the cloud is darker around the nucleus), this is because of the attraction of the negatively charged electron to the positively charged protons in the nucleus. The nucleus itself is quite interesting as well. In order to keep from being ripped apart by the intense electromagnetic forces pulling on it at all times, even stronger nuclear bonding forces are required. These force is known as the strong force and is one of the fundamental forces in our universe (along with gravity, electromagnetic forces, and weak forces).

Future Posts

Subsequent posts will cover: What are Atoms and Molecules?, Chemical Reactions, The Periodic Table, Electronic Configuration of Atoms, Chemical Bonding, and Molecular Geometry, and more.

Reference Figure: Periodic Table

Source
Reference

Constants and Conversions List

Source for Additional Constants

Works Cited

If you like my work, please consider giving me a follow: steemit.com/@justtryme90. I am a PhD holding biochemist with a love for science. My future science blog posts will cover a range of topics in the biology/chemistry fields.

Thank you for your support of my work!

Chemistry Lesson: Part 2 (Factor Labeling, Conversions, and Intro to Atomic Theory)