chemical paper(exact three pages)(Your grade on the fifth essay will depend, to large measure, on the improvement demonstrated relative to the first four essays.)

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LECTURE 18 VOLTA’S BATTERY to COMBUSTION to FERTILIZER $C:\Users\kozak\AppData\Local\Microsoft\Windows\INetCache\Content.MSO\A39024A7.tmp$ In Lecture 2, two fundamental generalizations, called Conservation Laws, were introduced. The first was the First Law of Thermodynamics which states that the total energy of an isolated system is constant. Energy can be transformed from one form to another, but can be neither created nor destroyed. This Law was instrumental in understanding exothermic and endothermic, chemical (Lecture 15) and nuclear (Lecture 16) reactions. This lecture will focus on the Second Law of Thermodynamics. This law states that a quantity called the entropy of an isolated system can never decrease over time. The total entropy can remain constant in ideal cases where the system is in an equilibrium state, a steady state, or for processes which are reversible. However, in all spontaneous processes , the total entropy increases and the process is irreversible . Our understanding of non-equilibrium processes advanced significantly in the last half of the 20th century, and we now have a deeper understanding of the role of biochemical reactions in driving the energetics of living systems and are able to understand oscillating chemical reactions (see above). Importantly, the increase in entropy which accounts for the irreversibility of natural processes, is also at the heart of our understanding of the asymmetry between past, present and future times, sometimes referred to as Time’s Arrow.

Let us start with a simple example, dropping a few grains of table salt in a glass of water. In crystalline NaCl, the NA+ ions and the Cl− ions are in a highly ordered, state. When salt dissolves, the ions are free to move around in the water. They are in a much more random and disordered state than before. In the dissolved state, water molecules are held around the ions as water of hydration. These water molecules are in a more ordered state than before because they are now confined to the immediate environment of the ions. Therefore, the dissolving of a salt involves both disordering and ordering processes. The disordering processes are usually dominant, so the overall effect is an increase in the disorder of the system.

How does an ionic substance like KCl dissolve in water? - Quora The representation of NaCl (above) shows ions on the surface of a small crystal.

In Lecture 2, we displayed the crystal structure of NaCl (see below).

As can be seen, “inside” the crystal, each Na+ ion is surrounded by six Cl- ions, but on the surface, a Na+ ion has only five nearest neighbors. Surface ions are less tightly bound than those in the interior of the crystal and, in the random motion that takes place in water, can be dislodged from the surface of the crystal. Each ion is immediately surrounded by water molecules. For a Na+ ion, the negative “end” of a water molecule provides a “hydration shell,” and for the Cl- ion it is the positive end of H2O, a consequence of Coulomb’s Law (see Lecture 3). Calculations show that the energy with which a surface ion is bound to the surface is approximately the same as the stabilization energy of an ion surrounded by a hydration shell of water molecules. The “bottom line” is that the dissolution of salt in water is, in fact, being driven by the change from an ordered to more disordered state of the system, i.e., by a change (an increase) in the entropy. The change in entropy S, like the change in energy E, depends only on the initial and final state of the system, ΔS = Sfinal – Sinitial. A positive value of ΔS indicates that the final state is more disorganized than the initial state. We are all familiar with what happens when an air freshener (or perfume) is sprayed into a room. Within minutes we can smell the scent in every nook and cranny. Attractions and repulsions between molecules in a gas are much weaker than attractions and repulsions between ions in a crystal. The expansion of a gas in an evacuated space is a spontaneous process, driven by an increase in randomness of molecules in the system, the process driven by an increase in the entropy.

Entropy of a system is related to its disorder. The quantitative relationship of entropy to disorder was first established by the Austrian physicist Ludwig Boltzmann (1844-1906). Boltzmann reasoned that the disorder of a particular state of the system, and thus its entropy, is related to the number of possible arrangements of molecules in the state. Boltzmann showed that entropy of an isolated system equals a constant times the natural logarithm of the number of possible arrangements of atoms or molecules in the system: S = k ln W W is the number of possible arrangements in the system, and k is a constant known today as Boltzmann’s constant: k = 1.38 x 10-23 J/oK (K is temperature in units of Kelvin) A pure crystalline solid at the absolute zero of temperature (K = 0o , C = − 273.15o ) is assumed to have only one arrangement of atoms or molecules; that is, W=1. If W=1, then S = k ln 1 = 0.

