Volume 2
2.12. HEAT OF REACTION. RATES OF CHEMICAL REACTIONS AND CHEMICAL EQUILIBRIUM

Exothermic and Endothermic Reactions • Rates of Chemical Reactions. Guldberg and Waage's Law of Mass Action. Rate Constant. Concentration and Temperature as Key Factors Influencing Reaction Rate • Catalysis • Reversible Reactions and Chemical Equilibrium Again. Bodenstein's Experiments • Le Chatelier's Principle. Synthesis of Ammonia • Exercises
2.12.1. Exothermic and Endothermic Reactions. In Volume 1, we learned that a chemical reaction is exothermic if it releases energy as it occurs. A reaction that consumes heat from the surroundings is referred to as endothermic.

Usually the energy produced in a chemical transformation is released in the form of heat. The neutralization reaction that occurs on mixing concentrated solutions of NaOH and HCl produces enough heat to be felt by the hand on touching the reaction flask. A complex exothermic reaction is used by the bombardier beetle for defense, as shown in this excellent video.

In some other exothermic reactions, the energy is released not only in the form of heat but also in the form of light. Examples of such reactions include the burning of natural gas, wood, and other organic materials in air, as well as explosions. The thermite reaction (Fe2O3 + 2 Al = 2 Fe + Al2O3) that we touched on earlier is a spectacular highly exothermic process that generates large amounts of heat and light. A fascinating case of light emitted by chemical reactions occurring in living organisms is bioluminescence (Figure 2-127). Another form of energy that some reactions can generate is electricity. Without these reactions there would have been no batteries that are broadly used in our everyday life.
Figure 2-127. Railroad worms (left; source) and fireflies (right; source) are examples of living creatures that use exothermic chemical reactions to emit light.


Endothermic reactions require external energy to occur. In most instances, heating the reagents is needed for an endothermic reaction to take place. For example, coal gasification, C + H2O = CO + H2, is an endothermic process that is conducted at 750 oC and above. For weakly endothermic reactions, just the warmth of the surrounding air is enough to get started. One such reaction is that of baking soda (NaHCO3) with dilute acetic acid (vinegar). This video shows a drop in the temperature of a mixture of vinegar and baking soda as the two react.

Energy released or absorbed in chemical reactions is expressed in energy units (kilojoules, kJ or kilocalories, kcal) per mole. The heat of a given reaction, also known as enthalpy, can be included in the chemical equation. In our course, we will use the symbol Q for heat of reaction. Importantly, the value of Q is negative for exothermic reactions and positive for endothermic reactions, as illustrated by the two examples presented in Figure 2-128. Without going into details, just remember the counterintuitive way the positive and negative signs are used for the energy produced and consumed in a chemical reaction. It might help to think about the energy released in an exothermic chemical transformation as a loss from the system comprising certain atoms and molecules upon their rearrangement into products. Hence the negative sign for the Q. The positive sign for the Q means that the energy is gained in an endothermic chemical transformation that absorbs energy as the reagents are converted to products.
Figure 2-128. The heat (Q) released in an exothermic reaction is presented with a minus sign. The heat (Q) consumed in an endothermic reaction is presented with a plus sign.


2.12.2. Rates of Chemical Reactions. Guldberg and Waage's Law of Mass Action. Rate Constant. Concentration and Temperature as Key Factors Influencing Reaction Rate. Different reactions occur at different rates. Neutralization of a strong alkali (such as KOH) with a strong acid (such as HCl) occurs within the time of mixing. Likewise, precipitation of BaSO4 on addition of Na2SO4 to a solution of BaCl2 is instantaneous. The oxidation of sucrose (sugar) with KMnO4 in the presence of NaOH, however, is noticeably slower, taking minutes as shown in this video. The time frame for some chemical transformations such as the rusting of iron in the air and the decomposition of hydrogen peroxide in dilute aqueous solutions can be months and even years. The rate of a chemical reaction is determined by measuring the decrease in reactant concentration or the increase in product concentration in a unit of time. A widely used reaction rate unit is M/sec, where M is molarity.

