Na(s) -----> Na+(aq) + e-
and the concomitant reduction of something. Since diatomic hydrogen gas is evolved, and, as an element, the H atoms in the molecule share the electron pair perfectly (covalently), there's no ionic character to H:H with which we might assign charges to the "ions." So each H atom in H:H has a formal charge of ZERO. But that wasn't the case in water; there the oxygen gained the lion's share of both bonding pairs. Oxygen in water has a formal charge of 2- which means each H atom in water must bear a formal charge of +1.So, during the reaction with sodium, it is the hydrogen which picks up an electron and has its formal charge reduced from +1 to zero:
2 H2O(l) + 2 e- -----> 2 OH-(aq) + H2(g)
Those formal charges on the atoms in water give rise to many of its interesting properties. They make the molecule polar (since it's bent) and hold water together as a liquid at temperatures where other molecules of similar molar mass have long since vaporized. The inter-water attractions (of the +1 H for a neighbor's -2 O atom) are called "hydrogen bonds" and hold not only water together by your DNA as well! Indeed "hydrogen bonds" (not the hydrogen bond, H:H) occur between and molecules with (a) ionic hydrogen (that let's out hydrocarbons) and (b) a strong electronegative atom; so HF exhibits strong hydrogen bonding as well.But what's left behind in that sodium-water reaction is a very caustic, corrosive, alkaline solution with hydroxyl (negatively charged OH) ions in it. Those common words, caustic, alkaline, refer to what chemists call a base. When chemists say a solution is basic, they don't mean it's "fundamental" but rather that it's alkaline or caustic; that means it is likely to be awash in those negative hydroxyl ions. Common bases include ammonia water and lye.
What of the opposite sort of corrosion, an acidic one? There it's not hydroxyl ions in abundance but rather the other half of water, so to speak: cationic (positive charged) hydrogen ion, that is, a proton. However, free protons are never found in aqueous solutions; they latch onto water molecules right away. But where do they come from...molecules eager to lose a proton to just about anything. Such molecules are called strong acids, symbolized by HA.
HA(aq) + H2O(l) -----> H3O+(aq) + A-(aq)
The "A" there stands for any Anion interested in ridding itself of the proton. And the proton? It latched onto the water to become that
H3O+ (aq), Hydronium ion
It is the quantitative production of hydronium ion which marks strong acids in aqueous solutions; no HA(aq) survives this reaction...which can be very exothermic, "heat releasing."But acids can exist in nonaqueous solutions too. All that's required is that HA be eager to donate a proton to something which is eager to accept it. According to Brønsted-Lowry definitions, the proton donor is the acid and the proton acceptor is the base.
That means that the water molecule which accepts a proton from a strong acid is acting as a base. But if aqueous solutions can be basic as well as acidic, water must act differently there. In fact, strong bases, B:, accept...no, demand...protons from water as
B:(aq) + H2O(l) -----> BH+(aq) + OH-(aq)
So in basic solutions, water donates protons and must be considered the acid! It's all relative; in some pairing, the acid is whichever molecule most readily gives up a proton to one which will accept it. Bear in mind, that proton transfer need not occur in chemical reactions, but if it does, that's the hallmark of an acid-base reaction type.One of the strongest acids is perchoric acid (it's also an enormously strong oxidizer):
HClO4(aq) + H2O(l) -----> H3O+(aq) + ClO4-(aq)
And the oxygen in sodium oxide is one of the strongest bases:
Na2O(s) + H2O(l) -----> 2 Na+(aq) + 2 OH-(aq)
Now true, that O atom didn't end up as a cation, but it would've had a negative 2 charge but for the hydrogen atom it stole; so it ends up as another (singly) negative hydroxyl ion.Since we take the trouble to characterize acids as STRONG, it implies that there might be WEAK ones as well. Indeed there are many more weak than strong acids. The same holds true for bases. The weakness stems from incomplete proton transfer; i.e., the acid/base's reaction in water is not quantitative. Many of the parent molecules are left in solution at equilibrium. Such reactions are indicted with double-headed arrows which we'll approximate here as "<=====>" So a typical weak base example might be
NH3(aq) + H2O(l) <=====> NH4+(aq) + OH-(aq)
where that equilibrium lies far to the left; so the arrows would be more descriptive if they read
->
<------
signifying the weakness of the reaction and of the base.The archtypal weak acid is acetic acid (vinegar) where
CH3-C=0 (aq) + H2O(l) <=====> CH3-C=O + H3O+ (aq)
\ \
O-H O-
and again the equilibrium lies to the left.
