Equilibrium chemicals

The equilibrium between dinitrogen tetroxide (colorless) and nitrogen dioxide (brown color) gases favor the formation of the latter as temperature increases (from bottom to top). The models show dinitrogen tetroxide and nitrogen dioxide molecules. 

3 The Relationship Between Chemical Kinetics and Chemical Equilibrium 

Equilibrium is a state in which there are no observable changes as time goes by. When a chemical reaction has reached the equilibrium state, the concentrations of reactants and products remain constant over time, and there are no visible changes in the system. However, there is much activity at the molecular level because reactant molecules continue to form product molecules while product molecules react to yield reactant molecules. This dynamic equilibrium situation is the subject of this chapter. Here we wUl discuss different types of equilibrium reactions, the meaning of the equihbrium constant and its relationship to the rate constant, and factors that can disrupt a system at equilibrium. 

Liquid water in equilibrium with its vapor in a closed system at room temperature. 

Few chemical reactions proceed in only one direction. Most are reversible, at least to some extent. At the start of a reversible process, the reaction proceeds toward the formation of products. As soon as some product molecules are formed, the reverse process begins to take place and reactant molecules are formed from product molecules. Chemical equilibrium is achieved when the rates of the forward and reverse reactions are equal and the concentrations of the reactants and products remain constant. 

Every chemical reaction can proceed in either direction, /A + + gY, even if it goes in one 

Every reversible reaction has an equilibrium expression in which K, the equilibrium constant, is defined in terms of molar concentrations (mol/L) as indicated by the square brackets  

Products favored X, is large Reactants favored X, is small 

The equilibrium must be shifted to the right, the side of the equilibrium where CHjCOOC.H, exists. This is achieved by any combination of the following adding C,H,OH, adding CH,COOH. removing H,0, removing CH,COOC,H,. 

Problem 3.15 Summarize the relationships between the signs of AW. T A5, and AG, and the magnitude of X,., and state whether a reaction proceeds to the right or to the left for the reaction equation as written. 

1 The Equilibrium Condition 13.5 Applications of the Equilibrium 13.6 Solving Equilibrium Problems 

The Characteristics of Chemical Constant Treating Systems That Have Small 

2 The Equilibrium Constant Reaction Quotient 13.7 Le ChStelier s Principle 

3 Equilibrium Expressions Involving Calculating Equilibrium Pressuresand The Effect of a Change in Concentration 

The equilibrium in a sak water aquarium must be carefully maintained to keep the sea life heakhy. (BorissoslDreamstime.com) 

To be in equilibrium is to be in a state of balance. A tug of war in which the two sides pull with equal force so that the rope does not move is an example of a sfaf/c equilibrium, one in which an object is at rest. Equilibria can also be dynamic, whereby a forward process and the reverse process take place at the same rate so that no net change occurs. 

A saturated solution coo (Section 13.2) of an ionic compound in contact with undissolved crystals of the same compound is a good example of dynamic equilibrium. The rate at which ions leave the surface of the crystals and enter the solution (dissolution) is equal to the rate at which ions leave the solution and become part of the solid (crystallization). Hence the concentration of ions in solution and the amount of undissolved solid do not change with time. 

A saturated solution is one of many instances of dynamic equilibrium that we have already encountered. Vapor pressure ooo (Section 11.5) is another example of dynamic equilibrium. The pressure of a vapor above a liquid in a closed container reaches equilibrium with the liquid phase, and therefore stops changing when the rate at which molecules escape from the liquid into the gas phase equals the rate at which molecules 

1 THE CONCEPT OF EQUILIBRIUM We begin by examining reversible reactions and the concept of equilibrium. 

2 THE EOUILIBRIUM CONSTANT We define the equilibrium constant based on rates of forward and reverse reactions, 

Chemical-equilibrium analysis allows finding the maximum achievable per-pass conversion and the composition of the reaction mixture at equilibrium. Accordingly, it may suggest measures for improving both conversion and selectivity. 

Gibbs free-energy minimization offers an elegant computational manner without the need to specify the stoichiometry. In addition, phase equilibrium is accounted for. Since occasionally the method fails, considering explicit reactions is safer. 

Chemical equilibrium plays an important role in the operation of chemical plants where manufacturers want to optimize the amount of product. This can be done by adjusting reaction conditions such as temperature, pressure, and concentrations of reacfanfs used, all key factors fhaf affecf chemical equilibrium. 

17-1 Basic Concepts 17-2 The Equilibrium Constant 17-3 Variation of with the Form of the Balanced Equation 17-4 The Reaction Quotient 17-5 Uses of the Equilibrium Constant, Kc 17-6 Disturbing a System at Equilibrium Predictions 

17-10 Relationship Between/Cp and Kc 17-11 Heterogeneous Equilibria 17-12 Relationship Between AG and the Equilibrium Constant 

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Chemical equilibrium is associated with chemical equations. To explain it, let us consider the following hypothetical chemical reaction  

In this equation, s represents a solid and g a gas AG of this reaction is the driving force. It can be shown that  

AG° is the free energy change when the reactants and the products are in their standard states. K, the equilibrium constant for the reaction, is given by Equation 18.45. 

