Initial testing was performed in stagnant electrolyte to test the electrochemical and thermal behavior in the absence of flow. The thermal response of both electrodes to 2 seconds of applied voltage is shown in Fig. 3. As expected, the anode experiences a momentary temperature decrease, and the cathode experiences a concomitant increase. Both temperature deviations are of shorter duration than the applied voltage pulse. While the short duration of cooling could be due to either thermal equilibration in the small gap between electrodes or the decreasing reaction current over time due to concentration polarization of the electrolyte (Fig. S2), concentration polarization is likely the larger contributor in this case because τ sand < τ thermal . As larger driving potentials are applied, the magnitude of the heating effect at the cathode grows faster than the cooling effect at the anode, as expected for less reversible heat pump operation. At very high overpotential, the cooling effect at the anode is overwhelmed by heating due to activation and transport losses even before τ sand , and long before heat can propagate across the inter-electrode space.

Figure 3 Temperature response of the anode (cold electrode) and cathode (hot electrode) upon application of various overpotentials in stagnant electrolyte. For this configuration τ sand ≈ 0.3 s while τ thermal ≈ 1.4 s, so it is likely that the cooling pulse duration was limited more by ion concentration polarization than by thermal equilibration between the electrodes. Full size image

As shown in Fig. 4, electrolyte flow dramatically altered the thermal response of the cell. In these tests, different overpotentials were applied and the temperature response monitored with and without electrolyte flow. The applied current and overpotential were also recorded and were used to compare the estimated to the theoretical cooling power. Both peak and maximum steady-state cooling were achieved at higher overpotential with flowing electrolyte than with stagnant electrolyte. Flowing electrolyte also allowed for some steady-state refrigeration (albeit delivered at lower η than the peak value), which was neither expected nor observed in the cell with stagnant electrolyte. However, the magnitude of the peak cooling was not significantly larger with flow than without.

Figure 4 Temperature response of the cold electrode (bottom) and hot electrode (top) upon application of various overpotentials in both stagnant (left panels) and flowing (right panels) electrolyte. Flowing electrolyte enables continuous cooling operation, but it does not significantly change the measured peak temperature depression. Full size image

The similarity of the peak cooling values for stagnant and flowing electrolyte is attributable to the scaling of both i and U with v1/2 (see derivation in SI). In other words, higher flow rates allow for higher cooling power by advection of reactant towards (and Joule heating away from) the cooling electrode but simultaneously allow better heat removal from the cooling electrode. Therefore, electrolyte flow does not appreciably change the transient ΔT.

Comparison of the estimated and theoretical cooling power density with electrolyte flow illustrates a similar dynamic. Since the electrode is porous and has a three-dimensional architecture, it is most useful to present this as a volumetric power density, where the volume is measured using the nominal area and the thickness of the electrode. As shown in Fig. 5, electrolyte flow allows for favorable partitioning of the temperature rise due to the cell’s irreversibility; less than 50% of the total overpotential applied to the cell results in undesired temperature rise at the cooling electrode. For a typical Butler-Volmer transfer coefficient α b−v ≈ 0.5, the expectation in stagnant electrolyte is that approximately 50% of the applied activation overpotential will manifest as an undesirable temperature increase on the cooling electrode. The anticipated cooling in this case is described by the brown line in Fig. 5. That the measured cooling power was higher than this power indicates the benefit of electrolyte flow in sweeping joule heated electrolyte away from the cooling junction. This effect diminished and the total kinetic losses regressed towards an equal partition at higher currents, as concentration polarization rather than ohmic losses in the electrolyte become dominant20. While the observed partitioning of overpotential losses onto the heating electrode could in theory be due to a kinetic asymmetry in the reaction, (i.e. a Butler-Volmer transfer coefficient α b−v ≠ 0.5), this asymmetry would not continue to scale with higher applied overpotentials and is inconsistent with prior measurements of α b−v for Fe(CN) 6 3−/4− kinetics21.

Figure 5 Estimated volumetric cooling power at different driving currents with 1 mm/s electrolyte flow, compared to theoretically achievable cooling power with no losses (black) and with measured irreversibility evenly split between the two electrodes (brown). The solid lines are calculated based on i and η input to the cell while individual points are based on the thermal signal at the IR microscope. Electrolyte flow from the cold to the hot electrodes allows less than 50% of kinetic losses in the cell to manifest as a temperature rise on the cooling electrode, enabling a more powerful refrigeration effect than would be possible in a cell with stagnant electrolyte. Full size image

The combination of CaF 2 - backed electrodes and infrared microscopy used in this experiment proved its worth for visualizing thermal aspects of electrochemical processes. Thermal microscopy provided much better spatial resolution than was required for these measurements, and might prove useful for future authors investigating localized heat generation in different parts of electrochemical cells, or even for parallelized electrocatalytic screening22. Temperatures were successfully measured on a variety of porous and nonporous materials, including metallic electrodes (Pt and Pd). While high emissivity is desirable for temperature measurements, it is not always easy to achieve. Our best results for low-emissivity electrodes were achieved by spin-coating a thin layer of photoresist as an emitter layer directly on the infrared-transparent substrate prior to electrode deposition. Even better results might be obtainable by adding an emissivity-enhancing dissolved or colloidal species to the electrolyte itself, or by utilizing thinner porous electrodes.

