This section is a continuation of the previous, Basic Physics, topic and assumes that you have already read the previous section. The following discussion is not meant to be comprehensive or overly technical but strives to explain some basic chemistry topics relevant to understanding water chemistry. There is some generalization and simplification.
Electrons in an atom occupy what a physicist would call energy levels and a chemist would describe as electron shells. These energy levels/electron shells are only found at certain very definite distances from the nucleus of an atom, distances which are unique to a particular element. The electrons cannot be at any other distances which is because the electrons have both particle and wave-like properties (all matter does but it is only at the very small scale of an electron that the wave-like properties of matter normally become important – quantum physics). The particular energy level/electron shell that an electron occupies corresponds to a certain amount of energy; the farther away from the nucleus that the energy level is, the greater the amount of energy that the electron in that level has.
In Physics Basics, you learned that when an electron changes energy levels, it either gains energy (it moves to an energy level farther from the nucleus) or emits (electromagnetic) energy (it moves to a level closer to the nucleus). In this section, you will learn that a given energy level, or as the chemists would say, an electron shell/subshell can only hold a certain number of electrons. This is important because an atom can be thought of as trying to do two different things at the same time: balancing electric charge and trying to arrange it so that all of its electron shells and subshells are either empty or full of electrons (Yes, this is anthropomorphising an atom which can’t ‘want’ anything but it makes for a better, more easily understood story).
The lowest energy level/electron shell (closest to the nucleus) can only hold two electrons. If an atom at rest has more than two electrons, any additional electrons will occupy higher (more distant) energy levels. An atom is said to be ‘at rest’ if all of its electrons are at the lowest possible energy level/electron shell. If one or more of the electrons are not at the lowest possible levels, the atom is said to be excited.
The second lowest energy level/shell can hold eight electrons in two subshells and so on. Chemistry gets into considerable detail as to the sequence and capacity of electron shells and subshells but for a cursory understanding, it is only necessary to know whether the lower shells/subshells are filled to capacity with electrons or, if not, what the shortage or overage of electrons is in the highest level/subshell with electrons. The electrons in that last, unfilled subshell are sometimes known as valence electrons.
For those of you who want a more detailed explanation, it goes like this: Shown below is an ideal, simplistic image of an atom with a total of 60 electrons which means, since this is an atom, that it also has 60 protons (element number 60) which makes it an atom of Neodymium. There is a problem with this image but it illustrates some important points, one of which is that not everything on the internet is correct.
In this image you can see four shells (different colors) and that within a shell, there are one or more subshells. The first, lowest energy (closest to the nucleus) shell has only one subshell, the ‘s’ subshell. As you move outward, each successive shell adds one additional subshell: a ‘p’ subshell, a ‘d’ subshell, and an ‘f’ subshell. There is a number in each subshell which is the number of electrons that that type of subshell can hold:
The Nomenclature column shows what looks like a letter with an exponent but the number is not an exponent; it simply indicates how many electrons are in that subshell. In the table above, all of the subshells are full. You can indicate which shell that subshell is in by putting the number of the shell in front of the letter for the subshell.
Earlier it was mentioned that an atom ‘wants’ to do two things: maintain a balanced electrical charge and have all of its subshells either completely full or completely empty. Oxygen can accomplish the latter by making bonds (of which more later) with another atom(s) which have a subshell with only one or two electrons. Water is a great example. Two hydrogen atoms can donate (they are electron donors) their lone electrons to an oxygen atom (an electron acceptor). However, there is now a charge imbalance so the two hydrogens must remain closely associated (bonded with) the oxygen atom to balance the electrical charges; three atoms become a compound, a water molecule.
If, however, an atom on its own already has completely filled or empty subshells, it has no need to form bonds with any other atoms and so, doesn’t. Such elements are known as the noble (they don’t bond with other atoms) gases and make up the 18th column, farthest to the right, of the Periodic Table of the Elements.
