by: Beichen Liu
When I first joined the Gebbie Group last winter, I often found myself at a loss when friends and family asked me to explain the research field that I was joining. How do I describe a field that I personally was new to in a succinct and understandable way to people who were unfamiliar with electrochemistry? How do I explain what ionic liquids even were, much less how they can be applied to electrochemistry? Ionic liquids were something I learned about not even months before! And truthfully, they were something that I only had a basic working knowledge of when I joined the group – and that was all because of the material that Matt gave me and my group members throughout the course of the advisor-selection process.
As I dug deeper into the scientific literature, I realized how expansive the design space for these ionic compounds were. There were so many interesting types and kinds of ionic liquids that covered a wide range of properties and characteristics. It is my hope that through this post, and many more to come, that I can make some of these concepts and ideas more accessible to readers who are curious but might not have the background or time to wade through the countless papers on this topic.
What Are Ionic liquids?
On a very basic level, ionic liquids, or ILs, are just salts! If you’re picturing table salt, then you’re definitely on the right track. A salt is a compound that contains a positively-charged cation and a negatively-charged anion. In general terms, these positive and negative charges balance out, making the overall compound neutral in charge. Regular table salt, NaCl, is made up of the Na+ cation and the Cl– anion.
A key characteristic of ionic liquids is that they exist in the liquid phase, which is different from many of the traditional salts you’d encounter in a typical chemistry lab. Those salts, like NaCl or KCl, are solid at room temperature and only melt at extremely high temperatures. At these temperatures (800.7 °C for NaCl, 770 °C for KCl)1,2, these salts are considered “molten salts.” To turn table salt into a liquid form at room temperature, you would need to dissolve it in water. But this makes table salt into an ion solution, not an ionic liquid. ILs are liquid at more moderate temperatures without needing to be dissolved in another solvent.
So what makes ILs different? It all comes down to the structure of the cations and anions. Traditional salts have relatively small cations and anions, which means that they can pack together in a very structured and ordered manner. Their small sizes also mean that they have strong interactions with each other. Consider again table salt, made of Na+ and Cl– ions – they can form this very nice lattice shape:
Example of lattice structure of NaCl. Here, there is a regular repeating structure of Na+ cations and Cl- anions in the lattice. (Image from Wikimedia Commons)
The tight packing of the ions means that this configuration is very energetically favorable – to disrupt this structure requires a LOT of energy. While there are definitely many more factors at play, this is one of the main reasons why traditional salts have such high melting temperatures – you need to add a lot of energy to the system to make the ions want to leave their lattice positions.
Cations and anions for ionic liquids are bulkier and more irregular in shape than traditional salts, making packing much harder. (Image from Gebbie Lab)
Ionic liquids, on the other hand, are made up very oddly-shaped ions! In many of the ILs used in research, the cation is usually a large, misshapen ion and the anion can range from small ions like Cl-, to larger ones like bistriflates (with chemical formula: C₂F₆NO₄S₂). These large and irregular ions make forming a lattice structure very difficult. Imagine trying to arrange marbles into a container – you’ll see the marbles take on a regular repeating structure. Try to do the same with rubber duckies of various sizes – no matter how you try, you won’t get a nice repeating structure like with the marbles. This less optimal packing is the reason why ILs are liquid. Without the tight lattice in place, you need to add less energy to the system via heat to make the IL flow. In fact, there are some ionic liquids that remain liquid even below the freezing temperature of water!
It’s much harder to get any sense of order for rubber duckies in a box than for marbles. This is the same concept for traditional salts and ionic liquids – try as you might, the ducks just won’t line up nicely!
A Brief History of Ionic Liquids
While ionic liquids have seen renewed interest in the past decade or two, its history actually extends back to the 19th century, when a “red oil” was observed during a Friedels-Crafts reaction3. Then in 1914, an alkylammonium nitrate salt was made that had a melting temperature of just 12 °C!3,4 In the 1960s, Dr. John Yoke of Oregon State University mixed copper (I) chloride with alkylammonium chlorides and discovered that the resulting solution formed a liquid. In the 1970s, Dr. Jerry Atwood from the University of Alabama added aluminum alkyls with salts and called the resulting product “liquid clathrates.”3
But the ILs that we think of today have their origin in an effort to find a replacement for the LiCl-KCl molten salt electrolytes that were being used for thermal batteries. These molten salt electrolytes, with melting temperatures of 355 °C (relatively low by salt standards), were often incompatible with the materials used in those batteries, resulting in issues with other devices nearby3.
1-ethyl-3-methylimidazolium tetrachloroaluminate (EMIM AlCl4), an example of chloroaluminate ionic liquids, which made up the first generation of ILs. (Image from PubChem)
In 1963, Major (Dr.) Lowell A. King at the US Air Force Academy looked into a class of chloroaluminates, which were a kind of molten salt. In fact, in many of the research papers published around this time, you may find that the term “molten salt” was often used for ILs. It wasn’t until later that “ionic liquids” overtook “molten salts” as the preferred descriptor. These chloroaluminates were a mixture of alkali halides and aluminum chlorides with much lower melting temperatures. For instance, NaCl-AlCl3 melted at just 107 °C!3
While chloroaluminates were a very attractive alternative, a key drawback was their reactivity with water, which was a serious safety hazard. In the 1990s, Dr. Mike Zaworotko looked to solve this issue of water reactivity by water-stable anions.3 From then on, scientists have gradually improved on the stability of ionic liquids through many different methods, which added to the ever-growing list of ionic liquids. Today, imidazolium-based ILs are one of the most popular classes of ionic liquids5.
An example of an imidazolium cation. This is 1-ethyl-3-methlyimidazolium since it has an ethyl group at the 1 position and a methyl group at the 3 position on the imidazolium ring. (Image from Gebbie Lab)
Why Are Ionic Liquids So Unique?
Ionic liquids have such a wide range of properties that can be tuned to fit different applications ranging from biological, to electrochemical, to catalytic research! There are, however, a few characteristics that many, if not all, ILs share. They are what make ILs such an attractive compound.
As stated before, ionic liquids have low melting temperatures. Those that remain liquid at or below room temperature are referred to as “room-temperature ionic liquids”, or RTILs. This property then lends itself well to many synthesis applications. For one, ILs can also dissolve both organic and inorganic compounds. Combined with the fact that they generally have very low vapor pressures, which means that they are very unlikely to evaporate, and are not flammable6, ILs make very enticing candidates for use as green solvents. Traditional organic solvents often give off volatile organic compounds as they are being used, which can be harmful for the environment as well as those who are working near the reactors. Even more, they tend to be expensive to use – in many synthetic processes, some of the reactants or products can remain in the organic solvent and since removal is often difficult, new solvent is needed to replace the old, unusable batch. Ionic liquids can make the separation process easier and less wasteful, but that’s a blog post for another time!
Given the vast design space of ionic liquids as a class of compounds, there are a wide range of uses for ILs. In the Gebbie Group for example, we’re interested in discovering what happens at the interface between an electrode and an electrolyte with a particular interest in its implications for sustainable energy and catalysis. Other groups use ionic liquids in biological systems, synthetic applications, or consumer-oriented products. It’s almost certain that researchers have at least attempted to apply ionic liquids in many of the current research fields.
Applications and Challenges
While the uses of ILs are almost too numerous to list comprehensively, many of the popular applications are in separations, battery electrolyte development, and electrocatalysis. The properties of ILs lend themselves well to use in separating chemicals from one another. For example, ILs are pretty good at solubilizing CO27-9 and since these ionic liquids do not release volatile organic compounds, it’s a very feasible idea to pass syngas through an ionic liquid to remove common impurities like CO2. In the field of battery electrolyte development, lots of effort is being put towards developing electrolytes that prevent dendrite formation and increase device performance10,11. It’s a hot topic among electrochemists especially as energy storage and sustainability are becoming important issues. Similarly, the use of ionic liquids in electrocatalysis has seen growing interest because it could mean improved catalytic efficiency and direct production of some value chemicals that can then be used to make products important to consumers12-15. Even further, topics like CO2 reduction to make carbon-based derivatives can be used to sequester CO2 and reduce the levels of the greenhouse gas in the atmosphere.
As with all ongoing research, however, there are many challenges to using ionic liquids. For one, the underlying mechanisms and explanations for why they behave the way that they do are unclear. Many research groups, like our own, are working to understand why they deviate so much from classical understanding of liquids. For instance, some recent articles have tried to look into why different compounds diffuse through ILs in manners that are not well described using conventional theories.16 Even more, the matter of what happens at interfaces in electrochemical cells that allow for ILs to be good cocatalysts is also not well-understood. On an industrial level, the cost of IL synthesis is currently still too high for implementation in the production of consumer goods and chemicals.
These challenges, however, are part of the reason why ILs are so exciting! By applying various characterization techniques like Raman spectroscopy, infrared spectroscopy, scanning electron microscopy (SEM), electrochemical characterization methods using potentiostats, etc. with computational tools like molecular dynamics (MD) simulations and density functional theory (DFT), we can really delve deep and work systematically to unravel the mysteries behind ionic liquids. And along the way, we can also hopefully share our findings and encourage others to consider their uses and applications.
If you’re interested in learning more, here are some useful (and free) resources to check out:
https://www.electrochem.org/dl/interface/spr/spr07/spr07_p38.pdf – a really nice, but more technical overview of what ILs are
https://www.organic-chemistry.org/topics/ionic-liquids.shtm – Gives some more examples of different kinds of ILs
https://www.ncbi.nlm.nih.gov/pmc/articles/PMC5988633/ – A very in-depth review of ionic liquids, their history, and applications by Dr. Tom Welton. Highly recommended if you want a more comprehensive understanding.
https://cen.acs.org/materials/ionic-liquids/time-ionic-liquids/98/i5 – A really interesting contemporary example of dissolving cellulose for sustainable clothing!
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U.S. National Library of Medicine, pubchem.ncbi.nlm.nih.gov/compound/Sodium-chloride.
2 “Potassium Chloride.” National Center for Biotechnology Information. PubChem Compound Database,
U.S. National Library of Medicine, pubchem.ncbi.nlm.nih.gov/compound/Potassium-
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