Lecture Notes from CHM 1341
5 July 1996

Sulfur

Sulfur is an ancient element...not in the sense that it's been around any longer than the other heavy elemental ejecta from local dying stars, but rather that, being found in its native (elemental) form, sulfur was known to our distant ancestors. In fact, "sulvere" comes from the Sanskrit, a long dead language of a prebiblical civilization in the Middle East.

In a region with any volcanic activity, it'll be impossible to miss! Take, for example, Yellowstone National Park; the earth and the geyser pools (above) are yellow because of their native sulfur content.

Being an atom in the same Group VI as oxygen, sulfur is even the "oxygen of choice" for an entire ecology of life in the deep oceanic rifts where volcanic "black smokers," so named due to their metal sulfide content, spew superheated water and sulfur onto an ocean floor upon which no light of day ever falls. Life has evolved even here to take advantage of this "opportunity," cheerfully replacing oxygen's chemical magic with sulfur's! This persistence and prospering in the face of apparent adversity makes one sanguine about the prospects for Life off of the earth!

Sulfur also hits the top of the charts with chemical companies. Not counting metals, water, and salt, sulfuric acid is produced and consumed as the highest volume industrial chemical (40 Gkg in 1990). One of its prime functions is to acidify mineral phosphates to phosphoric acid and intermediate oxoanions (see below) for fertilizers. And it does this because concentrated sulfuric acid is one of the most powerful acids known.

The element exists as a crown of 8 sulfur atoms to a sulfur molecule. These molecules stack together in a crystalline form called orthorhombic at room temperature. (A rhombus is 3d diamond shape; it has all its sides inclined or "tilted.") At slightly higher temperatures, the packing in the crystal changes to monoclinic (only one direction inclined). But it becomes truly tacky (pun intended) at temperatures above its melt because the 8 atom molecules polymerize into chains when (quasi) stabilized by quenching. (Quenching just involves rapid cooling as when a blacksmith quenches a red-hot horseshoe in water.) The resultant plastic sulfur has the consistency and resilence of putty until it reverts to the 8 atom crown. Although S₈ would be correct, it would unnecessarily complicate equations. So we'll follow the convention of using only S when we mean this molecular solid.

The /\/\/\/\/ bonding pattern in this crown of 8 sulfurs preserves the angular binding preference we know sulfur must have from its binding or electronic configuration:

     AX2E2   or   [Ne] 3s2 3p4   or   four sp3 orbitals (2 with lone pairs)

The same argument would lead us to expect that phosphorus, with one lone pair and three unpaired electrons in those tetrahedral hybrid orbitals would find its elemental molecular bonding well satisfied with one P atom at each corner of a tetrahedron and bonded to each of the other 3. So P₄ is, not surprisingly, the molecular formula for (white) phosphorus.

We've already been exposed to a great many sulfur compounds...no, I mean in this course. And many involve literal oxidation (with oxygen) to form the progression:

           O2          O2           O2
     S2- ------>  S  ------>  SO2 ------>  SO3
   sulfide      sulfur       sulfur       sulfur
                            dioxide      trioxide
     .. 2-        .              ..      ..     ..
    : S:         :S:       : S = O      :O = S = O:
      ¨           ·         ||   ¨          ||
                           : O:            : O:

While the sulfide and sulfur Lewis structures look OK, the di- and tri-oxide show 10 and a dozen electrons, respectively, around the sulfur! This violates Lewis's octet rule and stretches credulity of the 3s and 3p capacities. You may be confident of one thing: oxygen could never do this! It obeys the octet rule strictly because 2s and 2p and their 8 electrons complete the n=2 shell. Not so for sulfur.

Being in the n=3 shell, sulfur has orbitals 3s, 3p, and 3d; it just doesn't have enough electrons to populate the 3d unless many atoms rich in electrons (like oxygen) approach to bind. Then sulfur pulls carbon's trick and goes it two better. Instead of hybridizing merely sp³, sulfur can hybridize sp³d² (octahedral).

Those six equivalent hybrids, pointing up and down each of the Cartesian axes (X, Y, and Z), and lo, room is found for a dozen electrons. We'll soon see that sulfur is not alone in this ability; it's a possibility for any atom above Period 2.

The German notion of sauerstoff (oxygen), the acid-making atom, is borne out in spades here since those sulfur oxides hydrolyze (react with water) to produce the weak sulfurous and the strong sulfuric acids:

                             H
                            /
   H                      :O:
    \   ..                 |
    :O - S = O:       :O = S = O:
     ¨   |   ¨         ¨   |   ¨
        :O:               :O:
        /                 /
       H                 H

   Sulfurous acid      Sulfuric Acid

(Note the 10 e^- about S in H₂SO₃ comes from sp³d¹ not sp³d² hybridization.)

Since oxygen is so electronegative, more O atoms draw electrons away from the O-H bonds, making them vulnerable to ionization...that is, more acidic. It's a trend seen in all the oxoacids. Of course, as a proton leaves, both electrons in the bond that once held it are retained by the anion...in this case, the oxoanion. This reduces that electron withdrawing tendency somewhat, stabilizing the remaining acidic hydrogens somewhat. So polyprotic oxoacids get less and less acidic as their hydrogens ionize away.

So H₂SO₄ -----> H⁺ + HSO₄^- (hydrogensulfate or "bisulfate" ion)

but

HSO₄^- <====> H⁺ + SO₄^2- (sulfate ion) INCOMPLETE REACTION

That is, hydrogensulfate ion is only weakly acidic while sulfuric acid is violently acidic...and oxidizing...and dessicating (a strong dehydrating agent). Triple threat that. Sulfuric's dessicating power is so extreme that it is strongly inadvisable to dilute concentrated sulfuric acid by adding water! The energy released by sulfuric's violent hydration will splatter the solution...on the unlucky chemist. Instead one dilutes conc. sulfuric by adding it SLOWLY to water; then if the mixture splatters, at least it's already dilute!

Dangerous as it is, sulfuric is occasionally used as a drying agent, dessicating the air above it and driving water out of hydrated compounds like the beautiful blue copper(II) sulfate which crystallizes with 5 "waters of hydration" (sounds redundant, doesn't it?) in the crystal structure.

CuSO₄ · 5H₂O (blue) -----> CuSO₄(white) + 5 H₂O trapped in H₂SO₄(conc)

Acting as a dessicant, the conc. sulfuric becomes diluted. Acting as an acid, it becomes neutralized to sulfate ion. But acting as an oxidizer, sulfuric has to become reduced...that is, accept electrons and change its electronic state. But that latter is just the reverse of the sequence with which we started these notes.

Thus, when oxidizing a substance, sulfur trioxide (or its acid) becomes reduced to sulfur dioxide (or its acid). How can we tell that represents a reduction? Well, there's fewer oxygen atoms in the molecule; so in some sense sulfur as the dioxide gets to hang on to more electrons than sulfur as the trioxide.

All this can be placed on more trustworthy footing with the introduction of oxidation numbers next.

Chris Parr
University of Texas at Dallas
Programs in Chemistry, Room BE3.506
P.O. Box 830688  M/S BE2.6 (for snailmail)
Richardson, TX  75083-0688

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