The branch of chemistry that deals with rates of chemical transformations is called chemical kinetics. Studying reaction rates is beneficial for a number of reasons. One is that the information obtained in such studies often reveals important details about how a particular reaction occurs. These details, including the dependence of reaction rate on various factors help us develop means to speed up important chemical transformations that are too slow and slow down those that are too fast. It is not surprising that all industrially important chemical processes have been most scrupulously studied by methods of chemical kinetics.

The rate of a chemical reaction can be influenced by a number of factors, including reagent concentration, temperature, pressure, light, and the presence or absence of a catalyst. Watch this educational and entertaining video for a discussion of how various factors affect reaction rates. The first and foremost factor that obviously should and does influence reaction rates is reagent concentration. In order for molecules to react with one another, they must collide. The probability of collision between molecules is higher at higher concentrations. Pioneered by two Norwegian scientists, Cato M. Guldberg (1836-1902) and Peter Waage (1833-1900), the law of mass action states that the rate of a chemical reaction is proportional to the product of concentrations of the reagents. For a general chemical reaction,

a A + b Bproducts

(where A and B are the reagents and a and b are the coefficients),

the rate V is expressed as:

V = k [A]m [B]
n

where

* [A] and [B] are molar concentrations of the reagents A and B, respectively;
* m and n are the powers the concentrations are raised to; and
* k is the coefficient of proportionality that is called the rate constant.

Some important comments are due. The powers n and m are known as the orders of the reaction. We say that the reaction is nth order in reagent A and mth order in reagent B. The sum (n + m) is called the overall reaction order. Alas, it is not uncommon to see in some textbooks the statement that the powers n and m are equal to the coefficients a and b, respectively, of the balanced equation a A + b B → products. Although in some instances n = a and m = b, this is not always the case. The values of the powers n and m can be determined only experimentally. Reactions are known where the order in a reagent involved is fractional. Some other reactions display zero order in one of the reagents. The physical meaning of zero order in a reagent is that the reaction rate is independent of the concentration of that reagent (any number to the power of zero is 1). The experimental determination of the overall reaction order and the order in each of the reagents often sheds light on reaction mechanism, the sequence of elementary steps involved in a chemical transformation.

The rate constant k does not depend on reagent concentrations. As is clear from the expression for V above, k is the rate of the reaction when the concentration of both reagents A and B equals 1 mol/L (1 raised to any power is 1). To compare intrinsic rates of two or more different chemical reactions, their k values should be used as the reference, not rates V, which are often measured at randomly selected concentrations of reacting substances.

While being independent of reagent concentration, the rate constant k of a chemical reaction is temperature-dependent. Molecules move and collide nonstop. Not every collision, however, results in chemical change. The colliding particles must possess certain energy, known as activation energy, in order for the reaction to happen. At higher temperatures, molecules move faster, thereby colliding more frequently and with higher energy, which results in faster reaction.

The dependence of k and, consequently, V on temperature is expressed by the Arrhenius equation. Although the Arrhenius equation is beyond the scope of this course, it is useful to remember the simple rule of thumb, sometimes referred to as van't Hoff's rule: For every 10 oC rise in temperature, the reaction rate increases by a factor of approximately 2 to 4.
Digression. Now we know that temperature and reagent concentration are the two key factors influencing rates of chemical reactions occurring in solution. The higher the concentrations of reacting molecules, the faster the reaction. The higher the temperature, the faster the reaction. Yet in some rare cases, faster reaction rates are observed at lower temperatures! This may happen when one of the reagents is a gas that is only moderately soluble in the reaction solution. Recall that the solubility of gases is lower at higher temperatures. As we raise the temperature, the k value increases, thereby entailing faster reaction. At the same time, the gaseous reagent (partially) escapes from the reaction solution due to its diminished solubility at higher temperatures. The lowered concentration of the gas slows down the reaction. If the decelerating lower concentration effect overpowers the accelerating higher k temperature effect, the overall slower reaction rate is observed. Deceleration of a chemical reaction with an increase in temperature may also stem from a tricky interplay between reaction rate and the position of some chemical equilibria involved in the transformation. One such example is the oxidation of nitric oxide with oxygen in the gas phase, 2 NO + O2 = 2 NO2, which slows down as the temperature is increased.
There are homogeneous and heterogeneous reactions. A homogeneous reaction occurs within a homogeneous mixture, such as in solution or in a mixture of gases. The neutralization reaction between a water-soluble acid and a water-soluble base in water is a homogeneous reaction. Exchange reactions between water-soluble salts in aqueous solutions are homogeneous reactions, too. Note that for a reaction to be considered homogeneous, only the reagents must be in the same phase, and not necessarily the products. For example, the reaction AgNO3 + NaCl = NaNO3 + AgCl↓ is homogeneous, despite the fact that the AgCl product precipitates out as the reaction occurs, thereby making the originally homogeneous reaction mixture heterogeneous. What counts for a homogeneous reaction is that all of the reagents are in solution.

