chemistry

Slide 1: 

SUBMITTED TO: Mr. J.L. AGARWAL PGT(CHEMISTRY) SUBMITTED BY : BHANU PRATAP CLASS : XI - “B” ROLL NO. : 11237 CHEMISTRY PROJECT WORK

Slide 2: 

CHEMICAL EQUILIBRIUM “Old Chemists Never Die; they just reach EQUILIBRIUM!”

Slide 3: 

CHEMICAL EQUILIBRIUM In a chemical reaction, chemical equilibrium is the state in which the concentrations of the reactants and products have not yet changed with time. It occurs only in reversible reactions, and not in irreversible reactions. The reaction rates of the forward and reverse reactions are generally not zero but, being equal, there are no net changes in the concentrations of the reactant and product. This process is called dynamic equilibrium.

Slide 4: 

Start with NO2 Start with N2O4 Start with NO2 & N2O4 Equilibrium favors the reactant side

Slide 5: 

Concept of equilibrium

Slide 6: 

Concept of equilibrium Chemical equilibrium occurs when a reaction and its reverse reaction proceed at the same rate.

Slide 7: 

DYNAMIC EQUILIBRIUM The net result of a dynamic equilibrium is that no change in the system is evident. Equilibrium is dynamic since product is constantly made (forward reaction), but at the same rate it is consumed (reverse reaction).

Slide 8: 

DYNAMIC EQUILIBRIUM Evaporation Open System (No Equilibrium) Evaporation Liquid Gas (No Equilibrium) (No Equilibrium) Liquid Gas Liquid Gas (Equilibrium)

Slide 9: 

DYNAMIC EQUILIBRIUM Ag + + Cl - AgCl (s) Chemical Equilibrium Rate of Precipitation = Rate of Dissolving HC2H3O2 (aq) H + + C2H3O2 - Rate of dissociation (ionization) = Rate of Association HC2H3O2 H + H + C2H3O2 - C2H3O2 - HC2H3O2 Cl - Ag + Ag + Cl - AgCl (s)

Slide 10: 

EXPERIMENT Orange Red HC2H3O2 H+ + C2H3O2- Blue Pink [CoCl4]2- + 6 H2O (l) Co(H2O)62+ + 4 Cl- White Colorless Solution NH4Cl (s) NH 4+ + Cl- You can actually “see” the equilibrium shift!

Slide 11: 

HENRY’S LAW The solubility of a gas in a liquid depends on temperature, the partial pressure of the gas over the liquid, the nature of the solvent and the nature of the gas. The most common solvent is water. Carbonated beverages are an example of Henry's law in everyday life. The dissolved carbon dioxide stays in solution in a closed pop bottle or can where the partial pressure of carbon dioxide was set at a high value during bottling. When the can or bottle is opened the partial pressure of CO2 is much lower and the dissolved carbon dioxide will gradually escape from the pop. When the new low partial pressure equilibrium is established the soda will be "flat" . This loss of dissolved carbon dioxide will happen faster for warm soda than for cold.

Slide 12: 

HENRY’S LAW The illustration shows that if the pressure is doubled then the concentration of dissolved gas will double. The dissolving process for gases is an equilibrium. The solubility of a gas depends directly on the gas partial pressure. At equilibrium the number of molecules leaving the gas phase to enter the solution equals the number of gas molecules leaving the solution. Gas solubility is proportional to the gas partial pressure. If the temperature stays constant increasing the pressure will increase the amount of dissolved gas.

Slide 13: 

EQUILIBRIUM CONSTANT Equilibrium Constant - When the rates of the forward and reverse reactions are equal, the system is “at equilibrium” and the reaction quotient = equilibrium constant. Consider, If we start with a mixture of nitrogen and hydrogen (in any proportions), the reaction will reach equilibrium with a constant concentration of nitrogen, hydrogen and ammonia. However, if we start with just ammonia and no nitrogen or hydrogen, the reaction will proceed and N2 and H2 will be produced until equilibrium is achieved.

Slide 14: 

EQUILIBRIUM CONSTANT

Slide 15: 

EQUILIBRIUM CONSTANT No matter the starting composition of reactants and products, the same ratio of concentrations is achieved at equilibrium. For a general reaction the equilibrium constant expression is where Kc is the equilibrium constant.

Slide 16: 

EQUILIBRIUM CONSTANT ONLY THE TEMPERATURE IS THE FACTOR WHICH EFFECT THE EQUILIBRIUM CONSTANT.

