Chemical Equilibrium in Chemical Reactions.
defination of chemical equlibrium
the chemical reaction is the condition which occurs when the concentration of reactants and products participating in a chemical reaction exhibit no net change over time. Chemical equilibrium may also be called a "steady state reaction." This does not mean the chemical reactions has necessarily stopped occurring, but that the consumption and formation of substances have reached a balanced condition.
The quantities of reactants and products have achieved a constant ratio, but they are almost never equal. There may be much more product or much more reactant.
DYNAMIC EQUILIBRIUM
Dynamic equilibrium occurs when the chemical reaction continues to proceed, but a number of products and reactants remain constant. This is one type of chemical equilibrium.
WRITING THE EQUILIBRIUM EXPRESSION
The equilibrium expression for a chemical reaction may be expressed in terms of the concentration of the products and reactants. Only chemical species in the aqueous and gaseous phases are included in the chemical equation because the concentrations of liquids and solids does not change. For the chemical reaction:
jA + kB → lC + mD
The equilibrium expression is
K = ([C]l[D]m) / ([A]j[B]k)
K is the equilibrium constant
[A], [B], [C], [D] etc. are the molar concentrations of A, B, C, D etc.
j, k, l, m etc. are coefficients in a balanced chemical equation
FACTORS THAT AFFECT CHEMICAL EQUILIBRIUM
First, consider a factor that does not affect equilibrium: pure substances. If a pure liquid or solid is involved in equilibrium, it is considered to have an equilibrium constant of 1 and is excluded from the equilibrium constant. For example, except in highly concentrated solutions, pure water is considered to have an activity of 1.
Another example is solid carbon, which may be form by the reaction of two carbom monoxide molecules to form carbon dioxide and carbon.
Factors that do affect equilibrium include:
- Adding reactant or product or a change in concentration affects equilibrium. Adding reactant can drive equilibrium to the right in a chemical equation, where more product forms. Adding product can drive equilibrium to the left, as more reactant forms.
- Changing the temperature alters equilibrium. Increasing temperature always shifts chemical equilibrium in the direction of the endothermic reaction. Decreasing temperature always shifts equilibrium in the direction of the exothermic reaction.
- Changing the pressure affects equilibrium. For example, decreasing the volume of a gas system increases its pressure, which increases the concentration of both reactants and products. The net reaction will see to lower the concentration of gas molecules.
Le Chatelier's principle -
may be used to predict the shift in equilibrium resulting from applying a stress to the system. Le Chatelier's principle states that a change to a system in equilibrium will cause a predictable shift in equilibrium to counteract the change. For example, adding heat to a system favors the direction of the endothermic reaction because this will act to reduce the amount of heat
CHEMICAL EQUILIBRIUM
Synopsis:
• Based on the direction of their occurrence Chemical reactions are two types.
Irreversible reactions: in these reactions reactants are converted into products and products cannot be
converted into reactants.
• There are unidirectional as they occur in one direction. i.e.
Reactants → products.
• These are denoted by single arrow mark.
• These reactions almost go for completion i.e. reactants are almost completely convert into products.
• Precipitation – ionic reactions, explosive reactions, strong acid – strong base neutralisation reactions,
combustion- reactions are irreversible.
Eg :1) 2KClO3(s)→2KCl(s) + 3O2(g)
2) NH4NO2(s) → N2(g) + 2H2O(g)
3) C2H5OH(l) → 3O2(g) → 2CO2(g) + 3H2O(g)
4) 2Mg(s) + O2(g) → 2MgO (s)
5) HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
6) H2(g) + F2(g) → 2HF(g)
7) H2(g) + Cl2(g) → 2HCl(g)
Reversible reactions :
• A reaction is said to be reversible, if both the forward and the backward reactions are taking place
simultaneously under the given experimental conditions.
• Reactants giving rise to products is known as forward reaction.
• Products giving rise to reactants is known as reverse reaction (or) backward reaction.
• Reversible reactions are represented by writing a pair of half headed – arrows pointing in opposite directions
in between the reactants and products.
Reactants products
• A reverse reaction does not go to completion.
• Most of the reversible reactions are carried in the closed vessels.