Heterogeneous reactions occur in heterogeneous mixtures where reacting substances are located in different phases. For example, the glowing of white phosphorus (P4) in air deals with the oxidation reaction, as follows.

4 P + 5 O2 = 2 P2O5

This reaction is heterogeneous because it takes place at the interface separating the solid phosphorus and the O2 in the air. Watch this video to see a small lump of white phosphorus undergoing the highly exothermic oxidation in air. Note the white fumes emerging around the surface of the lump. These fumes are the P2O5 product that instantly reacts with moisture in the air to give phosphoric acid, P2O5 + 3 H2O = 2 H3PO4. As the demonstration shows, eventually the phosphorus lump catches fire. This is due to the heat produced at the beginning of the oxidation that causes the temperature to rise, thereby speeding up the reaction. The ignition is not immediate however, but rather takes some time to take place. In contrast, the solid white phosphorus that crystallizes out on evaporation of its solution in carbon disulfide (CS2) catches fire right away, as shown in another video. In both cases, the reagents are the same, the reaction is the same, and the reaction conditions are the same. Why is it that in the second experiment the phosphorus self-ignites instantly, whereas in the first one it takes much longer before the phosphorus catches fire?

The difference in the phosphorus oxidation rates observed in the two experiments stems precisely from the heterogeneous nature of the reaction that occurs on the surface. The reaction is much faster in the second experiment because the tiny crystals of phosphorus produced on the evaporation have a much larger surface area than the compact lump. On the larger surface, more P atoms get in contact with O2, which makes the reaction faster. For the same reason, air-oxidation of an iron nail (rusting) is a slow process, whereas a superfine Fe powder self-ignites on contact with the oxygen of the air, as can be watched here. Likewise, the sparks produced on striking iron against flint are also due to self-ignition of the tiny particles of the metal chipped off as a result of the strike. We can conclude that grinding up a solid chemical for a heterogeneous reaction will result in faster rates due to the larger surface area where the chemical transformation occurs.

2.12.3. Catalysis. There are many chemical reactions that, while yielding valuable products, cannot be utilized due to their unacceptably slow rates. A powerful remedy for many slow reactions is catalysis. As briefly discussed earlier, catalysis is the phenomenon of acceleration of a chemical reaction by a small amount of a substance called a catalyst. In many books and articles, it is stated that a catalyst can be recovered unchanged at the end of the reaction. More modern alternative definitions of a catalyst, however, omit this condition.

A spectacular example of catalysis is the effect of platinum metal on the reaction between hydrogen and oxygen, 2 H2 + O2 = 2 H2O. A mixture of hydrogen and oxygen could be stored at room temperature virtually unchanged for thousands of years (as long as it is kept away from an open flame). The reaction between H2 and O2 would still occur, but at such a slow rate that no sign of any chemical change could be detected. Things change dramatically, however, once a small amount of a fine platinum powder is added to a mixture of H2 and O2. As demonstrated in this video, about 30 seconds after some platinum is thrown into a bottle filled with hydrogen and oxygen, the mixture explodes. The violent reaction between H2 and O2 is prompted by the catalytic effect of platinum.
Digression. The ability of Pt metal to catalyze the reaction of hydrogen with the oxygen of the air is used in Döbereiner's lamp, the smart lighter invented in the 1820s by the famous German chemist Johann Wolfgang Döbereiner (1780-1849). This video shows how a simple Döbereiner's lamp can be built and how it works. (Note the right way the H2SO4 solution is prepared prior to the experiment. At stirring with a glass rod, the acid is added to water, not the other way around.) The diluted H2SO4 is then used for the reaction with zinc to generate hydrogen, Zn + H2SO4 = ZnSO4 + H2↑. The H2 produced flows to the nozzle at the top of the lighter, where it gets in contact with the air and the Pt gauze that initiates and catalyzes the combustion.
There are different types of catalysis and many distinct modes of action of various catalysts. In our course, we will touch on only one example of a particular mechanism of catalysis. The example deals with the the synthesis of ammonia, which will be discussed in detail at the end of this section.

