Value of Dn | Relation between Kp and Kc | Units of Kp | Units of Kc |
0 | Kp = Kc | No unit | No unit |
>0 | Kp > Kc | (atm)Dn | (mole l–1)Dn |
<0 | Kp < Kc | (atm)Dn | (mole l–1)Dn |
\[\Delta n=0\,;\,\,{{K}_{p}}={{K}_{c}}\] | \[\Delta n<0\] ; \[{{K}_{p}}<{{K}_{c}}\] | \[\Delta n>0;\ {{K}_{p}}>{{K}_{c}}\] | ||
\[\underset{(g)}{\mathop{{{H}_{2}}}}\,\]+ \[\underset{(g)}{\mathop{{{I}_{2}}}}\,\] ? \[\underset{(g)}{\mathop{2HI}}\,\] | \[\underset{(g)}{\mathop{{{N}_{2}}}}\,+\underset{(g)}{\mathop{3{{H}_{2}}}}\,\]? \[\underset{(g)}{\mathop{2N{{H}_{3}}}}\,\] | \[\underset{(g)}{\mathop{2S{{O}_{2}}}}\,+\underset{(g)}{\mathop{{{O}_{2}}}}\,\]?\[2\underset{(g)}{\mathop{S{{O}_{3}}}}\,\] | \[\underset{(g)}{\mathop{PC{{l}_{_{5}}}}}\,\]?\[\underset{(g)}{\mathop{PC{{l}_{3}}}}\,+\underset{(g)}{\mathop{C{{l}_{2}}}}\,\] | |
Initial mole | 1 1 0 | 1 3 0 | 2 1 0 | 1 0 0 |
Mole at Equilibrium | (1–x) (1– x) 2x | (1–x) (3–3x) 2x | (2–2x) (1–x) 2x | (1–x) x x |
Total mole at equilibrium | 2 | (4 – 2x) | (3 – x) | (1 + x) |
Active masses | \[\left( \frac{1-x}{V} \right)\] \[\left( \frac{1-x}{V} \right)\] \[\frac{2x}{V}\] | \[\left( \frac{1-x}{V} \right)\] \[3\,\left( \frac{1-x}{V} \right)\] \[\left( \frac{2x}{V} \right)\] | \[\left( \frac{2-2x}{V} \right)\] \[\left( \frac{1-x}{V} \right)\] \[\left( \frac{2x}{V} \right)\] | \[\left( \frac{1-x}{V} \right)\] \[\left( \frac{x}{V} \right)\] \[\left( \frac{x}{V} \right)\] |
Mole fraction | \[\left( \frac{1-x}{2} \right)\] \[\left( \frac{1-x}{2} \right)\] \[\frac{2x}{2}\] | \[\frac{1-x}{2\,\left( 2-x \right)}\]\[\frac{3}{2}\left( \frac{1-x}{2-x} \right)\]\[\frac{x}{(2-x)}\] | \[\left( \frac{2-2x}{3-x} \right)\] \[\left( \frac{1-x}{3-x} \right)\,\,\ \ \left( \frac{2x}{3-x} \right)\] | \[\left( \frac{1-x}{1+x} \right)\] \[\left( \frac{x}{1+x} \right)\] \[\left( \frac{x}{1+x} \right)\] |
Partial pressure |
\[p\,\left( more...
Factors which Change the State of Equilibrium: Le-Chatelier's Principle.
Le-Chatelier and Braun (1884), French chemists, made certain generalizations to explain the effect of changes in concentration, temperature or pressure on the state of system in equilibrium. When a system is subjected to a change in one of these factors, the equilibrium gets disturbed and the system readjusts itself until it returns to equilibrium. The generalization is known as Le-Chatelier's principle. It may stated as:
“Change in any of the factors that determine the equilibrium conditions of a system will shift the equilibrium in such a manner to reduce or to counteract the effect of the change.”
The principle is very helpful in predicting qualitatively the effect of change in concentration, pressure or temperature on a system in equilibrium. This is applicable to all physical and chemical equilibria.
(1) Effect of change of concentration : According to Le-Chatelier's principle, “If concentration of one or all the reactant species is increased, the equilibrium shifts in the forward direction and more of the products are formed. Alternatively, if the concentration of one or all the product species is increased, the equilibrium shifts in the backward direction forming more reactants.”
Thus,
Increase in concentration of any of the reactants \[\underset{equilibrium\,\,to}{\mathop{\xrightarrow{Shifts\,\,the}}}\,\] Forward direction
Increase in concentration of any of the products \[\underset{equilibrium\,\,to}{\mathop{\xrightarrow{Shifts\,\,the}}}\,\] Backward direction
(2) Effect of change of temperature : According to Le-Chatelier's principle, “If the temperature of the system at equilibrium is increased (heat is supplied), the equilibrium will shift in the direction in which the added heat is absorbed. In other words, the equilibrium will shift in the direction of endothermic reaction with increase in temperature. Alternatively, the decrease in temperature will shift the equilibrium towards the direction in which heat is produced and, therefore, will favour exothermic reaction.”
Thus,
Increase in temperature \[\underset{in\,\,the\,\,direction\,\,of}{\mathop{\xrightarrow{Shifts\,\,the\,\,equilibrium}}}\,\] Endothermic reaction
Decrease in temperature \[\underset{in\,\,the\,\,direction\,\,of}{\mathop{\xrightarrow{Shifts\,\,the\,\,equilibrium}}}\,\] Exothermic reaction
(3) Effect of change of pressure : Pressure has hardly effect on the reactions carried in solids and liquids. However, it does influence the equilibrium state of the reactions that are carried in the gases. The effect of pressure depends upon the number of moles of the reactants and products involved in a particular reaction. According to Le-Chatelier's principle, “Increase in pressure shifts the equilibrium in the direction of decreasing gaseous moles. Alternatively, decrease in pressure shifts the equilibrium in the direction of increasing gaseous moles and pressure has no effect if the gaseous reactants and products have equal moles.”
