 # Chapter 16: Chemical Equilibrium- General Concepts WHAT IS EQUILIBRIUM?

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Chapter 16: Chemical Equilibrium- General Concepts WHAT IS EQUILIBRIUM?

The decomposition of N 2 O 4 (g) into NO 2 (g). The concentrations of N 2 O 4 and NO 2 change relatively quickly at first, but eventually stop changing with time when equilibrium is reached.

The equilibrium mixture is independent of whether we start on the “reactant side” or the “product side” The equilibrium between N 2 O 4 and NO 2.

The same equilibrium composition is reached from either the forward or reverse direction, provided the overall system composition is the same.

There is a simple relationship among the concentrations of the reactants and products for any chemical system at equilibrium Consider the equilibrium:

Four experiments to study the equilibrium among H 2, I 2, and HI gases. Different amounts of the reactants and products are placed in a 10.0 L reaction vessel at 440 o C where the gases establish equilibrium. When equilibrium is reached, different amounts of reactants and products remain.

This average value is called the reaction quotient, Q

The reaction can be evaluated at any concentrations At equilibrium (and 440 o C) for this reaction the reaction quotient has the value 49.5 (a unitless number) This relationship is called the equilibrium law for the system

The value 49.5 is called the equilibrium constant, K c, and characterizes the system For chemical equilibrium to exist, the reaction quotient Q must be equal to the equilibrium constant K c Consider the general chemical equation

The exponents in the mass action expression are the same as the stoichiometric coefficients At equilibrium The form is always “products over reactants” raised to the appropriate powers

Various operations can be performed on equilibrium expressions –Changing the direction of equilibrium – when the direction of an equilibrium is reversed, the new equilibrium constant is the reciprocal of the original

–Multiplying the coefficients by a factor – when the coefficients in an equation are multiplied by a factor, the equilibrium constant is raised to a power equal to that factor

–Adding chemical equilibria – when chemical equilibria are added, their equilibrium constants are multiplied

The gas law can be used to write the equilibrium constant in terms of partial pressures Equilibrium constants written in terms of partial pressures are given the symbol K p

The size of the equilibrium constant gives a measure of how the reaction proceeds General statements can be made about the equilibrium constant (either K c or K P )

The magnitude of K and the position of equilibrium. A large amount of product and very little reactant at equilibrium gives K>>1 (large K). When, approximately equal amounts of reactant and product are present at equilibrium. When K<<1, mostly reactant and very little product are present at equilibrium.

The two different forms of the equilibrium constants can be related

In a homogeneous reactions, all the reactants and products are in the same phase Heterogeneous reactions involve more than one phase For example the thermal decomposition of sodium bicarbonate (baking soda) Heterogeneous reactions can come to equilibrium just like homogeneous systems

If NaHCO 3 is placed in a sealed container, homogeneous equilibrium is established The equilibrium law involving pure liquids and pure solids can be simplified

–For a pure liquid or solid, the ratio of amount of substance to volume of substance is constant The concentration of a substance in a solid is constant. Doubling the number of moles doubles the volume, but the ratio of moles to volume remains the same.

The equilibrium law for a heterogeneous reaction is written without concentrations terms for pure solids or pure liquids. The equilibrium constants found in tables represent all the constants combined

According to Le Châtelier’s principle: If an outside influence upsets an equilibrium, the system undergoes a change in the direction that counteracts the disturbing influence and, if possible, returns the system to equilibrium We can consider some common “stresses” –Adding or removing a product or reactant The equilibrium shifts to remove reactants or products that have been added The equilibrium shifts to replace reactants or products that have been removed

–Changing the volume Reducing the volume of a gaseous reaction causes the reaction to decreases the number of molecules of gas, if it can Moderate pressure changes have a negligible effect on reactions involving only liquids or solids –Changing the temperature Increasing the temperature shifts a reaction in a direction that produces an endothermic (heat- absorbing) change Decreasing the temperature shifts a reaction in a direction that produces an exothermic (heat- releasing) change

–Catalysts have no effect on the position of equilibrium Catalysts change how fast a system achieves equilibrium, not the relative distribution of reactants and products –Adding an inert gas at constant volume If the added gas cannot react with any reactants or products it is inert towards the substances in the equilibrium No concentration changes occur, so Q still equals K and no shift in equilibrium occurs

Equilibrium calculations can be divided into two main categories: 1)Calculating equilibrium constants from known equilibrium concentrations or partial pressures 2)Calculating one or more equilibrium concentrations or partial pressures using the known value of K c or K P Consider the decomposition of N 2 O 4

Calculating the equilibrium constant this way is easy

More commonly, you will have a set of initial conditions and an equilibrium constant If a K P describes the system, equilibrium will usually be described in terms of partial pressures If a K c describes the system, equilibrium will usually be describe in terms of concentration (molarity, mol/L) The Initial, Change, Equilibrium or “ICE” table is a useful way to summarize the problem

–Example: Ethyl acetate, CH 3 CO 2 C 2 H 5, is produced from acetic acid and ethanol by the reaction At 25 o C, K c =4.10 for this reaction. Suppose 0.100 mol of ethyl acetate and 0.150 mol of water are placed in a 1.00 L reaction vessel. What are the concentrations of all species at equilibrium? ANALYSIS: Use an ICE table and the equilibrium constant to find the concentrations.

This can be solved by putting it in quadratic form:

Negative concentrations are not allowed, so A similar procedure can be used to calculate partial pressures using K P

Sometime simplifications can be made Example: Nitrogen and oxygen react to form nitrogen monoxide with K c =4.8x10 -31. In air at 25 o C and 1 atm, the N 2 concentrations and O 2 are initially 0.033 M and 0.00810 M. What are the equilibrium concentrations? ANALYSIS: The equilibrium constant is very small, very little of the reactants will be converted into products

Substituting: [N 2 ]=0.033-x=0.033 M [O 2 ]=0.00810-x=0.030810 M [NO]=2x=1.60x10 -17 M