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Chapter 6: Equilibrium Chemistry
Transcript of Chapter 6: Equilibrium Chemistry
Reversible Reactions and Thermodynamics
Solving Equilibrium Problems
Equilibrium Constants for Chemical Reactions
Reactions are thought of as proceeding in a forward and reverse direction. In other words:
When the rate of the forward reaction balances the reverse, it reaches a steady state
We say that the reaction has reached equilibrium
The amount of products and reactant will not change unless the equilibrium is disturbed
The direction in which a reaction moves in order to reach equilibrium is determined by the change in Gibbs Free Energy, G. If delta G is negative, the reaction moves in the forward direction (as written).
Delta G is given by:
Where reaction quotient, Q is:
At equilibrium, delta G = 0 and the reaction quotient becomes the equilibrium constant, K.
Thermodynamics and Equilibrium Chemistry
For a general equation:
When an equation is reversed:
If a reaction is the sum of two or more reactions, the individual equilibrium constants can be multiplied.
Manipulating Equilibrium Constants
Defined as two of more soluble chemicals reacting to form a solid product (precipitate)
The ability of a precipitate to dissolve is called the solubility product constant, Ksp
Using the Bronsted-Lowry definition an acid-base reaction involves the transfer of a proton (H+) from an acid to a base.
Strong acids and bases are considered to react complete (no equilibrium exists)
Strong Acid-Base Reactions
Weak acids and bases do not fully react with the solvent
The acid dissociation constant defines the strength of the acid
The base dissociation constant defines the strength of the base
Weak Acid-Base Reactions
Water can act as a acid and base
pH is defined as the –log[H3O+]
For a neutral solution:
For a conjugate acid base pair:
Autoionisation of Water and pH
Reactions involving the transfer of electrons
Oxidation involves the loss of electrons
Reduction involves a gain in electrons
The relative strengths of oxidation or reduction agents are found from their reduction potentials. A large reduction potential suggests a species with a large potential to gain electrons and remove an electrons from another species (therefore a good oxidizing agent)
The potential of a reaction at standard conditions is given by:
The potential is related to the equilibrium constant, K, according to the equation:
At non-standard conditions, we find the potential using the Nernst equation:
Reduction Potential and Equilibrium Constant
The equilibrium position of a reaction will move in order to relieve an applied stress
For example if a reactant is added to a reaction at equilibrium, it will produce more products in order to relieve the stress
The value of the equilibrium constant does not change
La Chatelier’s Principle
F- is the dominant species
When pNH3 is between 3.31 and 3.91
ICE Table Approach
Write the balanced chemical equation
Write the equilibrium constant expression
Draw an ICE (Initial, Change, Equilibrium) Table
Insert concentrations, using x for unknowns
Insert equilibrium concentrations into the equilibrium constant expression
Apply assumptions (if any)
Solve for x
A Systematic Approach
Write all relevant reactions and equilibrium constant expressions
Count unique species appearing in the equilibrium constant expressions. If the number of unknowns is great than the number of equilibrium expressions, add a
Combine equations and solve for one unknown making appropriate assumptions when possible
Check any assumptions
Reactions involving the formation of coordinative covalent bond between a metal ion and a ligand
A ligand is a species with one or more lone pairs of electrons
The product is called the complex
Complex formation is characterized by the formation constant, Kf.
Sn2+ & Fe2+
Sn4+ & Fe2+
Sn4+ & Fe3+
A statement of conservation of matter
The products must contain the same number and type of atoms as the reactants
e.g. when HF ionizes, the number of moles of F- and HF produced must equal the initial moles of HF. For a constant volume:
The charge of cations must equal the charge of anions
In other words, the solution must have an overall neutral charge
Equation must include every ion in solution.
E.g. for a solution of magnesium fluoride:
A buffer solution is one that resists a change.
For example, an acid/ base buffer will resist changes in pH due to addition of acid or base to the solution
Buffers can also exists for other types of reactions
Buffer problems can be solved using the ICE table approach or by applying the Henderson-Hasselbach Equation
Be aware that this equation includes assumptions.
Solving Buffer Problems
Buffer capacity is the ability of a buffer to resist a change in pH when adding a strong acid or base
The buffer capacity depends on the concentrations of the buffer components
Acid-base buffers are made by combining known amounts of acid and conjugate base or by titrating an acid or base in order to produce the conjugate
Ladder diagrams can be used to illustrate buffers
The diagram shows a) an acid-base buffer b) a complexation buffer and c) a redox buffer with the buffer regions indicated by grey boxes
If an equilibrium occurs in a solution containing other ions (not common ions) the measured concentration will be different to the calculated value
e.g. for the reaction:
In a potassium nitrate solution, the potassium and nitrate ions form an ionic atmospheres around the silver and iodate ions causing drag and slowing down the reverse reaction, shifting the equilibrium and making the silver iodate more soluble
The “effective concentration” of silver and iodate ions is called the
Equilibrium constants are really in terms of activities not concentrations
Activity is related to concentration by the
At infinite dilution the activity coefficient is one, and activity equals concentration
The activity coefficient can be calculated using the extended Debye-Huckel equation
is given by: