knc2133 chap 4 chemical equilibrium
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8/3/2019 KNC2133 Chap 4 Chemical Equilibrium
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Chapter 4
Chemical Equilibrium
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Learning ObjectivesCalculation of equilibrium constant.Activity and activity coefficients.
Systematic approach to equilibriumcalculations.Mass balance and charge balance
equations.
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Chemical Reactions: The Rate Concept
• Consider the chemical reaction
aA + bB cC + dD
• The rate forward is equal to
ratef = k f [A]a[B]b
where ratef is the rate of the forward reaction andk f is the rate constant
• Kf is dependent on factors such as temperature
and the presence of catalysts
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• The rate reverse is equal to
rateb = k b[C]c[D]d
where rateb is the rate of the reverse reaction and k b is
the rate constant
• For a system equilibrium, the forward and reverse
rates are equal
k f [A]
a
[B]
b
= k b[C]
c
[D]
d
• The molar equilibrium constant will be[ ] [ ]
[ ] [ ]
c d
f
a b
b
k C D K
A B k
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A and B disappearing
Equilibrium Concentration
C o n c e n t r a t i o n
Time
C and D appearing
0
Note:
The equilibrium constantdoes not provide any info.
on how fast the reaction
will occur
• K can be evaluated empirically by measuring the
concentration of A, B, C and D at equilibrium
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Fig. 6.1. Progress of a chemical reaction.
The rate of the forward reaction diminishes with time, while that of thebackward reaction increases, until they are equal.
A large K means the reaction lies far to the right at equilibrium.
©Gary Christian, Analytical Chemistry, 6th Ed. (Wiley)
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Equilibrium constants may be written for dissociations, associations, reactions
or distributions.
Equilibrium Reaction Equilibrium Constant
Acid-base
dissociation
HA + H2O H3O+ +
A-
Ka, acidity constant
Solubility MXMn+ + An- Ksp, solubility product
Complex formation Mn+ + aLb- MLa(n-
ab)+Kf, formationconstant
Reduction-oxidation Ared + Box Aox + Bred Keq, reaction eq.
constantPhase distribution AH2O Aorganic KD, distribution
coefficient
Types of Equilibria
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Gibbs Free Energy and Equilibrium Constant
• The tendency of a reaction to occur is defined
thermodynamically from its change in enthalpy (∆H) and
entropy (∆S).
• A system always favor lower energy and increased randomness,that is lower enthalpy and higher entropy.
• The combined effect is given by the Gibbs free energy, G:
G = H – TS
• The change in energy of the system at constant T:
ΔG = ΔH – TΔS
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• At standard condition (1 atm, 298 K)
ΔGo
= ΔHo
– TΔSo
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When stress is applied to a system at
chemical equilibrium, the
equilibrium will shift in a directionthat tends to relieve or counteract
that stress.
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What are the factors (stress) ?
• Temperature (T)
Affects the individual rate constants.
the equilibrium constant Also affects the free energy.
Depends on the magnitude of the heat of reaction of the
system.
Q: for endothermic reaction, in which direction the reaction
will occur if heat is added?
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• Pressure
Larger influence for reaction in gaseous phase
An increase of pressure will favor a shift in the
direction that results in a reduction in the volume of the
system. Negligible effect for reactions in liquid phase.
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• Concentration
Does not affect the value of equilibrium constant.
Affects the position of equilibrium for the reaction.
Consider:
3I- + 2Fe3- I3- + 2Fe2+
What will happen if Fe2+ is removed from the reactor?
What will happen if it is added to the reactor?
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• Catalysts
Catalysts do not affect the equilibrium constant or the
position of equilibrium.
• Completeness of Reactions
For quantitative analysis, equilibria should be at least
99.9% to the right for precise measurements.
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Chemical Equilibrium
Review of PrinciplesChemical reactions are never
“complete” Chemical reactions proceed to a state
where ratio of products to reactants isconstant
NH3 + HOH NH4+ + OH-
[NH4+][OH-]/[NH3][HOH] = Kb
o
If Kb << 1 (little ionization)
H2SO4 + HOH H3O+ + HSO4- [H3O+][HSO4
-] / [H2SO4][HOH] = KaIf Ka
>> 1 (mostly ionized)
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Chemical Equilibrium
Equilibrium
is not reached instantaneously can be approached from either direction is a dynamic state amounts of reactants/products can be changed
by “mass action”
(adding/ deleting products/reactants) HCO3
- + H+ CO2(g) + HOH
Ke = [CO2][HOH]/[HCO3-][H+]
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Chemical Equilibrium
Equilibrium Constants 2 A + 3 B C + 4 D
Ke = [C][D]4/[A]2[B]3
Concentrations [ ] : molar for solutes partial pressures (atm) for gases [1.0] for pure liquid, solid, or solvent
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Equilibrium constants for dissociating or
combining species
When a substance dissolves in water, it will often
partially or completely dissociate or ionize.
• Weak electrolytes: partially dissociate
Ex: acetic acid
• Strong electrolytes: completely dissociate
Ex: hydrochloric acid
The dissociation of weak electrolytes or the solubility of
slightly soluble substances can be quantitatively described
by equilibrium constant.
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Consider the dissociating species AB,
ABA + B
The equilibrium constant can be written as:
[ ] [ ]
[ ] eq
A B
K AB
The larger the Keq, the
greater will be the
dissociationSome species dissociate stepwise,
A2BA + AB
ABA + B
1
2
[ ][ ]
[ ]
A AB K
A B
2
[ ] [ ]
[ ]
A B K
AB
Overall:
Keq = K1.K2
= [A]2[B]/[A2B]
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Activity and Activity Coefficients
Activity of an ion,ai = Ciƒi
Ci = concentration of the ionƒi = activity coefficient ( @ Ci < 10-4M )= 1
Ionic Strength,
= ½ CiZi2
)Zi = charge on each individual ion
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Activity and Activity Coefficients
Calculation of Activity CoefficientsDebye-Huckel Equation:
-log ƒi = 0.51Zi2 ½
i½
i = ion size parameter in angstrom (Å)1 Å = 100 picometers (pm, 10-10 meters)
Limitations: singly charged ions = 3 Å
-log ƒi = 0.51Zi2 ½
½
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Systematic Approach to EquilibriumCalculations
How to Solve Any Equilibrium Problem?
1. Write balanced chemical reactions2. Write equilibrium constant expressions
3. Write all mass balance expressions4. Write the charge balance expression5. Equations >= Chemical Species solpossible6. Make assumptions where possible7. Calculate answer 8. Check validity of assumptions
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