The same thing can be done with fluorite (calcium difluoride) to recover diatomic fluorine gas, and the rest of the halogens could be recovered likewise were there concentrated natural sources of their halides. But while fluorite and table salt are plentiful, the bromides and iodides are found only in low concentrations in seawater. Some sea plants, like kelp, concentrate iodide, but it's simpler sacrificing some of our hard-won chlorine gas to oxidize the sea's bromides and iodides directly. This works because of the electronegativity decline down Group VII; chlorine has a greater propensity to attract and hold an extra electron than does either bromine or iodine and so will oxidize both bromide and iodide to their neutral molecules.
In all of these redox (oxidation/reduction) reactions, such as the one from the cover of Gillespie show at left, oxidation involves wresting electrons from a compound (changing it into another into the bargain) and simultaneous reduction, depositting those electrons on a willing receptor molecule (usually changing that compound too). The very name "redox" implies that those processes are inextricably linked; after all, you can't buy bottles of free electrons...they're going to attach themselves to something!So that reaction pictured involves the fiery, spontaneous oxidation of lithium metal in water with simultaneous reduction of one of water's protons to free hydrogen. Actually, the reaction produces such great heat that the liberated hydrogen immediately catches fire in air, its flame tinged lilac with electronically excited lithium atoms...an aurora in a teacup, as it were. The initial redox reaction is
2 Li(s) + 2 H2O(l) -----> 2 Li+ (aq) + 2 OH- (aq) + H2(g)followed instantaneously by
2 H2(g) + O2(g) ------> 2 H2O(l) in the air.That Li metal lost an electron to become its cation is obvious. That hydrogen gained it takes modest explanation since we don't see hydrogen cation anywhere there. But we know that the electronegativity difference between hydrogen and oxygen means that in forming water, the oxygen as gained (more than its share of the bonding) electrons and has thus been "reduced," so hydrogen must have lost (its fair share of the bonding) electrons, having been oxidized. The lithium-water reaction reverses the process to the extent that it liberates at least one of water's hydrogens, so that liberated hydrogen must have been reduced in the process.
The halogens are terrific oxidizers; that's their primary function in the chemical industry and society in general. But they are terribly corrosive (as a consequence); so free halogens are used only in chemical plants. Our day-to-day halogen oxidations (like bleaching laundry) take place with tamed versions of these beasts. For example,
Cl2(g) + H2O(l) <=====> HOCl + H3O+ + Cl-produces HOCl (hypochlorous acid) whose neutralized salt, NaOCl (sodium hypochlorite) is familiar to all as the prime ingredient in "bleach." Note that in its production, one of the chlorines has already been reduced to chloride ion. That should jar you because something must've been oxidized at the same time...and it was...the other chlorine! In diatomic chlorine, both atoms were neutral. Then in that reaction, one picked up a negative charge (becoming the chloride ion), and the other got stuck in HOCl where, formally at any rate, we imagine O is still double negative, so its Cl must (formally) have a plus one charge! It's not a cation, of course, but it's part of a neutral molecule which for its neutralilty needs Cl to adopt a +1 formal charge. (We'll see a species next lecture where a Cl has a formal +7 charge!)
So there's some bite left in that HOCl acid or its neutralized anion since
ClO- (aq) + H2O(l) + 2 e- -----> Cl- (aq) + 2 OH- (aq)represents the reduction hypochlorite must undergo when it oxidizes the compounds staining your clothes.
Note that we've written a reaction there which supposes that you can buy bottles of electrons...but that's impossible. This is therefore only HALF of what's actually going on; so it's called a half reaction...how unimaginatively descriptive. The other half reaction mate to this would be an oxidation, such as
2 I- (aq) ------> I2(s) + 2 e-Hypochlorite no longer has the guts (chemists would call that "electromotive force") to oxidize bromide to bromine the way chlorine itself can. But that pair of reactions isn't forbidden by the tables of electromotive force. So since by good fortune, both half reactions involve the same number of electrons, adding them together cancels the electrons exactly, and we get
ClO- (aq) + 2 I- (aq) + H2O(l) -----> Cl- (aq) + I2(s) + 2 OH- (aq)Notice in all of the last 3 equations, balancing them involved not only making sure that the atoms are conserved (neither created nor destroyed) but also that charge is conserved as well. We get paid back for going to this extra effort when trying to balance redox equations; the half reactions are much easier to balance than trying to write the whole reaction down correctly in one go! Yet another example of the Roman Emporer Caesar's favorite dictum: "Divida et imperata" ("divide and conquer").
Indeed the only real irritation associated with redox species comes in the casual nomenclature used to describe them. To try to straighten that all out, imagine that we've two species, Y and Z. Suppose that the formal charge on Y is n before but n+1 after the reaction while that on Z was m before but m-1 after. What happened was that a single electron was tossed from Y to Z. So
e-
Yn+1 - - - - - - - > Xm-1
was oxidized transfer was reduced
was where oxidation occured was where reduction occured
was the "reducing agent" was the "oxidizing agent"
was the "oxidizer"
The confusion, obviously, is that words beginning "oxidi..." and "reduc..." appear BOTH places. Not a problem if you notice the grammatical tense of how they're used...some active and some passive! So an "oxidizer" causes oxidation in another species by being itself "reduced." Get it? "Oxidizer" is active (doing to) while "reduced" is passive (being done to). With that application of a rule you'd hoped never to see outside of Rhetoric, redox nomenclature is a snap.