At any temperature above zero, the atoms acquire energy, more arrangements are possible, and so W >1 and S>0. This is referred to as the Third Law of Thermodynamics. Boltzmann made many other contributions to science, particularly in the area of statistical mechanics, which is the derivation of the bulk properties of large collections of atoms and molecules using the laws of probability.

Boltzmann strongly believed in the existence of atoms which, strange as it seems today, was a controversial viewpoint in physics at the beginning of the 20th century. The reason atomic theory was not universally accepted can be understood by recalling the classical expression for kinetic energy: KE = ½ mv2 Velocity v is the distance d divided by the time t. Inserting this into the above expression for the KE, KE = ½ m (d/t)2 Notice that you get the same value for the KE whether you use +t or –t. Thus, there is no difference in whether you have time t going “forward” or “backward.” This property is called time reversal invariance. If, as everyone believed, motion is correctly defined by classical Newtonian physics, then a system of gaseous atoms following Newton’s laws could not exhibit the spontaneous, irreversible expansion of a gas as found in Nature. Therefore, serious scientists rejected the atomic theory of matter. Boltzmann’s life had a tragic ending. In poor health and unable to endure continual intellectual attacks on his beliefs, he committed suicide on September 5, 1906. Ironically, it was only a few years later that the work of J.J. Thomson, Robert Millikan and Ernest Rutherford led to the acceptance of atoms. See Lecture 14. Incidentally, Boltzmann’s most famous student at the University of Vienna where he was professor of theoretical physics, was Lisa Meitner, whose contribution to the understanding of nuclear fission was brought out in Lecture 16. The reconciliation of Newtonian mechanics with spontaneous processes in Nature, starting from an atomic description of matter, is a challenging theoretical problem. One of the most important contributors to the understanding of this important problem was Ilya Prigogine, with whom I studied for two years at the Universite libré de Bruxelles in Belgium after receiving my doctorate at Princeton University. I give below a short summary of his contributions to the understanding of entropy. Prigogine (1917-2007) was born in Moscow a few months before the Russian Revolution of 1917, into a Jewish family. His father was a chemical engineer at the Imperial Moscow Technical School ; his mother was a pianist. Because the family was critical of the new Soviet system , they left Russia in 1921. They first went to Germany and then, in 1929, to Belgium , where Prigogine received Belgian nationality in 1949. Ilya Prigogine 1977c.jpg Prigogine began studying entropic processes in the years just before and during the Second World War. This led him to propose in 1947 the principle of minimum entropy production which states that the steady state of an irreversible process, i.e., a state in which the thermodynamic variables are independent of the time, is characterized by a minimum value of the rate of entropy production. Following the War, he began to study the work of Boris Pavlovich Belousov (1893 – 1970), a Russian chemist who discovered in the 1950s an oscillating chemical reaction, now called the Belousov–Zhabotinsky reaction . The BZ reaction is one of a class of reactions that serve as a classical example of non-equilibrium thermodynamics , resulting in the establishment of a nonlinear chemical oscillator . The reactions are important in that they show that chemical reactions do not have to be dominated by equilibrium thermodynamic behavior. These reactions are far from equilibrium and remain so for a significant length of time and evolve chaotically . In this sense, they provide an interesting chemical model of non-equilibrium biological phenomena. An essential aspect of the BZ reaction is its so called "excitability." Under the influence of stimuli, patterns develop in what would otherwise be a perfectly homogeneous, quiescent solution. A snapshot in time of the nature of such patterns is shown at the top of this Lecture. Prigogine made a connection being patterns driven by the chemical energy released in chemical reactions, and the patterns observed in fluids when thermal energy is imparted, which result in Rayleigh-Bénard convection. This a type of natural convection , occurring in a horizontal layer of fluid heated from below. The fluid develops a regular pattern of convection cells known as Bénard cells. See the convection patterns below. $C:\Users\kozak\AppData\Local\Microsoft\Windows\INetCache\Content.MSO\4032239D.tmp$

