In addition to rate, chemical equilibrium between the reactants and products limits many catalytic systems. As shown in Figure , industrially important reactions including ammonia synthesis, water-gas shift chemistry, dry reforming of methane, methanol synthesis, and alkane dehydrogenation all exhibit Gibbs free energy of reaction (and associated equilibrium constants, K i ) at standard conditions limiting high overall conversion . To overcome these inherent limitations, equilibrium between reactants and products can be manipulated via reaction conditions; ammonia synthesis is operated at high pressure , while water-gas-shift chemistry is conducted in staged reactors of varying temperature . Other strategies for overcoming equilibrium limitations combine a second reaction or separation phenomenon to drive a reaction. For example, alkane dehydrogenation is combined with oxidation , ammonia synthesis is combined with absorption , and carbon-carbon coupling chemistries such as cycloaddition can be combined with dehydration . The driving force for chemical reaction is the difference in Gibbs free energy between reactant and product, with a reaction system stabilizing at equilibrium defined as the two states existing at equal Gibbs free energy. The two strategies for manipulating reactions are based on this definition; change the conditions and equilibrium, or deprive the system of one of the two components (reactant or product) to unbalance the free energy distribution (i.e., Le Chatelier's principle) . A third strategy for driving physical systems away from equilibrium is the application of work; added energy can perturb a physical system to a steadystate condition different from equilibrium such as the case of electrocatalytic water splitting or electrocatalytic ammonia synthesis . Alternatively, a catalyst surface with dynamicallyoscillating adsorbate binding energy provides work to a catalytic reaction to raise the energy of adsorbed surface species .