Thus,
Increase in pressure \[\underset{in\,\,the\,\,direction\,\,of}{\mathop{\xrightarrow{Shifts\,\,the\,\,equilibrium}}}\,\] Decreasing gaseous moles
Decrease in pressure \[\underset{in\,\,the\,\,direction\,\,of}{\mathop{\xrightarrow{Shifts\,\,the\,\,equilibrium}}}\,\] Increasing gaseous moles
(4) Effect of volume change : We know that increase in pressure means decrease in volume, so the effect of change of volume will be exactly reverse to that of pressure. Thus, “decreasing the volume of a mixture of gases at equilibrium shifts the equilibrium in the direction of decreasing gaseous moles while increasing the volume shifts the equilibrium in the direction of increasing more...
The Le-Chateliers principle has a great significance for the chemical, physical systems and in every day life in a state of equilibrium.
(1) Applications to the chemical equilibrium
(i) Synthesis of ammonia (Haber’s process)
\[\underset{1\ vol}{\mathop{{{N}_{2}}}}\,+\underset{3\ vol}{\mathop{3{{H}_{2}}}}\,\] \[\rightleftharpoons \] \[\underset{2\ vol}{\mathop{2N{{H}_{3}}}}\,+23kcal\] (exothermic)
(a) High pressure \[(\Delta n<0)\]
(b) Low temperature
(c) Excess of \[{{N}_{2}}\] and \[{{H}_{2}}\]
(d) Removal of \[N{{H}_{3}}\] favours forward reaction.
(ii) Formation of sulphur trioxide
\[\underset{2\ vol}{\mathop{2S{{O}_{2}}}}\,+\underset{1\ vol}{\mathop{{{O}_{2}}}}\,\] \[\rightleftharpoons \] \[\underset{2\ vol}{\mathop{2S{{O}_{3}}}}\,+45\ kcal\] (exothermic)
In the following reversible chemical equation.
\[A\] \[\rightleftharpoons \] \[yB\]
Initial mole 1 0
At equilibrium (1–x) yx x = degree of dissociation
Number of moles of \[A\] and \[B\] at equilibrium \[=1-x+yx=1+x(y-1)\]
If initial volume of 1 mole of A is V, then volume of equilibrium mixture of \[A\] and \[B\] is,\[=[1+x(y-1)]V\]
Molar density before dissociation,
\[D=\frac{\text{molecular}\ \text{weight}}{\text{volume}}=\frac{m}{V}\]
Molar density after dissociation, \[d=\frac{m}{[1+x(y-1)]V}\];\[\frac{D}{d}=[1+x(y-1)]\] ; \[x=\frac{D-d}{d(y-1)}\]
\[y\] is the number of moles of products from one mole of reactant. \[\frac{D}{d}\] is also called Van’t Hoff factor.
In terms of molecular mass,\[x=\frac{M-m}{(y-1)\,m}\]
Where \[M=\] Initial molecular mass,
\[m=\] molecular mass at equilibrium
Thus for the equilibria
(I) \[PC{{l}_{5(g)}}\] \[\rightleftharpoons \] \[PC{{l}_{3(g)}}+C{{l}_{2(g)}},y=2\]
(II) \[{{N}_{2}}{{O}_{4(g)}}\] \[\rightleftharpoons \] \[2N{{O}_{2(g)}},\ y=2\]
(III) \[2N{{O}_{2}}\] ? \[{{N}_{2}}{{O}_{4}},\ y=\frac{1}{2}\]
\[\therefore \] \[x=\frac{D-d}{d}\] (for I and II) and \[x=\frac{2(d-D)}{d}\] (for III)
Also \[D\times 2=\] Molecular weight (theoretical value)
\[d\times 2=\] Molecular weight (abnormal value) of the mixture.
Substances, which allow electric current to pass through them, are known as conductors or electrical conductors. Conductors can be divided into two types,
(1) Conductors which conduct electricity without undergoing any chemical change are known as metallic or electronic conductors.
(2) Conductors which undergo decomposition (a chemical change) when an electric current is passed through them are known as electrolytic conductors or electrolytes.
Electrolytes are further divided into two types on the basis of their strengths,
(i) Substances which almost completely ionize into ions in their aqueous solution are called strong electrolytes. Degree of ionization for this type of electrolyte is one i.e., \[\alpha \approx 1\].
For example : \[HCl,\ {{H}_{2}}S{{O}_{4}},\ NaCl,\ HN{{O}_{3}},\ KOH,\ \]\[NaOH,\ \] \[HN{{O}_{3}},AgN{{O}_{3}},\ CuS{{O}_{4}}\] etc. means all strong acids, bases and all types of salts.
(ii) Substances which ionize to a small extent in their aqueous solution are known as weak electrolytes. Degree of ionization for this types of electrolytes is \[\alpha <<<1\].
For example :\[{{H}_{2}}O,\ C{{H}_{3}}COOH,\ N{{H}_{4}}OH,\ \]HCN, \[Liq.\ S{{O}_{2}}\], \[HCOOH\] etc. means all weak acids and bases.
Current Affairs CategoriesArchive
Trending Current Affairs
You need to login to perform this action. |