Eg : 1) H2(g) + I2(g) ⇔ 2HI(g)
2) PCl5(g) ⇔ PCl3(g) + Cl2(g)
3) 2NO2(g) ⇔ N2O4(g)
4) N2(g) + O2(g) ⇔ 2NO(g)
5) 2SO2(g) + O2(g) ⇔ 2SO3(g)
6) CaCO3(s) ⇔ CaO(s) + CO2(g)
7) CH3COOH(l) + C2H5OH(l) ⇔ CH3COOC2H5(l) H2O(l)
Equilibrium State :
• The stage or state at which the rate of forward reaction is equal to the rate of the reverse reaction in a
reversible reaction is known as the equilibrium stage or state.
• The chemical equilibrium is considered as dynamic equilibrium because the forward and the reverse reactions Chemical Equilibrium continue to take place simultaneously at this stage also Equilibrium is established in
i) a reversible reaction ii) in a closed vessel
• In the beginning in a reversible reaction the rate of forward reaction is more since the concentration of
reactants is more.
• As time proceeds the rate of forward reaction decreases as the concentrations of reactants decreases.
• In the beginning in a reversible reaction the rate of backward reaction is absolutely zero because the
concentration of products is zero.
• AS time proceeds the rate of backward reaction increases since the concentrations of products also increases.
• At one stage the rate of forward reaction becomes equal to rate of backward reaction and no further change
occurs in the concentration of reactions or products.
• At equilibrium, the concentration of reactants and products may not be equal but they remain constant.
Characteristics of chemical equilibrium :
• The rate of the forward reaction is equal to the rate of the reverse reaction.
• The concentrations of the reactants and the products remain unchanged with time.
• The observable properties such as pressure, concentration, density, colour also remain unchanged with time.
• The attainment of chemical equilibrium can be recognised by the constancy in some macroscopic potteries
like pressure, concentration, density, color etc.
• The equilibrium is dynamic in nature. Both the forward reaction and the reverse reaction continue to take place
simultaneously with equal rates.
• A catalyst does not alter the state of equilibrium or the composition of the chemical substances of the reaction
at the equilibrium. It only speeds up the attainment of the equilibrium.
• Chemical equilibrium can be established from either side of the reversible reaction.
• Chemical equilibrium can be homogeneous or heterogeneous and also ionic or molecular.
• The factors such as pressure, concentration temperature presence of inert gas influence the position of the
equilibrium.
• At equilibrium, the value of Gibbs free energy change (ΔG) is zero (ΔG = 0).
• At equilibrium, ΔS is maximum.
• The equilibrium does not tell us how long it takes for a reaction to attain equilibrium.
• Once equilibrium is reached, it continues forever until the conditions like pressure, temperature, concentration
etc, are altered.
• At equilibrium the concentration of reactants may be equal or less or more than the concentration of products.
Law of mass action :
• This laws stated by C.M. Guldberg and P. Wage in 1863.
• This law gives the relation between the rate of a reaction and the concentration of the reactants.
• The rate of a chemical reaction at a temperature at any instant is proportional to the product of the active
masses of the reactants.
• This law is applicable to all reactions i.e. reversible and irreversible occurring in the gas phase or in the liquid
phase.Chemical Equilibrium
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• aA + bB ⇔ cC + dD, the equilibrium constant.
Kc= Kf/Kb [C]c[D]d/[A]a[B]b
Kf = forward reaction rate constant.
Kb = backward reaction rate constant.
• The equilibrium constant
KC=product of the concentration of reactants/product of the concentration of products.
Partial pressure of the gas = mole fraction of gas × total pressure.
K=KF/KB=Pcc.PDd/PAa.PBb
• product of partialpressures of reac tants
prodctof patial pressures of products Kp =
kc= equilibrium constant in terms of molar concentration.
kp=equilibrium constant in terms of partial pressure.
Active mass =no.of moles/ volume in litres
Active mass is considered for gas or liquid.
• The active mass of a solid is unity whatever may be its mass.
Types of chemical equilibrium:
• Based on the physical states of substances equilibrium is of two types.
1) Homogeneous equilibrium : All the reactants and products are present in same physical state. i.e same
phase.
Eg :1) 2SO2(g)+ O2⇔ 2SO3(g)
2) N2(g) + 3H2(g) ⇔ 2NH3(g)
3) CH3COOC2H5(l) + H2O(l)⇔ CH3COOH(l) + C2H5OH(l)
4) CH3COOH-(l) ⇔ +H+
2) Heterogeneous equilibrium : Reactants and products are in different physical states or different phase.