P represents partial pressure. This equation is also known as mass action expression. At equilibrium, there is no change in free energy. Therefore, Equation 18.44 reduces to  

And K is given in terms of standard free energy change as  

2 Chemical Equilibrium Forward and Reverse Reactions Mathematical Relationships 

The Equilibrium (Mass Action) Expression Gas Phase Equilibria Kp vs. Kp Homogeneous and Heterogeneous Equilibria Numerical Importance of the Equilibrium Expression Mathematical Manipulation of Equilibrium Constants Reversing the Chemical Equation Adjusting the Stoichiometry of the Chemical Reaction Equilibrium Constants for a Series of Reactions Units and the Equilibrium Constant 

When Cases Are Present Effect of a Change in Temperature on Equilibrium Effect of a Catalyst on Equilibrium 

The Bronsted-Lowry Theory of Acids and Bases The Role of Water in the Brensted-Lowry Theory Weak Acids and Bases 

1 Relationship between Equilibrium Constant, Kp/po and Gibbs Energy Change, A G° 

In general If AT 1 products dominate and reaction moves towards completion If K 1 reactants dominate and reaction tends not to occur 

If AG /kJmol-1 0 this does not mean that the reaction cannot occur because 

Many industrial processes actually have AG°/kJ mol-1 values which are positive. A G° has the following attributes  

The more negative A G°, the more likelihood there is that the reaction will occur of its own accord. However, reactions may be too slow (in the absence of a catalyst) despite having a large negative free energy change. For example  

At the upper left end of the table K is very small in magnitude.  

At the upper left end of the table the concentration of the products is small compared with the concentration of the reactants At the upper left end of the table AG /kJ moP is large and positive  

As we descend the table the equilibrium shifts such that the concentrations of the products become larger and the concentrations of the reactants become smaller. 

Homogeneous Equilibria Equilibrium Constants and Units Heterogeneous Equilibria The Form of K and the Equilibrium Equation Summary of Rules for Writing Equilibrium Constant Expressions 

The Efifect of a Catalyst Summary of Factors That May Affect the Equilibrium Position 

Chemical Equilibrium Chemical Equilibrium describes the state in which the rates of forward and reverse reactions are equal and the concentrations of the reactants and products remain unchanged with time. This state of dynamic equilibrium is characterized by an equilibrium constant. Depending on the nature of reacting species, the equilibrium constant can be expressed in terms of molarities (for solutions) or partial pressures (for gases). The Equilibrium constant provides information about the net direction of a reversible reaction and the concentrations of the equilibrium mixture. 

Factors That Affect Chemical Equilibrium Changes in concentration can affect the position of an equilibrium state—that is, the relative amounts of reactants and products. Changes in pressure and volume may have the same effect for gaseous systems at equilibrium. Only a change in temperature can alter the value of equilibrium constant. A catalyst can establish the equilibrium state faster by speeding the forward and reverse reactions, but it can change neither the equilibrium position nor the equilibrium constant. 

Interactivity Determining Extent—Concentration from Equilibrium Expression (15.3) 

Every system in chemical equilibrium, under the influence of a change of any one of the factors of equilibrium, undergoes a transformation... [that produces a change]... in the opposite direction of the factor in question. 

5 Fleterogeneous Equilibria Reactions Involving Solids and liquids 661 

6 Calculating the Equilibrium Constant from Measured Equilibrium Concentrations 662 

9 Le Ch telieFs Principle How a System at Equilibrium Responds to Disturbances 677 

The double arrows in this equation indicate that the reaction can occur in both the forward and reverse directions and can reach chemical equilibrium. We encountered this term in Chapters 11 and 12, and we define it more carefully in the next section. 

1 The Equilibrium Constant Governs the Concentration of Reactants and Products at Equilibrium 512 

2 The Equilibrium Constant Can Be Used to Predict the Direction and Equilibrium Concentrations of a Chemical Reaction 524 

3 The Equilibrium Constant for a Reaction Can Be Determined from the Standard Gibbs Energy Change 531 

4 The Response of an Equilibrium System to a Change in Conditions Can Be Determined Using Le Chatelier s Principle 536 

Keeping fish in an aquarium requires maintaining an equilibrium among the living organisms and the water. 

Reversible Reactions Rates of Reaction Chemical Equilibrium Le Chatelier s Principle 

The FREZCHEM model is a chemical equilibrium model. For a reaction such as gypsum dissolution 

What about chemical reactions Why do chemical systems eventually reach equilibrium The answer is analogous to that for the rock There are balanced forces acting on the chemical species in the system. These forces are actually energies—chemical potentials of the different chemical species involved in the system at equilibrium. The next section introduces chemical equilibrium in those terms. 