Starting from the principle of minimum energy production, Prigogine and his students analyzed these patterns theoretically. Dissipative structures was the name given by Prigogine for the complex structures created by irreversible processes. The Nobel Prize in Chemistry 1977 was awarded to Ilya Prigogine "for his contributions to non-equilibrium thermodynamics, particularly the theory of dissipative structures." Below is a photograph of the invited members of the Solvay Conference in Chemistry, organized in 1978 following the award of the Nobel Prize. Prigogine is in the front row; yours truly is in the last row. A Window to the World | SpringerLink

In the first decade of the 20th century when Thomson, Millikan and Rutherford were carrying out experiments leading to our modern understanding of the atom, experiments were being carried out in Berlin by Fritz Haber (1868 – 1934), a German chemist, on the synthesis from nitrogen gas (N2) and hydrogen gas (H2) of ammonia (NH3), a “Molecule that Changed the World.” In the first billion years of the Earth's existence, there was intense volcanic activity. The early atmosphere was mostly carbon dioxide (CO2), with little or no oxygen. There were smaller proportions of water vapor (H2O), ammonia (NH3) and methane (CH4). Our present atmosphere containing ~20% oxygen is due to the evolution of life on earth. The Haber process, also called the Haber–Bosch process, is an synthetic process for producing ammonia, and is the main industrial procedure for the production of ammonia today. Before the development of the Haber process, ammonia had been difficult to produce on an industrial scale, with early methods all being highly inefficient. The Haber-Bosch process converts atmospheric nitrogen (N2) to ammonia (NH3) by a reaction with hydrogen (H2) using a metal catalyst under high temperatures and pressures. The Nobel Prize in Chemistry 1918 was awarded to Fritz Haber "for the synthesis of ammonia from its elements."

Although the Haber process is mainly used to produce fertilizer today, during the first World War, it provided Germany with a source of ammonia for the production of explosives , compensating for the Allied Powers ' trade blockade on Chilean saltpeter . Saltpeter is the compound potassium nitrate with the chemical formula KNO₃. Saltpeter, or Guano (sometimes called “white gold”) is the accumulated excrement of seabirds and bats along the central los Molles coast in Chile. $C:\Users\kozak\AppData\Local\Microsoft\Windows\INetCache\Content.MSO\DB60D43.tmp$ The excreta, carcasses and egg shells cover large areas which tend to cover the entire bedrock of a colony and build up to large deposits when precipitation fails to wash the debris back into the ocean. Especially in dry and hot climates, these deposits can be massive and become rock-like during the aging process. The material of such deposits is called “guano,” a word that derives from the original Peruvian language Quechua; guano means “dung to fertilizer.”

As a manure , guano is a highly effective fertilizer due to its exceptionally high content of nitrogen , phosphate and potassium : key nutrients essential for plant growth.

KNO₃. is an ionic salt of potassium ions K⁺ and nitrate ions NO₃⁻, and is therefore an alkali metal nitrate. It also occurs in nature as a mineral, niter. Because of its ready solubility in water, niter is most often found in arid environments and often in conjunction with other soluble minerals like halides , iodates , borates , gypsum , and rarer carbonates and sulfates. Niter is a source of nitrogen, from which it derives its name. A major source of the sodium nitrate mineral nitratine (NaNO3) ("Chile saltpeter") is the Atacama Desert in Chile. It is also a readily available source of the nitrate anion (NO3−), which is useful in reactions carried out on industrial scales for the production of fertilizers , pyrotechnics and smoke bombs , glass and pottery enamels , preservatives (esp. meats), and solid rocket propellant .

As a fertilizer, NaNO3 is found in country stores across the Midwest. The substance is a strong oxidizer and its heat of reaction with combustibles (such as fuel oil, used in rural areas to heat homes and barns) may cause ignition. NaNO3 decomposes explosively when heated to temperatures above 1000 oF. 2 NH4NO3(s) + C → 2 N2(g) + 4 H2O(g) + CO2(g) + energy The Oklahoma City bombing was a domestic terrorist truck bombing on the Alfred P. Murrah Federal Building in Oklahoma City, Oklahoma , United States, on April 19, 1995. Perpetrated by Americans Timothy McVeigh and Terry Nichols , using fertilizer and fuel oil, the bombing happened at 9:02 am and killed at least 168 people, injured more than 680 others, and destroyed one third of the building. The blast destroyed or damaged 324 other buildings within a 16-block radius, shattered glass in 258 nearby buildings, and destroyed or burned 86 cars, causing an estimated $652 million worth of damage.

A view of the destroyed Alfred P. Murrah Federal Building from across the adjacent parking lot, two days after the bombing. The importance of the Haber process is that it yields ammonia (NH3), the starting reagent in producing fertilizers and gunpowder.

The first use led to the possibility of producing crops which saved from starvation literally millions of people since the synthesis was discovered in the first years of the 20th century. See the following Lecture on Agriculture and the Green Revolution. The second use led to a compound which resulted in the human destruction of the great wars of the 20th century.

The Haber Process combines nitrogen from the air with hydrogen derived mainly from natural gas (methane) into ammonia. The reaction is reversible and the production of ammonia is exothermic. N2(g) + 3 H2(g) 2 NH3(g)

A flow scheme for the Haber Process looks like this:

Experimental Conditions

The catalyst

The catalyst is actually slightly more complicated than pure iron. It has potassium hydroxide (KOH) added to it as a promoter - a substance that increases its efficiency.

The pressure

The pressure varies from one manufacturing plant to another, but is always high. (200 atmospheres).

Recycling

At each pass of the gases through the reactor, only about 15% of the nitrogen and hydrogen converts to ammonia. (This figure also varies from plant to plant.) By continual recycling of the unreacted nitrogen and hydrogen, the overall conversion is about 98%.

The proportions of nitrogen and hydrogen

The mixture of nitrogen and hydrogen going into the reactor is in the ratio of 1 volume of nitrogen to 3 volumes of hydrogen.

Avogadro's Law: Equal volumes of gases at the same temperature and pressure contain equal numbers of molecules. This means that the gases are going into the reactor in the ratio of 1 molecule of nitrogen to 3 of hydrogen.

That is the proportion demanded by the equation.

In some reactions an excess of one of the reactants is used. You do this if it is particularly important to use up as much as possible of the other reactant - if, for example, it was much more expensive. That doesn't apply in this case.

There is always a down-side to using anything other than the equation proportions. If you have an excess of one reactant there will be molecules passing through the reactor which can't possibly react because there isn't anything for them to react with. This wastes reactor space - particularly space on the surface of the catalyst.

The Temperature

Equilibrium considerations

You need to shift the position of the equilibrium as far as possible to the right in order to produce the maximum possible amount of ammonia in the equilibrium mixture.

The forward reaction (the production of ammonia) is exothermic.

According to Le Chatelier's Principle, this will be favored if you lower the temperature. The system will respond by moving the position of equilibrium to counteract this - in other words by producing more heat.

In order to get as much ammonia as possible in the equilibrium mixture, you need as low a temperature as possible. However, 400 - 450°C isn't a low temperature!

Rate considerations

The lower the temperature you use, the slower the reaction becomes. A manufacturer is trying to produce as much ammonia as possible per day. It makes no sense to try to achieve an equilibrium mixture which contains a very high proportion of ammonia if it takes several years for the reaction to reach that equilibrium.

You need the gases to reach equilibrium within the very short time that they will be in contact with the catalyst in the reactor.