The consequence of all this is that the last subshell in the actual configuration has only four electrons in it, not fourteen (it is not a misprint); neodymium has a partially filled subshell which is why it isn’t a noble gas. That part of the real configuration in parentheses is the configuration of Xenon (element 54) which is a noble gas; all of its subshells are either completely full or empty. The electron shell image of neodymium is now seen to be an oversimplification. The real world is more complicated (and interesting), leading to the transition elements (not discussed) that make up so much of the periodic table and multiple oxidation numbers (discussed later).
Electrons can and do move between energy levels but it requires a change in energy. In order to move to a higher energy level, the electron must absorb enough energy to do that. If there should be an opening available at a lower energy level, the electron will drop down to it but will emit energy characteristic of the difference in energy levels. That absorbed or emitted energy is electromagnetic radiation which includes everything from radio waves, through microwaves, infrared, visible light, ultraviolet, x-rays, to gamma radiation.
All of the radiation above, carried by photons (do not confuse them with protons), is part of the electromagnetic spectrum; the difference is in the wavelength and energy of the radiation (by the way, all of those radiations move at the speed of light). Radio waves at the low end of the spectrum have very long wavelengths and little energy. At the high end, gamma radiation has an extremely short wavelength and carries enormous energy. If it helps to visualize it, you could think of wavy electromagnetic radiation as being like a spring. Compress (shorten the wavelength) of a spring and it has more energy (don’t take this analogy too far).
An electron can be bounced up to a higher energy level by absorbing select wavelengths of electromagnetic energy which match the difference in energy levels. When that happens, the atom is said to be excited. If the absorbed energy is great enough, the electron might be bounced completely free of the atom at which point the atom is no longer an atom (the charges no longer balance) and the atom becomes what is known as an ion. Any radiation less energetic than visible light cannot ionize an atom (note that that includes microwaves). Visual light is borderline but not normally described as ionizing radiation. Plants, however, can set up special conditions that will allow visible light to ionize atoms, which is what photosynthesis is all about; solar panels do something similar (the photoelectric effect). Ultraviolet light, x-rays, and gamma radiation do have enough energy to ionize atoms and are, therefore, described as being ionizing radiation. Ionizing an atom in a chemical compound can break bonds. Break the wrong bond in an organic compound in a cell and it can lead to cancer. Incidentally, heat and chemical reactions also break bonds.
It works the other way too. When an electron drops down to a lower energy level, it will emit photons of electromagnetic radiation. The bigger the drop (the greater the difference in energy levels), the shorter the wavelength and the greater the energy of the emitted photon. The electron does not have to drop from a high energy level to the lowest possible energy level in one big drop but can drop in a series of steps down sub-shells, each step resulting in the emission of photons whose energy and wavelength are directly related to the energy difference between the levels. The sum of the energy of each individual step should equal the energy of one big drop. Should an electron be able to drop from a very high level to a very low level in one big drop, it might even emit gamma radiation (which is what happens in beta decay, described in the previous topic, Basic Physics). UV light (ultraviolet or black light) can be used to excite the atoms of some fluorescent materials. When the electrons drop back down in several steps, they may emit several photons of visible light (they fluoresce). Some laboratory water-testing equipment uses these principles to identify what and how much of something is dissolved in water.
It was earlier mentioned that atoms, by definition, have balanced electric charges; the number of protons = the number of electrons. There is, however, another important balance. Since this discussion is not really a strictly technical one, I shall take the liberty of anthropomorphizing the atom for the sake of a better story (atoms don’t ‘want’ anything). Remember those energy levels/electron shells? An atom ‘wants’ (is more stable if) its energy levels are either completely filled to capacity with electrons or have no electrons at all. So, an atom is trying to do two things at once: maintain a balanced electric charge and have completely filled or empty electron shells.
The first, lowest energy level can hold two electrons. Helium, element #2, has two electrons and so is completely satisfied both with respect to charge balance and electron shells. Neon, element #10, has ten electrons, two for the first shell and eight for the second shell; it, too is completely satisfied in both respects. You might recognize where this is going by looking at the right-most column of the periodic table of the elements, the column of the noble gasses. These elements are ‘noble’ because they don’t form compounds (yes, there are a few rare exceptions). These ‘noble’ elements are completely satisfied both with respect to balanced charge and filled/empty electron shells – they don’t need to interact with (form compounds) with anything else to achieve a balance, but all the other elements do.
