Chapter 18.1 : The Nature of Chemical Equilibrium


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Chapter 18.1 : The Nature of Chemical Equilibrium

  1. 1. Chapter 18.1
  2. 2. 1. Define chemical equilibrium. 2. Explain the nature of the equilibrium constant. 3. Write chemical equilibrium expressions and carry out calculations involving them.
  3. 3. • Theoretically, every reaction can proceed in two directions, forward and reverse. • Essentially all chemical reactions are considered to be reversible under suitable conditions. • A chemical reaction in which the products can react to re-form the reactants is called a reversible reaction.
  4. 4. • A reversible chemical reaction is in chemical equilibrium when the rate of its forward reaction equals the rate of its reverse reaction and the concentrations of its products and reactants remain unchanged. • A state of dynamic equilibrium has been reached when the amounts of products and reactants remain constant. • Both reactions continue, but there is no net change in the composition of the system. • The chemical equation for the reaction at equilibrium is written using double arrows to indicate the overall reversibility of the reaction.  22HgO( ) 2Hg( ) O ( )s l + g
  5. 5. • Many chemical reactions are reversible under ordinary conditions of temperature and concentration. • They will reach a state of equilibrium unless at least one of the substances involved escapes or is removed from the reaction system. 1. When the products of the forward reaction are favored, there is a higher concentration of products than of reactants at equilibrium. • The equilibrium “lies to the right”
  6. 6. 2. When the products of the reverse reaction are favored, there is a higher concentration of reactants than of products at equilibrium. • the equilibrium “lies to the left” 3. In other cases, both forward and reverse reactions occur to nearly the same extent before chemical equilibrium is established. • Neither reaction is favored, and considerable concentrations of both reactants and products are present at equilibrium.
  7. 7. • products of the forward reaction favored, lies to the right 2 2 32SO ( ) O ( ) 2SO ( )g + g g  – 2 3 2 3 3H CO ( ) H O( H O ( ) HCO ( )aq + l) aq + aq  – 2 3 2 3 3H SO ( ) H O( ) H O ( ) HSO ( )aq + l aq + aq Neither reaction is favored • products of the reverse reaction favored, lies to the left
  8. 8. • Initially, the concentrations of C and D are zero and those of A and B are maximum. • Over time the rate of the forward reaction decreases as A and B are used up. • The rate of the reverse reaction increases as C and D are formed. • When these two reaction rates become equal, equilibrium is established.  A B C Dn m x y
  9. 9. • After equilibrium is reached, the individual concentrations of A, B, C, and D undergo no further change if conditions remain the same. • A ratio of their concentrations should also remain constant. • The equilibrium constant is designated by the letter K.  A B C Dn m x y x y n m K [C] [D] [A] [B] 
  10. 10. • The constant K is independent of the initial concentrations. • K is dependent on the temperature of the system. The Equilibrium Constant • The numerical value of K for a particular equilibrium system is obtained experimentally. • If K is equal to 1 at equilibrium, there are roughly equal concentrations of reactants and products.
  11. 11. The Equilibrium Constant, continued • If the value of K is small, the reactants are favored. • A large value of K indicates that the products are favored. • Only the concentrations of substances that can actually change are included in K. • Pure solids and liquids are omitted because their concentrations cannot change.
  12. 12. • The equilibrium constant, K, is the ratio of the mathematical product of the concentrations of substances formed at equilibrium to the mathematical product of the concentrations of reacting substances. • Each concentration is raised to a power equal to the coefficient of that substance in the chemical equation. • The equation for K is sometimes referred to as the chemical equilibrium expression.
  13. 13. • The rate of the reaction between H2 and I2 vapor in a sealed flask at an elevated temperature can be followed by observing the rate at which the violet color of the iodine vapor diminishes. • The color fades to a constant intensity but does not disappear completely because the reaction is reversible. • Hydrogen iodide decomposes to re-form hydrogen and iodine. • The constant color achieved indicates that equilibrium exists among hydrogen, iodine, and hydrogen iodide.
  14. 14. The net chemical equation for the reaction is 2 2H ( ) I ( ) 2HI( )g + g g K 2 2 2 [HI] [H ][I ]  • The value for K is constant for any system of H2, I2, and HI at equilibrium at a given temperature. The following chemical equilibrium expression is • At 425°C, the equilibrium constant for this equilibrium reaction system has the average value of 54.34.
  15. 15. • The balanced chemical equation for an equilibrium system is necessary to write the expression for the equilibrium constant. • Once the value of the equilibrium constant is known, the equilibrium constant expression can be used to calculate concentrations of reactants or products at equilibrium.
  16. 16. An equilibrium mixture of N2, O2 , and NO gases at 1500 K is determined to consist of 6.4  10–3 mol/L of N2, 1.7  10–3 mol/L of O2, and 1.1  10–5 mol/L of NO. What is the equilibrium constant for the system at this temperature? 2N ( ) ) 2NO( )2g + O (g g K 2 2 2 [NO] [N ][O ]  Given: [N2] = 6.4  10–3 mol/L [O2] = 1.7  10–3 mol/L [NO] = 1.1  10–5 mol/L Solution: The balanced chemical equation is Unknown: K The chemical equilibrium expression is
  17. 17. K 5 2 3 3 5(1.1 10 mol/L) (6.4 10 mol/L)(1.7 10 mol/ 1.1 10 L)          Reactants are favored