For a chemical reaction occurring in a closed system, species that have some initial chemical identity ( reactants ) change to some different chemical identity ( products ). In the previous chapter, we made the point that the Gibbs energy is dependent on the amount of any substance, and defined the chemical potential as the change in the Gibbs energy with respect to amount  

Because G varies with each it should be no surprise that during the course of a chemical process, the total Gibbs energy of the entire system changes. 

We now define the extent as a measure of the progress of a reaction. If the number of moles of the ith chemical species in the system at time t = 0 is n, o the extent is given by the expression 

The following reaction is set up with the initial amounts of each substance listed below. 

1 To study the effect of addition of solutions ofFe (NOy), KSCN, SnCli or KNO to the equilibrium system obtained by mixing aqueous solutions of Fe(N03) i and KSCN 

When aqueous solutions of Fe(N03)3 and KSCN are mixed, the following three equilibria exist simultaneously  

The consequent decrease in the concentration of Fc will cause FcSCN to dissociate so that K( remains constant. Dissociation of FeSCN will result in fading away of the reddish-brown colour. 

To about 40 ml of distilled water taken in a beaker, add 4 drops of Fe(N03)3 solution. Add KSCN solution dropwise with shaking until the solution acquires a light reddi.sh-brown colour. Mix well and pour equal volumes of the solution into 5 clean test tubes. Mark the tubes as A, B, C, D and E and set aside tube A for colour comparison. Add 10 drops of the solution of Fe(N03)3 to tube B, KSCN to 

Tube Initial colour Solution added (10 drops) Colour intensity 

One of the main problems of electrochemistry is the study of conditions concerning the conversion of chemical energy into electrical and vice versa. Quantitative relations between both forms of energy are based on exact thermodynamical laws which can bo applied, if the process occurs reversibly. These laws, therefore, enable to determine the amount of energy to be gained or spent in a reversible course of a reaction, if the thermodynamical properties of the reacting system are known. 

FUNDAMENTALS OK THERMO DYNAMICS If a chemical process under consideration, expressed by a common equation 

Similarly the free energy of a system after reaction can be expressed as  

It has been already said that the condition of the spontaneity of a reaction is the possible transfer of a certain quantity of energy to the surroundings, the maximum of which being equivalent at constant pressure and temperature to the decrease in the free energy of the system. A characteristic feature of all thermodynamically feasible and spontaneously occurring processes is therefore expressed as follows  

In electrochemistry all galvanic cells and storage batteries belong to systems satisfying this condition. 

In contrast to chemical kinetics, which focuses on the rates of chemical reactions, chemical equilibrium focuses on the final state of reactions. It is defined as the state in which the concentrations of all the components reach a steady-state condition and no further changes occur macroscopically. That means there is no tendency toward changes in molecular concentrations or ion abundances, although they do change microscopically (dynamic equilibrium state). One should keep in mind that chemical equilibria must involve both forward and reverse reactions whereas chemical kinetics only concerns forward reactions. Contrary to the simple reaction represented in Equation 10.1, the equilibrium state involves both the forward and the reverse processes, for example  

In the case of processes occurring in a reactor, the molecules can be delivered directly via transfer line systems or picked up by carrier gases to the inlet of a mass spectrometer. If the reaction occurs in ambient conditions, the molecules can be sampled by atmospheric pressure ion sources. If the ionization, transmission, and detection efficiencies of the reactants and the products of Equation 10.6 are available, their concentrations can be estimated from 

Once the equilibrium constant is obtained, the standard Gibbs free energy (AG°) of the reaction can be computed using the following equation  

Notably, the change in with temperature can also be estimated by monitoring spectral features as a function of reaction temperature, T. Since AG° = AH -TAi, the standard enthalpy change (A//°) and standard entropy change (A5°) of the studied reaction can be derived  

The association of solvation molecules to an ion is called nucleation reactions. It is normally conducted under low vacuum conditions to facilitate ion-molecule interactions. 

The two opposing processes in the dehnition of equilibrium may be chemical reactions. Imagine placing two chemicals that react with other, A and B, into a reaction vessel. Let us say that the products of the reaction are 

As the reaction proceeds, C and D begin to form, and their concentrations increase while the concentrations of A and B decrease. Now let us say that C and D also react with each other and that A and B are the products. 

This means that the reverse of reaction shown in Equation 11.4 occurs at the same time. After a period of time, imagine that these two opposing chemical reactions occur not only at the same time but also at same rate. When this happens, chemical equilibrium occurs. Chemical equilibrium thus refers to two opposing chemical reactions occurring at the same time and at the same rate with no net change. Such a reaction is written with a double arrow ( ). 

