So let's knock out a few basic concepts to jump-start your education. First the ultra-quick version:
Chemistry studies how atoms share or exchange electrons. Of roughly 100 kinds of atoms, a few—twelve, to be exact—have one or two "loose" electrons that are easy to strip off, while another twelve have room for one or two more, and will easily plunder those loose electrons. Some others can either gain or lose three, four, or even five electrons. The rest typically share electrons. Chemistry is learning all the ways this can happen, and which elements behave in which fashion.For more, read on. We begin with Electrons.
ElectronsElectrons are particles that make up the outer "skin" and "flesh" of atoms. What we usually mean when we say "chemistry" is properly "electron chemistry". There is also nucleon chemistry, plus other subdisciplines such as crystal chemistry and organic chemistry. The odd thing is, you first have to know a little about nucleon chemistry to get a framework to learn electron chemistry.
Nucleons and ElementsPerhaps you have heard that there are 92 "natural" elements, or maybe, as I wrote above, that there are "about 100 elements". There are actually 90 elements called "naturally occurring". That is because, although the heaviest natural element is Uranium, #92, the elements numbered 43 and 61 are not found in nature, for reasons we'll soon get into.
Nucleons are the particles that make up the nucleus: Protons and Neutrons. The number of protons in a nucleus determine what element it belongs to. For a nucleus to be stable (and the "what for" about this is a major subject of nucleon chemistry) there need to be neutrons present also. Only one element has no neutrons in its nucleus, Hydrogen. An atom of hydrogen, the simplest and lightest element, has one proton and one electron, and nothing more. Every other kind of nucleus has at least one neutron, and with only one exception, the number of neutrons is at least as large as the number of protons.
The main item of nucleon chemistry that you must know is that the Atomic Number is the number of Protons. The term Atomic Number is used everywhere. It is also extremely useful to understand that radioactivity expresses the tendency for certain combinations of protons and neutrons to break apart in one way or another. A very few kinds of "unstable" nuclei are nearly stable and last for millions or billions of years. Uranium is one of these.
Nuclei of elements #43 (Technetium) and #61 (Promethium) are always unstable, in every variety, no matter how many or how few neutrons are in there with the protons. In this case, "unstable" means having a half-life short enough that every single atom of these elements that may have existed billions of years ago when Earth was formed, has decayed. Half-life is another very useful term, though mainly in nucleon chemistry. For a bunch of any specific, unstable kind of nucleus, the half-life is the time it takes for half of them to decay. Lots of uranium (originally produced when big stars blew up billions of years ago) is still here because its half-life is about 4.7 billion years.
The fundamental tool for understanding electron chemistry is a table in order of Atomic Number, that is arranged according to how electrons pack together in each kind of element: the Periodic Table.
Periodic TableGet ready for it! I am about to explain this monstrosity:
The columns are arranged the way they are because elements in a column have similar chemical behavior. Down the left side, for example, the six elements Li, Na, K, Rb, Cs, and Fr all have similar chemical behavior because the outermost electron is "loose" and easily lost to more acquisitive elements. Hydrogen is special; though it can both lose and gain an electron, it also participates in a third kind of sharing bond we'll describe later.
Each row represents an electron shell, which fills from left to right. The rightmost column, topped by Helium (element #2) contains all the elements with a completely filled shell. This is the group of elements with the easiest chemistry: They don't participate in chemical reactions! But right next to them we find F, Cl, Br, I, At, and the "artificial" element currently called Uus (Un-Un-Septium, a fake Latin term for 117). They all have an outermost shell that is nearly filled, but is ready to grab an electron from another element that has a "loose" one available.
The rows are different lengths because the shells have different capacities. It takes some learning in quantum physics to comprehend what electrons are doing (as much as that may be possible!). Here is the simple explanation:
- Electrons come in pairs.
- The first shell is filled by a single pair, thus Helium has a filled shell. This filled shell is the core of all heavier elements.
- The shells of all elements other than Hydrogen and Helium have sub-shells.
- The sub-shells were discovered by spectroscopy, and are called, for historical reasons, s, p, d, and f.
- Sub-shells increase by odd numbers of electron pairs:
- p has 3, so s+p = 4 pairs or 8 electrons.
- d has 5, so s+p+d = 9 pairs or 18 electrons.
- f has 7, so s+p+d+f = 16 pairs or 32 electrons.
- Shells 2 and 3 have s+p only; 4 and 5 also have d (thus the lower-middle block); and 6 and 7 also have f (shown as the extra stuff below the main table).
- The placement of the rows shows that the d sub-shell fills before the p sub-shell, and the f sub-shell fills before d.
