Lecture 14. Chemistry of Groups I, II, III and IV (cyclic polyether) ethereal oxygens

Group 1: Hydrogen (H&S chapter 9). Hydrogen is the simplest element, consisting of a single proton and electron. It has a reasonably high electro- negativity, which means that it forms covalent bonds with carbon which has a similar electronegativity of 2.5. Once it has ionized to form a proton, it has no remaining electrons, and, in theory has an ionic radius of zero. In fact, it is never a bare proton, and always retains some electron density, but still has a very small size where it is formally cationic. This gives it a very high charge density, and the proton is a very strong Lewis acid.

The proton as a Lewis acid: The strength of the proton as a Lewis acid can be seen in its affinity for some Lewis bases: H + (aq) + OH - (aq) = H 2 O(l) log K = 14.0 H + (aq) + F - (aq) = HF(aq) log K = 3.2 H + (aq) + NH 3 (aq) =NH 4 + (aq) log K = 9.22 It is, however, fairly hard in the HSAB sense, so that we find that it has high affinity for hard ligands such as F -, OH -, and NH 3, but virtually no affinity for soft halide ions such as Cl -, Br -, or I -.

The hydronium ion: The small proton in aqueous solution forms the linear two-coordinate [H 2 O­H­OH 2 ] + ion, shown below (left). This low coordination number is expected from the small size of the proton. [H(OH 2 ) 2 ] +.4H 2 O H+H+

The hydride anion, H - The electron affinity of H is high enough that it can form a negative hydride anion, H -, where it achieves the He electron configuration. However, H - is a very strong Lewis base. Thus, H - displaces OH - from water to give H 2 and OH -, showing that it is a stronger base than OH - : H - (aq) + H 2 O =OH - (aq) + H 2 (g) The saline (salt-like) hydrides are formed by the group 1 and 2 metal ions, e.g. NaH or CaH 2. Here the hydride ion resembles an F - anion in salts such as NaF, and has about the same ionic radius as F -. Many other more electronegative metals form more covalent hydrides, such as [Al 2 H 6 ] or transition metal ions that have covalent bonds to H, as will be discussed later.

Hydride as a ligand: H H H H H H bridging hydride Al H [Al 2 H 6 ] [Fe(CO) 4 H 2 ] H

Hydrogen bonding: One of the most important properties of the proton is its ability to form H-bonds when attached to more electronegative donor atoms such as F, O, or N. An H-bond is judged to be present when the separation between the two atoms forming the H-bond is less than the sum of the van der Waal’s radii. For H-bonds typical X-H-X distances between X atoms are: O- -H- -O 2.76 Å F- - H- -F 2.55 Å N- -H- -N 3.00 Å Figure 2. Hydrogen bonding of four water molecules around a central water molecule.

Group 1: The Alkali Metals: Li, Na, K, Rb, Cs (H&S Chapter 10). Li + Na + K + Rb + Cs + 1Å1Å

The alkali metals are very reactive, and react violently with water to give the metal hydroxide and H 2 gas. The standard reduction potentials are very negative in accord with this: Li + (aq) +e =Li(s) E o = V Li (s)+ H 2 O = Li + (aq) + OH - (aq) + H 2 (g) Because of their low charge and large size, the ability of the group 1 metal ions to form complexes in solution is limited. Thus, the metal hydroxides are completely ionized to give metal cations and hydroxide ions. They are therefore strong bases.

The coordination numbers increase with increasing metal ion size: Metal ion:Li + Na + K + Rb + Cs + Ionic radius (Å): Coord. No.: Figure 3. A four coordinate complex of Li + is seen (left) with four THF (tetrahydrofuran) molecules attached to the Li +.

Crown ethers and Cryptands The low electronegativity of the alkali metals means that they are very hard in the HSAB sense, and their chemistry is largely that of being bound to the hard oxygen donor atoms, as seen for [Li(THF) 4 ] + above. The most important aspect of their chemistry is their ability to bind to crown ethers and cryptands. The crown ethers were discovered in 1967 by Charles Pedersen when he was working at DuPont. These are cyclic polyethers called macrocycles (‘large cycles’). Some examples of crown ethers and cryptands are shown below (Figure 4):

Crown ethers and Cryptands Figure 4. Cryptands and crown ethers.

The important aspect of the crown ethers was that these complexed alkali metal cations in solution. Up until that time it was considered that the alkali metal ions had very little ability to form complexes in aqueous solution. This was important, because ion channels in cell membranes allowed K + and Na + to pass through selectively, and the properties of the crown ethers suggested how this might be achieved. The striking feature of crown ethers was their ability to complex alkali metal ions selectively on the basis of their size. Figure 5. The D 3d conformer of the free 18-crown-6 ligand, and its complex with K +, showing how well the K + cation fits into the cavity of the ligand.

Thus, the log K 1 values for 18-crown-6 with alkali metal ions vary in aqueous solution as shown below. The diagram shows that 18-crown-6 has a definite preference for the K + ion. This can be understood by looking at a space-filling drawing (Figure 5) of 18-crown-6, and how the K + cation can fit into the cavity in the ligand. Figure 6. Variation in log K 1 for 18-crown-6 complexes as a function of metal ion radius for alkali metal ions.

Cryptands: The cryptands were developed by Jean-Marie Lehn, and have a three-dimensional cavity. The complexes they form with group 1 and 2 metal ions are thermodynamically much more stable than those formed by crown ethers. Cryptand-2,2,2

K + cryptand-2,2,2 complex: K+K+ cryptand

The Alkali Earth Metals (group 2). (H&S Chapter 11) The alkali earth metal ions resemble the alkali metal ions in having a low electronegativity, and being very hard in the HSAB classification. The big difference, though, is their charge, which makes them stronger Lewis acids. The effect of charge on log K 1 for hard metal ions with EDTA, all having an ionic radius of about 1.0 Å, makes this point (see next slide for Ca EDTA complex): Metal ion:Na + Ca 2+ La 3+ Th 4+ Ionic radius (Å): log K 1 (EDTA):

We thus find that the metal ions in Group 2 are much better at complexing with ligands than are those in Group 1. Being hard, complexing of Group 2 cations is confined largely to oxygen donors, and to nitrogens, more so where the nitrogen donors are part of a ligand that also has some oxygen donors, such as in EDTA. EDTA

The alkali earth metal ions Ca 2+, and particularly Sr 2+, and Ba 2+ are large enough to fit well into the cavities of crown ethers and cryptands, and actually form more stable complexes than large alkali metal ions. Thus, we can compare log K 1 values with some crown ethers and cryptands for Ba 2+ and K +, which are almost identical in size: Ligand: 18-crown-6 15-crown-5 cryptand-222 log K 1 (K + ): log K 1 (Ba 2+ ): Thus, even with these ligands, the charge on the metal ion has an effect on complex stability.

Group 3. B, Al, Ga, In, Tl. (H&S Chapter 12). In group 3 the electronegativity of the metals is getting a bit higher, and the heavier metals Ga, In, and Tl are actually post-transition elements (they are close to Au), so have much higher electronegativity and a very different chemistry from B and Al. They form trivalent cations that form very strong complexes: Metal ion:Al(III) Ga(III)In(III) Tl(III) ionic radius (Å): log K 1 (OH) log K 1 (EDTA): increasing electronegativity

The Tl(III) ion is stabilized by complexation with ligands, and is an extremely powerful Lewis acid. Because of its high electronegativity, Tl(III) is classified as soft in HSAB, as reflected by its log K 1 values with halide ions: Metal ion:Al 3+ Ga 3+ In 3+ Tl 3+ log K 1 (F - ): log K 1 (Cl - ): HARD ← → SOFT

The inert pair effect in Thallium(I): For the first time we have to consider the inert pair effect. Thus, for Tl, the most stable oxidation state is not Tl(III) but Tl(I). The Tl(I) ion has an ionic radius of 1.50 Å, and so resembles K + and Rb + to some extent in its chemistry. It does have some tendency towards covalence (it is soft), and so forms many complexes where it is bound to soft donors such as S. At right is seen the complex of Tl(I) with the sulfur-donor macrocycle 9-ane-S 3. position of lone pair Figure 8. Structure of the Tl(I) complex with the S-donor macrocycle 9-ane-S 3.

Boron Boron is very different in its chemistry from the other members of the group. While they all have preferred coordination numbers of 6, with occasional higher coordination numbers of 7 or 8, boron always has a coordination number of four or less. Thus, B(III) in aqueous solution exists as B(OH) 3 (aq) at lower pH, and is too acidic to ever be protonated to yield a B 3+ (aq) ion. At higher pH (9.1) a water coordinated to B(OH) 3 (aq) ionizes to yield the borate anion: B(OH) 3.OH 2 ( aq ) = [B(OH) 4 ] - ( a q) + H + ( aq ) pK a = 9.1 “boric acid”borate anion This behavior is readily understood in terms of the small size (ionic radius = 0.11 Å) and high charge on B(III).

B(III) forms compounds of considerable covalency, with electronegativity = 2.0, and forms a reasonably stable hydride, as in Li[BH 4 ], lithium borohydride. Here we have a T d [BH 4 ] - anion, which is used in organic chemistry as a mild reductant. The chemistry of the boranes, those compounds involving boron and hydrogen, is enormous. The structure of [B 2 H 6 ] is shown below. Figure 9. B 2 H 6, showing the bridging H-atoms, which donate electron density to the adjacent B atom.

Group 4. C, Si, Ge, Sn, Pb. (H&S Chapter 14). Here the group valency is four. The electronegativity of the elements has risen quite high, with the C atom having an electronegativity of 2.5. None of these elements forms an M 4+ cation in solution. Carbon forms the CO 3 2- and HCO 3 - anions at higher pH, and at lower pH (<6) breaks up to form CO 2 (g). Si and Ge form many compounds with a coordination number of four, such as SiCl 4 or GeCl 4. They also readily expand their coordination numbers to six, as in complexes such as [SiF 6 ] 2- and [GeF 6 ] 2-.

The inert pair effect in Pb(II) and Sn(II): The high electronegativity of these elements leads to a strong inert pair in Sn and Pb. For Sn both the Sn(IV) abd Sn(II) state are relatively stable. For Pb, the Pb(IV) state is of rather low stability. Important Pb(IV) compounds are PbO 2, which is important in the lead/acid battery, and Pb(CH 2 CH 3 ) 4 (tetraethyl lead), which used to be added to gasoline to prevent ‘knock’ (premature ignition on compression). The lead/acid battery works on the cell: PbO 2 (s) + 4 H + (aq) +Pb(s) = 2 Pb 2+ (aq) + 2 H 2 O E o = +1.2 V

The Pb(II) and Sn(II) ions display a sterically active inert pair, which means that in structures of the complexes of these cations, there is usually a gap in the coordination geometry which is occupied by the lone pair. This resembles the structure of NH 3 as predicted by VSEPR, where the structure is derived from a tetrahedron, with one site occupied by the lone pair. This is seen in the structures below of the [SnCl 3 ] - and the [Pb(C 6 H 5 ) 3 ] - anions: [SnCl 3 ] - [Pb(C 6 H 5 ) 3 ] - lone pair Cl Sn phenyl group Pb