N2 + 3 H2 = 2 NH3

This reaction is credited with feeding 40-50% of the current world population because it makes ammonia that is paramount for the industrial productions of fertilizers. Watch this animation that models the mechanism of catalysis of the reaction between N2 and H2 on the surface of specially activated iron metal. In the cartoon, the blue spheres depict the N atoms and white spheres the H atoms. First, the exceptionally strong N≡N triple bond of a molecule of N2 is split into two N atoms by the Fe atoms (grey spheres) on the surface. Molecules of H2 then can react with the resultant nitrogen atoms bonded to the iron atoms to produce NH3.

Catalysis is used to manufacture not only fertilizers, but also petrochemical fuels, plastics, dyes, pharmaceuticals, crop protection agents, and even certain types of food and supplements such as margarine and vitamins.

2.12.4. Reversible Reactions and Chemical Equilibrium Again. Bodenstein's Experiments. In 1899, the outstanding German physical chemist Max Ernst August Bodenstein (1871-1942) published his experimental study that since then has become a classic for generations of chemists. In his work, Bodenstein investigated rates of the reaction between iodine (I2) and hydrogen (H2) that leads to the formation of hydrogen iodide (HI).

H2 + I2 = 2 HI

Picture a mixture of H2 (1 mol) and I2 (1 mol) that is confined in a sealed glass reactor and heated to a particular temperature between 300 and 500 oC. Although at room temperature iodine is a purplish-black solid, within the used temperature and pressure range I2 is a purple gas. The other reactant, H2, and the product of the reaction, HI, are both colorless gases at room temperature and above. Thus, under the reaction conditions used, all three substances involved, H2, I2, and HI, form a homogeneous gaseous mixture.

As the mixture of H2 and I2 is heated to a high temperature, the originally deep purple color gradually fades due to the consumption of I2 (Figure 2-129). After a certain period of time, however, the color fading stops. At that point, the reaction mixture is not completely colorless but pale purple, which suggests that a small amount of I2 (and, consequently, the same molar amount of H2) are still present in the mixture. No matter how much longer the mixture is kept at that temperature, the color stays the same.
Figure 2-129. Color change progress in the reaction between H2 and I2.


In a parallel experiment, an identical vessel is filled with 2 mol of HI, the product of the reaction between H2 and I2. The vessel is sealed and heated to exactly the same temperature as was used in the first experiment. Although HI is colorless, soon the gas in the vessel turns pale purplish in color (Figure 2-130). This suggests that the HI is decomposing to hydrogen and iodine, 2 HI = H2 + I2. As the heating continues, the color deepens until it reaches exactly the same depth and shade as that finally observed in the previous experiment (Figures 2-129 and 2-130).
Figure 2-130. Color change progress observed on heating of HI at the same temperature as was used for the reaction between H2 and I2 (Figure 2-129).


We can continue heating the two mixtures at the same temperature for as long as we want, but the color will stay the same (Figure 2-131).
Figure 2-131. A mixture of HI, H2, and I2 of exactly the same composition is formed as a result of heating a mixture of H2 (1 mol) and I2 (1 mol) or heating HI (2 mol) at the same temperature.


These experiments indicate that two reactions occur simultaneously when a mixture of hydrogen and iodine is heated. One is the formation of HI.

H2 + I2 = 2 HI

The other is the decomposition of the HI product back to iodine and hydrogen.