Slide 17: 

EQUILIBRIUM CONSTANT IN TERMS OF PRESSURE If KP is the equilibrium constant for reactions involving gases, we can write: KP is based on partial pressures measured in atmospheres. We can show that, PA = [A](RT)

Slide 18: 

EQUILIBRIUM CONSTANT IN TERMS OF PRESSURE PA = [A](RT) This means that we can relate Kc and KP: where Dn is the change in number of moles of gas. It is important to use: Dn = ngas(products) - ngas(reactants)

Slide 19: 

EQUILIBRIUM CONSTANT IN TERMS OF PRESSURE = Kp

Slide 20: 

reaction quotient (Qc) The reaction quotient (Qc) is calculated by substituting the initial concentrations of the reactants and products into the equilibrium constant (Kc) expression. IF, Qc > Kc system proceeds from right to left to reach equilibrium Qc = Kc the system is at equilibrium Qc < Kc system proceeds from left to right to reach equilibrium

Slide 21: 

Relationship between Kc and Kp Relationship between concentration and pressure obtained from the ideal gas law. Recall PV = nRT or Substitute for P in equilibrium expression. Consider the reaction: aA + bB  cC + dD

Slide 22: 

MAGNITUDE OF EQUILIBRIUM CONSTANT The equilibrium constant, K, is the ratio of products to reactants. Therefore, the larger K the more products are present at equilibrium. Conversely, the smaller K the more reactants are present at equilibrium. If K >> 1, then products dominate at equilibrium and equilibrium lies to the right. If K << 1, then reactants dominate at equilibrium and the equilibrium lies to the left.

Slide 23: 

MAGNITUDE OF EQUILIBRIUM CONSTANT

Slide 24: 

Law of chemical equilibrium The Law of Chemical Equilibrium is a relation stating that in a reaction mixture at equilibrium, there is a condition (given by the equilibrium constant, Kc) relating the concentrations of the reactants and products. For the reaction:aA(g) + bB(g) ↔ cC(g) + dD(g), Kc = [ C ]c·[ D ]d / [ A ]a·[ B ]b

Slide 25: 

Homogeneous equilibrium Homogenous equilibrium applies to reactions in which all reacting species are in the same phase. Kp = Kc(RT)Dn Dn = moles of gaseous products – moles of gaseous reactants = (c + d) – (a + b)

Slide 26: 

Homogeneous equilibrium

Slide 27: 

Heterogeneous equilibrium When all reactants and products are in one phase, the equilibrium is homogeneous. If one or more reactants or products are in a different phase, the equilibrium is heterogeneous. Consider: The concentration of a solid or pure liquid is its density divided by molar mass.

Slide 28: 

Heterogeneous equilibrium

Slide 29: 

Heterogeneous equilibrium Neither density nor molar mass is a variable, the concentrations of solids and pure liquids are constant. For the decomposition of CaCO3: We ignore the concentrations of pure liquids and pure solids in equilibrium constant expressions. The amount of CO2 formed will not depend greatly on the amounts of CaO and CaCO3 present.

Slide 30: 

Reaction sequence As with DH in thermo. Kc can be determined from a reaction sequence. Consider the reactions to the right. The third reaction is the result of the sum of the other two (called a reaction sequence..

Slide 31: 

Application of equilibrium constant 1. To predict the extent of reaction

Slide 32: 

1. To predict the extent of reaction

Slide 36: 

2. To predict the direction of reaction

Slide 37: 

2. To predict the direction of reaction

Slide 38: 

There are three possible values of this ratio when it is compared with the value of Kc. When ratio=Kc According to the law of mass action, there is no shifting of reaction and there will be no change in the concentration of reactants and products and the system is already at equilibrium. When ratio > Kc                               In this condition the reaction will shift in the backward direction to achieve equilibrium state. At equilibrium quantity of product will decrease and the quantity of reactants will increase. When ratio < Kc In this condition the reaction will shift in forward direction to achieve equilibrium state. At equilibrium quantity of product will increase and the quantity of reactants will decrease.

Slide 39: 

Le Châtelier’s Principle In chemistry, Le Chatelier's principle, also called the Chatelier's principle, can be used to predict the effect of a change in conditions on a chemical equilibrium. The principle is named after Henry Louis Le Chatelier and sometimes Karl Ferdinand Braun who discovered it independently.

Slide 40: 

Le Châtelier’s Principle **“If a system at equilibrium is disturbed by a change in temperature, pressure, or the concentration of one of the components, the system will shift its equilibrium position so as to counteract the effect of the disturbance.” OR **“When a system at equilibrium is placed under stress, the system will undergo a change in such a way as to relieve that stress and restore a state of equilibrium.”

