Ionic Liquids - ACS Publications - American Chemical Society

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Chapter 18

Significance of Cations in Ionic Liquids Chemistry

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Keith E. Johnson, Li Xiao, and Gordon Driver Department of Chemistry and Biochemistry, University of Regina, 3737 Wascana Parkway, Regina, Saskatchewan S4S 0A2, Canada

There is a tendancy to classify ionic liquids either as individual solvents (LiCl-KCl vs. NaCl-KCl, NaNO -KNO vs. LiNO -KNO ) or in terms of the anions. In fact, a compromise is appropriate. We shall point out how the cation choice influences the chemistry, including the electrochemistry, of an anion in an ionic liquid. Some examples follow. The effect of the cation on the structure of liquid alkali halides is exemplified by the coordination chemistry of dissolved Ni(II) as well as by diffraction studies. Raman spectroscopy demonstrates the cation influence on the fundamental equilibrium 2AlCl --> Cl + Al Cl in alkali chloroaluminates. When the electrochemical reduction of NO in alkali nitrates is examined, we observe product dependence that can be traced back to metal + oxygen reactions. Regarding the Lewis acidity of room temperature ionic liquids, it has been shown that the cracking of alkanes and the dissolution of heavy oil are influenced by the cation present with a given haloaluminate. 3

3

3

3

-

4

-

-

2

7

-

3

230

© 2002 American Chemical Society

In Ionic Liquids; Rogers, R., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2002.

231 This is a cautionary tale. We will attempt to point out the relationship of the old "molten salts" research to the new "ionic liquids" research and to demonstrate the strengths and weaknesses of classifying ionic liquids by the nature of the anions - a situation quite pronounced for the older literature but not to be ignored for ambient temperature systems.

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W h a t is an Ionic L i q u i d ? The glib definition is a liquid composed totally of ions with the forces overwhelmingly coulombic - a relaxed simple ionic crystal such as K G . Indeed, most ionic liquids are salts but there are intriguing situations such as the salt H g C l which has very low conductivity in the liquid state and the molecule PC1 which turns out to be ionic (PC1 PC1 *) in the solid. There are also ionic liquids such as emimAlCl and its mixtures with emimCl where there is evidence of considerable hydrogen bonding between the emim and the anions (emim=lethyl-3-methyl-lif-imidazolium). Then we have systems in which ionization is incomplete or significant ion pairing occurs. One might arbitrarily define an ionic liquid as one in which more than half the "molecules" are ionized or one could be even less precise and talk about a liquid in which the chemistry is primarily determined by the ions even though they are in a minority e.g. the 28 mol% K O H of the Stuart electrolyzer (7). We are then concerned with the chemistry of electrolyte solutions where the solvent itself is molecular but ionizing. 2

5

+

4

6

4

+

Simple Ionic Salts as Solvents In this section we will discuss a variety of examples of medium to high temperature ionic liquids in which the cations are obviously far more than spectator ions viz. basic physieoehemical properties and structure, electronic spectra of dissolved transition metal ions and some thermodynamic properties and electrode reactions of nitrates.

Basic Physieoehemical Properties and Structure Over 40 years ago, X-ray diffraction studies showed that the melting of simple alkali halides produced liquids with alternating anions and cations and average first coordination numbers showing the effects of radius ratios (2). Neutron diffraction is commoner for this work now and details beyond the first "sphere" can be well matched by theoretical calculations (3). We must be

In Ionic Liquids; Rogers, R., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2002.

232 careful in assigning any meaning or precise structure to a given set of numbers. We also can learn of structural changes in a series of liquids from measurements of viscosity and conductivity. These two quantities are expected to move in opposite directions with a change of one ion component although the charge aspect of conductivity need not always correlate with size.

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Electronic Spectra of Dissolved Transition Metal Ions A n examination of the UV-visible spectra of the 3d ions in the single solvent L i C l - K C l (4) indicated that Cr(III) is octahedrally coordinated, Fe(II), Co(II) and Mn(II) tetrahedrally coordinated, Cu(II), Cr(II) and Ni(II) take up distorted tetrahedral geometry and Ti(III), V(III) and V(II) exist in equilibrium mixtures of octahedral and tetrahedral sites (but apparently not in 5-coordination). The situation with Ni(II) was studied in considerable detail ( N i C l in L i C l - K C l at different temperatures (5), N i C l in 5 individual alkali chlorides (6), N i C l in various C s C l - N i C l (7), CsCl-ZnCl (8) and L i C l - K C l (9) mixtures) with the conclusions that in most liquid chlorides the N i ion is found in both weldefined tetrahedral (T) and less-defined octahedral (O) sites; it was suggested that larger cations form the shell around the Τ sites and the smaller cations e.g. L i , polarize the chlorides of the Ο sites, with increasing temperature favoring Τ site formation and occupation. Figure 1 shows a selection of these spectra compared with that for tetrachedral N1CI4 " in an organic salt (10). For detailed discussions see (11). 2

2

2

2

2

2 +

+

2

Thermodynamic Properties and Electrode Reactions of Nitrates Measurements of the e.m.f.s of cells of the type A g / A g N 0 / A g N 0 ( N ) in M N 0 / A g where the mole fraction Ν is given by 10~ > M 04

+

20 ™ 2

N0 " + 2

20ΗΓ

2

>

N0

>

-

2

2

»

2

H 0 2M

2

0 "

2

> 2N0 ~ 20H-

2

2

8. 9.

2

O '

2

2

+

V2N2 +

+ N0 " 3

>

N0

2

O

2 -

2

30 "

> Pt(NO)

3+

2

+ 20 "

+ V2O2

Voltammetric studies of alkali nitrates coupled with chemical analysis of the products of coulometry show the reduction of nitrate to nitrite at —1.5 V vs. A g / A g (13, 14, 15). When L i or N a are the counterious (as pure salts or eutectic components) this process is inhibited by oxide film formation (activation energy -60 kJ mol" ) and one observes a sharp current increase at ~-3 V with dissolution of a platinum electrode (14, 16). It was shown that in dry liquids this first electrode reduction to oxide and nitrite (Eqn. 3) is followed by chemical oxidations of oxide to peroxide (for M=Li, Na, or K ) (Eqn. 4) and to superoxide (Eqn. 5) (for M=Na and especially K ) (17, 18, 19). In addition, the oxides L i 0 and Na2Û are sparingly soluble in the liquid nitrates, thus inhibiting the electrode reactions by film formation (Eqn. 8). Only traces of N O 2 are seen at the cathode. If water is present (its solubility is ~10" M ) then a catalytic nitrate reduction is seen before the direct process and some peroxide (and superoxide) is destroyed. Alkali metal ion reduction does indeed occur at very negative potentials but leads to the corrosion of most electrode materials (Equation 11 is a feasible process). If we recall the oxygen chemistry of the alkali metals the differences between the nitrates make sense: L12O, N a 0 and K 0 are the favored products of the M/O2 reactions. +

+

+

1

2

3

2

2

2

The electrochemical reduction of CIO4" in liquid L1CIO4 (20) similarly leads to the fomation of insoluble L12O but it is possible to build up a film of L i behind the L12O at more negative potentials and this L i can react E X P L O S I V E L Y with the bulk perchlorate! l

In Ionic Liquids; Rogers, R., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2002.

235 Haloaluminate and Hydrogen Halide Systems Haloaluminate systems (ZX/AIX3 ) are controlled by equilibria of the type illustrated in Table II for X ^ C l . Theoretical calculations (21), which apply to the G A S P H A S E at 298 K , gave a value for the equilibrium constant Κ for reaction 5, the fundamental one, of 7 χ 10 . Experimental estimates of this quantity with the necessary inclusion of the counter ions, Z , are ~10 for liquid N a salts at 598 Κ (22) and ~10 -10 for liquid emim salts at 298 Κ (23). The magnitude of K , even at 598 K , allows for the clear division of systems of this type into Lewis basic (Xzci XA!C1 ) & Lewis acidic ( X A I C I Z C l ) systems with their divers chemistries. Hydrogen halides in these systems in turn behave as Brônsted acids (in Lewis basic liquids) or superacids (in Lewis acid liquids). For the pure hydrogen halides there is evidence, both theoretical and experimental (24), for the formation of species H X " (n = 1-4) and H X (m = 1,2) while within the haloaluminate liquids polyhalahydrogenate (l-)anions and haloaluminate - hydrogen halides have been observed or proposed (25, 26). These adducts are further examples of the established line of hydrogen-bonded anions QHR" with Q = halide and R = haloaluminate. Table II includes examples of these equilibria for chlorides. 32

7

+

+

16

18

>

+

m

>

Downloaded by UNIV OF BATH on October 30, 2014 | http://pubs.acs.org Publication Date: July 25, 2002 | doi: 10.1021/bk-2002-0818.ch018