It was mentioned that acids and bases, defined as donors and acceptors of protons (Brønsted), need not be confined to aqueous solution. As a counterexample consider our one-time, sometime, never-never lecture demo:
NH3(g) + HCl(g) -----> NH4Cl(s)
or even
NH3(l) + HCl(l) -----> NH4Cl(in NH3 solution)
The shapes of some of these species are interesting and easy to come by. We can take advantage of their isoelectronic counterparts to determine it. For exampleNH4+ is isolectronic with CH4 and so must be AX4 tetrahedral H3O+ is isoelectronic with NH3 and so must be AX3E trigonal pyramidal OH- is isoelectronic with HF and so must be AX1E3linearWhat's truly unfortunate about these chemical symbols is that many of them appear to place the overall charge on the wrong atom. Chemists don't worry about this sort of thing since they understand that an ion's charge refers to the entire molecule not just the atom which happens to be written last in the formula. So the positive charge on hydronium doesn't reside on the oxygen as written but rather is shared by all of the hydrogens. Ditto hydroxyl ion; that negative charge doesn't really belong on the hydrogen. It's just habit that causes the hydroxyl species to be written with the oxygen first, but in all ionic species, the ionic charge comes last.
Sometimes to make matters more clear, the formulas will be written with square brackets to indicate the entire species carries the charge:
[NH4]+ [H3O]+ [OH]-
Finally, we have introduced some new and potentially marvelous species in this lecture. For example, the strong acid
HClO4 to which we should add
HNO3 and H2SO4 this latter reacting in water as
H2SO4(aq) + H2O(l) -----> H3O+(aq) + HSO4-(aq)
That final species is instead a weak acid:
HSO4-(aq) + H2O(l) -----> H3O+(aq) + SO42-(aq)
What (should be) marvelous about these critters is that many of them violate the heck out of the octet rule. The nitrate ion is OK:
..
:O:
\ .. _
N=O: show only 8 electrons around N
/ ¨
:O:
¨
if you remember that the double bond "=" is worth two pair (or 4 electrons). What nitrate appears to violate is not the octet rule but rather nitrogen's normal valency. If we count the "formal charges" in this ion, and we formally assign negative 2 to each oxygen, they're worth -6, but the ion's showing -1 so the nitrogen must balance this off with a formal +5 charge. This is not the expected +3 for nitrogen which makes us expect that it would much rather be reduced (gain some electrons); not surprisingly then, nitrate ion is a good oxidizer. It'll even oxidize ammonia; so the ammonia nitrate molecule ought to be unstable. It is what brought down Oklahoma City's Federal Building (and almost the entire town of Texas City earlier this century).But more alarming is the perchlorate ion whose Lewis structure is
.. _
:O:
.. | ..
O=Cl=O
¨ || ¨
O:
¨
where that double bond to the south of Cl was intended. Adding up the electrons around Cl gives us not 8 but rather 14! Cl and other atoms in and beyond Period 3 have the capacity to tap unused valence electron orbitals beyond those available to Periods 1 and 2 elements. (More about that later.) Moreover, the same "formal charge" count here gives Cl a charge of +7 rather than the -1 it would prefer. Not surprisingly then, perchlorate is a VIOLENT oxidizer; it figures prominently in fireworks.
.. 2-
For completeness: :O:
.. | ..
O=S=O
¨ | ¨
:O:
¨