Eg : 1) CaCO3(s)⇔ CaO(s) + CO2(g)
2) NH4HS(s)⇔ NH3(g)+ H2S(g)
3) Fe(s)+ 4H2O(g)⇔ Fe3O4(s) + 4H2(g)
• Relationship between kp and kc:
kp = kc (RT)Δn
R = gas constant, T = absolute temperature
Δn = change in number of moles
= nP - nR (no.of moles of gaseous products – no.of moles of gaseous reactants )
case (i) if nP = nR, Δn = 0, kp = kc
Eg. H2+ I2 ⇔ 2HI
CASE (ii) if nP > nR, Δn = +ve, kp > kc
Chemical Equilibrium
PCl5 ⇔ PCl3 + Cl2
iii) If nP < nR, Δn = -ve, kp < kc
N2 + 3H2 ⇔ 2NH3
• Units of equilibrium constant :
Unit of kc = (mol. lit–1)
Δn
Unit of kp = (atmosphere) Δn
EX-
i) CaCO3(s) CaO(s) + CO2(g)
kc = [CO2] ; kc = mol. lit–1
kp = CO2 P ; kp = atm
Characteristics of equilibrium constant : (kp or kc)
• The value of k depends on the nature of the reaction.
• The value of k will be a constant for a given reaction at a given temperature.
• The value of k depends on temperature of reaction.
• The value of k is independent of concentration and pressure.
• The value of k is independent of presence of catalyst and presence of inert gas.
• The value of k depends on stoichiometry of the equation.
• The value of k depends on mode of writing the equilibrium reaction.
Lechattelier’s principle and its application to industrial processes ;
• The effect of change of pressure, concentration and temperature on equilibrium was studied by Henry Lewis
Lechatterlier in 1885 and F. Brawn. So this is called Lechattelier- Brawn Principle or simply Lechattelier’s
principle.
• If a system at equilibrium is subjected to a stress, the system shifts the equilibrium in such a way to reduce or
nullify the stress.
• Effect of concentration : Increase in concentration of reactants or decrease in concentration of products
favours the shift of equilibrium towards products side and rate of forward reaction is increased.
Increase in the concentration of the products or decrease in the concentration of reactants favours the shift ofChemical Equilibrium
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equilibrium towards the reactant side and rate of backward reaction is increased.
Effect of pressure :
• Pressure has no effect on equilibrium if Δv or Δn=0, (np= nr).
Eg : H2(g) + I2(g) ⇔ 2HI(g)
• Pressure has effect on equilibrium if Δv≠0 or Δn≠0, (np≠ nr). When pressure increases, equilibrium shifts in a
direction of decrease of volume or towards less mole number and vice vessa.
Eg: N2(g) + 3H2(g) ⇔ 2NH3(g)
When pressure increases, equilibrium shifts towards right and forward reaction rate increases.
• Pressure does not show any marked effect on equilibrium reactions taking place in the solution phase or in the
solid phase.
• Effect of temperature : increase of temperature of the equilibrium system favours endothermic reactions, and
decrease of temperature of the equilibrium system favours exothermic reactions.
• Effect of Catalyst : Catalyst has no net effect on equilibrium. It helps the system to attain equilibrium at a
faster rate by increasing the rate of forward as well as backward reaction to the same extent.
Eg:(1) Synthesis of ammonia by Haber’s process:
N2(g) + 3H2(g) ⇔ 2NH3(g) + heat ΔH = – 92.0 K.J
Favourable conditions for high yield of NH3:
High pressure:200atm, Catalyst : Fe as catalyst
Low temperature : 773 K, Promoter: small amount of molybdenum or Al2O3 and K2O.
(2) Manufacture of H2SO4 by the contact process.
2SO2(g) + O2(g) ⇔ 2SO3(g) + Heat; ΔH = –189 K.j
• Favourable conditions for higher yield of SO3.
I) High pressure : 1.5 – 1.7 atm.
ii) Low temperature : 673 k
iii) Catalyst : V2O5 or platinised asbestos
N2(g) + O2(g)⇔ 2NO(g) – heat
i) High temperature ii) No effect of pressure
Melting of ice :
H2O(s)+ heat⇔ H2O
I) high temperature ii) high pressure
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