All the elements from 93 to 118 have been produced in nuclear reactors and particle accelerators. With element #118, the seventh shell is filled, so once elements #119 and greater are produced, an eighth shell will begin to fill. This is expected to have a new sub-shell, usually called g. It can contain 9 electron pairs. It is likely that the g sub-shell will begin to be filled with element #121, but we will only know this for sure if element #121, or a heavier one, has a long enough half-life so the electron arrangement can be studied before the whole sample decays away.
BondingWhen one atom takes control of the loose electron given up by a different atom, or when atoms share electrons, we talk of a chemical bond. To discuss this, a version of the Periodic Table with different highlighting will be helpful:
You know that term "alkali"? It refers to substances that neutralize acids. The two columns of elements at the left, in lavender and blue coloring, are called the Alkali Metals (lavender) and the Alkaline Earth Metals (blue). The ones with an odd atomic number have one loose electron, and the even ones have two loose electrons. They participate in compounds that tend to be alkaline; in some cases, the compounds are so caustic they will remove your skin.
Now, at the far right, as I mentioned above, the elements in the last column do not combine chemically with others. A few very extreme experiments have been done to force them into unstable chemical compounds. We call them the Noble Gases. They, and four other elements in which the lettering is dark green colored, are gases at "room temperature", defined for chemists as 25°C or 77°F.
The elements in the next column, with beige coloring, are called Halogens. "Halogen" is from the Latin word for "salt". They like to glom onto loose electrons. Any of these reacted with hydrogen will form a strong acid, but when paired with one of the Alkali Metals or an Alkali Earth Metal, they form stable salts. Two of them are usually gases, one is a liquid (Br, with dark blue letters), and the rest are solids. They are a major part of a group also called Non-Metals.
Hydrogen plus the other elements in orange coloring are the rest of the Non-Metals. In element form, solidified at low temperature in the case of Nitrogen and Oxygen, they are insulating solids that look like soft ceramics. While Oxygen and those below it tend to snatch two loose electrons whenever possible, they also participate in the sharing bond I mentioned earlier.
The elements with brown coloring are called Semi-Metals. In element form, they are semiconductors, and one in particular, Si or Silicon, forms the basis for most electronic circuits. The lime green colored elements are Metals that are either semiconductors by themselves, or form semiconductors when alloyed with Semi-Metals.
All the rest of the elements in the main part of the table are colored light yellow, and are Metals. The top row of them, from Scandium to Zinc, are the Transition Metals. "Transition" refers to their similar chemistry. They all have a filled s sub-shell and an empty p sub-shell, and from 1 to 10 electrons in the d sub-shell, which is "hidden" beneath the filled s sub-shell. However, those two outermost electrons can act as loose electrons to combine with Non-Metals or Oxygen, and frequently one of the d electrons will also do so. Thus, they have more complicated chemistry than those to the extreme right or left. The three pale yellow rows below behave a lot like the Transition Metals, but it is harder and harder to get them to react. In particular, Platinum and Gold (Pt and Au) are very resistant to participating in chemical activity, as are the elements directly beneath them, though those are radioactively unstable and are very short-lived.
The Transition Metals are useful to living things in various amounts, usually quite small amounts. Even Iron (Fe), the most abundant metal in our bodies, is present as 4-6 grams in an adult human, or less than 1/100 of a percent. The heavier metals are called "heavy metals", particularly in medicine, because they are all toxic. Lead (Pb) is the most familiar toxic metal.
Ionic BondsThe shift of one or more electrons between strong "electron donors" such as Li or Ca, and "electron acceptors" such as Se or Cl, produces an Ionic Bond. This kind of bond is strong in the pure solid, but is pulled apart in water to dissolve salts such as LiCl, CaBr2, or MgSe. However, salts with S or Se are poorly soluble compared to salts with Halogen elements "on the right". In water solution, the elements that have lost electrons are + ions, and those that have accepted electrons are - ions.
Covalent BondsElectron sharing in which two atoms form a strong bond to fill their outermost shell produces mainly insoluble compounds held together by Covalent Bonds. The Non-Metals, when in elemental form, usually exist as paired atoms sharing one or more electrons. The simplest example is ordinary Hydrogen:
Here the electrons are shown as dots. The shared electrons satisfy the s sub-shell of both atoms.
Most elements can participate in covalent bonds. The most versatile is Carbon, which has 4 outer electrons, and thus room for 4 more. It prefers to share a covalent bond in 4 directions. This makes it the most versatile in its chemistry, and a huge discipline, Organic Chemistry, is the study of carbon chemistry. Where a chemist who studies inorganic chemistry will become familiar with thousands or tens of thousands of chemical compounds, the number of organic compounds so far known exceeds 50 million.