Soil 240 Chapter 1.1 Chemical Equilibria and Thermodynamics of Reactions PDF

Chapter 1 CHEMICAL EQUILIRIA AND THERMODYNAMICS OF REACTIONS 1.1. Chemical Equilibrium That stage in a chemical reaction when there is no further tendency for the composition of the reaction mixture to change. there is a state of balance between the concentrations of the reactants...

Chapter 1 CHEMICAL EQUILIRIA AND THERMODYNAMICS OF REACTIONS 1.1. Chemical Equilibrium That stage in a chemical reaction when there is no further tendency for the composition of the reaction mixture to change. there is a state of balance between the concentrations of the reactants and products  the concentrations of the reactants and products remain the same BUT the reactions don’t stop rate at which the rate at which the reactants change = products change to into products the original reactants Consider the hypothetical reversible reaction: A + B = AB If A and B are mixed they will react to produce AB. A sample of pure AB on the other hand will decompose to form A and B.  [AB] concentration [A] or [B] - - te time Figure 1.1 Curve showing changes in concentrations of materials for the reaction A + B = AB  Consider the synthesis of ammonia: N2(g) + 3H2(g) = 2NH3(g) The reaction reaches a steady state before all the reactants are consumed  H2 H2 concentration NH3 N2 NH3 N2 (a) (b) time Figure 1.2. Changes in the concentration of Nitrogen, Hydrogen and Ammonia with time Law of Mass Action The rate of a chemical reaction is directly proportional to the concentrations of the reacting substances In 1863, Norwegian chemists Cato Maximilian Guldberg and Peter Waage introduced the concept of equilibrium constant Consider the reaction: A + B = AB Rate of the forward reaction: Ratef α [A] and [B] Ratef = k1[A][B] where k1- specific rate constant of the forward reaction Rate at the reverse reaction is: Rater α[AB] Rater = k2[AB] where k2- specific rate constant of the reverse reaction At equilibrium, these two rates are equal and therefore: Ratef = Rater k1[A][B]=k2[AB] rearranging, k1 [AB] = k2 [A][B] [AB] K = => equilibrium constant [A][B] In general, for the reaction: aA + bB = cC + dD the equilibrium constant expression is: [C]c [D]d K = [A]a [B]b The convention is: substances on the right-hand side of the equation are written at the top of the K expression those on the left-hand side are written at the bottom IMPORTANT: Indicate phases of reactants and products i.e. gas (g), liquid (l), solid (s), aqueous (aq) Concentration of solids and pure liquid are constant Concentration of water is practically constant in dilute aqueous solutions Examples: Fe2O3 (s) + 2CO2 (g)= 2FeCO3 (s) + ½ O2 (g) 1/2 PO K = 2 2 PCO 2 Fe2O3 (s) + 3H2O(l) = 2Fe2+ + 6OH- (aq) (aq) K = [Fe 2+ ]2 [OH − ]6 Types of equilibria: A homogeneous equilibrium has everything present in the same phase. include reactions where everything is a gas, or everything is present in the same solution. example: Haber Process N2(g) + 3H2(g) = NH3(g)  A heterogeneous equilibrium has things present in more than one phase. include reactions involving solids and gases, or solids and liquids example: heating of CaCO3 CaCO3(s) = CaO(s) + CO2(g) Kinds of equilibrium constants: depends upon the units in which reactants and products are expressed Concentration constant, KC – units are expressed as concentration of substances, [ ] moles/liter Disadvantage: they change with ionic strength must be used in system of the same ionic strength in which they are determined ionic strength corrections must be applied  Kp for reactions involving gases where the concentrations of the reacting gases are expressed as partial pressures Example: 2NH3(g) = N2(g) + 3H2(g) PN2 P3 H2 Kp = P 2NH3 Kc and Kp: Since for an ideal gas: PV = nRT and concentration (molarity, M = mole/liter) is: M = n/V Then P = MRT Where R – gas law constant T – temperature, K  Kp = Kc (RT) Δng where Δng = number of moles of gaseous products minus the number of moles of gaseous reactants  Activity constants, Ka – expressed as activities, ( ) Significance: They can be calculated from thermodynamic data. They are true constants that hold for solutions of all ionic strengths Disadvantage: many reactants and products consist of specific ionic or molecular species whose activities are difficult or impossible to measure. For this reason, concentration is often used. ▪ General Significance of K: If K>>1, the products dominate the reaction mixture K