Hydrogen (almost all of it is protium) can achieve a shell balance by either getting rid of its one electron or acquiring a second electron, thereby making its first shell either completely empty or completely full; it usually gets rid of its electron instead of accepting a second electron. But now the charge is out of balance; what was once an atom or molecule becomes an ion. If the ion has an excess positive charge, it is a cation. If it has a negative charge, it is an anion. The hydrogen ion, with one excess positive charge (the single proton in its nucleus), is a cation although it is usually just referred to as a hydrogen ion with the understanding that it does have a positive charge.
Oxygen, with a total of eight electrons, is two electrons short of filling its second sub-shell and so tries to acquire two more electrons, throwing its charge out of balance. Since hydrogen usually gives up its one electron; it is an electron donor. Oxygen tries to take two extra electrons; it is an electron acceptor. Triple up an oxygen with two hydrogens and all three can attain a shell balance; each hydrogen donates one electron to the oxygen. But in doing that, the oxygen now has two excess negative charges (the extra electrons) and the hydrogens one positive charge each (a lack of an electron). Since opposite charges attract, the hydrogens become bonded to the oxygen (the charges are balanced) and the three atoms form a compound (water, of course). As a compound, all charges are balanced and all electron shells are either full or empty; the three atoms, now a compound, become a water molecule.
Once again, atoms have, by definition, a balanced charge. So, too, do molecules but the charges do not have to be equally distributed over the entire molecule, they just have to be balanced as a whole. Molecules that have areas with a slight excess (usually referred to as ‘partial’ charges) of a negative charge and other areas with a slight excess of a positive charge are said to be polar. Water is an excellent example of a polar molecule. The angle between the two hydrogen bonds to the oxygen is about 106°; the front end near the hydrogens has a slight positive charge as opposed to the back end of the oxygen which has a slight negative charge. This means that the front end of one water molecule is weakly attracted to the back end of another water molecule, forming what are known as hydrogen bonds which are much weaker than ionic or covalent bonds. Nevertheless, these hydrogen bonds are why water is a liquid at temperatures and pressures other similar molecules are gasses.
Weak links form between the oppositely charged polar ends of a polar covalent molecule.
This explains why the boiling and freezing points of water are so high for the small size of the molecule.
The oxidation state of carbon 1 in glucose is + 1, carbon 6 is – 1, and all of the other carbons are at 0. To see how the oxidation state of the carbons is determined, consider C1 (carbon 1). It has four bonds, one with C2, another with –OH, with –H, and one with –O–. All of the carbon bonds are covalent bonds which means that, while the carbon will share electrons, it will not completely let go of them or completely accept them; its partners remain tightly linked to the carbon atom.
The bond between C1 and C2 is perfectly equal in that an electron is equally shared between the carbon atoms so this bond contributes 0 to the oxidation number of those two carbon atoms. In the –OH bond, the oxygen ‘wants’ the shared electron more than C1 does so, although C1 will not ‘let go’ of the electron, that electron will associate more closer with the oxygen atom. This contributes a minus charge to the oxygen and a plus charge to C1. The link with –H is the opposite; the hydrogen ‘pushes’ an electron closer to the carbon atom, giving C1 a negative charge and the hydrogen a positive charge. The link with –O– adds a positive charge to C1 and a negative charge to that oxygen.
The total charge or oxidation number of C1 then, is the sum of the charges from the individual bonds or 0 + (–1) + (+1) + (+1) = +1. All of the oxygens are at –2 and the hydrogens have an oxidation number of +1. The sum of the oxidation numbers of all of the elements in a compound must equal 0 which, as you can see in the glucose molecule, they do even though individual carbons have different oxidation numbers.
A summary of general rules, oxidation states, and more examples follows: