3.8: Redox Examples

Introduction

Many of the useful energetic reactions around us are examples of redox reactions. These include reactions such as:

the combustion of methane (\(\ce{CH4 +2O2 -> CO2 + 2H2O}\)); many of the energy consuming and energy producing reactions in biological organisms; and the reactions in electrochemical energy sources such as batteries and fuel cells. A few examples are summarized in this section.

Oxidative phosphorylation and photophosphorylation

Electron Transport/Oxidative Phosphorylation

Mitocohondria are called the power plants of the cell because most of a cell’s ATP is produced there, in a process referred to as oxidative phosphorylation. There are two parts to the process. The first is electrons from reduced energy carriers, such as NADH and FADH2, enter an electron transport system via protein complexes containing iron. As seen in the figure \(\PageIndex{1}\), electrons move from one complex to the next, not unlike the way they might move through an electrical circuit. The energy from the oxidation processes in this electron transport chain is used to pump protons and sodium ions out of the inner region of a mitochondrion, creating a concentration gradient and a membrane potential. The extra protons outside then flow down this gradient through the ATP Synthase complex to drive the production of ATP.

Figure \(\PageIndex{1}\): Cartoon of a mitochondrion, the oxidations occurring in the in the electron transport chain (I, II, III, IV) used to pump H⁺ (also Na⁺) out of the inner "Matrix" and the conversion of ADP to ATP by ATP Synthase using the energy from the membrane potential and the proton gradient.

When electrons pass through complexes I, III, and IV, protons are moved from the mitochondrial matrix (inside of mitochondrion) and deposited in the intermembrane space (between the inner and outer membranes of the mitochondrion). The effect of this redistribution is to increase the electrical and chemical potential across the membrane. The energetics of the oxidations are¹:

Complex I: \(\ce{H+ + NADH + Q -> NAD+ + QH2}\), E⁺ = + 0.42 V, ∆_rG⁺ = -81 kJ/mol

Complex II: \(\ce{FADH2 + Q -> FAD + QH2}\), E⁺ = + 0.32 V, ∆_rG⁺ = -62 kJ/mol

Complex III: \(\ce{QH2 + 2Fe^{3+}(Cyt c) -> Q + 2Fe^{2+}(Cyt c) + H2O}\), E⁺ = + 0.15 V, ∆_rG⁺ = -29 kJ/mol

Complex IV: \(\ce{2Fe^{2+}(Cyt c) + 2H+ (1/2)O2 -> Q + 2Fe^{3+}(Cyt c) + H2O}\), E⁺ = + 0.56 V, ∆_rG⁺ = -108 kJ/mol

This is clearly enough energy to convert more than 1 ADP to ATP in the reaction:

\[ \ce{ADP^{3-} + HPO4^{2-} + H3O^{+} -> ATP^{4-} + H2O} \quad \Delta_r G^{+} = +31 kJ/mol\nonumber\]

The energy actually available for this last reaction by moving a mole of H⁺ ions across the proton gradient and the membrane potential is:

\[\Delta G_m = RTln\frac{[H^{+}]_{in}}{[H^{+}]_{out}}+F\Delta \phi\label{EQ:Gm}\]

Measurements suggest typical values are \(\Delta\phi \approx 0.14 V\) and \(\Delta pH \approx -1\)  (\(\frac{[H^{+}]_{in}}{[H^{+}]_{out} }\approx 10\)). Putting these values into equation \(\ref{EQ:Gm}\) yields ∆G_m ≈ 19 kJ/mol. So at least 2 moles of protons must be transferred across the membrane to generate the 31 kJ/mol necessary to create a mole of ATP.

In summary, one can think of the electron transport system as charging the battery for oxidative phosphorylation by pumping protons out of the mitochondrion. The intact inner membrane of the mitochondrion keeps the protons out, except for those that re-enter through ATP Synthase. The ATP Synthase allows protons to re-enter the mitochondrial matrix and harvests their energy to make ATP.

References

1. P. Atkins, J. de Paula, Physical Chemistry for the Life Sciences, 2nd Ed. W.H. Freeman and Company, New York, 2011, p. 208.

Photophosphorylation

Photophosphorylation to make ATP is found only in cells that carry out photosynthesis. This process is similar to oxidative phosphorylation in several ways. A primary difference is the ultimate source of the energy for ATP synthesis. In oxidative phosphorylation, the energy comes from electrons produced by oxidation of biological molecules. In the case of photosynthesis, the energy comes from the light of the sun.

Figure \(\PageIndex{6}\): Photophosphorylation.

Photons from the sun interact with chlorophyll molecules in reaction centers in the chloroplasts of plants or membranes of photosynthetic bacteria. A schematic of the process is shown above. The similarities of photophosphorylation to oxidative phosphorylation include:

• an electron transport chain
• creation of a proton gradient
• harvesting energy of the proton gradient by making ATP

with the help of an ATP synthase. Some of the differences include:

• the source of the electrons – \(\ce{H2O}\) for photosynthesis versus \(\ce{NADH/FADH2}\) for oxidative phosphorylation
• direction of proton pumping – into the thylakoid space of the chloroplasts versus outside the matrix of the mitochondrion
• movement of protons during ATP synthesis – out of the thylakoid space in photosynthesis versus into the mitochondrial matrix
• nature of the terminal electron acceptor – \(\ce{NADP^{+}}\) in photosynthesis versus \(\ce{O2}\) in oxidative phosphorylation.

Figure \(\PageIndex{7}\): Electron movement in photosynthesis.

Electron Transport in Chloroplasts vs. Mitochondria

In some ways, the movement of electrons in chloroplasts during photosynthesis is opposite that of electron transport in mitochondria. In photosynthesis, water is the source of electrons and their final destination is \(\ce{NADPH}\). In mitochondria, \(\ce{NADH/ FADH2}\) are electron sources and \(\ce{H2O}\) is their final destination.

How do biological systems get electrons to go both ways? It would seem to be the equivalent of going to and from a particular place while always going downhill, since electrons will move according to potential. The answer is the captured energy of the photons, which elevates electrons in photosynthesis to an energy where they move “downhill” to their \(\ce{NADPH}\) destination in a Z-shaped scheme. The movement of electrons through this scheme in plants requires energy from photons in two places to “lift” the energy of the electrons sufficiently. Last, it should be noted that photosynthesis actually has two phases, referred to as the light cycle (described above) and the dark cycle, which is a set of chemical reactions that captures \(\ce{CO2}\) from the atmosphere and “fixes” it, ultimately into glucose. The dark cycle is also referred to as the Calvin Cycle.

"Batteries"

The term "battery" is used in common speech to refer to any device that produces electrical energy via redox reactions. As chemists we differentiate between galvanic cells (a single electrochemical cell) and a battery (a connected collection of electrochemical cells). In this section we will concentrate on what happens in a single galvanic cell.

Single use batteries

A common primary battery is the dry cell, which uses a zinc can as both container and anode (“–” terminal) and a graphite rod as the cathode (“+” terminal). The Zn can is filled with an electrolyte paste containing manganese(IV) oxide, zinc(II) chloride, ammonium chloride, and water. A graphite rod is immersed in the electrolyte paste to complete the cell. The spontaneous cell reaction involves the oxidation of zinc:

\[\ce{Zn(s) -> Zn^{2+}(aq) +2 e^{-}} \tag{anode} \]

and the reduction of manganese(IV)

\[\ce{2 MnO2(s) + 2 NH4Cl(aq) + 2 e^{-} -> Mn2O3(s) + 2 NH3(aq) + H2O(l) + 2 Cl^{-}} \tag{cathode} \]

which together yield the cell reaction:

\[\ce{2 MnO2(s) + 2NH4Cl(aq) + Zn(s) -> Zn^{2+}(aq) + Mn2O3(s) + 2 NH3(aq) + H2O(l) + 2 Cl^{-}} \tag{cell} \]

The voltage (cell potential) of a dry cell is approximately 1.5 V. Dry cells are available in various sizes (e.g., D, C, AA, AAA). All sizes of dry cells comprise the same components, and so they exhibit the same voltage, but larger cells contain greater amounts of the redox reactants and therefore are capable of transferring correspondingly greater amounts of charge. Like other galvanic cells, dry cells may be connected in series to yield batteries with greater voltage outputs, if needed.

Figure \(\PageIndex{2}\): A schematic diagram shows a typical dry cell.

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Alkaline batteries (Figure \(\PageIndex{3}\)) were developed in the 1950s to improve on the performance of the dry cell, and they were designed around the same redox couples. As their name suggests, these types of batteries use alkaline electrolytes, often potassium hydroxide. The reactions are

\begin{align*}
& \text { anode: } \quad && \ce{Zn(s) + 2OH^{-}(aq) -> ZnO(s) + H2O(l) + 2e^{-}} \\[4pt]
& \text {cathode: } \quad && \ce{2 MnO2(s) + H2O(l) + 2 e^{-} -> Mn2O3(s) + 2 OH^{-}(aq)} \\[4pt]
\hline &\text { cell: } \quad && \ce{Zn(s) + 2 MnO2(s) -> ZnO(s) + Mn2O3(s)} \quad \quad \quad E_{\text {cell }}=+1.43\,\text{V}
\end{align*}

An alkaline battery can deliver about three to five times the energy of a zinc-carbon dry cell of similar size. Alkaline batteries are prone to leaking potassium hydroxide, so they should be removed from devices for long-term storage. While some alkaline batteries are rechargeable, most are not. Attempts to recharge an alkaline battery that is not rechargeable often leads to rupture of the battery and leakage of the potassium hydroxide electrolyte.

Figure \(\PageIndex{3}\): Alkaline batteries were designed as improved replacements for zinc-carbon (dry cell) batteries.

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Rechargable batteries

Nickel-cadmium, or NiCd, batteries (Figure \(\PageIndex{4}\)) consist of a nickel-plated cathode, cadmium-plated anode, and a potassium hydroxide electrode. The positive and negative plates, which are prevented from shorting by the separator, are rolled together and put into the case. This is a “jelly-roll” design and allows the NiCd cell to deliver much more current than a similar-sized alkaline battery. The reactions are

\[\begin{align*}
& \text { anode: } \quad && \ce{Cd(s) + 2OH^{-}(aq) -> Cd(OH)2(s) + 2 e^{-}} \\[4pt]
& \text {cathode: } \quad && \ce{NiO2(s) + 2 H2O(l) + 2 e^{-} -> Ni(OH)2(s) + 2OH^{-}(aq)} \\[4pt]
\hline &\text { cell: } \quad && \ce{Cd(s) + NiO2(s) + 2 H2O(l) -> Cd(OH)2(s) + Ni(OH)2(s)} \quad \quad \quad E_{\text {cell }}\approx+1.2 \,\text{V}
\end{align*} \nonumber \]

When properly treated, a NiCd battery can be recharged about 1000 times. Cadmium is a toxic heavy metal so NiCd batteries should never be ruptured or incinerated, and they should be disposed of in accordance with relevant toxic waste guidelines.

Figure \(\PageIndex{4}\): NiCd batteries use a “jelly-roll” design that significantly increases the amount of current the battery can deliver as compared to a similar-sized alkaline battery.

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Lithium ion batteries (Figure \(\PageIndex{5}\)) are among the most popular rechargeable batteries and are used in many portable electronic devices. The reactions are

\[\begin{align*}
& \text { anode: } \quad && \ce{LiCoO2 -> Li_{1-x}CoO2 + x~ Li^{+} + x~ e^{-}} \\[4pt]
& \text {cathode: } \quad && \ce{x~ Li^{+} + x~e^{-} + x~ C6 -> x~ LiC6} \\[4pt]
\hline &\text { cell: } \quad && \ce{LiCoO2 + x~ C6 -> Li_{1-x}CoO2 + x~ LiC6} \quad \quad \quad E_{\text {cell }}\approx+3.7\,\text{V}
\end{align*} \nonumber \]

The variable stoichiometry of the cell reaction leads to variation in cell voltages, but for typical conditions, x is usually no more than 0.5 and the cell voltage is approximately 3.7 V. Lithium batteries are popular because they can provide a large amount current, are lighter than comparable batteries of other types, produce a nearly constant voltage as they discharge, and only slowly lose their charge when stored.

Figure \(\PageIndex{5}\): In a lithium ion battery, charge flows as the lithium ions are transferred between the anode and cathode.

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The lead acid battery (Figure \(\PageIndex{6}\)) is the type of secondary battery commonly used in automobiles. It is inexpensive and capable of producing the high current required by automobile starter motors. The reactions for a lead acid battery are

\[\begin{align*}
& \text { anode: } \quad && \ce{Pb(s) + HSO4^{-}(aq) -> PbSO4(s) + H^{+}(aq) + 2 e^{-}} \\[4pt]
& \text {cathode: } \quad && \ce{PbO2(s) + HSO4^{-}(aq) + 3 H^{+}(aq) + 2 e^{-} -> PbSO4(s) + 2 H2O(l)} \\[4pt]
\hline &\text { cell: } \quad && \ce{Pb(s) + PbO2(s) + 2 H2SO4(aq) -> 2 PbSO4(s) + 2 H2O (l)} \quad \quad \quad E_{\text {cell }}\approx +2\,\text{V}
\end{align*} \nonumber \]

Each cell produces 2 V, so six cells are connected in series to produce a 12-V car battery. Lead acid batteries are heavy and contain a caustic liquid electrolyte, H₂SO₄(aq), but are often still the battery of choice because of their high current density. Since these batteries contain a significant amount of lead, they must always be disposed of properly.

Figure \(\PageIndex{6}\): The lead acid battery in your automobile consists of six cells connected in series to give 12 V.

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 Fuel cells

A fuel cell is a galvanic cell that uses traditional combustive fuels, most often hydrogen or methane, that are continuously fed into the cell along with an oxidant. (An alternative, but not very popular, name for a fuel cell is a flow battery.) Within the cell, fuel and oxidant undergo the same redox chemistry as when they are combusted, but via a catalyzed electrochemical that is significantly more efficient. For example, a typical hydrogen fuel cell uses graphite electrodes embedded with platinum-based catalysts to accelerate the two half-cell reactions:

Figure \(\PageIndex{7}\): In this hydrogen fuel cell, oxygen from the air reacts with hydrogen, producing water and electricity.

\[\begin{align*}
& \text { anode: } \quad && \ce{2 H2(g) -> 4 H^{+}(aq) + 4 e^{-}} \\[4pt]
& \text {cathode: } \quad && \ce{O2(g) + 4 H^{+}(aq) + 4 e^{-} -> 2 H2O (g)} \\[4pt]
\hline &\text { cell: } \quad && \ce{2 H2(g) + O2(g) -> 2 H2O (g)} \quad \quad \quad E_{\text {cell }}\approx+1.2\,\text{V}
\end{align*} \nonumber \]

These types of fuel cells generally produce voltages of approximately 1.2 V. Compared to an internal combustion engine, the energy efficiency of a fuel cell using the same redox reaction is typically more than double (~20%–25% for an engine versus ~50%–75% for a fuel cell). Hydrogen fuel cells are commonly used on extended space missions, and prototypes for personal vehicles have been developed, though the technology remains relatively immature.

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