2 HI = H2 + I2

Both processes are therefore reversible, as can be expressed by using the double arrow symbol for the equation.

H2 + I2 ⇄ 2 HI

When the reaction reaches the state of no net change in the concentrations of the reactants (H2 and I2) and the product (HI), the chemical reaction is at equilibrium or, as we can also say, the system has equilibrated. The equilibrium is dynamic, meaning that both the direct and reverse reactions continue. At equilibrium, however, the formation of HI and its decomposition occur nonstop at equal rates, thereby making the reagent to product ratio constant.

2.12.5. Le Chatelier's Principle. Synthesis of Ammonia. It is often important, if not critical, to shift the equilibrium of a reversible chemical transformation toward the desired product(s). Obviously, to shift a chemical equilibrium toward the product is to make the forward reaction faster than the backward reaction. A system at equilibrium may be controlled by using guidelines from Le Chatelier's principle. This principle, devised by, and subsequently named after, the French chemist Henry Louis Le Chatelier (1850-1936) states that if a change in reaction conditions (temperature, concentration, pressure, etc.) is applied to an equilibrated chemical transformation, the equilibrium will shift in such a way as to reduce the effect of the change.

First, let us see how Le Chatelier's principle works for the just discussed simple reversible reaction between hydrogen and iodine.

H2 + I2 ⇄ 2 HI

At equilibrium, the rate of the forward reaction equals the rate of the backward reaction. As a result, the concentrations of H2, I2, and HI remain the same in time. If we add more H2 or I2 or both to the reaction mixture, the concentration of the reagent(s) will increase. According to Le Chatelier's principle, the system will respond to this change by speeding up the rate of the forward reaction to "fight back" the new, higher concentration of the reagent(s). The reagents will react faster and the equilibrium will shift to the HI product or, as chemists often say, to the right, because the HI is on the right side of the equation. Conversely, if we add extra HI to the mixture, the equilibrium will shift to the left to counteract the change by increasing the HI decomposition rate.

Now we will learn how Le Chatelier's principle is applied to the synthesis of ammonia from nitrogen and hydrogen.

N2 + 3 H2 ⇄ 2 NH3

This reaction is run on a humongous scale scale, consuming more than 1% of all energy produced world-wide and, as mentioned above, feeding up to one-half of the current world population. Given the exceptionally large production scale, it is not surprising that the industrial synthesis of ammonia has been most thoroughly optimized. Le Chatelier's principle played a pivotal role in making the ammonia synthesis process as efficient and productive as it is today.

Before we proceed to the application of Le Chatelier's principle to the ammonia synthesis, it is worth to reemphasize one important point. The reversible reaction between N2 and H2 is catalytic. In the absence of a catalyst, the reaction is not feasible at all, as it does not occur at any appreciable rate. However, no catalyst can change the position of a chemical equilibrium because a catalyst always speeds up both the forward and reverse reactions equally. This fundamental rule applies to all reversible reactions. The reaction between N2 and H2 is no exception. On the surface of the ammonia synthesis catalyst molecules of NH3 are formed faster and decomposed back to N2 and H2 faster, too.

Now we can discuss how Le Chatelier's principle can be applied to the problem of shifting the equilibrium N2 + 3 H2 ⇄ 2 NH3 to the right in order to maximize the yield of ammonia. Let us consider possible changes in reaction conditions one by one, to see how they should be adjusted to favor the formation of NH3. All of the essential information that is needed for our discussion is presented in Figure 2-132.
Figure 2-132. Synthesis of ammonia from hydrogen and nitrogen.


1. Temperature. The higher the temperature, the faster the rate of a chemical transformation. The reaction we are considering is no exception. Should the process be run at as high a temperature as possible?

The answer is "No!" Our reaction is exothermic (Figure 2-132), meaning that heat is produced as the formation of NH3 from N2 and H2 takes place. As follows from Le Chatelier's principle, if we raise the temperature, the system will respond to relieve the stress by shifting the equilibrium toward the reverse process (decomposition of NH3) that is endothermic and that absorbs the heat.

So, Le Chatelier's principle suggests that in order to shift the equilibrium to the product, the reaction should be run at as low a temperature as possible. Yet, reaction rate considerations point to the need to raise the temperature in order to make the process faster. Are we between a rock and a hard place here? In a way, yes. But we can find a compromise by optimizing the temperature to make it high enough for sufficiently good reaction rates yet not too high to spare the NH3 product from too much decomposition. The optimal temperature conventionally used for the industrial ammonia synthesis is between 400 and 500 °C.

2. Pressure.
The reagents (N2 and H2) and the product (NH3) are gases. Gases are compressible. Using Le Chatelier's principle, we can predict how a higher pressure will affect the equilibrium. Figure 2-132 shows that 1 mol of N2 and 3 mol of H2 (4 mol in total) give rise to 2 mol of NH3. Since 1 mol of any gas occupies the same volume under identical conditions, the forward reaction is accompanied by the reduction in volume by a factor of 2. Therefore, applying pressure to the reaction mixture should shift the equilibrium toward the product in accord with Le Chatelier's principle: the stress from a higher pressure applied to the gaseous reaction mixture will be relieved by lowering its volume upon product formation. In industry, the process is run at a very high pressure of 150-250 atmospheres.

3. Change in concentration.
Le Chatelier's principle suggests that removal of ammonia from the reaction mixture should favor the formation of NH3. Indeed, lowering the concentration of NH3 in the mixture would prompt the system to shift the equilibrium to the right. The boiling point of ammonia is -33 oC, much higher than those of H2 (-253 oC) and N2 (-196 oC), which makes it possible to separate the NH3 produced in the reaction by cooling the mixture. But how is it possible to keep the mixture at 400-500 °C in order for the reaction to occur rapidly enough and, at the same time, cool it down to condense and separate the ammonia product? Sure these two operations cannot be done simultaneously within the same reactor. In the industrial process, a mixture of N2 and H2 is passed through a reactor containing the catalyst and maintained at 400-500 °C (Figure 2-133). After the pass, the gaseous mixture containing about 15% NH3 just formed is run through cooling coils in a separate unit to condense out the ammonia. The liquid NH3 is drained and transferred to a receiving tank, whereas the separated N2 and H2 gases are pumped back into the reactor for further conversion to ammonia. Simultaneously, the atmosphere in the reactor is continuously replenished with a fresh mixture of N2 and H2.
Figure 2-133. Schematic representation of the industrial setup for ammonia synthesis (source).


Besides the applications of Le Chatelier's principle to the ammonia synthesis, other interesting things can be learned from the industrial setup for running this exceptionally important reaction. One is that the process (Figure 2-133) is continuous, meaning that the product, NH3, is made nonstop. This type of chemical production is usually more economical and efficient as compared with batch manufacturing, where a chemical reaction is run over a finite period of time to make the desired product in a finite quantity. Another important feature of the ammonia synthesis process is the smart energy conservation design. Find the heat exchanger in Figure 2-133 to see how energy is saved in the process. The heat of the just produced hot gaseous mixture of NH3, H2, and N2 is used to warm up the recovered H2/N2 mixture after the separation of NH3 before cycling it back into the catalytic reactor. The hot water produced in the cooling condenser is also utilized.
Digression. Two Nobel Prizes were awarded for the ammonia synthesis process. One went to Fritz Haber (1868-1934) for the development of the catalytic reaction and the other to Carl Bosch (1874-1940) for its commercialization. Not surprisingly, the process is called the Haber-Bosch process.

Fritz Haber was and still remains a controversial figure. Given the unmatched benefit of the ammonia synthesis to mankind, the importance of Haber's scientific developments is hard to overestimate. Yet Haber was also the father and a strong advocate of chemical warfare. In addition to being a professor and the most powerful and influential scientist in Germany in the 1910-20s, Haber held the rank of captain in the German army during World War I. Unlike Haber, his wife Clara (née Immerwahr) was a pacifist and a particularly strong opponent of chemical warfare. Not being able to bear the shame of her husband's central role in the development and deployment of chemical weapons, Clara Haber killed herself using her husband's service pistol. She committed suicide upon her husband's return from Belgium, 10 days after the first successful gas attack by the Germans in Ypres on 22 April 1915. Just hours after Clara shot herself in the Habers' home garden in Berlin, Fritz Haber left for the Eastern Front to deploy gas against the Russians.