As with the vapor pressure example, equilibrium will not occur if any of the chemicals, reactants, or products escape or are removed from the container. 

PURPOSE OP EXPERIMENT Study the properties of a system at chemical equilibrium, and determine the value of the equilibrium constant. 

When reactants are first mixed and held at a given temperature, the concentrations of the reactants decrease and the concentrations of the products increase. However, since the products can also be converted back to the reactants, you have opposing changes. With time and under a given set of conditions, you reach a point where the concentrations of all species remain constant, even though the reactions in both directions continue. The chemical system is then said to have attained a state of chemical equil-brium. This state persists as long as the conditions remain constant. 

In this experiment you will study qualitatively a complex ion equilibrium, and you will determine the equilibrium constant for the formation of FeNCS and see if the value stays the same as you change the concentration. The net ionic equation that describes the reaction is 

The following expression for the equilibrium constant describes the state of equilibrium. 

In this experiment you will study the properties of the system at chemical equilibrium, and you will determine the equilibrium constant for the reaction by using either a Visual Method or an Instrumental Method. These methods are based on the following general ideas. The intensity of the color of a solution of FeNCs2+ will depend on the concentration of this ion in the solution and the depth of the solution through which you look. In the Visual Method you will compare a solution of known concentration with a 

As previously mentioned, the Gibbs free energy of a chemical system is minimized at equilibrium. At equilibrium, the forward and reverse reactions of a chemical reaction, such as given in Eq. [1-8], are necessarily occurring at the same rate, and the change in Gibbs free energy is zero  

For the moment, consider only reactions involving chemicals dissolved in water. The preceding equations can be combined with the definition of the reaction quotient, Eq. [1-10], to define an equilibrium constant, K, that applies to the final expected chemical composition of the system  

Note that molar concentrations must be used in the expression for K and that K equals the reaction quotient Q only at equilibrium (Fig. 1-8). Equation [1-12], the expression of the equilibrium state of a reaction, is known as the mass action law and is a direct consequence of the minimization of Gibbs free 

TABLE 1-1 Equilibrium Constants and Standard Free Energy Changes for Some Common Environmental Reactions0 

CHjCOOH Acetic acid H+ Hydrogen ion + ch3coo- Acetate ion 1 00 6.5 

The most important basis of analytical chemistry is the theory of chemical equilibrium. For ca. 350 years chemists have been performing chemical operations with the intention obtaining information about chemical composition. This means, first of all, utilizing the laws and the relationships describing chemical equilibria. 

Chemical equilibrium is a dynamic equilibrium. In a system which is in equilibrium, reactions do not stop, even if no movement is visible for an external observer. When molecules react, they form new products, and the products are decomposed again, and whereas a certain amount of products is formed, simultaneously an equal amount of the original reactants is generated as a result o the consumption of the products. 

For a theoretical description of chemical equilibrium and to derive its inherent laws, there exist two fundamentally different models, namely the thermodynamic approach and the kinetic approach. Both approaches result in the same mathematical relationships. 

For quantitation of mixtures of substances, the following quantities are important  

X the mole fraction (molar fraction). This denotes the munber of moles of dissolved substance wb as a proportion of the total munber of moles in a solution ( a + b. where the index B denotes the solvent), x = nA+m a the activity a = / c is a concentration where the activity coefficient f is some kind of correcting factor . Activity is measured using the same units as concentration c. For low values of c, approximately a c and/ 1. 

In addition to the determination of enthalpies of reaction as a function of temperature and pressure, thermodynamics allows us to calculate the equilibrium conversion for single or complex reversible reactions at given conditions (temperature, pressure, and initial composition). 

At constant temperature and pressure, chemical equilibrium is reached when the Gibbs energy shows a minimum. To describe the change of the Gibbs energy with temperature, pressure, and composition, the following fundamental equation can be applied  

The chemical potential in the standard state is identical with the molar Gibbs energy of formation Thus, the part on the left-hand side of Eq. (12.17) is id( ntical with the standard Gibbs energy of reaction Ag. The part 

This means that the value of the equilibrium constant at 25 C can directly be calculated from the tabulated standard Gibbs energies of formation. For a large number of compounds, the required standard Gibbs energies of formation at 25 C and 1 atm can be found in factual data banks or data compilations for the different states of aggregation (liquid, solid, and hypothetical ideal gas state) (3). In the various process simulators, usually the hypothetical ideal gas state is used. 

If no standard Gibbs energies for a given compound can be found, the value in the ideal gas state can be estimated with the help of group contribution methods (see Section 3.1.5). 

For choosing the conditions of a chemical process, we shonld know the thermodynamic properties and particularly the conditions of the chemical equilibrium. Before determining the reaction kinetics, it is necessary to verify if the reaction is thermodynamically possible. The pressure and temperature conditions are important to calculate the conversion of a reversible or irreversible reaction. For reversible reactions, we need to determine the chemical equilibrium constant, which is temperature dependent. With this constant, it is possible to predict what is the maximum equilibrium conversion of a reversible reaction. Therefore, the reversibility of the reaction imposes serious limitations. 

