Unit 7: Equilibrium

19 terms. Say what each one means, then open the card to check.

• Dynamic equilibrium

  The state of a reversible reaction in a closed system when the forward and reverse reactions run at the same rate, so every concentration stays constant.

  Example. In a sealed bottle of soda the dissolved carbon dioxide and the gas above it keep trading molecules, yet the pressure reading no longer changes.

• Equilibrium constant (K)

  A number, fixed at a given temperature, equal to the ratio of product to reactant concentrations (or pressures) at equilibrium, each raised to its coefficient.

  Example. A reaction with K near $10^{9}$ ends almost entirely as products; one with K near $10^{-9}$ barely starts.

• Equilibrium expression

  The formula for K: products over reactants, each concentration raised to the power of its coefficient in the balanced equation.

  Example. For $2\text{NO}(g) + \text{O}_{2}(g) \rightleftharpoons 2\text{NO}_{2}(g)$, $K_c = \frac{[\text{NO}_{2}]^2}{[\text{NO}]^2[\text{O}_{2}]}$.

• Kc and Kp

  $K_c$ is written with equilibrium molarities; $K_p$ is written with equilibrium partial pressures of the gases. Both describe the same equilibrium.

  Example. For $\text{CO}(g) + \text{Cl}_{2}(g) \rightleftharpoons \text{COCl}_{2}(g)$, $K_p = P_{\text{COCl}_2} / (P_{\text{CO}} P_{\text{Cl}_2})$.

• Heterogeneous equilibrium

  An equilibrium with species in more than one phase. Pure solids and pure liquids are left out of K because their concentrations do not change.

  Example. For $\text{CaCO}_{3}(s) \rightleftharpoons \text{CaO}(s) + \text{CO}_{2}(g)$, $K_p = P_{\text{CO}_2}$ only.

• Changing K with the equation

  Reversing an equation inverts K, multiplying its coefficients by n raises K to the nth power, and adding equations multiplies their K values.

  Example. If K = 5.0 for a reaction, its reverse has K = 0.20, and the reaction with every coefficient doubled has K = 25.

• ICE table

  A table of Initial concentrations, the Change in each set by the coefficients, and the Equilibrium concentrations, used to find K or unknown amounts.

  Example. If 0.30 M of $\text{A}_2$ forms 2A and 0.10 M of $\text{A}_2$ reacts, the rows read 0.30, then minus 0.10, then 0.20 M for $\text{A}_2$, and 0.20 M for A.

• Reaction quotient (Q)

  The same ratio as K but built from the concentrations or pressures present at any moment, equilibrium or not.

  Example. A flask of pure reactants has Q = 0, since every product concentration in the numerator is zero.

• Q compared with K

  If Q is less than K the net reaction runs forward; if Q is greater than K it runs in reverse; if Q equals K the system is at equilibrium.

  Example. With K = 10 and Q = 2, more products form until Q rises to 10.

• Le Chatelier’s principle

  When a change disturbs a system at equilibrium, the net reaction shifts in the direction that partly undoes the change.

  Example. Removing a product as it forms keeps pulling the reaction forward, which raises the yield.

• Temperature and K

  Temperature is the one change that alters K. Raising the temperature increases K for an endothermic reaction and decreases K for an exothermic one.

  Example. Many salts dissolve endothermically, so their solubility equilibrium constant grows as the water is warmed.

• Volume and pressure changes

  Shrinking the volume of a gas mixture raises every partial pressure, and the net reaction shifts toward the side with fewer moles of gas.

  Example. Compressing a mixture for $2\text{CO}(g) + \text{O}_{2}(g) \rightleftharpoons 2\text{CO}_{2}(g)$ favors $\text{CO}_2$, two moles of gas against three.

• Inert gas at constant volume

  Adding an unreactive gas to a rigid container raises the total pressure but leaves each partial pressure, and so Q, unchanged, so there is no shift.

  Example. Two rigid flasks hold the same equilibrium gas mixture, but one also holds 2 atm of unreactive nitrogen; both contain the same amounts of the reacting gases.

• Catalyst and equilibrium

  A catalyst speeds up the forward and reverse reactions equally, so equilibrium is reached sooner but its position and K do not change.

  Example. A metal surface lets a gas mixture reach equilibrium in minutes instead of days, with the same final amounts.

• Solubility product (Ksp)

  The equilibrium constant for a slightly soluble ionic solid dissolving into its ions; the solid is left out of the expression.

  Example. For $\text{PbI}_{2}(s) \rightleftharpoons \text{Pb}^{2+}(aq) + 2\text{I}^{-}(aq)$, $K_{sp} = [\text{Pb}^{2+}][\text{I}^{-}]^2$.

• Molar solubility

  The moles of a solid that dissolve per liter to make a saturated solution, written s; $K_{sp}$ is found from s using the ion ratio.

  Example. For a salt of formula $\text{MX}_3$, the ion concentrations are s and 3s, so $K_{sp} = s(3s)^3 = 27s^4$.

• Common-ion effect

  A salt is less soluble in a solution that already holds one of its ions, because the extra ion shifts the dissolving equilibrium back toward the solid.

  Example. Silver bromide dissolves less in a sodium bromide solution than in pure water.

• Q compared with Ksp

  For a solution of ions, if the ion product Q exceeds $K_{sp}$ a precipitate forms; if Q is below $K_{sp}$ the solution is unsaturated and more solid can dissolve.

  Example. Mixing two clear solutions that give an ion product above $K_{sp}$ makes the mixture turn cloudy with solid.

• pH and solubility

  A salt whose anion is a base, such as a hydroxide, fluoride or carbonate, dissolves more in acid, because the acid removes that anion from the equilibrium.

  Example. Limestone, mostly calcium carbonate, is slowly eaten away by acidic rain.