The global energy production is influenced by the growing shortage of traditional fossil fuels and the threat of disturbing the ecological balance due to climate change. Conferences like COP21 and COP26 set goals to reduce greenhouse gases such as CO₂. These COPs proclaimed the leading role of renewable energy in solving climate problems. The wider use of hydrogen was advocated ^[1][2][1,2] through, e.g., hydrogen-enriched natural gas [3], the direct reduction of iron ore [4], the use of hydrogen in the minerals’ processing [5], and the production of CO₂-based synthetic natural gas [6], among others. It is important to assess the ways that hydrogen production can help to meet the renewable energy goals.

Since there are no significant resources of hydrogen on earth, hydrogen is referred to as a secondary energy carrier, made from a primary energy source (nowadays, ~95% from fossil fuels) ^[1][3][1,3]. The global transition to hydrogen as an energy carrier accompanies the rise in required “green” energy.

In 2020, the annual production of hydrocarbons reached 14 billion tons of oil-equivalent [7]. To affect the global CO₂ emissions, at least 10% of hydrocarbons should be replaced by renewable energy resources and/or by hydrogen. If considering hydrogen as a major potential renewable energy carrier, roughly 1 billion tons of H₂ produced from non-fossil sources are required to accomplish this. In 2020, only 87 million tons of hydrogen were produced worldwide [8], mainly used in ammonia production and oil refining. This is more than 10 times below the required quantity to replace 10% of hydrocarbons.

The main method currently considered in producing green hydrogen is electrolysis using electricity from renewable sources (wind turbines, photovoltaics). It theoretically takes about 50 kWh to produce electrolysis H₂ at an overestimated efficiency of 80% [1], without taking other losses into account. The annual global energy production from renewable sources (excluding hydro-energy) was estimated to be 2800 TWh in 2019, which would translate into 56 million tons of H₂ annually, if these sources were solely used to produce H₂. This is lower than the current hydrogen production. In addition, the production cost of the renewable-based H₂ is two to four times higher than when using traditional petrochemical pathways, and it is not deemed practical to expand these renewable technologies to meet the goal of 1 billion tons of H₂ [7].

The global electricity production of hydropower in 2019 was around 38 Exajoules (10⁴ GWh) [9], which allows an annual production of 200 million tons of H₂ if no electricity is sent to the power grid, which is only three times the current production. Since hydropower is already realized for over 20% of its potential, it is therefore unable to solve the problem [9]. The same reason applies to nuclear energy, where the energy production is only 25 Exajoules (even less than hydropower). With the often negative public attitude towards nuclear energy and the limited reserves of Uranium, nuclear power will not sustain a future for hydrogen production.

According to the arguments above, fossil hydrocarbons are deemed to remain a valid source for hydrogen production in the near future. However, the production of hydrogen with Steam Methane Reforming (SMR) results in the emission of 10 kg CO₂/kg H₂, and such hydrogen is classified as “gray hydrogen” according to ecological standards [1]. It is environmentally unattractive and does not contribute to the reduction of CO₂ emissions. An additional process to make this technology “cleaner” is carbon capture and storage (CCS). The major disadvantage of this process is that the total cost rises because the process requires more energy. CCS would not be sufficient to have an environmental impact. With the rise of the hydrogen cost to almost the double, it is not a viable solution [10].

2. Hydrogen Production by Water Splitting in Thermal Redox Systems

An earlier rpapesearchr by the researcheauthors ^[11][12] assessed and classified different thermo-chemical water splitting redox reactions. A redox system ranking was developed by applying multiple assessment criteria, including energy efficiency, H₂ yield, conversion, cyclic operation, production cost, and environmental and safety impacts. Several very high temperature reactions (≥1000 °C) were not taken into consideration, e.g., involving. metal–metal oxides/hydroxides, perovskites, or doped ceria. At such high temperatures, concentrated solar energy would be the potential heat source, but at the expense of having to use extravagant construction materials for the solar receiver-reactor. Selected systems should operate at lower temperatures, preferably below 1000 °C for the reactor wall temperature. Deng et al. ^[11][12] finally selected 4 out of 24 redox reactions that meet the systems’ selection priorities. These systems included Mn₃O₄/MnO/NaMnO₂, MnFe₂O₄, U₃O₈/UO₂CO₃ and ZnO/Fe₃O₄/ZnFe₂O₄ redox reactions. The U₃O₈ system was discarded for nuclear hazard reasons, and the ZnO system scored 30% lower than both other redox reactions and was hence eliminated from a priority investigation. Both the MnFe₂O₄ and Mn₃O₄/MnO/NaMnO₂ redox processes are of a multi-step nature, where water is decomposed into H₂ and O₂ via a medium-temperature two-step (MnFe₂O₄) or a medium-temperature four-step process (Mn₃O₄), respectively, that forms a closed cycle ^[11][12]. This process was previously fostered as having a high potential ^[12][13] and involves four reaction steps, as shown below.

3. MnO_x/Na₂CO₃ Cycle

Xu et al. ^[12][13] comprehensively studied the Mn₃O₄/MnO/Na₂CO₃ system. Theoretical considerations were validated by experiments on a small scale (200 mg) in a thermogravimetric analyzer. The Mn-based reduction and oxidation reactions required an operating temperature in excess of 850 °C. This four-step redox cycle is given by the equations below. $$3{\mathrm{Na}{2}_{}{\mathrm{CO}{3}_{} \left(\mathrm{s}\right)+2{\mathrm{Mn}{3}_{}{\mathrm{O}}_4\left(\mathrm{s}\right)\to 4{\mathrm{NaMnO}{2}_{}\left(\mathrm{s}\right)+2{\mathrm{CO}{2}_{}\left(\mathrm{g}\right)+2\mathrm{MnO}\left(\mathrm{s}\right)+{\mathrm{Na}{2}_{}{\mathrm{CO}{3}_{} \left(\mathrm{s}\right)}_{}}_{}}_{}}_{}}_{}}_{}}_{}$$ $$2\mathrm{MnO}\left(\mathrm{s}\right)+{\mathrm{Na}{2}_{}{\mathrm{CO}{3}_{} \left(\mathrm{s}\right)+{\mathrm{H}}_2\mathrm{O}\left(\mathrm{v}\right)\to {\mathrm{H}}_2\left(\mathrm{g}\right)+{\mathrm{CO}{2}_{}\left(\mathrm{g}\right)+2{\mathrm{NaMnO}{2}_{}\left(\mathrm{s}\right)}_{}}_{}}_{}}_{}$$ $$6{\mathrm{NaMnO}{2}_{}\left(\mathrm{s}\right)+{\mathrm{ayH}{2}_{}\mathrm{O}\left(\mathrm{v}\right)+(3+\mathrm{b}){\mathrm{CO}{2}_{}\left(\mathrm{g}\right)\to 3{\mathrm{Na}{2}_{}{\mathrm{CO}{3}_{} \left(\mathrm{s}\right)+{\mathrm{aH}{\mathrm{x}}_{}{\mathrm{MnO}{2}_{}·{\mathrm{yH}{2}_{}\mathrm{O}\left(\mathrm{s}\right)+{\mathrm{bMnCO}{3}_{} \left(\mathrm{s}\right)+{\mathrm{cMn}{3}_{}{\mathrm{O}}_4\left(\mathrm{s}\right)}_{}}_{}}_{}}_{}}_{}}_{}}_{}}_{}}_{}}_{}$$ $${\mathrm{aH}{\mathrm{x}}_{}{\mathrm{MnO}{2}_{}·{\mathrm{yH}{2}_{}\mathrm{O}\left(\mathrm{s}\right)+{\mathrm{bMnCO}{3}_{} \left(\mathrm{s}\right)\to (2-\mathrm{c}){\mathrm{Mn}{3}_{}{\mathrm{O}}_4\left(\mathrm{s}\right)+{\mathrm{ayH}{2}_{}\mathrm{O}\left(\mathrm{v}\right)+{\mathrm{bCO}{2}_{}\left(\mathrm{g}\right)+0.5{\mathrm{O}}_2\left(\mathrm{g}\right)}_{}}_{}}_{}}_{}}_{}}_{}}_{}$$ The results of Xu et al. ^[12][13] exhibited a more than 90% yield for both the H₂ and O₂ evolution for five successive redox cycles. The key recyclability feature of the Mn-based system was demonstrated to be the complete in/out shuttling of Na⁺ of the manganese oxides, where the reactions to form α-NaMnO₂ are thermodynamically favorable. It was demonstrated that the Na⁺ extraction is enhanced by the mobility of Na⁺ in the layered structure when intercalated by water and is promoted by CO₂, thus driving the reaction equilibrium to a mixture of Mn₃O₄, and MnCO₃. After thermal reduction at 850 °C, Mn₃O₄ is formed, and this closes the thermo-chemical cycle. The cycles consist of multiple stages, operating from 50 °C to 850 °C ^[12][13][13,14]. Dissolved Na₂CO₃ must be recovered from the aqueous solution before the needed re-heating to 850 °C, and it will impose a considerable energy cost. The wide range of the temperatures in the different steps requires adequate heat management and makes the process more complex. The possible loss of volatile Na species at high temperature can be responsible for the progressive loss of process cyclability ^[14][15]. The hydrolysis reaction, however, drives the formation of Na-birnessite. Bayón et al. demonstrated that the surface area of Mn₃O₄ is the main factor, while the crystalline domain size has a secondary influence ^[15][16]. Xu et al. ^[16][17] performed the multistep reactions on 200 mg Mn₃O₄/Na₂CO₃ below 850 °C and demonstrated that the rates of H₂ releasing and Na⁺ extraction depend on the redox properties of metals in ferrite-based oxides and the facility of intercalating alkali cations. Bayón et al. ^[13][14] also studied the feasibility of the MnO/Na₂CO₃ cycle, with a maximum conversion of 47% and a productivity of 20.1 μmol H₂ min⁻¹ g⁻¹. Alonso et al. ^[17][18] studied the Mn₂O₃/Mn₃O₄/MnO cycles at >1835 K using solar-driven thermogravimetry and derived the overall kinetic rate equations.