1.For the reaction 2H2O (g)↽−−⇀2H2 (g)+O2 (g) the equilibrium concentrations were found to be [H2O]=0.250 M, [H2]=0.580 M, and [O2]=0.750 M. What is the equilibrium constant for this reaction?
Initial Concentration (M) 0.010 0.010 0 Change in Concentration (M) - x - x + 2 x Equilibrium Concentration (M) 0.010 - x
Change in Concentration (M) - x - x + 2 x Equilibrium Concentration (M) 0.010 - x 0.010 - x 0 + 2 x Substitute the expressions for the equilibrium concentrations into the equilibrium expression and solve for "x".
The binding of oxygen by hemoglobin (Hb), giving oxyhemoglobin (HbO2), is partially regulated by the concentration of H3O+and dissolved CO2in the blood. Although the equilibrium is complicated, it can be summarized as HbO2(aq)+H3O +(aq)+CO 2(g)⇌CO2−Hb−H ++O 2(g)+H2O(l) (a) Write the equilibrium constant expression for this reaction.
Figure 13.2A mixture of NO2and N2O4moves toward equilibrium. Colorless N2O4reacts to form brown NO2. As the reaction proceeds toward equilibrium, the color of the mixture darkens due to the increasing concentration of NO2.
When hydrogen reacts with gaseous iodine, heat is evolved. H2(g)+I2(g)⇌2HI(g) ΔH=−9.4kJ(exothermic) Because this reaction is exothermic, we can write it with heat as a product. H2(g)+I2(g)⇌2HI(g)+heat Increasing the temperature of the reaction increases the internal energy of the system. Thus, increasing the temperaturehastheeffectofincreasingtheamountofoneoftheproductsofthisreaction.Thereactionshiftstothe lefttorelievethestress,andthereisanincreaseintheconcentrationofH2andI2andareductionintheconcentration ofHI.Loweringthetemperatureofthissystemreducestheamountofenergypresent,favorstheproductionofheat, and favors the formation of hydrogen iodide. Whenwechangethetemperatureofasystematequilibrium,theequilibriumconstantforthereactionchanges. Lowering the temperature in the HI system increases the equilibrium constant: At the new equilibrium the concentrationofHIhasincreasedandtheconcentrationsofH2andI2decreased.Raisingthetemperaturedecreases the value of the equilibrium constant, from 67.5 at 357 °C to 50.0 at 400 °C. Temperature affects the equilibrium between NO2and N2O4in this reaction N2O4(g)⇌2NO2(g) ΔH=57.20kJ The positive ΔHvalue tells us that the reaction is endothermic and could be written heat+N2O4(g)⇌2NO2(g) At higher temperatures, the gas mixture has a deep brown color, indicative of a significant amount of brown NO2
Figure 13.8(a) The test tube contains 0.1MFe3+. (b) Thiocyanate ion has been added to solution in (a), forming the red Fe(SCN)2+ion. Fe3+(aq)+SCN−(aq)⇌Fe(SCN)2+(aq). (c) Silver nitrate has been added to the solution in (b), precipitating some of the SCN−as the white solid AgSCN. Ag+(aq)+SCN−(aq)⇌AgSCN( s). The decrease in the SCN−concentration shifts the first equilibrium in the solution to the left, decreasing the concentration (and lightening color) of the Fe(SCN)2+. (credit: modification of work by Mark Ott)
The reaction quotient is equal to the molar concentrations of the products of the chemical equation (multiplied together)overthereactants(alsomultipliedtogether),witheachconcentr ationraisedtothepowerofthecoefficient ofthatsubstanceinthebalancedchemicalequation.Forexample,thereactionquotientforthereversiblereaction 2NO2(g)⇌N2O4(g) is given by this expression:
In the early 20th century, German chemist Fritz Haber (Figure13.9) developed a practical process for convertingdiatomicnitrogen,whichcannotbeusedbyplantsasanutrient,toammonia,aformofnitrogenthat is easiest for plants to absorb.
Aswelearnedduringourstudyofkinetics,acatalystcanspeeduptherateofareaction.Thoughthisincrease inreactionratemaycauseasystemtoreachequilibriummorequickly(byspeedinguptheforwardandreverse reactions), a catalyst has no effect on the value of an equilibrium constant nor on equilibrium concentrations. Theinterplayofchangesinconcentrationorpressure,temperature,andthelackofaninfluenceofacatalystona chemicalequilibriumisillustratedintheindustrialsynthesisofammoniafromnitrogenandhydrogenaccordingto the equation N2(g)+3H2(g)⇌2NH3(g) Alargequantityofammoniaismanufacturedbythisreaction.Eachyear,ammoniaisamongthetop10chemicals, by mass, manufactured in the world. About 2 billion pounds are manufactured in the United States each year. Ammoniaplaysavitalroleinourglobaleconomy.Itisusedintheproductionoffertilizersandis,itself,animportant fertilizer forthegrowthofcorn,cotton,andothercrops.Largequantitiesofammoniaareconvertedtonitricacid, whichplaysanimportantroleintheproductionoffertilizers,explosives,plastics,dyes,andfibers,andisalsousedin the steel industry.