## 2 Fundamentals of Heterogeneous Catalysis  
Chemical reactions underpin societal development, with advances in chemical technologies yielding essential products, such as pharmaceuticals, vaccines, fertilizers, fuels, food additives, and so on [27]. The efficiency gains driving modern scientific and technological progress are particularly linked to catalytic processes. Catalysts are defined as substances that accelerate chemical reactions without changing the standard Gibbs free energy (\(\Delta\text{G}_{\text{r}}^{\circ}\)). As we know, in a chemical reaction, the activation energy is the energy barrier that reactants must overcome to form products, and it is reduced by catalysts. This reduction allows a greater proportion of reactant molecules to attain the transition state at a given temperature, thereby increasing the reaction rate [28]. Notably, catalysis contributes to over 35% of the global GDP, and it is a key component in approximately 80% of all manufactured goods [29], owing to the fact that catalysts accelerate reactions, reducing energy consumption, and increasing the yield of desired products, thereby enabling the industrial-scale production of critical materials [30; 31; 32]. Importantly, while catalysts speed up the attainment of equilibrium, they do not alter the position of the equilibrium itself. Thus, the \(\Delta\text{G}_{\text{r}}\) is a thermodynamic parameter that indicates the difference in free energy between reactants and products under standard conditions, thereby determining reaction spontaneity (\(\Delta\text{G}^{\circ}<0\)) or non-spontaneity (\(\Delta\text{G}^{\circ}>0\)). Given that catalysts do not alter the intrinsic energy states of the reactants or products, the overall \(\Delta\)G\({}^{\circ}\) remains constant [33]. Consequently, while catalysts enhance the reaction rate by lowering the activation energy, they do not affect the thermodynamic favorability of the reaction [34]. Therefore, catalysis represents a highly relevant and multidisciplinary field, finds applications in chemical synthesis, materials engineering, surface science, and physical chemistry. Its versatility enables the optimization of chemical reactions in numerous contexts, leading to innovative and efficient solutions. Catalysts can be classified as homogeneous and heterogeneous, where, in the first type of catalysis, the reactants and the catalyst are in the same phase; in its turn for heterogeneous catalysis, the reactants and the catalyst must be in different phases [35]. For instance, in biological systems, enzymes function as natural homogeneous catalysts that facilitate vital biochemical reactions necessary for life, operating under the same principle of lowering activation energy without altering \(\Delta\)G\({}^{\circ}\)[33]. Therefore, catalysts are important because they allow reactions to have an alternative mechanism, faster than the one that occurs in their absence. Although the catalyst participates in the mechanism reaction, at the end of the reaction cycle, it can be regenerated and used again in the same or some other chemical process [36; 37]. In principle, the catalyst would not be consumed at the end of a catalytic cycle; however, undesirable parallel reactions, the presence of 'catalytic poisons', or bad operating conditions can lead to loss of catalytic activity [38]. In this review article, we will emphasize heterogeneous catalysts. Heterogeneous catalysis is pivotal in many industrial processes, producing various chemicals, fuels, and materials, while enhancing reaction rates and selectivity [39; 40]. In this way, there has been a search for more specific, reusable catalysts that generate high-quality products without the need for so much purification; hence, to meet these requirements, the use of heterogeneous catalysts is pursued. Historically, the first heterogeneous catalytic reaction was studied by [PERSON] in 1778 when attempting to dehydrate ethanol by active clays. Later, in 1796, [PERSON] was the first to use metallic catalysts for the dehydrogenation of ethanol.