Reactions Not at Equilibrium

Concentration - Concentration Plots to Represent Simple Equilibria Effect on Equilibria of Changing Concentrations 

A G° is the free energy difference between products and reactants (each in their standard states at pressure, P° and 298.15 K). This equation does not refer to the actual reaction (41.2) (Frame 41) at equilibrium (except where Kp/po = 1 and A G° = 0). A G refers to the difference in Gibbs energy between products and reactants at other concentrations/pressures not corresponding to equilibrium. 

When AG = 0 the reaction is then at equilibrium and the concentrations (or partial pressures) of reactants and products are then those that appear in the equilibrium constant expression. 

In Frames 23 and 24 we introduced phase diagrams which comprised of a series of lines on a pressure, P, temperature, T, grid, which related to the conditions in which either solid was in equilibrium with liquid, liquid in equilibrium with gas or solid in equilibrium with gas. At any point with coordinates (T, P) which was situated on these lines, the two relevant phases then co-existed at equilibrium. 

In the case of simple equilibria we can draw similarly informative concentration versus concentration diagrams. Consider a simple reaction  

Although the reaction represented here is a gas reaction we do not need to use partial pressures, pi, to represent the equihbrium constant, (= K) we can equally well use gaseous concentrations, Ci (see for example Frame 40, equation (40.6)). 

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The change in free energy of a reaction not at equilibrium (AG) is given by... 

The value of the equilibrium constant for a reaction makes it possible to judge the extent of reaction, predict the direction of reaction, and calculate equilibrium concentrations (or partial pressures) from initial concentrations (or partial pressures). The farther the reaction proceeds toward completion, the larger the value of Kc. The direction of a reaction not at equilibrium depends on the relative values of Kc and the reaction quotient Qc, which is defined in the same way as Kc except that the concentrations in the equilibrium constant expression are not necessarily equilibrium concentrations. If Qc Kcr net reaction goes from left to right to attain equilibrium if Qc > Kc/ net reaction goes from right to left if Qc = Kc/ the system is at equilibrium. 

Activity product For a reaction not at equilibrium, the activity product (Q) at a given temperature and pressure is derived from the ratio of the activities of the products to the activities of the reactants. For the following reaction ... 

The equilibrium constant describes only equilibrium conditions. For reactions not at equilibrium, a similar equation gives information about the reaction ... 

Generally inhibitors are competitive or non-competitive with substrates. In competitive inhibition, the interaction of the enzyme with the substrate and competitive inhibitor instead of the substrate can be analysed with the sequence of reactions taking place as a result, a complex of the enzyme-inhibitor (El) is formed. The reaction sets at equilibrium and the final step shows the product is formed. The enzyme must get free, but the enzyme attached to the inhibitor does not have any chance to dissociate from the El complex. The El formed is not available for conversion of substrate free enzymes are responsible for that conversion. The presence of inhibitor can cause the reaction rate to be slower than the ordinary reaction, in the absence of the inhibitor. The sequence of reaction mechanisms is ... 

In the various laboratory studies when the outlet gas composition was not at equilibrium, it was observed that the steam-to-gas ratio (S/G) significantly affected the hydrogen leakage while the carbon monoxide still remained low. On the assumption that various reactions will proceed at different rates, a study was made to determine the effect of S/G on the reaction rate. The conditions for this test are presented in Table VII the findings are tabulated in Table VIII. 

A mechanical system, typified by a pendulum, can oscillate around a position of final equilibrium. Chemical systems cannot do so, because of the fundamental law of thermodynamics that at all times AG > 0 when the system is not at equilibrium. There is nonetheless the occasional chemical system in which intermediates oscillate in concentration during the course of the reaction. Products, too, are formed at oscillating rates. This striking phenomenon of oscillatory behavior can be shown to occur when there are dual sets of solutions to the steady-state equations. The full mathematical treatment of this phenomenon and of instability will not be given, but a simplified version will be presented. With two sets of steady-state concentrations for the intermediates, no sooner is one set established than the consequent other changes cause the system to pass quickly to the other set, and vice versa. In effect, this establishes a chemical feedback loop. 

PHI = 0.10 bar. (a) Calculate the reaction quotient, (b) Is the reaction mixture at equilibrium (c) If not, is there a tendency to form more reactants or more products ... 

If a system is not at equilibrium, which is common for natural systems, each reaction has its own Eh value and the observed electrode potential is a mixed potential depending on the kinetics of several reactions. A redox pair with relatively high ion activity and whose electron exchange process is fast tends to dominate the registered Eh. Thus, measurements in a natural environment may not reveal information about all redox reactions but only from those reactions that are active enough to create a measurable potential difference on the electrode surface. 

The surface is a very minor part of the entire Earth, but being heated from the interior and energised by the Sun s radiation can not only maintain an approximately constant temperature but it can allow activated chemical reactions that create especially energised organic compounds not at equilibrium These processes gave rise to life, apparently starting in the aqueous phase. 

The reaction of X with S must be fast and reversible, close to if not at equilibrium with concentration of S. It can be that there is an intermediate step in which X binds to a protein kinase (a protein which phosphorylates other proteins mostly at histidine residues in bacteria) using phosphate transferred from ATP. It then gives XP which is the transcription factor, where concentration of S still decides the extent of phosphorylation. No change occurs in DNA itself. Here equilibrium is avoided as dephosphorylation involves a phosphatase, though changes must be relatively quick since, for example, cell cycling and division depend on these steps, which must be completed in minutes. We have noted that such mechanical trigger-proteins as transcription factors are usually based on a-helical backbones common to all manner of such adaptive conformational responses (Section 4.11). 

Just by looking at the value of AG, you can determine which way a reaction goes. If AG < 0, the reaction goes to the right. If AG > 0, the reaction goes to the left. And, if AG = 0, the products and reactants are of exactly the same free energy (note that this does not mean that the products and reactants are at the same concentration), and the reaction is at equilibrium. 

Many systems are not at equilibrium. The mass action expression, also called the reaction quotient, Q, is a measure of how far a system is from equilibrium and in what direction the system must go to get to equilibrium The reaction quotient has the same form as the equilibrium constant, K, but the concentration values put into Q are the actual values found in the system at that given moment. 

Initially, both electrodes are at equilibrium. Since the anode has accumulated electrons and the cathode has depleted electrons, electrons begin to flow from electrode from the anode to the cathode. The thermodynamic driving force for the electron flow is the electrode potential difference, which for the fuel cell reaction is 1.23 Y at standard conditions. In addition to electron flow, H + ions produced at the anode diffuse through the bulk solution and react at the cathode. The reaction is able to continue as long as H2 is fed at the anode and 02 at the cathode. Hence, the cell is not at equilibrium. The shift in electrode potential from equilibrium is called the overpotential (>/). 

In this section, you iearned that the equiiihrium constant, iCc, is a ratio of product concentrations to reactant concentrations. You used concentrations to find K, and you used K to find concentrations. You aiso used an ICE table to track and summarize the initial, change, and equiiihrium quantities in a reaction. You found that the value of Kc is small for reactions that reach equilibrium with a high concentration of reactants, and the value of IQ is large for reactions that reach equilibrium with a low concentration of reactants. In the next section, you will learn how to determine whether or not a reaction is at equilibrium, and, if it is not, in which direction it will go to achieve equilibrium. 

The system is not at equilibrium. The reaction will proceed by moving to the left. 

Various samples were analyzed. The concentrations are given in the table below. Decide whether each sample is at equilibrium. If it is not at equilibrium, predict the direction in which the reaction will proceed to establish equilibrium. 

The galvanic cell pictured in Figure 7.1 is not at equilibrium. If switch S is closed, electrons will spontaneously flow from the zinc (anode) to the copper (cathode) electrode. This flow will continue imtil the reactants and products attain their equilibrium concentrations. If switch S is opened before the cell reaches equilibrium, the electron flow will be interrupted. The voltmeter would register a positive voltage, which is a measure of the degree to which the redox reaction drives electrons from the anode to the cathode. Since this voltage is a type of energy that has the potential to do work, it is referred to as a redox potential or cell potential, denoted as... 

The equilibrium constant for a dissociation reaction is usually referred to as the dissociation constant ( d), and the products of that dissociation are then written in the numerator. For reactions not at equilbrium, the value for the product of the concentrations of all products divided by the product of the concentrations of all reactants is frequently referred to as the mass action ratio. See Association Constant Dissociation Constant Mass Action Ratio... 

This principle has only limited application to reactions that are not at equilibrium. Furthermore, the Principle of Microscopic Reversibility does not apply to reactions commencing with photochemical excitation. See also... 

If the three gases in the reaction were at equilibrium and you then increased the carbon monoxide concentration, some Bt2 would combine with added CO to produce COBr2 and thereby minimize the increase in CO. Alternatively, if you decrease the CO concentration, some COBtj would decompose to produce CO and Br2 and thereby minimize any decrease in CO. Notice how the concentrations of all constituents shift to counteract the imposed change in a single substance. Of course, this shift does not affect the value of the equilibrium constant. Only a change in temperature can do that. 

The phases are certainly not at equilibrium, (i. e. vapor-liquid, liquid-liquid, vapor-liquid-liquid equilibrium), until all reactions have proceeded to completion. Therefore, mass transfer can be assumed to occur across the interfaces for all time t. 

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