Slide 41: 

Le Châtelier’s Principle

Slide 42: 

Le Châtelier’s TRANSLATED: “When you take something away from a system at equilibrium, the system shifts in such a way as to replace some what you’ve taken away.” “When you add something to a system at equilibrium, the system shifts in such a way as to use up some of what you’ve added.”

Slide 43: 

Consider the production of ammonia As the pressure increases, the amount of ammonia present at equilibrium increases. As the temperature decreases, the amount of ammonia at equilibrium increases. Le Châtelier’s Principle

Slide 44: 

Le Châtelier’s Principle

Slide 45: 

Le Châtelier’s Principle Removing SO2, O2

Slide 46: 

Le Châtelier’s Principle

Slide 47: 

Le Châtelier’s Principle

Slide 48: 

Le Châtelier’s Principle

Slide 49: 

What Happens When More of a Reactant Is Added to a System?

Slide 50: 

THE HABER PROCESS “The transformation of nitrogen and hydrogen into ammonia (NH3) is of tremendous significance in agriculture, where ammonia-based fertilizers are of utmost importance.”

Slide 51: 

THE HABER PROCESS If H2 is added to the system, N2 will be consumed and the two reagents will form more NH3.

Slide 52: 

THE HABER PROCESS This apparatus helps push the equilibrium to the right by removing the ammonia (NH3) from the system as a liquid.

Slide 53: 

LECHATELIER EXAMPLE #1 **A closed container of ice and water is at equilibrium. Then, the temperature is raised. Ice + Energy  Water The system temporarily shifts to the _______ to restore equilibrium. right

Slide 54: 

LECHATELIER EXAMPLE #2 **A closed container of N2O4 and NO2 is at equilibrium. NO2 is added to the container. N2O4 (g) + Energy  2 NO2 (g) The system temporarily shifts to the _______ to restore equilibrium. left

Slide 55: 

LECHATELIER EXAMPLE #3 **A closed container of water and its vapor is at equilibrium. Vapor is removed from the system. water + Energy  vapor The system temporarily shifts to the _______ to restore equilibrium. right

Slide 56: 

LECHATELIER EXAMPLE #4 **A closed container of N2O4 and NO2 is at equilibrium. The pressure is increased. N2O4 (g) + Energy  2 NO2 (g) The system temporarily shifts to the _______ to restore equilibrium, because there are fewer moles of gas on that side of the equation. left

Slide 57: 

Le Châtelier’s Principle System starts at equilibrium. A change/stress is then made to system at equilibrium. Change in concentration Change in volume Change in pressure Change in Temperature Add Catalyst System responds by shifting to reactant or product side to restore equilibrium.

Slide 58: 

1.Change in concentration

Slide 59: 

1.Change in concentration Add Add Change Shifts the Equilibrium *Increase concentration of product(s) *Decrease concentration of product(s) *Increase concentration of reactant(s) *Decrease concentration of reactant(s) left right right left

Slide 60: 

2. Change in volume and pressure Only a factor with gases Change Shifts the Equilibrium Increase pressure Side with fewest moles of gas Decrease pressure Side with most moles of gas Decrease volume Increase volume Side with most moles of gas Side with fewest moles of gas

Slide 61: 

2. Change in volume and pressure As volume is decreased pressure increases. Le Châtelier’s Principle: if pressure is increased the system will shift to counteract the increase.

Slide 62: 

3. Change in temperature

Slide 63: 

3. Change in temperature The equilibrium constant is temperature dependent. For an endothermic reaction, H > 0 and heat can be considered as a reactant. For an exothermic reaction, H < 0 and heat can be considered as a product. Adding heat (i.e. heating the vessel) favors away from the increase: if H > 0, adding heat favors the forward reaction, if H < 0, adding heat favors the reverse reaction.

Slide 64: 

3. Change in temperature Removing heat (i.e. cooling the vessel), favors towards the decrease: if H > 0, cooling favors the reverse reaction, if H < 0, cooling favors the forward reaction. Consider for which DH > 0. Co(H2O)62+ is pale pink and CoCl42- is blue.

Slide 65: 

4. Effect of catalyst A catalyst lowers the activation energy barrier for the reaction. Therefore, a catalyst will decrease the time taken to reach equilibrium. A catalyst does not effect the composition of the equilibrium mixture.

Slide 66: 

4. Effect of catalyst Catalysts increase the rate of both the forward and reverse reactions.