The compromise

400 - 450°C is a compromise temperature producing a reasonably high proportion of ammonia in the equilibrium mixture (even if it is only 15%), but in a very short time.

The Pressure

Equilibrium considerations

Notice that there are 4 molecules on the left-hand side of the equation, but only 2 on the right.

According to Le Chatelier's Principle, if you increase the pressure the system will respond by favoring the reaction which produces fewer molecules. That will cause the pressure to fall again.

In order to get as much ammonia as possible in the equilibrium mixture, you need as high a pressure as possible. 200 atmospheres is a high pressure, but not amazingly high.

Rate considerations

Increasing the pressure brings the molecules closer together. In this particular instance, it will increase their chances of hitting and sticking to the surface of the catalyst where they can react. The higher the pressure the better in terms of the rate of a gas reaction.

Economic considerations

Very high pressures are very expensive to produce on two counts.

You have to build extremely strong pipes and containment vessels to withstand the very high pressure. That increases your capital costs when the plant is built.

High pressures cost a lot to produce and maintain. That means that the running costs of your plant are very high. The compromise

200 atmospheres is a compromise pressure chosen on economic grounds. If the pressure used is too high, the cost of generating it exceeds the price you can get for the extra ammonia produced.

The Catalyst

Equilibrium considerations

A catalyst has no effect whatsoever on the position of the equilibrium. Adding a catalyst doesn't produce any greater percentage of ammonia in the equilibrium mixture. Its only function is to speed up the reaction.

Rate considerations

In the absence of a catalyst the reaction is so slow that virtually no reaction happens in any sensible time. The catalyst ensures that the reaction is fast enough for a dynamic equilibrium to be set up within the very short time that the gases are actually in the reactor.

Separating the ammonia

When the gases leave the reactor they are hot and at a very high pressure. Ammonia is easily liquefied under pressure as long as it isn't too hot, and so the temperature of the mixture is lowered enough for the ammonia to turn to a liquid. The nitrogen and hydrogen remain as gases even under these high pressures, and can be recycled. The Haber process produces about 500 million tons (453 billion kilograms) of fertilizer every year. This fertilizer helps to feed about 40% of the world's population, and is one of the pillars of the Green Revolution. See the following, final Lecture.

Volta’s battery, the BZ reaction described earlier, and ammonia synthesis are examples of a general class of reactions called oxidation-reduction ( or Redox ) reactions. Redox reactions are characterized by the actual or formal transfer of electrons between chemical species , most often with one species (the reducing agent) undergoing oxidation (losing electrons) while another species (the oxidizing agent) undergoes reduction (gaining electrons). The chemical species from which the electron is removed is said to have been oxidized, while the chemical species to which the electron is added is said to have been reduced. Note that acid-base reactions and combustion (oxidation) reactions, already known and used in industrial processes in the 19th century, are understood today in terms of two subatomic particles, the proton in acid-base reactions and the electron in redox reactions.

Oxidation is the loss of electrons (or an increase in the oxidation state) of an atom, an ion , or of certain atoms in a molecule .

Reduction is the gain of electrons (or a decrease in the oxidation state) of an atom, an ion, or of certain atoms in a molecule.

Many reactions in organic chemistry are redox reactions due to changes in oxidation states but without distinct electron transfer. For example, during the combustion of wood with molecular oxygen, the oxidation state of carbon atoms in the wood increases and that of oxygen atoms decreases as carbon dioxide and water are formed. The oxygen atoms undergo reduction, formally gaining electrons, while the carbon atoms undergo oxidation, losing electrons. Thus oxygen is the oxidizing agent and carbon is the reducing agent in this reaction.

Although oxidation reactions are commonly associated with the formation of oxides from oxygen molecules, oxygen is not necessarily included in such reactions, as other chemical species can serve the same function.