A number of steps could have been taken to achieve a greater cooling effect but were forgone in this work for experimental simplicity and ease of visualizing results. These include tighter electrode spacing for lower total ohmic losses16, a supporting electrolyte for greater ionic conductivity20, high surface-area catalyst loading, flow-through rather than flow-past electrode configuration, a multi-stage design (Fig. S3), and a much higher flow rate to eliminate concentration polarization on the cooling electrode. The relatively sluggish flow rates used in this work reflect the goal of elucidating the effect of flow, rather than leading to the maximum possible cooling power. The flow-past electrode configuration was similarly chosen to facilitate lower-noise IR thermography; extensive work in the flow battery community suggests that a flow-through configuration would in fact better reduce concentration polarization23. Additionally, while thermodynamic calculations indicate that cooling can be achieved up until a product/reactant concentration C oxidized /C reduced = e−αF/R (see derivation in SI), the use of a reference electrode could aid future researchers in localizing overpotential and thus in formulating the ideal ratio of reduced and oxidize species in the refrigerant. We expect that these measures will be taken in future work.

An important aspect of optimizing future electrochemical refrigeration systems will be materials design. While the Fe(CN) 6 3−/4− redox couple was chosen in this work based on its track record of use in thermogalvanic systems, there is a wide parameter space open for electrochemical refrigerants with the right combination of high standard entropies of reduction ΔS, low activation barrier E a for reduction (if α > 0) or oxidation (if α < 0), low specific heat C p , and a high capacity of entropy carriers C (in mol/kg), which equates to high total solubility for dissolved species. These properties are the primary determinants of the achievable cooling effect ΔT real in a system in which heating due to concentration polarization and ohmic losses in the electrolyte are managed by forced convection.

$$\Delta {T}_{real}=\frac{[\Delta S(\frac{J}{mol\cdot K})T(K)-{E}_{a}(\frac{J}{mol})]\times C(\frac{mol}{kg})}{{c}_{p}(\frac{J}{kg\cdot K})}$$ (1)

Based on a Buckingham Pi analysis, we propose the following dimensionless figures of merit for electrochemical refrigerants:

$$Y=\frac{\Delta S(\frac{J}{mol\cdot K})\cdot C(\frac{mol}{kg})}{{c}_{p}(\frac{J}{kg\cdot K})}$$ (2)

and

$${Q}_{g/b}=\frac{\Delta S(\frac{J}{mol\cdot K})T(K)}{{E}_{a}(\frac{J}{mol})}$$ (3)

Y gives the ratio of redox cooling potential to sensible heat energy stored in the refrigerant and is also described in the literature on electrochemical energy harvesting11. Y expresses the thermodynamic reality; YT = ΔT max, the maximum cooling that could be achieved adiabatically by this refrigerant given totally reversible operation and a very long residence time near the electrode surface (or in a staged design, see SI). By contrast, Q g/b expresses the kinetic reality as the ratio between the “good” and “bad” thermal signatures per carrier; only refrigerants with Q g/b > 1 can demonstrate a cooling effect in practice.

Redox reaction entropy is generally attributed to rearrangements of molecular structure and solvation24. Marcus theory, however, correctly predicts that large molecular and solvation rearrangements disfavor electron transfer20. As a result, one might expect a fundamental tradeoff between reversibility and reaction entropy that limits Q g/b . Due to solvent effects in particular, larger dissolved molecules should demonstrate generally lower E a and ΔS whereas smaller molecules should demonstrate higher ΔS and E a . In this respect, a “happy medium” for Q g/b might be hard to find. To date, the best Q g/b values have been found in coordinated metal redox couples, which to some extent avoid this kinetic tradeoff with entropies of reduction that are based more on solvent rearrangement than on reorganization of the redox centers themselves24.

Table 1 illustrates a different tradeoff. Pure substances such as water require no solvent by definition and so have high Y ratios and correspondingly high ΔT max relative the species’ thermopower. However, redox of these small molecules requires inner-sphere electron transfer, which tends to proceed slowly20, leading to a poor ratio Q g/b . By contrast, many coordinated metal species are only moderately soluble, and the C p contribution of the excess solvent is reflected in low ΔT max. However, these species undergo comparatively fast single-electron transfer and frequently have high ΔS leading to high Q g/b . An important step in the future will be to identify the trick that allows an electrochemical refrigerant to circumvent this apparent tradeoff.

Table 1 Properties of potential redox refrigerants. Full size table

The redox flow battery community has spent decades screening electroactive species for solubility, stability and reversibility23, which can now be leveraged to identify electrochemical refrigerants. Further inspiration can be found in the strikingly similar set of compounds found in biological vascular systems. Hemovanadin (vanadium); hemoerythrin, chlorocruorin, and hemoglobin (iron); and hemocyanins (copper) all point to the enormously tunable redox properties of an appropriately coordinated metal center25. Laboratory results suggest that the standard entropy of reduction is perhaps as tunable as the redox potential24. The inexhaustive list of Table 1 contains only a few potential liquid-phase refrigerants and completely omits other electrochemical transformations (solid dissolution, hydriding, intercalation, solution-precipitation reactions, redox of slurries, use of non-aqueous media etc.) that may be of great interest in future work.

Interestingly, due to the difference in reduction entropies of the V4+/5+ and V2+/3+ redox couples, a vanadium redox flow battery can already be considered to be producing a beneficial electrochemical cooling effect during charging. However, due to the high rates at which these batteries are charged, it is unlikely that a net cooling effect is ever observed23. A more likely scenario is that the battery is discharged during cool evening hours and recharged during the heat of the day. In this case, the entropy changes align to slightly increase the equilibrium voltage during discharging and to slightly decrease it during recharging. This energy harvesting effect likely results in a small but measurable boost in the batteries’ cycle efficiency. Other mooted grid-scale energy storage systems such as Zn-air, Fe-air, and regenerative hydrogen fuel cells would also benefit from this effect.