Plus One Chemistry Notes Chapter 7 Equilibrium is part of Plus One Chemistry Notes. Here we have given Kerala Plus One Chemistry Notes Chapter 7 Equilibrium.
|Text Book||NCERT Based|
|Category||Plus One Kerala|
Kerala Plus One Chemistry Notes Chapter 7 Equilibrium
In a reversible reaction, a stage is reached when the concentrations of reactants and products stay unchanged. This state is called as the state of equilibrium.
At equilibrium, the concentrations of reactants and products may or nay not be equal but must remain unchanged with time.
The mixture of reactants and products in the equilibrium state is called an equilibrium mixture.
- Irreversible Reactions. The reactions in which products donot react to form reactants back are called irreversible reaction.
eg., AgNo3(aq)+NaCI(aq) → AgCI(s)+NaNo3(aq)
- Reversible Reactions. The reactions in which products react to form reactants back are called reversible reaction.
eg., H2(S) + l2 (g) → 2HI (g)
Equilibrium in Physical Processes
Consider a perfectly insulated thermos flask containing some ice and water at 273 K. Since the flask is insulated, there will be no exchange of heat between its contents and the surroundings. It is seen that as long as the temperature remains constant, there is no change in the mass of ice and water. This represents an equilibrium state between ice and water.
Consider a closed vessel connected to a manometer. The water vapour present in the vessel is first removed by placing some drying agent such as anhydrous calcium chloride in it. The drying agent is then removed. Now the level of mercury in both the limbs of the manometer will be same. Introduce some water into the vessel and allow to stay at room temperature.
Now water starts evaporating. A pressure will gradually develop within the vessel due to the formation of water vapour. The change of pressure can be easily measured from the manometer. As evaporation continues, the pressure goes on increasing and the level of mercury in the right limb of the manometer starts rising. After sometime it is observed that pressure becomes constant. This shows that the quantity of water vapour is not increasing any more, although liquid water is still present in the vessel. This indicates that a state of dynamic equilibrium has been attained between liquid water and water vapour.
Equilibrium Involving Dissolution of Solid or Gas in Liquids
Solids in Liquids. In a saturated solution, a dynamic equilibrium exists between the solute molecules in the solid state and in the solution: the rate of dissolution of sugar = rate of crystallisation of sugar.
Gases in Liquids. This equilibrium is governed by Henry’s law, Which states that the mass of a gas dissolved in a given mass of a solvent at any temperature is proportional to the pressure of the gas above the solvent.
Dynamic Equilibrium. It represents a permanent situation where two exactly opposite changes takes place at the same rate. At equilibrium both forward and backward reactions take place at the same rate does not stop. Hence called dynamic equilibrium.
Law of Chemical Equilibrium
The quantitative relationship between the amount of the reactants and products in equilibrium at a constant temperature was first established by two Norwegean chemists G.M Guldberg and P. Waage in 1864. This equilibrium equation ius called law of chemical equilibrium.
Law of Mass action
It states that “at a constant temperature and pressure, the rate of a chemical reaction is directly proportional to the product of the molar concentrations of the reactants each raised to the power equal to the corresponding stoichiometric coefficient which appear in the balanced chemical equation”.
Equilibrium constant, K
It is defined as the ratio of the product of the equilibrium concentrations of the products to that of the reactants with each concentration term raised to the power equal to the number of moles of that substance in the balanced equation.
Homogeneous & Heterogeneous Equilibria
In a homogeneous system, all the reactants and products are in the same phase.
Equilibrium in a system having more than one phase is called hetrogeneous equilibrium.
Predicting the Direction of the Reaction
The equilibrium constant is also used to find in which direction the reaction mixture of reactants and products will proceed. For this, we have to calculate the reaction quotient (Q) and compare with the equilibrium constant (K). the concentrations of the species in Qc are not necessarily equilibrium values.
- If Q > K, the reaction will proceed in the direction of the reactants (i.e., reverse reation).
- If Q < K, the reaction will proceed in the direction of the products (i.e., forward reaction).
- If Q = K, the reaction mixture is already at equilibrium.
Application of Equilibrium constant
a. Prediction of extent of reaction
Larger the value of K, greater is the extent to which the reactants are converted into the products.
- If K>103, the product predominates over reactants at equilibrium. If K is very large the reaction proceeds nearly to completion
- If K<103, the product predominates over product. If K is very small the reaction proceeds to very small extend or rarely.
- If K is in the range of 10-3 to 103, appreciable concentrations of both reactants and products are present.
b. Prediction of direction of a reaction
The reaction quotient in terms of concentration or concentration Quotient (Q) with molar concentrations (Qc) and with partial pressures (Qp) is related to the equilibrium constant (K) for the reaction.
For a general reaction, aA + bB cC +dD
- If Qc > kc, the reaction will proceed in the direction of reactants (reverse reaction)
- If Qc < kc, the reaction will proceed in the direction of the products (forward reaction)
- If Qc= K, the reaction is at equilibrium.
Le Chatelier’s Principle (Factors affecting equilibria)
It states that “If we change either pressure, temperature or concentration in a system in equilibrium the system adjusts itself to cancel the effect of the constraint.”