Redox reactions can occur relatively slowly, as in the formation of rust , or much more rapidly, as in the case of burning fuel. There are simple redox processes, such as the oxidation of carbon to yield carbon dioxide (CO2) or the reduction of carbon by hydrogen to yield methane (CH4), and more complex processes such as the oxidation of glucose (C6H12O6) in the human body. The processes of oxidation and reduction occur simultaneously and cannot happen independently of one another, a feature that also characterized acid-base reactions, as was pointed out in Lecture 17. The oxidation alone and the reduction alone are each called a half-reaction , because two half-reactions always occur together to form a whole reaction. This brings us to an important point about Redox reactions. In balancing oxidation-reduction reactions, the resulting chemical equation must be balanced both with respect to mass and charge. The constraint on mass is imposed by the Law of Conservation of Mass, the second on charge is a consequence of the Law of Conservation of Charge. See Lecture 3.

When writing half-reactions, the gained or lost electrons are included explicitly in order that the half-reaction be balanced with respect to electric charge. The electrons cancel out when the half-reactions are combined to make the net chemical equation .

Though sufficient for many purposes, these general descriptions are not precisely correct. Although oxidation and reduction properly refer to a change in oxidation state , the actual transfer of electrons may never occur. The oxidation state of an atom is the “fictitious” charge that an atom would have if all bonds between atoms of different elements were 100% ionic. Thus, oxidation is defined as an increase in oxidation state, and reduction as a decrease in oxidation state. In practice, the transfer of electrons will always cause a change in oxidation state, but there are many reactions that are classed as "redox" even though no electron transfer occurs (such as those involving covalent bonds). As a result, simple half-reactions cannot be written for the individual atoms undergoing a redox process.

In redox processes, the reductant transfers electrons to the oxidant. Thus, in the reaction, the reductant or reducing agent loses electrons and is oxidized, and the oxidant or oxidizing agent gains electrons and is reduced. The pair of an oxidizing and reducing agent that are involved in a particular reaction is called a redox pair. A redox pair (or couple) is a reducing species and its corresponding oxidizing form, e.g., Fe+2 and Fe+3. Substances that have the ability to oxidize other substances (cause them to lose electrons) are said to be oxidative or oxidizing and are known as oxidizing agents , oxidants, or oxidizers. That is, the oxidant (oxidizing agent) removes electrons from another substance, and is thus itself reduced. And, because it "accepts" electrons, the oxidizing agent is also called an electron acceptor . Oxygen is the quintessential oxidizer.

Substances that have the ability to reduce other substances (cause them to gain electrons) are said to be reductive or reducing and are known as reducing agents , reductants, or reducers. The reductant (reducing agent) transfers electrons to another substance, and is thus itself oxidized. And, because it donates electrons, the reducing agent is also called an electron donor . Electron donors can also form charge transfer complexes with electron acceptors.

Reductants in chemistry are very diverse. Elemental metals , such as lithium, sodium, magnesium , iron , zinc , and aluminum , are good reducing agents. These metals donate or give away electrons relatively readily. Hydride transfer reagents, such as NaBH4 and LiAlH4 , are widely used in organic chemistry , primarily in the reduction of carbonyl compounds to alcohols . Another method of reduction involves the use of hydrogen gas (H2) with a palladium , platinum , or nickel catalyst . These catalytic reductions are used primarily in the reduction of carbon-carbon double or triple bonds.

A redox reaction is the force behind an electrochemical cell like the Galvanic cell shown below. The battery is made out of a zinc electrode in a ZnSO4 solution connected with a wire and a porous disk to a copper electrode in a CuSO4 solution.

In this type of reaction, a metal atom in a compound (or in a solution) is replaced by an atom of another metal. For example, copper is deposited when zinc metal is placed in a solution containing Cu+2 ions, e.g. copper (II) sulfate :

Zn(s)+ CuSO4(aq) → ZnSO4(aq) + Cu(s)

In this reaction, zinc metal displaces the copper(II) ion from copper sulfate solution and thus liberates free copper metal. The reaction is spontaneous and releases 213 kJ per 65 g of zinc because relative to zinc, copper metal is lower.

