How long will substances remain in equilibrium

However, the rate of attainment of equilibrium in this case is very slow at low temperatures [Eq. See also: Ammonia. A typical liquid-phase equilibrium is the deprotonation of acetic acid in water, reaction 23 and Eq. See also: Acetic acid ; Ionic equilibrium. These are usually studied at constant pressure, because at least one of the phases will be a solid or a liquid. The imposed pressure may be that of an equilibrium gaseous phase, or it may be an externally controlled pressure.

In describing such systems, as for all reaction systems, the nature of each phase must be specified. In the examples to follow, s, l, g, aq, and soln identify solid, liquid, gaseous, aqueous, and nonaqueous phases, respectively. For solutions or mixtures, the composition is needed, in addition to the temperature and pressure, to complete the specification of the system; the nature of any solvent also must be specified.

In the equilibrium shown as 25 , For pure water, the equilibrium constant is approximately equal to the vapor pressure of water at least at low pressures. When a small amount of solute is added, thereby decreasing the mole fraction of solvent, the vapor pressure p must be lowered to maintain equilibrium Raoult's law. The effect of the total applied pressure P upon the vapor pressure p of the liquid is given by the Gibbs-Poynting equation Here V l and V g are the molar volumes of liquid and vapor, respectively.

The equilibrium vapor pressure will increase as the total pressure is increased both V l and V g are positive and the activity of the liquid increases with pressure. If the external pressure is applied to a solution by a semipermeable membrane, then an applied pressure can be found which will restore the vapor pressure or activity of the solvent to its standard state value.

For a solid, such as barium sulfate, BaSO 4 , which dissolves in water as shown in Eq. When the solid state is pure, its activity is unity at kilopascals. If the solid is extremely finely divided, then its activity is greater than unity; with this increase in the activity of the solid state, the solubility must increase to maintain equilibrium.

On the other hand, inclusion of foreign ions in the crystal lattice solid solution formation lowers the activity of the solid state. When a gas, such as CO 2 , is dissolved in a liquid, its equilibrium with the gas phase is as shown in Eq. Equation 31 , Henry's law, represents the equilibrium constant for this equation, where m is the molality of dissolved CO 2.

When the gas dissociates in the liquid, as in reaction 32 , Eq. Similarly, when a solute distributes itself between two immiscible phases, that is, nA soln 1 A n soln 2 , the equilibrium constant takes the form of Eq. The equilibrium concentrations c i in Eq.

See also: Extraction. When components form immiscible phases at equilibrium, each condensed phase will be a saturated solution; complete immiscibility is impossible in principle, because the chemical potential of any component must be the same in all phases. The separation of a liquid system into two liquid phases is a manifestation of the nonideality of the solutions. For practical purposes, however, many solids may be regarded as immiscible because of the stringent requirements associated with formation of solid solutions.

In reactions involving condensed and immiscible phases, for example, reaction 35 , In such a case the Q term in Eq. Such is the case in transition phenomena, or melting-freezing phenomena. From the phase-rule viewpoint, the system in reaction 35 has only one degree of freedom if reaction is possible, so once a value of P or T is chosen, the remaining variable T or P is fixed by nature at equilibrium.

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Chemical equilibrium Article by: Rock, Peter A. Key Concepts Hide In a kinetic sense, chemical equilibrium occurs in a chemical reaction when the rates in its forward and reverse directions are equal, so that the concentrations of the reactant and product substances do not change with time. In a thermodynamic sense, chemical equilibrium is the condition in which there is no tendency for the composition of the system to change.

The chemical potential is the tendency of a substance to enter into chemical or physical change; for example, a reaction will be spontaneous left to right when the total chemical potential of the reactants is greater than that of the products.

When equilibrium is reached, the total chemical potentials of products and reactants become equal. For reactions at constant temperature and pressure, the difference in chemical potentials becomes equal to the Gibbs energy change. The decrease in Gibbs energy represents the maximum net work obtainable from the process; when no more work is obtainable, then the system is at equilibrium.

Chemical potential Thermodynamics attributes to each chemical substance a property called the chemical potential, which may be thought of as the tendency of the substance to enter into chemical or physical change. See also: Potentials The importance of the chemical potential lies in its relation to the affinity or driving force of a chemical reaction.

The equilibrium position can be changed by changing the reaction conditions. How quickly an equilibrium is reached depends upon:. The table summarises the effects of these factors on the equilibrium position and the time taken to reach equilibrium. Describe the change in reaction conditions which increases the rate of reaching equilibrium, but does not change the position of equilibrium position. Adding a suitable catalyst.

Reaching equilibrium - Higher When a reversible reaction happens in a closed container, it reaches a dynamic equilibrium. At equilibrium: the forward and backward reactions are still happening the forward and backward reactions have the same rate of reaction the concentrations of all the reacting substances remain constant they do not change The equilibrium position of a reversible reaction is a measure of the concentrations of the reacting substances at equilibrium.

If, for example, C is removed in this way, the position of equilibrium would move to the right to replace it. If it is continually removed, the equilibrium position shifts further and further to the right, effectively creating a one-way, irreversible reaction. This only applies to reactions involving gases, although not necessarily all species in the reaction need to be in the gas phase.

A general homogeneous gaseous reaction is given below:. Pressure is caused by gas molecules hitting the sides of their container. The more molecules in the container, the higher the pressure will be.

The system can reduce the pressure by reacting in such a way as to produce fewer molecules. In this case, there are three moles on the left-hand side of the equation, but only two on the right. By forming more C and D, the system causes the pressure to reduce. Increasing the pressure on a gas reaction shifts the position of equilibrium towards the side with fewer moles of gas molecules. If this mixture is transferred from a 1.

Because the volume is increased and therefore the pressure reduced , the shift occurs in the direction that produces more moles of gas. To restore equilibrium the shift needs to occur to the left, in the direction of the reverse reaction. The equilibrium will move in such a way that the pressure increases again. It can do that by producing more gaseous molecules. In this case, the position of equilibrium will move towards the left-hand side of the reaction. In this case, increasing the pressure has no effect on the position of the equilibrium.

Because there are equal numbers of molecules on both sides, the equilibrium cannot move in any way that will reduce the pressure again. Again, this is not a rigorous explanation of why the position of equilibrium moves in the ways described. A mathematical treatment of the explanation can be found on this page.

Three ways to change the pressure of an equilibrium mixture are: 1. Add or remove a gaseous reactant or product, 2. Add an inert gas to the constant-volume reaction mixture, or 3. Change the volume of the system. To understand how temperature changes affect equilibrium conditions, the sign of the reaction enthalpy must be known. Assume that the forward reaction is exothermic heat is evolved :.

In this reaction, kJ is evolved indicated by the negative sign when 1 mole of A reacts completely with 2 moles of B. For reversible reactions, the enthalpy value is always given as if the reaction was one-way in the forward direction.

The back reaction the conversion of C and D into A and B would be endothermic, absorbing the same amount of heat. The main effect of temperature on equilibrium is in changing the value of the equilibrium constant. It is not uncommon that textbooks and instructors to consider heat as a independent "species" in a reaction.