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Thursday, February 18, 2016

Chemistry for Preppers: Molecules

First, I need to give a word of warning: Chemistry follows rules, but there are almost always exceptions to those rules. I will try to point out exceptions as they occur to me, but my memory is not perfect and I may miss some. If I make a mistake, please let me know so I can correct it; my ego is not so large that I am unwilling to admit that I can make mistakes.

I studied Chemistry for three years in college, but that was almost 30 years ago; my interest and desire to learn have kept me dabbling in chemistry for many years, and the various jobs I've had over those many years have taught me more than college did.

I am also struggling with subscript and superscript in the blog format, without resorting to writing the whole mess in HTML. If I ever figure out how to properly represent molecular formulas, I will. Until then you'll have to put up with the mess along with me.

Recap
Last week I covered atoms, which are the smallest pieces that elements can be divided into and still retain their physical and chemical properties. Start adding or subtracting protons and you're creating a different element; add or subtract neutrons and you're creating isotopes; electrons bounce around between atoms creating ions. All three will act differently, physically and chemically, than the atom you started with. Atoms are the base of chemistry, but molecules are what we work with.

Electron Shells
When atoms get together and share electrons, they create molecules. I mentioned electron shells briefly last week, and how there are a series of shells around the atomic nucleus with each shell capable of holding a set number of electrons. The shells tend to (almost always, but there are exceptions) fill completely before electrons start orbiting in the next higher shell. Thus, the outer shell is the one with the “available” electrons or “available” spaces for electrons, since it is the one that is not full (except for the Noble Gasses, with their full outer shells which don't combine with anything under normal circumstances). These outer electrons are called valence electrons, and are the ones shared in the “bonds” that tie atoms together into molecules. As you go down the Periodic Table, the outer electron shells hold more electrons and it becomes possible for atoms to have several natural valence states, or be polyvalent (more on that when I get to naming and nomenclature).

Ions
If electrons from one atom are stripped from the outer shell and take up orbit around another atom it is called an ionic bond, because both sides have had their electron/proton balance upset and have become electrically charged. Electrons carry a negative charge, so the atom that gains an electron will become more negatively charged and is known as an anion, while the donor atom will become more positively charged and is known as a cation. My simple way of remembering which is which is, “I like cats, so cats are positive and therefore cations are positively charged”. Those of you who don't like cats will have to use the “an-” part of anion and remember that it means “minus” or “negative”. Anhydrous = without (minus) water, anarchy = without rule, anaerobic = without air, etc.

Bonding
Since opposites attract, the positive and negative ions will stick together like magnets. 

Ionic bonds are fairly weak and easily broken; dissolving an ionic molecule in a solvent will create a solution of ions. Table salt is a good example of an ionic molecule: NaCl is the chemical formula for table salt, sodium chloride, and if you add NaCl to water, the molecule splits into Na (+) and Cl(-) ions. The Cl atoms have stripped an electron from the Na atoms, ionizing both of them.

If electrons are shared equally in pairs between the components of a molecule, it is called a covalent bond. Covalent bonds are strong and stable. Water is a good example of a covalent bond: H2O is the formula for water, two hydrogen atoms bound to a single oxygen atom, and these bonds require an external source of energy to break. Quite a few of the elements exist in a molecular form due to covalent bonds -- Oxygen atoms tend to buddy up and you'll find them as O2 in nature, and Sulfur (or sulphur) tends to gang up in groups of 8 such that the common yellow powder form of sulfur is known as S8.

There is a third type of bond, known as a hydrogen bond, that is peculiar to the unique state of hydrogen only having one electron to share but needing two to fill its first and only electron shell. This one is more applicable to the interactions between molecules than those between the atoms that make up molecules. Hydrogen bonds are very weak.

Depiction
If there are more than one bond in a molecule, the bonds will also have an angle between the bonds. If you look up H2O, you'll generally find a molecule that looks like a Mickey Mouse head, with the oxygen atom as Mickey's face and the hydrogen atoms being his ears. This shape is set by the angle of the O-H bonds (104.5 degrees) to each other, which is why you'll often see complex molecules pictured as bizarre tinker-toy constructions. Bond angles are determined by the number of valence electrons available and how many are actually being shared.

If you're looking through a chemistry text and see a molecule represented like H-O-H then you're seeing a simple Lewis Structure. Each of the bars (there may be more than one) between the atoms represents two electrons that are being shared in a bond. Unshared or “lone pair” electrons may be shown as dots around the atomic symbols. This system shows the electrons being shared, but not the bond angle. This is an older system that is being replaced with 3D graphical representations in order to relay more details about a molecule's structure.


Next week I will be introducing the various categories of chemicals, with reduction/oxidation chemistry, stoichiometry, and reaction potentials in the near future. My CRC Handbook of Chemistry and Physics must be packed away, so I bought an older version on eBay for less than $10, shipped. I'm going to need that for some of my references, and it's handy to have around for the conversion factors in the back (about 40 pages of fine print, how to convert any common unit of measure into any other) and some of the history of chemistry.

The Fine Print


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