By means of thermodynamics, it is possible to predict whether a chemical reaction occurs and determine its composition. The thermodynamic equilibrium conversion represents the maximum conversion that can be achieved, regardless of the catalyst and reaction rates. However, the rates and conversions depend only on temperature, 

When a reaction occurs at constant temperature and pressure, it proceeds spontaneously varying in the direction of the increase in entropy. Once the equilibrium is reached, this entropy does not increase further. Consequently, from the first law of thermodynamics, the total change of the free energy of the system is always negative for any spontaneous reaction and zero at the equilibrium. 

The variation of free energy with temperature and pressure for an open system will be  

The chemical potential for ideal gases expressed as a function of the partial pressnre is defined by  

We have seen that a reaction will proceed just as long as AG is negative, although the reaction may take place extremely slowly. Further, a sufficient criterion for the state of equilibrium, where no net reaction in either direction occurs, is that AG=0. 

Consider in more detail the hydrogenation of ethylene the equation is  

Suppose the initial partial pressures to be 1 bar for each reactant. (The partial pressure of one component of a mixture of gases is defined as the pressure it would exert if it were alone in the available space. For an ideal mixture of perfect gases, the total pressure is the sum of the various partial pressures.) At 25°C, AG° for this reaction is -100.4 kJ/mol, and so the driving force is considerable, and reaction takes place over the catalyst with great vigour. The temperature of the reaction vessel is kept constant. As the reaction occurs, the partial pressures of ethylene and hydrogen drop, and their reduced active mass , or 

Let dn(C2H4), dn(H2) and dn(C2H6) be the changes in the amount of moles of C2H4, H2 and C2H6. A useful variable is the extent of the reaction c, (Fig. 7.1). In this case we may write  

If =0 no reaction occurs, so there are only reactants. If t= at least one of the reactants has completely reacted. 

Some of the most valuable results we have obtained from chemical thermodynamics are those that relate the position of chemical equilibrium to the thermodynamic properties of the reactants and products. With the aid of statistical thermodynamics we can go one step further and relate the position of equilibrium to the masses, dimensions, and vibrational frequencies of the molecules involved. 

Consider the equilibrium between two chemical species A and B. Each possesses a series of energy states, the energies of which we assign ef and e when measured from the lowest levels of each species. That is e is the energy of the kth state of B measured from the lowest level of B, s , and sf is the energy of the jth level of A measured from e (see Fig. 9.8). 

We can also define other equilibrium constants in terms of partition functions. For an equilibrium 

Now — RT In KP = AG° (Section 4.11), where AG° is the change in the Gibbs free energy for a mole of reaction with all reactants and products in their standard states at a pressure of 1 atm. 

It can also be shown that Kct the equilibrium constant in terms of concentrations, is related to the partition functions evaluated per unit volume 

The same types of equations may be derived for evaporation, micellization, and so on. 3.7.3 Chemical Equilibrium 

Using Equation 3.29 for components in solution, this leads to the following expression for the equilibrium constant K(s c c /CaCb)  

Although CVD processes inherently involve rapid changes, it is useful to examine the limiting case of long reaction times for insights into the nature of the films that can be deposited. To do this, we examine the final equilibrium state for the reactions of interest, which will depend on the initial reactant gas composition and the final pressure and temperature. 

The problem we are addressing here is what is the gas phase composition of a mixture of gases under specified conditions of pressure and temperature, where as much time as is necessary is allowed for the gases to equilibrate If there is a change of phase as one proceeds from one equilibrium state to another (i.e., solid silicon film forming on the container walls), then this has to be accounted for as well. 

The logical joining point between the study of chemical kinetics and chemical thermodynamics is the point, in an elementary step, where the rate of the forward reaction is equal to the rate of the reverse reaction. This is the point of chemical equilibrium and is implied in the relationships of equations (1-22) and (1-23). While we cannot conduct a short course here on the thermodynamics of chemical equilibrium, some review will be useful. The reader may also wish to refer to a favorite text on thermodynamics for amplification of our condensed presentation. 

The criterion for chemical equilibrium in a single-phase, single-reaction system is stated in terms of the minimization of the free energy of the system. In terms of the reaction, then 

For present purposes we will state without further derivation or proof that this may be generalized to 

For nonideal gas mixtures describable by a fugacity function (ex equation of state or the principle of corresponding states)  

Now these expressions clearly define the composition-free energy relationships required to ensure adherence to the general equilibrium relationship of equation (1-138) and are useful for calculations if we know quantities such as G, 7, or fi- However they are still somewhat far removed from the number that is normally of primary interest to us, which is the equilibrimn constant. To get to this point in a practically useful way, we can define a reference or standard state for each species to define the partial molar free energy at a reference mol fraction, x°, the temperature of interest, T, and a pressure of 1 atm. Following through with a little bit of thermodynamics, here is the following sequence of equations along the road to an equilibrium constant  

There is one word of warning that is essential to remember thermodynamics is silent about the rates of reaction. All it can do is to identify whether a particular reaction mixture has a tendency to form products it cannotsay whether that tendency will ever be realized. We explore what determines the rates of chemical reactions in Chapters 6 through 8. 

One of the major goals of chemistry is to predict what will happen when various substances come into contact. Will a chemical reaction occur, or will the substances just exist side by side One to way to approach this problem is through the concept of chemical equilibrium, which is the focus of this chapter. 

No reaction proceeds to 100% conversion. The conversion of all reactions is controlled by equilibrium. Just as equilibrium controls the ratio of vapor and liquid and the composition of each phase, equilibrium also controls the relative amounts of products that can be converted or reactants that can be formed in the chemical reaction. 

One may think of an equilibrium reaction as two reactions occurring in parallel  

The amount of A that exists at equihbrium is controlled by how much B is present, and vice versa. In some cases, the equilibrium condition gives extremely high conversion, which approaches 100%. In these cases, we say that the reaction goes to completion and we assume that complete conversion can actually be achieved. However, other cases exist in which less than 100% conversion may be obtained as a result of the thermodynamic limitations. It is these cases with which we are concerned in this section. 

In terms of kinetics, a chemical equilibrium is reached when the rate of the forward reaction is equal to the rate of the reverse reaction. This is often described as a dynamic equilibrium and a double arrow ( = ) is usually drawn between reactants and products. 

In Chapter 4 (Section 4.6), we introduced the concept of acids and bases. We can now look at the concept of equilibrium associated with the dissociation of weak acids and weak bases in water. Remember that strong acids and strong bases dissociate in water, they completely dissociate and a forward reaction does not exist. 

It is assumed that the tendency of a molecular mixture to interact can be analyzed as a function of the chemical (quantum) potential energy field and some action variable that reflects mass ratios or amounts of substance. Spontaneous chemical change occurs as the chemical potential of a system decreases, i.e. while Ap 0, and ceases when Ap = 0, at equilibrium. The quantity here denoted by Ap, also known as the affinity, a of the system, is the sum over all molecules, reactants and products 

The action represents the mass ratio in terms of individual activities, i. e. 

To use the relationship of equation (3) in practice a logical procedure would be to define an initial state in terms of valence-state electronegativities. However, this procedure does not provide a fixed reference point for the potential function, since the valence state for each species occurs under its own characteristic conditions. 

It has been common practice to define a standard state at an action of unity, Ai = 1, at which point by definition A/q = A/re. Equation (3) by reference to this standard state becomes 

Equations (4) and (5) are well-known thermodynamic expressions. A reaction attains equilibrium when the affinity becomes zero, hence 

Another flow-type apparatus has been operated by Reichl et It was set up as a non-recycle flow still and, due to the short residence time, it was used to determine isoharic VLB data of thermally unstable components and of reactive mixtures. Two esterification systems (methyl formate and ethyl acetate) and one etherification system (fert-amyl methyl ether) were investigated. 

Arlt proposed a new type of static still. The temperature and pressure are recorded at a given binary start mixture over time. By extrapolation to the zero time, pressure and temperature for this point can be derived. The method works if the reaction half-times are longer than 10 min, which is a prerequisite for the use of reactive distillation coluiims (see above). 

The binary data of the mentioned apparatus are used to determine binary interaction parameters of G -models or of equations of state. 

In formulating and understanding the problems inherent in reactive distillation, it is necessary to take the chemical equilibrium into account. 

Starting from the fundamental equation of thermodynamics for one component  

Great Barrier Reef and other coral reefs are Paper mill on the Potomac River near Weslemport, Maryland, neutralizes acid mine 

Part of the North Branch of the Potomac River runs crystal clear through the scenic Appalachian Mountains, but it is lifeless—a victim of acid drainage from abandoned coal mines. As the river passes a paper mill and a wastewater treatment plant near Westemport, Maryland, the pH rises from an acidic, lethal value of 4.5 to a neutral value of 7.2, at which fish and plants thrive. This happy accident comes about because calcium carbonate exiting the paper mill equilibrates with massive quantities of carbon dioxide from bacterial respiration at the sewage treatment plant. The resulting soluble bicarbonate neutralizes the acidic river and restores life downstream of the plant.1 

Equilibria govern diverse phenomena from the folding of proteins to the action of acid rain on minerals to the aqueous reactions used in analytical chemistry. This chapter introduces equilibria for the solubility of ionic compounds, complex formation, and acid-base reactions. Chemical equilibrium provides a foundation not only for chemical analysis, but also for other subjects such as biochemistry, geology, and oceanography. 