3

x

3

n

n+l

m + 1

m

+

+

The nature of the cation Z , besides determining K , influences the struture and reactivity of the chloroaluminate liquids in several ways. The melting points follow from the lattice energies which are much lower for organic than for simple inorganic salts. Within the alkali metal series (27), it was shown by Raman spectroscopy that (i) within the AICI3-L1CI system, AICI4" and AI2CI6 persist to XAICI 0.75; (ii) higher chloroaluminate polymer formation increases in the order Li >

HC1 >

+

AI3CI10"

> HC1 ~

HC1

+

7

>

HC1

2

2

7

HC1

+

A1 C1 "

>

AICI3

+

2

2

6

>

2

+

2

AICI3

7

+

A1 C1

H C1 " 2

H C1

H C1 3

HC1 " + 2

>

+ 2

H C1

+

2

>

3

HCl

+

2

>

HCl

3

2HC1 " 2

QHÇIAICI3-

>

C1HC1A1 C1 " 2

6

> AICI4" + C l H C l A k C k "

Trialkylsulfonium salts have not been investigated over wide AICI3 mole fraction ranges - certain systems form liquids at ambient temperatures with excess aluminum halides or with hydrogen halides (XHX 0.5) (31, 32). These cations offer neither hydrogen bonding nor aromatic character and thus provide a useful comparison for some reactions. Simple pyridinium salts present other possibilités through their onium character. It has been shown that pyridinium chloride itself (M. Pt. 144°) can play the role of a Bronsted acid in cleaving alkyl aryl ethers (33), can form anions H^Cl^+j" with H C l and can form cations pyHClHpy in which the protons on nitrogen are N O T exchangeable (on the H N M R time scale) with those of >

+

l

In Ionic Liquids; Rogers, R., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2002.

237 H C\ \" (34). The species pyHClHpy may be thought of as H2C1 (Equation 10 of Table II) stabilized by pyridine. We have shown that acidic pyridinium chloroaluminate liquids (XAICI - X p y H C l ) - ^ long-chain alkanes, giving rise to branched C 4 - C 7 alkanes (see Figure 2) and lesser amounts of alkenes (35). These liquids are also capable of dissolving heavy oils (36) with a significant reduction in the higher molecular weight (hence higher boiling) components (Figure 3). This superacidic behavior is also shown by the acidic haloaluminates of the other cations (emim , bmin , R3S , N-Bupy ) but their rates of reaction are less. Furthermore, although the presence of protons as C l H C l A ^ C l g " enhances the reactivity, it is not essential: cracking still occurs when this proton type is reduced to below electrochemical detection limits by M A C (methylaluminum sesquichloride) which does not attack pyH ! While basic haloaluminate mixtures do not crack alkanes, salts of HX2" (with Hammett acidities comparable to that of concentrated hydrochloric acid) show some reactivity but at too low a level at 25°C to distinguish cationic differences. In making comparisons of the reactivity of various haloaluminate systems, it has been customary to think in terms of first the mole fraction of AICI3 and then those of the anions present e.g. AICI4" and AI2CI7" in Lewis acidic mixtures. We have calculated the molarities of the ions for a number of mixtures (Table III) in order to ascertain whether concentration is a more appropriate quantity to consider in making such evaluations: clearly the concentrations of given anions are influenced by the sizes of the counter-cations. +

n

+

n+

>

crac

3

+

+

+

+

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+

Table III: Mole percentage and molarity of anions in different chloroaluminate ionic liquid systems: Density AlCh Ref M ici MW M cif Mole% (gmL- ) (mol/L) emimCl 50.00 1.2941 4.62 279.94 0 29 M W : 146.6 54.99 1.3209 0.95 3.32 309.51 60.01 1.3488 1.94 346.69 1.95 66.66 1.3888 413.20 0 3.36 Me SCl 66.66 1.40 378.78 0 3.69 38 MW:112.5 66.66 1.59 423.28 32 3.75 0 Me SBr* MW:157 50.00 332.34 1.4116 37 4.25 0 Et SBr* MW:199 54.99 1.4286 361.97 3.07 0.88 60.01 1.4455 399.01 1.81 1.81 66.66 1.4680 465.28 0 3.15 * For the mixed halogen systems we have ignored exchange and we quote molarities of A l C l B r " and A l C l B r " . Note: NaCl with a formula weight of 58.5 has a molarity of 26.6 at the temperature of 801°C. A

4

Ah

1

3

3

3

3

2

6

In Ionic Liquids; Rogers, R., et al.; ACS Symposium Series; American Chemical Society: Washington, DC, 2002.

238 4.60