Factors affecting Equilibria:
- Effect of concentration change
- Effect of pressure change
- Effect of inert gas addition at constant volume or pressure
- Effect of temperature change
- Effect of catalyst
Ionic Equilibrium in Solution
Michael Faraday classified the substances into two categories based on their ability to conduct electricity. One category of substances conduct electricity in their aqueous solutions and are called electrolytes while the other do not and are thus, referred to as non-electrolytes
2NO(g) + O2 (g) 2NO2(g)ΔH=-117kJ
i. Predict the effect of an increase in concentration of NO on the equilibrium concentration of NO2
ii. Predict the effect of pressure decrease as a result of increased volume on the equilibrium concentration of NO2
2NO(g) + O2(g) 2NO2(g) ΔH=-117kJ
i. If we increase the concentration of NO, the rate of forward reaction will increase, i.e., more NO2 will be formed.
ii. Decrease in pressure will favour back-ward reaction, i.e., less NO2 will be formed.
Electrolytes and Non-electrolytes
Substances which conduct electricity either in the molten or in the aqueous state and simultaneously undergo chemical change, eg., all acids, bases and salts.
Substances which do not conduct electricity and do not undergo any chemical decomposition on the passage of an electric current, eg., benzene, CCI4, sugar etc. Faraday further classified electrolytes into strong and weak electrolytes.
It is a substance which ionizes almost completely in the aqueous solution,
eg., HCI, NaOH.
It is a substance which ionizes to a small extent in aqueous solution,
eg., CH3COOH, NH4OH.
Arrhenius Concept of Acids and Bases
Acid is a substance which dissociate to give H+ ion in aqueous solution.
Base is a substance which dissociate to give OH’ ion in aqueous solution,
eg., HCI, H2SO4, CH3COOH etc. are acids which give free H+ ions in aqueous solution.
Limitations of Arrhenius concept
- It is applicable only to aqueous solution.
- Many compounds which act as acid or base do not contain H+ ion or OH ion. eg., CO2 (acid), NH3 (base).
Bronsted – Lowry Concept of Acids and Bases
- Acid is a proton donor.
- Base is a proton acceptor.
Conjugate acid-base pair
The base formed by the loss of proton from an acid is called conjugate base of the acid, eg., CL is the conjugate base of the acid HCI.
The acid formed by gain of proton by a base is called conjugate acid of the base.
acid = conjugate base + H+
base + H+ = conjugate acid
Thus each conjugate acid has one extra proton and each conjugate base has one proton less.
eg., H3O+ is the conjugate acid of the base H2o.
Limitation of Bronsted – Lowry concept
- It is only applicable to aqueous and non- aqueous protic solvents
- It give stress on the proton only.
- It cannot account for the acidic (CO2, SO2) and basic (CaO, Na2O) nature of oxides which do not contain hydrogen.
Lewis Concept of Acids and Bases
Acid is electron pair acceptor (electrophile) Base is electron pair donor (nucleophile)
eg., BF3+ : NH3 ⇒ BF3: NH3
Ammonia-borontrifluoride addition product
Limitation of Lewis concept
- It fails to account for the relative strength of acid and bases, as it is not based on ionization.
- It stress on the electron transfer, which should be quite fast
- Many Lewis acid-base reactions are slow.
The pH Scale
Hydronium ion concentration,in molarity is more conveniently expressed on a logarithmic scale known as the pH scale.
The pH of a solution is defined as the negative logarithm to base 10 of the acitivity (aH+) of hydrogen ion.
- solution pH < 7
- For basic solution pH > 7
- For neutral solution pH = 7
Common ion effect in the Ionization of Acids and Bases
Common ion effect may be defined as the suppression of the dissociation of a weak electrolyte (weak acid or weak base) by the addition of some strong electrolyte containing a common ion.
Solubility Product Ksp
It may be defined as the product of ionic concentration in a saturated solution of a sparingly soluble salt at a particular temperature.
Applications of solubility product constant
1. Calculation of solubility.
Knowing Ksp of a sparingly soluble salt at any given temperature, we can calculate its solubility ‘s’ or vice versa.
2. Predicting the precipitation of a salt in reaction.
- In a saturated solution, ionic product = Ksp
- In an unsaturated solution, ionic product < Ksp. No precipitation occurs.
- In a supersaturated solution, ionic product < Ksp. Precipitation occurs.
3. In inorganic qualitative analysis. In qualitative analysis metal ions Cu2+ (group II cation),
Al3+ (group III), Zn2+ (group iv cation) etc. are precipitated. It is based on the difference in the solubility products of their sulphides.
Solubility equilibria of sparingly soluble salts
|Category I||Soluble||Solubility > 0.1m|
|Category II||Slightly soluble||0.01m<solubility < 0.1m|
|Category III||Sparingly soluble||Solubility<0.01m|
Common ion effect on solubility of ionic salts
The precipitation of soluble salts in pure state from their saturated solution is known as salting out.
In the saturated solution of NaCI, there exists an equilibrium,
It is a process in which salt react with water to form acid and base.
Degree of hydrolysis (h)
It is defined as the fraction of the salt which is hydrolysed by water.
Hydrolysis constant (Kn)
The equilibrium constant of hydrolysis reactions is called hydrolysis constant.
A buffer solution is one which can resist change in its pH value on the addition of small quantities of acid or base.
Buffer mixture of a weak acid and its salt with a strong base,
eg., CH3COOH and CH3COONa (pH = 4.75).
Buffer mixture of a weak base and its salt with a strong acid,
eg., NH4OH and NH4CI (pH = 9.25).
pH of buffer solution (Henderson equation)
Preparation of Acidic Buffer
To prepare a buffer of acidic pH we use weak acid and its salt formed with strong base. For the general case where the weak acid HA ionises in water,
This expression is known as Henderson- Hasselbalch equation.
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