In the thermodynamic derivation of the equilibrium constant, each quantity in Equation 6-2 is expressed as the ratio of the concentration of a species to its concentration in its standard state. For solutes, the standard state is 1 M. For gases, the standard state is I bar (= 105 Pa 1 atm = 1.013 25 bar), and for solids and liquids, the standard states are the pure solid or liquid. It is understood (but rarely written) that [A] in Equation 6-2 really means [A]/( 1 M) if A is a solute. If D is a gas, [D] really means (pressure of D in bars)/( 1 bar). To emphasize that [D] means pressure of D, we usually write Pn in place of [D. The terms of Equation 6-2 are actually dimensionless therefore, all equilibrium constants are dimensionless. 

For the ratios [A]/(l M) and [D]/(l bar) to be dimensionless, LA] must be expressed in moles per liter (M), and [D] must be expressed in bars. If C were a pure liquid or solid, the ratio [( /(concentration of C in its standard state) would be unity (1) because the standard state is the pure liquid or solid. If [Cl is a solvent, the concentration is so close to that of pure liquid C that the value of [C] is still essentially 1. 

Effect of the Surface Area on the Reaction Rate. Assemble an apparatus as shown in Fig. 45. Introduce 0.1 g of powdered metallic zinc into a 50-ml reaction flask and pour in 10 ml of a 20% sulphu- ric acid solution. Put the end of the gas-discharge tube under the burette and determine the volume of hydrogen that evolves during two to five minutes. Perform the experiment at a constant temperature, shaking the reaction mixture. Why  

Run the same experiment, using 0.1 g of zinc in the form of pieces or plates. Repeat similar experiments with magnesium, taking it in an amount of 0.04 g and using a 10% sulphuric acid solution instead of the 20% one. How does the surface area of the reactants. affect the rate of the chemical reaction proceeding in a heterogeneous medium  

Effect of Light on the Decomposition of Silver Chloride. Introduce 2 ml of a 0.1 W silver nitrate solution into a test tube and add dropwise such an amount of a saturated sodium chloride solution that will be sufficient for the complete precipitation of the silver as -a chloride. Rapidly filter out the precipitate, rinse it with water and put it on two watch glasses. Place one glass with the precipitate in a dark cupboard, and leave the other at a window in daylight. In an hour, compare the colour of the precipitates and explain the phenom- ena you observe. How does light affect the reaction After the experiment, put the precipitates containing a silver compound into he jars set aside for them. 

What reactions are called practically reversible ones  

What are the conditions of irreversibility and reversibility of chemical reactions  

The shaded areas represent the values of the different enthalpies of transformation. 

Customarily chemical equilibrium has very instructively been introduced by describing the underlying meaning of reversible and irreversible reactions. 

In many cases, it has been proved that the products of a reaction themselves react forming the original reactants once more. A general reaction of the type in which reactants R1 and R2 from products P1 and P2 at a given temperature is considered. This is expressed as a forward reaction by  

and P2 react to regenerate Rx and R2. This is expressed as a reverse reaction by Pi + P2 Ri + R2 

These two oppositely directed reactions are represented by a single equation as  

To save effort, we often write these two exactly opposite equations as one, with double arrows  

We call the reagents on the right of the chemical equation as it is written the products and those on the left the reactants, despite the fact that we can write the equation with either set of reagents on either side. 

With the reaction just above, if you start with a mixture of nitrogen and hydrogen and allow it to come to 500°C at 200-atm pressure, some nitrogen and hydrogen combine to form ammonia. If you heat ammonia to 500°C at 200-atm pressure, some of it decomposes to nitrogen and hydrogen. Both reactions can occur in the same vessel at the same time. 

If we change the conditions on a system at equilibrium such as the N2, H2, NH3 system at equilibrium, for example by changing the temperature, we can get some further net reaction. Soon, however, the system will achieve a new equilibrium at the new set of conditions. 

Le Chatelier s principle states that if a stress is applied to a system at equilibrium, the equilibrium will shift in a tendency to reduce that stress. A stress is something done to the system (not by the equilibrium reaction). The stresses that we consider are change of concentration(s), change of temperature, change of pressure, and addition of a catalyst. Let us consider the effect on a typical equilibrium by each of these stresses. 

The ultimate aim of the modern movement in biology is in fact to explain all biology in terms of physics and chemistry. 

If all reactants and products are in solution, then the equilibrium constant can assume several meanings. If the substance in solution is an acid or base that dissociates in solution, then the equilibrium constant becomes the dissociation constant. For example, some acetic acid molecules dissociate into hydrogen (H+) and acetate (C2H3O2) ions  

Similar equilibrium constants have been found and tabulated for solubilities and precipitations. 

