Environmental Hydrogen Concentration as a Novel Factor Determining Changes in Redox Potential

Abstract

Intracellular oxidation–reduction (redox) potential is a key factor regulating various physiological phenomena in the cell. Monitoring this potential change is therefore important for understanding physiological homeostasis in cells. Herein, we propose a new approach for the real-time, non-invasive estimation of the redox potential impacting biological metabolism and reactive oxygen species generation. Enzymes, specifically oxidoreductases, play a crucial role in catalyzing redox reactions by facilitating the transfer of electrons and hydrogen atoms between molecules. The redox potential of substrates, such as nicotinamide adenine dinucleotide, is determined by the ratio of its oxidized and reduced forms, while that of enzymes, such as succinate dehydrogenase, is determined using the reference electrode in protein-film voltammetry. Although the standard hydrogen electrode potential is defined as zero under standard conditions, the electrode potential of a reversible hydrogen electrode changes according to the ratio of the hydrogen ions (H⁺) and hydrogen gas (H₂) in the biological fluids, as a reference electrode. The pH is maintained at 7.4 ± 0.1 in the arterial blood and the H₂ that produced by the gut microbiota is measured in the endo-tidal breath for clinical diagnosis. The H₂ in the endo-tidal breath equilibrates arterial blood during gas exchange in the lungs, as well as in whole-body tissues, due to the systemic circulation. In this study, H₂ can be measured in the environmental gas compared to the atmosphere, and may serve as a novel factor for redox potential changes in redox enzymes, impacting biological metabolism and reactive oxygen species generation.

1. Introduction

The energetics of electron flow is determined by the reduction-oxidation (redox) potentials of organic and inorganic cofactors as determined by the protein environment [1]. For example, in aerobic respiration, oxygen (O₂) is the final electron acceptor and reduced to water (H₂O). In anaerobic respiration, the hydrogen ion (H⁺) is the final electron acceptor, and hydrogen (H₂) is produced by several members of the gastrointestinal tract microbiota [2]. Oxidoreductases can couple the oxidation and reduction of at least two substrates. The activity of those enzymes depends on both the substrate concentrations and electrochemical potential of the enzymes due to catalysis and electron transfer [3]. Since cofactors of the active center (e.g., iron, heme, and copper) can be covered by proteins, the influence of other redox substrates is limited, except the substance and product. In protein film voltammetry, some enzymes can be wired to electrodes and function in either direction of the reaction depending on the electrochemical driving force (Figure 1) [4].

The working electrode, with a film of redox enzyme, is placed in a solution-filled electrochemical cell. The addition of a reference and counter electrode yields the three-electrode set-up utilized in standard electrochemical measurements. Potentials are applied between the working and reference electrodes, and current flows through the counter electrode. The gas flow (e.g., O₂ and CO₂ concentration) is controlled.

Thermodynamics predicts only the value of the open circuit potential (OCP), where i = 0, such that the forward and reverse reactions proceed at exactly the same rate (Figure 2) [5].

In reversible redox catalysis, the ideal potential-current response displays a single sigmoidal wave that crosses the zero-current axis at the open circuit potential (OCP), marked by an open circle, and reaches a potential-independent limiting current at a relatively high overpotential. When the electrode potential is above the OCP, the forward reaction (i.e., oxidation) occurs, with the reaction rate (current) proportional to the electrode potential. In contrast, when the electrode potential is below the OCP, the reverse reaction (the reduction) proceeds, and the current is similarly proportional to the electrode potential. Finally, in electro-catalysis, as in simple electrochemistry, a potential-independent process eventually becomes rate-limiting at a high overpotential.

This potential equates the formal reduction potential of the substrate/product combination, which can be calculated for any concentration ratio using both the Nernst equation and published value [6]. Even in intracellular fluids comprising various redox substances, the redox potential of the active center of the redox enzyme exhibits a constant potential against the standard hydrogen electrode (SHE) because the substrate specificity of the redox enzyme prevents reactions with other solutes, and the substrate and product are in equilibrium. In the case of succinate dehydrogenase (SDH) in complex II of the respiratory chain, the potential-current wave exhibits a sigmoidal increase in succinate oxidation activity upon increasing the electrode potential, but the wave for fumarate reduction shows a decreased wave upon reducing the electrode potential [7]. Thus, the direction and rate of the catalytic reaction of the redox enzyme depend on alterations in the electrode potential of the active center. The redox potential of SDH exhibits weak dependence on pH (10 mV per pH unit), while the redox potential of the substrate/product couple decreases 60 mV per pH unit [8]. Nonetheless, there is no indicator of the alteration in the potential of the redox enzyme in vivo and it is unclear if the electrochemical potential of the redox enzyme has any physiological relevance.

2. Redox Potential of Water as the Environment of the Redox Enzyme

The environment of living organisms and cultured cells is neutral and aerobic. Water is in equilibrium with gaseous oxygen (O₂) and/or hydrogen (H₂), as well as hydrogen ions (H⁺) from oxygen reduction and hydrogen ion reduction [9].

4H⁺(aq) + O₂(g) + 4e⁻ ⇄ 2H₂O

(1)

2H⁺(aq) + 2e⁻ ⇄ H₂(g)

(2)

O₂ and H₂ evolution reactions occur when the electrode potentials are high and low, respectively. The range of electrochemical potential is called the electrochemical window. The Nernst equation links the equilibrium potential of an electrode, E, to its standard potential, E₀, and the concentrations or pressures of the reacting components at a given temperature, T.

$$\mathrm{E}={\mathrm{E}}_0-{\displaystyle \frac{\mathrm{R}\mathrm{T}}{\mathrm{n}\mathrm{F}}}\mathrm{l}\mathrm{n}({\displaystyle \frac{\left[reduced\right]}{\left[oxidaized\right]}})$$

(3)

where n is the number of electrons transferred, F is Faraday’s constant and R is the universal gas constant. As O₂ and/or H⁺ increase, their electrochemical potential increases. As H₂ increases, its electrochemical potential decreases. The redox potential of water is determined according to the Nernst equation: E = +0.815 V for 1 bar-O₂ and E = −0.414 V for 1 bar-H₂ at pH 7 and 25 °C. Thus, the gaseous O₂ and H₂ partial pressures above the equilibrium solution determine the electrochemical potential of the solution according to the Nernst equation and is employed to measure the partial pressure of dissolved gases utilizing a Clark electrode with a permeable membrane [10]. At 37 °C, the change in the O₂/H₂O redox potential combined with the dissolved O₂ tension (DOT) and H⁺ activity is determined by the Nernst equation in pure water [11] as follows:

$$∆\mathrm{E}=+15.4\times ∆\mathrm{l}\mathrm{o}\mathrm{g}\left(\mathrm{D}\mathrm{O}\mathrm{T}\right)-61.5\times ∆\mathrm{p}\mathrm{H}(\mathrm{m}\mathrm{V}).$$

(4)

The redox potential of cell-free media is a proportional logarithm of the DOT, with slopes of +21.7 and +34.7 mV/Δlog(DOT), which is greater than 15.4 mV, as predicted by the Nernst equation for pure water. The redox potential of hybridoma cells during exponential growth and at confluence was reported to be in the range of −130 to +70 mV for reduced and oxidized conditions at a constant DOT ranging from 3 to 300% [12]. Different O₂ concentrations not only influence the absolute redox values, but also involve oxidation of proteins and/or the generation of reducing agents, such as thiols. O₂ is essential for aerobic metabolism, and arterial blood O₂ tension is clinically measured to diagnose respiratory and/or circulatory function.

A cell is not at equilibrium, and there are weak couplings between various redox pairs, resulting in the establishment of different redox potentials for coexisting redox pairs in a cell [13]. One of the most common uses of electrochemistry is to measure hydrogen ion concentration of a solution. However, pH is more related to hydrogen ion activity than hydrogen ion concentration.

$$\mathrm{p}\mathrm{H}=-\mathrm{l}\mathrm{o}\mathrm{g}\left(a_{H^{+}}\right)$$

(5)

$${\mathrm{a}}_{{\mathrm{H}}^{+}}=\mathrm{f}\times \left[{\mathrm{H}}^{+}\right],$$

(6)

where ${\mathrm{a}}_{{\mathrm{H}}^{+}}$ is activity of hydrogen ions, f is activity coefficient, and [H⁺] is hydrogen ion concentration. However, since it is defined in terms of a quantity that cannot be measured by a thermodynamically valid method, Equation (4) can be only a notional definition of pH [14]. The U.S. National Institute of Standards and Technology has defined pH values in terms of the electromotive force existing between certain standard electrodes in specified solutions where the sample (X) and the standard pH solution (S) as follows;

$$\mathrm{p}\mathrm{H}\left(\mathrm{X}\right)=\mathrm{p}\mathrm{H}\left(\mathrm{S}\right)+{\displaystyle \frac{{\mathrm{E}}_{\mathrm{X}}-{\mathrm{E}}_{\mathrm{S}}}{\mathrm{z}}},$$

(7)

which depends upon a half-cell reaction of the hydrogen electrode (Equation (2)) according to the Nernst equation. The Nernst slope is calculated as follows;

$$\mathrm{z}={\displaystyle \frac{\mathrm{R}\mathrm{T}}{\mathrm{F}}}\times \mathrm{l}\mathrm{n}10={\displaystyle \frac{2.303\times \mathrm{R}\mathrm{T}}{\mathrm{F}}},$$

(8)

where R is the gas constant (R = 8.314 J K⁻¹ mol⁻¹), T is the temperature (K), and F is the Faraday constant (F = 9.6485 × 10⁴ C mol⁻¹). Glass electrodes sensitive to H⁺ are the most commonly employed sensors in chemistry and related disciplines. The formation of an electric potential difference on a thin glass membrane in contact with solution was observed similarly to the hydrogen electrode by Cremer in 1906 and used for the construction of a device that measures the acidity of solutions by Haber and Klemensiewicz in1909 [15]. The glass electrode is in fact a half cell with a membrane forming contact with the other half-cell: Ag|AgCl; KCl|glass membrane. In order to maintain homeostasis, the human body utilizes several physiological adaptations, one of which is maintaining the acid-base balance [16]. H⁺ (40 nM at pH 7.4) is equilibrated with phosphate (2 mM), bicarbonate (25 mM), and hemoglobin (15 g/dL) in the blood. Dissolved O₂ homeostasis is tightly regulated and relies mainly on hemoglobin. While a small fraction of O₂ (about 2%) is dissolved in plasma, the majority (approximately 98%) is carried by hemoglobin in red blood cells [17]. Therefore, it is not possible to use the pH nor dissolved O₂ as an index of electrochemical potential in the human body due to those buffering systems.

H₂ is an inert gas in mammalian cells due to the lack of hydrogenase enzyme. H₂ does not readily dissociate into H+ in pure water. However, it can react with water in the presence of a catalyst, such as platinum [18]. Although the pH is proportional to the redox potential, equilibrium behavior for the H⁺/H₂ couple has been verified at gold electrodes from H₂ partial pressures of 1 to about 6 × 10⁻⁴ atm [19]. Primary pH standard values can be deduced from electrochemical data from the cell without transference using the hydrogen gas electrode, known as the Harned cell [14]. Thus, environmental H₂ partial pressure appears to be a redox potential-specific and sensitive factor of the hydrogen electrode in the solution where pH is buffered. It is hypothesized that the environmental H₂ concentration is a novel factor determining changes in the redox potential of the water solution.

3. Measurement of the Reversible Hydrogen Electrode Potential as a Reference Electrode

The SHE serves as the fundamental reference element in electrochemical devices [20]. It consists of a platinum electrode and an acidic solution with the unit activity of protons (H⁺) (pH 0), through which H₂(g) supplied at a fugacity of 1.00 bar (f_H2 = 1.00 bar) passes, ideally in the form of small bubbles. However, if at least one of these parameters is not unitary, the SHE becomes a reversible hydrogen electrode (RHE) and its potential deviates from that of the SHE. The hydrogen electrode is based on the redox half-cell reaction (Equation (2)). The electrode potential was calculated from the activity of H⁺ and the fugacity of H₂ using the Nernst equation, as follows:

$$\begin{gathered} {\mathrm{E}}_{\mathrm{R}\mathrm{H}\mathrm{E}}={\mathrm{E}}_0-{\displaystyle \frac{\mathrm{R}\mathrm{T}}{2\mathrm{F}}}\mathrm{l}\mathrm{n}{\displaystyle \frac{{\mathrm{f}}_{{\mathrm{H}}_2}}{{\left({\mathrm{a}}_{{\mathrm{H}}^{+}}\right)}^2}}=0-{\displaystyle \frac{2.303\mathrm{R}\mathrm{T}}{2\mathrm{F}}}\times \log {\displaystyle \frac{{\mathrm{f}}_{{\mathrm{H}}_2}}{{\left({\mathrm{a}}_{{\mathrm{H}}^{+}}\right)}^2}} \\ =0-{\displaystyle \frac{2.303\mathrm{R}\mathrm{T}}{2\mathrm{F}}}\times \left[log\left(f_{H_2}\right)-2\times log\left(a_{H^{+}}\right)\right](\mathrm{V}), \end{gathered}$$

(9)

where $a_{H^{+}}$ is the activity of H⁺, and $f_{H_2}$ is the fugacity of H₂. The value of 2.303RT/2F, the Nernst slope, was calculated at 25 °C (T = 297.15 K) as follows:

$${\displaystyle \frac{2.303\mathrm{R}\mathrm{T}}{2\mathrm{F}}}={\displaystyle \frac{2.303\times 8.314\times 297.15}{2\times 9.6485\times {10}^4}}={\displaystyle \frac{0.0590}{2}}$$

(10)

In a liquid mixture, the fugacity of each component is equal to that of the vapor component in equilibrium with the liquid. H₂ is a trace gas that comprises only a minuscule portion of the atmosphere and has a mixing ratio of around 0.530 ± 0.006 ppm with seasonal change [21]. The surface layers of the world’s oceans are generally supersaturated with H₂, typically 2- to 5-fold (up to 15-fold) relative to the atmosphere, where H₂ is mainly generated by nitrogen fixation by cyanobacteria [22]. Mammalian cells lack functional hydrogenase genes, but H₂ is fermented by the gut microbiome and a part of H₂ is exhaled from the lungs by the systemic circulation [23]. The organs and tissues are saturated by H₂ in the alveolar air [24]. The mean breath H₂ concentration in healthy Japanese subjects was 13 ppm, which was 20-fold relative to the atmosphere [25]. The H⁺ concentration at pH 7.4 is 40 nM, whereas the dissolved H₂ concentration in the solution equilibrated with 13 ppm-H₂ is calculated as 10 nM using Henry’s law constants [Hcp = 7.8 × 10⁻⁴ (mol/(L∙bar))].

Application of the Nernst equation allows for explanation of the influence of the activity of H⁺(aq) and the fugacity of H₂(g) on the potential of the hydrogen electrode. This is achieved in two stages. In the first stage, the fugacity of H₂(g) is assumed to be a unit, while the activity of H⁺(aq) is varied between 10⁻¹⁴ and 1.00 M. The potential of RHE is calculated from Equations (2) and (8) as

$${\mathrm{E}}_{\mathrm{R}\mathrm{H}\mathrm{E}}=0-0.0590\times \mathrm{p}\mathrm{H}(\mathrm{V})\mathrm{a}\mathrm{t}25°\mathrm{C}.$$

(11)

A negative logarithm function is used to convert very small numbers into more manageable values. The concentration of H⁺ in any solution ranges from 1 mol to 10⁻¹⁴ mol per liter of solution and pH ranges between 0.00 and 14.00. The pH is defined by the electrode potential difference as compared to the standard pH solution for the quantitative measure of the acidity or alkalinity of an aqueous solution [17]. In the second stage, the activity of H⁺(aq) is a unit, while the fugacity of H₂(g) is varied between 0.53 × 10⁻⁶ (the atmosphere) and 1.00 bar. The fugacity of H₂ at low pressure is equal to the ratio of the partial pressure of H₂ (P_H2) over the standard pressure (P_H2/P⁰, P⁰ = 1.00 bar) and is equal to its concentration in the atmosphere above sea level (${\mathrm{C}}_{{\mathrm{H}}_2}$). The potential of RHE is calculated as

$${\mathrm{E}}_{\mathrm{R}\mathrm{H}\mathrm{E}}=0+{\displaystyle \frac{0.0590}{2}}\times {\mathrm{p}\mathrm{H}}_2(\mathrm{V})\mathrm{a}\mathrm{t}25°\mathrm{C}.$$

(12)

where pH₂ is defined as

$$\mathrm{p}{\mathrm{H}}_2=-\mathit{log}\left(f_{H_2}\right)=-\mathrm{l}\mathrm{o}\mathrm{g}\left({\displaystyle \frac{P_{H2}}{P^0}}\right)=-\mathit{log}\left(C_{H_2}\right),$$

(13)

and pH₂ varies between 0.00 and 6.28. The potential of the RHE is calculated using the pH of the solution, as well as the pH₂ measured with an H₂ gas sensor at a gaseous H₂ concentration in equilibrium with the solution. If both the fugacity of H₂ and the activity of H⁺ are not a unit, The potential of RHE is calculated from Equations (11) and (12) as

$${\mathrm{E}}_{\mathrm{R}\mathrm{H}\mathrm{E}}={\displaystyle \frac{0.0590}{2}}\times {\mathrm{p}\mathrm{H}}_2-0.0590\times \mathrm{p}\mathrm{H}(\mathrm{V})\mathrm{a}\mathrm{t}25°\mathrm{C}.$$

(14)

Instead of an SHE, which requires an acidic solution to be continuously supplied with H₂ gas as a reference electrode potential, the RHE is a half-cell with a well-defined and highly reproducible electrode potential [26].

4. Alterations in the Electrode Potential of the Active Site of Oxidoreductases

Prominent members of redox centers include cytochromes and cupredoxins with hemes and blue-copper centers, respectively. Together, these centers cover the entire range of reduction potentials, the electrochemical window, in biology [27]. The redox potential of the solution represents its mixed potential. The redox environment of a linked set of redox couples, such as those found in biological fluids, organelles, cells, or tissues, is determined by the sum of the products of the reduction potentials and reducing capacities of the linked redox couples that are present [28]. Historically, several approaches have been developed to estimate the actual cellular reduction potentials for the oxidized/reduced nicotinamide adenine dinucleotide and glutathione/glutathione disulfide ratios as solutes, as well as to determine whether a given substance undergoes reduction or oxidation [29,30]. The intracellular fluid contains various redox substances, as well as dissolved O₂, H₂ and H⁺ (pH) (Figure 3).

Redox enzymes are large and electronically insulating through most of their volume, which stabilizes the redox potentials of their active centers against the standard hydrogen electrode, even if the ratio of substrate/product is not constant. The solution contains many redox pairs with different redox potentials, e.g., nicotinamide adenine dinucleotide, glutathione, and thiols, as well as oxygen and hydrogen gases. The redox potential of the solution represents its mixed potential.

The electrode potential can be regarded as the potential difference between a point in the solid conductor (metal) and a point in the electrolytic solution [31]. In practical measurements, the electrode potential is measured as a relative value by introducing an additional electrode (i.e., a reference electrode). For example, dissolved O₂ is measured using a Clark electrode with a permeable membrane [32] and pH (H⁺) is measured using a glass electrode. Each oxidoreductase has a specific intrinsic redox potential (E°) of the active site relative to the SHE [33]. The potential of the redox center (E°′) is stable due to both the substrate specificity of the enzyme and the substrate/product equilibrium state. It is calculated as the potential difference relative to the redox potential of water (E water), rather than based on the ratio of the oxidized/reduced redox couple. Every redox center has different electrochemical potential in the solution as compared to water as a solvent (the mixed potential) (Figure 4).

$$\mathrm{E}°\prime =\mathrm{E}°-{\mathrm{E}}_{\mathrm{w}\mathrm{a}\mathrm{t}\mathrm{e}\mathrm{r}}$$

(15)

The electrode potential is the potential difference between a point in the solid conductor (the redox enzyme) and another point in the electrolytic solution (the Galvani potential). In practical measurements, the electrode potential is measured as a relative value by introducing an additional electrode (reference electrode). For example, it is used to measure the potential of a water containing dissolved oxygen, hydrogen, and hydrogen ions (pH), as well as to determine whether reduction or oxidation of a substance occurs in a potential-pH diagram. The electrode potential of a redox enzyme is determined by the redox potential difference between the redox enzyme’s active center and water.

The redox potential of water is determined by the relative concentrations of O₂ and H₂. When the O₂ tension decreases due to O₂ consumption and the environmental fluid becomes reduced, the electrode potential of the oxidoreductase increases, whereas when the O₂ tension increases due to exogenous O₂ supply and the environmental fluid becomes oxidized, the electrode potential of the oxidoreductase decreases. Reducing conditions resulted in the highest viable cells, monoclonal antibody concentrations, and thiol production in hybridoma cells [12]. Oxidant conditions resulted in a detrimental effect in all culture parameters, increasing the specific glucose, glutamine, and O₂ consumption rates, as well as inducing the apoptotic process.

The standard potential of any redox enzyme should be reported with respect to the SHE. The conversion of potential values between a RHE and the SHE can be easily accomplished using the below equation:

$$\mathrm{E}°=\mathrm{E}°\prime +{\mathrm{E}}_{\mathrm{R}\mathrm{H}\mathrm{E}}$$

(16)

where E° is measured versus the SHE; E°′ is the potential measured versus the RHE; and E_RHE is the potential of the RHE versus the SHE (Figure 5). When measurements are made at other H₂ pressures, the sum of the electrode potential of enzyme (E°′) and reversible hydrogen electrode potential (E_RHE) is equal to the sum of the electrode potential of enzyme (E°″) and reversible hydrogen electrode potential (E′_RHE).

$$\mathrm{E}°=\mathrm{E}°\prime +{\mathrm{E}}_{\mathrm{R}\mathrm{H}\mathrm{E}}=\mathrm{E}°″+{\mathrm{E}\prime }_{\mathrm{R}\mathrm{H}\mathrm{E}}$$

(17)

Changes in the electrode potential of the redox enzyme is generated by changes in the reversible hydrogen electrode potential as follows:

$$∆\mathrm{E}=\mathrm{E}°″-\mathrm{E}°\prime =-\left({{\mathrm{E}}^{\prime }}_{\mathrm{R}\mathrm{H}\mathrm{E}}-{\mathrm{E}}_{\mathrm{R}\mathrm{H}\mathrm{E}}\right)$$

(18)

Reversible hydrogen electrodes, which depend on the solution pH and hydrogen gas pressure, are widely used experimentally as a reference electrode. When the reversible hydrogen electrode potential changes with pH and hydrogen pressure, the electrode potential difference (ΔE) is determined by these changes in the reversible hydrogen electrode potential (ΔE_RHE).

When environmental H₂ tension increases and the environmental fluid becomes reduced in spite of pH buffer systems, the electrode potential of the redox enzyme will be considered. An increase in the H₂ partial pressure (environmental concentration) was shown to reduce the redox potential of a phosphate buffer solution according to the Nernst equation [25].

$$\mathrm{E}°\prime =\mathrm{E}°-{\mathrm{E}}_{\mathrm{R}\mathrm{H}\mathrm{E}}$$

(19)

The electrode potential of the redox enzyme will increase due to an increase in the H₂ partial pressure. Changes in the electrode potential can be calculated wherever pH is buffered as follows:

$$\begin{gathered} ∆\mathrm{E}=-∆{\mathrm{E}}_{\mathrm{R}\mathrm{H}\mathrm{E}}=-{\displaystyle \frac{2.303\mathrm{R}\mathrm{T}}{2\mathrm{F}}}\left({\mathrm{p}\mathrm{H}}_2\prime -{\mathrm{p}\mathrm{H}}_2\right) \\ =-{\displaystyle \frac{0.0590}{2}}\times \left({\mathrm{p}\mathrm{H}}_2\prime -{\mathrm{p}\mathrm{H}}_2\right)(\mathrm{V})\mathrm{a}\mathrm{t}25°\mathrm{C}. \end{gathered}$$

(20)

Therefore, under aerobic conditions, the lower limit of H₂ concentration is the atmospheric H₂ concentration, and an increase in H₂/a decrease in pH₂ leads to an increase in the electrode potential of the redox enzymes as follows at 37 °C (310.15 K) (Figure 6).

$$∆\mathrm{E}=-{\displaystyle \frac{0.0615}{2}}\times \left({\mathrm{p}\mathrm{H}}_2-6.28\right)(\mathrm{V})\mathrm{a}\mathrm{t}37°\mathrm{C}.$$

(21)

The H₂ concentration in the exhaled/inhaled breath ranges from 0.53 ppm to 4%. The pH₂ (negative logarithm of H₂ concentration) ranges from 6.28 to 1.4. The electrode potential difference against the atmosphere is proportional to the pH₂.

5. Environmental O₂ and/or H₂ Levels Change the Direction of Redox Reactions and Reactive Oxygen Species (ROS) Generation

Most (~98%) of the oxygen in blood is bound to hemoglobin in red blood cells (RBCs), while dissolved O₂ in the plasma and RBC water is about 2% of the total [17]. The solubility of O₂ is minimal (0.3 vol% at a O₂ tension of 100 mmHg), and O₂ is consumed at complex IV of the respiratory chain. In an animal model, 100% O₂ administration doubled the subcutaneous and large intestine intramural O₂ tension and increased the O₂ tension to 40% in the small intestine [34]. During hyperbaric O₂ conditions the magnetic resonance imaging (MRI)-derived extracellular O₂ tension in the brain was markedly lower than the arterial pO_2, but slightly higher than the cerebral venous pO₂ [35]. Cerebral venous blood is still not fully saturated (below 100 mmHg) up to 3.5 atmosphere absolute (ATA) of hyperbaric oxygen (2660 mmHg) [36]. This is because dissolved O₂ in the blood plasma is rapidly consumed, and O₂ is released from hemoglobin in RBCs, which results in a venous O₂ partial pressure below 100 mmHg. Hyperbaric oxygen (4.8 ATA) rapidly produces intracellular bioenergetic dysfunction and induces oxygen toxicity in human pulmonary cells [37].

Exogenous hypoxic conditions (0.2% O₂) in an in vitro study reduced mitochondrial O₂ concentration and reactive oxygen species (ROS) levels, whereas exogenous hyperoxic conditions (40% O₂, 300 mmHg) increased mitochondrial O₂ concentration and ROS levels [38]. Ischemia–reperfusion injury is tissue damage caused by ROS when the blood supply returns to tissues after ischemia or O₂ deprivation. The mean mitochondrial O₂ tension in rat liver was demonstrated to drop from approximately 30–40 mmHg to approximately 10 mmHg during ischemia [39]. The 15 min reperfusion time led to a complete alteration in the mitochondrial O₂ distribution, which flattened overall as compared to the baseline situation. The mitochondrial O₂ tension histograms revealed that both low- and high-O₂ regions were more prominently present, with some development toward higher O₂ during reperfusion. Above-normal O₂ levels in the blood can increase the production of ROS, potentially contributing to reperfusion injury [40]. Such a negative effect is supported by the Australian Air Versus Oxygen in Myocardial Infarction (AVOID) trial, which reported larger infarct sizes in patients with ST-segment elevation myocardial infarction (STEMI) who received O₂ compared to those who did not receive O₂ [41]. ROS are primarily produced by complexes I and III of the respiratory chain via reverse electron transport [42]. Since the direction of electron flow depends upon the electric potential gradient, relatively high mitochondrial O₂ tension would affect the redox potential of the active center in both complexes I and III. Both reducing and oxidant conditions were maintained in 1 l bioreactors through automatic control of the inlet gas composition at constant dissolved O₂ tension [12]. As culture redox potential increased, a linear decrease in the maximum concentration of viable cells, monoclonal antibodies, and thiols, as well as an increase in the various nutrient consumption rates was observed. Such a behavior indicates that a more oxidant environment is detrimental to culture performance. Instead of inlet O₂ tension, culture redox potential can also be controlled by manipulating other independent variables, such as pH or by the addition of different antioxidants [43].

For a long time, H₂ was thought to be a “biologically inert gas” which could not react with biomolecules under normal pressure. The biological effect of H₂ is much less understood compared to that of O₂. Interestingly, it has been reported that inhalation of H₂ gas reduced oxidative stress, the volume of cerebral infarction, and the size of myocardial infarction in a rat model [44,45]. When H₂ was added to mitochondria isolated from A549 cells, electron transport by succinate was reversed to forward electron transport, and ROS generated primarily from respiration was suppressed [46]. In complex I, environmental H₂ influence the direction of electron transport and ROS generation. H₂ decreases mitochondrial membrane potential by 11%. ROS generation in complex III is suppressed by mitochondrial membrane potential depletion [47]. Furthermore, in order to prevent ischemia–reperfusion injury, inhalation of 2–4% H₂ is required prior to injury. Since the H₂ concentration of the atmosphere is below one-millionth, 2–4% H₂ inhalation would compensate with fluctuations of the redox potential due to changes in mitochondrial O₂ tension during ischemia–reperfusion from 0–10 mmHg to 30–40 mmHg-O₂.

6. Fluctuations in Environmental H₂ Concentration and Metabolism

In the healthy human, H₂ is produced by the gut microbiota and exhaled according to the H₂ partial pressure gradient. In young women, breath H₂ was reported to exhibit an obvious circadian rhythm, increasing in the morning, decreasing to a minimum by 16:00, and then increasing again during the night [48]. Breath H₂ excretion depends not only on feeding and fasting patterns, but also on host organ function [49]. Consumption of a low-carbohydrate diet for 24 h prior to measurement reduces breath H₂ excretion [50]. Resistant starch and dietary fiber have been shown to increase breath H₂ excretion [51,52]. Higher breath H₂ concentrations during fasting have been reported in patients with exocrine pancreatic insufficiency and small intestinal bacterial overgrowth (SIBO) [53,54]. Body weight, body mass index, and visceral fat in patients with SIBO have been shown to be inversely correlated with H₂ production [55]. A previous in vivo study demonstrated that 4% H₂ inhalation for 1 h twice a day for 6 months reduced the body weight of rats and increased the level of metabolites such as glucose 6 phosphate dehydrogenase (G6PD) and 6-phosphogluconate dehydrogenase (Pentose Phosphate Pathway) and malic acid (TCA cycle and cytosol) involved in reactions that reduced nicotinamide adenine dinucleotide phosphate (NADP⁺) to NADPH [56]. In PFV the oxidative current of G6PD rapidly increased in the applied potential from +0.2 to +0.8 V vs. Ag/AgCl [57]. Glutathione reductase (GR) catalyzes the reduction of oxidized glutathione (GSSG) to reduced glutathione (GSH) using NADPH as the reducing cofactor, and the initial-velocity of GR reached saturation depending on the concentration of NADPH [58]. Although H₂ concentration in the breath varies during the day, short-term exposure to high concentrations of H₂ may influence redox enzyme reaction rates, leading to reduction of NADP⁺ and GSSG. Therefore, environmental H₂ in a gaseous state may be a novel approach for the electrode potential of redox enzymes to evaluate the kinetic characteristics and metabolic function. The importance of the enzyme kinetics involved must be addressed in future studies to fully grasp how reversible hydrogen electrode potential influence the cellular redox state and oxidative stress.

7. Conclusions

H₂ is fermented by the colonic microbiotas and exhaled from the lungs out of the body. A breath H₂ test noninvasively measures the end-tidal H₂ concentration, which is equilibrated within the tissues. When the H₂ partial pressure of the tissues is higher than that of the end-tidal H₂, H₂ will be exhaled to the alveoli. However, when the H₂ partial pressure of the tissues is lower than that of the end-tidal H₂, H₂ will be diffused to the tissues. The H₂ concentration is typically expressed as a percentage of hydrogen in the atmosphere, and as such, the breath H₂ concentration can be thought of as being equal to its partial pressure. The reversible hydrogen electrode potential depends on the logarithm of the gaseous H₂ partial pressure equilibrated with the solution and H⁺ activity in the solution. In this study, the negative logarithm of environmental H₂ concentration is introduced as an index of redox potential in human body as compared to the air in order to evaluate metabolic function and ROS generation.