Being a Jew, Haber lost everything and fled Germany shortly after Hitler came to power. Devastated and in very poor health, Haber died of heart failure in Basel in 1934, less than half a year after his escape from Nazi Germany. As one article says, "Fritz Haber's inventions both saved millions, and were eventually used to kill millions".

Carl Bosch, a brilliant chemist and chemical engineering genius, was the key figure in the first commercialization of the synthesis of ammonia. Unlike Haber, Bosch was not Jewish, so he did not have to escape the Third Reich. However, due to his criticism of Nazism and its policies, Bosch was demoted and eventually lost his positions. He developed strong alcohol dependence as he fell into despair and depression, and finally died in Heidelberg in April 1940.

This is a brief story of the two authors of the arguably the greatest innovation and most significant and important invention of the 20th Century.
2.12.6. Exercises.

1. A reaction is exothermic if (a) heat is released as it occurs; (b) heat is consumed as it occurs; (c) heating is required for the reaction to occur. Answer

2. Which of the two following reactions is exothermic and which endothermic?

(a) N2 + 2 O2 = 2 NO2 + Q

(b) MgO + CO2 = MgCO3 – Q

Answer

3. Which of the two reactions (a) 2 KClO3 = 2 KCl + 3 O2 - Q and (b) KСlO4 = KCl + 2 O2 + Q do you think should be more useful for making oxygen? Answer

4. For the reaction NO2 + CO = NO + CO2, the rate equation (rate law) is (a) = k [NO2] [CO]; (b) = k [NO2]n [CO]m. Answer

5. Chemical reactions always occur faster at higher concentrations of all of the reagents involved. True or false? Answer

6. Rate constant of a chemical reaction is (a) independent of reagent concentrations but depends on temperature; (b) depends on both temperature and concentrations of reagents; (c) does not depend on temperature. Answer

7. The rate of a chemical reaction was measured at 20 oC. One would expect the same reaction at 30 oC to occur (a) 10 times faster; (b) 10 times slower; (c) 2-4 times faster; (d) 30 times faster; (e) 1.5 times faster. Answer

8. Concentrations of reagents in calculations of reaction rates are expressed in (a) mass percent; (b) mol per liter; (c) either unit. Answer

9. A catalyst is a substance that (a) shifts the equilibrium of a chemical reaction toward the product; (b) speeds up both the forward and reverse reactions; (c) accelerates the direct reaction and slows down the reverse reaction; (d) does not affect the position of chemical equilibrium. Answer

10. Put Le Chatelier's principle in your own words.

11. A chemical process at equilibrium features a particular reagent-to-product ratio that (a) does not change in time regardless of whether reaction conditions are changed or not; (b) does not change in time so long as the reaction conditions stay the same; (c) is likely to change upon addition of a catalyst. Answer

12. Using Le Chatelier's principle, predict how the specified changes in reaction conditions will affect the following gas-phase transformations at equilibrium.

(a) 2 NO + O2 ⇄ 2 NO2 - Q [pressure is raised, temperature remains the same]

(b) N2 + 2 O2 ⇄ 2 NO2 + Q [both temperature and pressure are raised]

(c) 2 NH3 ⇄ N2 + 3 H2 + Q [pressure is lowered, temperature is raised]

(d) 2 NO2 ⇄ N2O4 - Q [both temperature and pressure are raised]

Answer

13. It is said everywhere that a catalyst has no effect on the position of chemical equilibrium. It is also said in 2.12.3 above that a mixture of H2 and O2 could be stored unchanged for thousands of years, which implies that the mixture is at equilibrium, 2 H2 + O2 ⇄ 2H2O, that is shifted to the left. Yet once a Pt catalyst is added to the mixture, H2O is produced in the violent reaction between the H2 and O2 gases. Is this not a clear indication that the Pt catalyst shifts the equilibrium to the right? Answer