Chemical concentrations are given in different units depending on the field of technology they come from. They can be converted from one to another, if need be, through suitable conversion equations given in Johnson (1999). Numerical values of equilibrium constants depend on the concentration units used, so the conventional concentrations are given in moles per liter. Other commonly used concentration units are  

11 mole fraction moles solute/moles solution 

Gaseous fuels can be made from coal In a number of ways. One way (described In the text) is to react coal with steam to produce a mixture of CO and H2, which then reacts by catalytic methanation to yield CH4,the major component of natural gas. 

The product mixture of CO and H2 (synthesis gas) is used to prepare a number of industrial chemicals. 

The processes of catalytic methanation and steam reforming illustrate the reversibility of chemical reactions. Starting with CO and H2 and using the right conditions, you can form predominantly CH4 and H2O. Starting with CH4 and H2O and using different conditions, you can obtain a reaction mixture that is predominantly CO and H2. An important question is. What conditions favor the production of CH4 and H2O, and what conditions favor the production of CO and H2  

As noted earlier, certain reactions (such as catalytic methanation) appear to stop before they are complete.The reaction mixture ceases to change in any of its properties and consists of both reactants and products in definite concentrations. Such a reaction mixture is said to have reached chemical equilibrium. In earlier chapters, we discussed other types of equilibria, including the equilibrium between a liquid and its vapor and the equilibrium between a solid and its saturated solution. In this chapter, we will see how to determine the composition of a reaction mixture at equilibrium and how to alter this composition by changing the conditions for the reaction. 

Many chemical reactions are like the catalytic methanation reaction. Such reactions can be made to go predominantly in one direction or the other depending on the conditions. Let us look more closely at this revCTsibility and see how to characterize it quantitatively. 

Addition or Removal of a Substance Changes in Volume and Pressure Changes in Temperature Catalysis 

The increase in the RBC in people who live at high altitudes or who sleep in hypoxic tents is the result of chemical equilibrium. 

In this chapter, you will learn what constitutes an equilibrium, what factors influence equilibrium, and how knowledge of equilibrium can be used to solve a variety of problems. 

Mountain climbers sometimes fall ill due to the lower oxygen content of air at high altitudes. Long-term exposure to an oxygen-poor environment causes the production of more hemoglobin. The additional hemoglobin facilitates the transport of oxygen to the body. 

We have previously treated chemical reactions as though they only proceed as written— from left to right—wherein reactants react and products form. In a reverse reaction, however, the products actually become the reactants, and vice versa. To avoid confusion, we will always refer to species on the left side of the equation as reactants and those on the right side as products—regardless of whether we are discussing the forward or reverse reaction. 

LEARNIH6 GOAL Describe how temperature, concentration, and cataiysts affect the rate of a reaction. 

3 In the following reaction, what happens to the number of collisions when more Br2(g) molecules are added  

5 How would each of the following changes affect the rate of the reaction shown here  

As of the 2009 edition of the WADA code, the use of hypoxic tents has not been harmed, but the issue likely will be revisited in the years to come. 

Student Annotation In Michael Crichton s 1990 novel Jurassic Park, frog DNA was used to repair ancient dinosaur DNA to facilitate cloning of the extinct animals. In the story, the scientists believed the cloned animals could not reproduce because the population was designed to be entirely female. However, some of the dinosaurs became males—something known to occur in the frogs from which the DNA for repair had been taken—and the population of dinosaurs grew out of control. 

At the end of this chapter, you will be able to solve a series of problems related to the detection of hypoxic RBC enhancement [ H Page 661]. 

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To describe this chemical equilibrium quantitatively, we use a chemical equilibrium constant... 

This new quantity Sv p, the negative of which De Bonder (1920) has called the affinity and given the symbol of a script A, is obviously the important thennodynamic fiinction for chemical equilibrium ... 

The approach outlined here will describe a viewpoint which leads to the standard calculational rules used in various applications to systems in themiodynamic (themial, mechanical and chemical) equilibrium. Some applications to ideal and weakly interacting systems will be made, to illustrate how one needs to think in applying statistical considerations to physical problems. 

In tills chapter we shall examine how such temporal and spatial stmctures arise in far-from-equilibrium chemical systems. We first examine spatially unifonn systems and develop tlie tlieoretical tools needed to analyse tlie behaviour of systems driven far from chemical equilibrium. We focus especially on tlie nature of chemical chaos, its characterization and the mechanisms for its onset. We tlien turn to spatially distributed systems and describe how regular and chaotic chemical patterns can fonn as a result of tlie interjilay between reaction and diffusion. 

You can investigate the energetics of chemical equilibrium by comparing the heats of formation of reactants and products. This produces one of the most useful results of a chemical calculation. The accuracy and reliability of the heats of formation depend on the method used (see Choosing a Semi-Empirical Method on page 148). 

ELECTROLYTES, EME, AND CHEMICAL EQUILIBRIUM TABLE 8.6 Solubility Product Constants Continued)... 

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