The ionic equation for this reaction is:

Zn + Cu2+ → Zn2+ + Cu

From two half-reactions , it is seen that the zinc is oxidized:

Zn → Zn2+ + 2 e−

And the copper is reduced:

Cu2+ + 2 e− → Cu

The term corrosion refers to the electrochemical oxidation of metals in reaction with an oxidant such as oxygen. Rusting , the formation of iron oxides , is a well-known example of electrochemical corrosion. Rust forms as a result of the oxidation of iron metal.

Common rust often refers to iron(III) oxide , formed in the following chemical reaction: 4 Fe + 3 O2 → 2 Fe2O3

The oxidation of iron(II) to iron(III) by hydrogen peroxide in the presence of an acid:

Fe2+ → Fe3+ + e−

H2O2 + 2 e− → 2 OH−

Overall equation:

2 Fe2+ + H2O2 + 2 H+ → 2 Fe3+ + 2 H2O

Oxidation is used in a wide variety of industries such as in the production of cleaning products and oxidizing ammonia to produce nitric acid , which is used in most fertilizers .

Redox reactions are the foundation of electrochemical cells , which can generate electrical energy or support electrosynthesis . Metal ores often contain metals in oxidized states such as oxides or sulfides, from which the pure metals are extracted by smelting at high temperature in the presence of a reducing agent. The process of electroplating uses redox reactions to coat objects with a thin layer of a material, as in chrome-plated automotive parts, silver cutlery, gold-plated jewelry , and galvanization .

**Many processes in Biology involve redox reactions. Perhaps the most important is the synergetic interplay between cellular respiration in animals and photosynthesis in plants.

Cellular respiration , for instance, is the oxidation of glucose (C6H12O6) to CO2 and the reduction of oxygen to water . The summary equation for cell respiration is:

C6H12O6 + 6 O2 → 6 CO2 + 6 H2O**

Photosynthesis involves the reduction of carbon dioxide into sugars and the oxidation of water into molecular oxygen.

6 CO2 + 6 H2O + light energy → C6H12O6 + 6 O2

Photosynthesis and cellular respiration are complementary, but photosynthesis is not simply the reverse of the redox reaction in cell respiration. Respiration, oxidizes sugars to produce carbon dioxide and water. As intermediate steps, the reduced carbon compounds are used to reduce nicotinamide adenine dinucleotide (NAD+) to NADH, which then contributes to the creation of a proton gradient , which drives the synthesis of adenosine triphosphate (ATP) and is maintained by the reduction of oxygen.

I close with an account of Fritz Haber after he developed his method for the synthesis of ammonia.

Haber greeted World War I with enthusiasm, joining 92 other German intellectuals in signing the Manifesto of the Ninety-Three in October 1914. Haber played a major role in the development of the non-ballistic use of chemical warfare in World War I , in spite of the proscription of their use in shells by the Hague Convention of 1907 (to which Germany was a signatory). He was promoted to the rank of captain and made head of the Chemistry Section in the Ministry of War soon after the war began. In addition to leading the teams developing chlorine gas (Cl2) and other deadly gases for use in trench warfare , Haber was on hand personally when it was first released by the German military at the Second Battle of Ypres (22 April to 25 May 1915) in Belgium , resulting in over 67,000 casualties. Haber also helped to develop gas masks with adsorbent filters which could protect against such weapons.

Haber defended gas warfare against accusations that it was inhumane, saying that death was death, by whatever means it was inflicted. During the 1920s, scientists working at his institute developed the cyanide gas Zyklon A , which was used initially as an insecticide , especially as a fumigant in grain stores. Haber was born in Breslau (now Wrocław, Poland), Prussia , into a well-off Jewish family, but later converted to Christianity. With the rise of Hitler in 1933, Haber (like Einstein) left Germany, staying briefly in Paris, Spain, and Switzerland. He was in extremely poor health during these travels, eventually suffering fatally from what was either a stroke or heart attack. In a terrible irony, Zyklon A was used in Hitler’s gas chambers, where a number of his relatives perished.