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RATE OF CHEMICAL REACTION AND CHEMICAL EQUILIBRIUM

RATE OF CHEMICAL REACTION AND CHEMICAL EQUILIBRIUM. CHEMICAL REACTIONS. A chemical change is called a chemical reaction. Iron rusts when kept exposed to humid air for a considerable time. Change of milk into curd. Burning of a paper. Combination of Nitrogen and Hydrogen to form Ammonia.

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RATE OF CHEMICAL REACTION AND CHEMICAL EQUILIBRIUM

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  1. RATE OF CHEMICAL REACTION AND CHEMICAL EQUILIBRIUM

  2. CHEMICAL REACTIONS • A chemical change is called a chemical reaction. • Iron rusts when kept exposed to humid air for a considerable time. • Change of milk into curd. • Burning of a paper. • Combination of Nitrogen and Hydrogen to form Ammonia

  3. SLOW CHEMICAL REACTIONS • The reactions which take place in a few minutes or more reactants are called slow reactions. • When Rusting of iron • Weathering of rocks • Generally the reactions between covalent compounds are slow • When ethyl alcohol and acetic acid are heated in the presence of a little concentrated sulphuric acid(catalyst), the reaction takes place in a few minutes and an ester called ethyl acetate is formed. This is a slow reaction.

  4. FAST CHEMICAL REACTIONS • The reaction which takes place immediately on mixing the reactants are called fast reactions. • Reaction between an Acid and a Base • HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) • 2Mg(s) + O2(g)  2 MgO(s)

  5. RATE OF A CHEMICAL REACTION • The rate of a chemical reaction means how fast the reaction is taking place by indicating how much of a reactant is consumed, or how much of a product is formed in a given time. • Therefore Reaction rate = Change in concentration of a reactant or product/ time taken

  6. UNIT OF RATE OF REACTION • The concentration is usually expressed in molarity M ( the number of moles of substance dissolved per litre of the solution) and time in seconds • Then unit of the rate of reaction is is moleL-1/s or molesL-1s-1

  7. REACTION RATE When acidified hydrogen peroxide (H2O2) is added to a solution of potassium iodide (KI) , iodine is liberated • H2O2(aq)+2 KI (aq) + H2SO4(aq) 2H2O(I)+ I2(aq)+ K2SO4(aq) • Here the concentration of iodine is zero initially. With the passage of time, it increases and the reaction solution becomes brownish. Concentration of iodine can be measured at different intervals of time by titration against sodium thiosulphate.

  8. REACTION RATE • If the concentration of iodine rises from 0 to 10-5mol L-1 in 10 seconds, we write ∆ [ I2] 10-5mol L-1 • Reaction rate = ----------- = -------------- ∆ t 10s = 10-6 mol L-1s-1 • Here, symbol ∆ represents a change and [I2], molar concentration of iodine.

  9. AVERAGE RATE • The rate of a reaction which is obtained by measuring the change in concentration of a reactant or product over a considerable period of time then it is called average reaction rate. • Suppose iodine( I2) is being evolved in a reaction as a product and we measure the change in its concentration over a time, say 10 seconds, so in this case • ∆ [I2] / ∆ t will give us an average reaction rate. • So Average rate of reaction = ∆ [I2]/ ∆ t

  10. INSTANTANEOUS RATE • The rate of a reaction at a particular instant of time is called instantaneous reaction rate. • It is the accurate estimate of reaction rate at any moment. • In order to determine the rate at a particular instant, one should make the time interval almost zero. i.e d[I2] • Instantaneous rate = -------------- dt • Here d[I2] denotes very small change in the conc. of I2 and dt denotes very small change in time

  11. RATE OF REACTION • One can think of expressing the rate in terms of change in concentration of H2O2 and KI is expressed as, ∆[H2O2] 1 ∆[KI] • Rate of reaction = ------------------ = - --- --------- ∆ t 2 ∆ t • Here negative sign indicates that concentration of H2O2 and KI will decrease with time. Here the rate of consumption of KI is twice the rate of consumption of H2O2 therefore, in order to make the two rates equal we divide it by two (stoichiometric coefficient) and write ½( ∆[KI]/ ∆ t)

  12. RATE OF REACTION • Cisplatin, Pt(NH3)2Cl2 is used in chemotherapy of cancer. It reacts with water and releases chloride ion bound to central platinum metal. The reaction is represented as Pt(NH3)2Cl2 + H2O  Pt(NH3)2Cl+ + Cl- • Here the conc. of cisplatin decreases with lapse of time but conc. of Cl increases.

  13. REACTION RATE CURVES Reaction Rate Graphs - Conc. of reactants(mol L-1) Conc of products (mol L-1) Time (s)  Time (s)  Here the conc. Of Product changes with time Here conc. of reactant changes with time

  14. QUESTIONS • Give an example of slow reaction. • Give an example of fast reaction. • Which reactions are slow and which are fast • Rusting of iron, Burning of Magnesium wire, setting of cement, Formation of coal, formation of sodium oxide • Fill in the blanks • Reactions between covalent compounds are usually ----- • Reactions between ionic compounds are very ------

  15. QUESTIONS • Define the rate of a chemical reaction ( in term products and reactants) • What is meant by average rate of reaction. • What is meant by instantaneous reaction rate. • Why do we put a minus sign in the formula for the rate of a reaction in terms of the change in conc. of a reactant but no such sign in case of product. • Write the unit for expressing i) the concentration of an aqueous solution ii) rate of a chemical reaction • Define molarity ? • Draw a graph to show the variation of conc. Of i) a reactant with time ii) a product with time. • Write the chemical name and formula of Cisplatin. For what purpose is it used. • A solution contains 0.021 mole of a substance in 3.0 L of the solution. What is its molarity ? • T he formula of H2O2 has been put within a square bracket as [H2O2]. What does it represent. • Consider the reaction • 2K

  16. ENERGY CHANGES IN A CHEMICAL REACTION • Most of the chemical reactions are accompanied by energy changes. The energy is either released or absorbed during a chemical reaction. In some reactions, energy is given in form of heat and in some cases, in form of light.

  17. EXOTHERMIC REACTION • Exothemic reaction – reactions which release heat • Ex- all combustion reactions. • CH4(g) + 2O2(g) CO2(g) +2H2O(l) +energy(890.4 kj mol-1) • 2AL(s) + Fe2O3(s) Al2O3(s) + 2 Fe(l)+ heat • N2 (g) + 3H2 (g) 2NH3 (g) + heat Fe

  18. ENDOTHERMIC REACTION • endothermic reactions- reactions which absorb heat. • Ba(OH)2.8H2O(s) + 2NH4SCN(s)+ heat Ba(SCN)2(aq) + 2NH3(aq) + 10H2O(l) • N2 (g) + O2(g) + heat -- 2 NO ( g )

  19. PHOTOCHEMICAL REACTION • Photochemical reactions – Reactions that take place in presence of light. For example, photosynthesis of plants. Photographic films which are coated with silver bromide (very small amount of silver iodide) undergo chemical reaction when exposed to sunlight. • CH4 (g) + Cl2 (g)  CH3Cl (g) + HCl(g) • 6CO2(g) + 6H2O (l)  C6H12O6(aq) + 6O2(g) • 2AgBr (s)  2Ag (s) + Br2(l)

  20. FACTORS AFFECTING THE RATE OF REACTION The following factors affect the rate of reaction. These are • Effect of Concentration • Effect of Temperature • Effect of Catalyst • Effect of light

  21. EFFECT OF CONCENTRATION • Increase in the concentration of reactants increases the number of molecules of reactants. There will be more collisions between the reactant molecules and the rate of reaction will increase. • When nitrogen oxide (NO) reacts with oxygen ( Which is a homogenous system) • 2NO (g) + O2 (g)  2 NO2 (g) , the reaction rate doubles when oxygen concentration doubles. But rate quadruples when the concentration of nitrogen monoxide is doubled.

  22. EFFECT OF TEMPERATURE • On increasing the temperature of reactants, the energy of the molecules increases. Due to increased energy the frequency of collision increases and finally the rate of reaction increases. • CaCO3 (s) + 2 HCl (aq)  CaCl2 (aq) + CO2 (g) + H2O (l) . In this reaction the evolution of CO2 is slow but if we increase the temperature the evolution of gas is more rapid indicating that the rate of reaction increases on increasing the temperature.

  23. EFFECT OF CATALYST • A catalyst is a substance which increases the rate of a chemical reaction but there is no change in the catalyst itself. In the presence of a catalyst a reaction takes place at a faster rate and at lower temperature. • Iron is used as a catalyst in the manufacture of ammonia. Iron catalyst increases the rate of reaction between nitrogen and hydrogen to form ammonia. • N2 (g) + 3H2(g) 2 NH3 (g) + heat • Similarly , the reaction of SO2 and O2 gives SO3 in presence of nitrogen monoxide (NO), which acts as a catalyst. 2 SO2 (g) + O2 (g) 2SO3 ( g) NO (g)

  24. EFFECT OF LIGHT • Photosynthesis and photography both involve light sensitive reactions. • In our daily life, a coloured cloth gradually fades when dried in sunlight. • Dyes of the clothes undergo chemical change and fade gradually. • Plants prepare food in presence of sunlight.

  25. REVERSIBLE REACTIONS • The reaction in which the product can react under any conditions to give back reactants are called reversible reactions or a reaction which takes place in a forward as well as reverse direction both, is called reversible reaction. It is represented by putting a double arrow between the reactants and products. One arrow pointing towards the products and other towards reactants.

  26. Examples of reversible reactions • 2H2 + O2 2H2O • N2 + O2 2NO • CaCO3 CaO + CO2 • N2 + 3H2 2NH3

  27. IRREVERSIBLE REACTIONS • Those chemical reactions in which the products cannot recombine to give back reactants are called irreversible reactions. • 2 Na(s) + 2 H2O (l)  2 NaOH ( aq) + H2 (g) • Here the products, sodium hydroxide and hydrogen can’t recombine to give back the reactants sodium and water.

  28. REVERSIBLE PROCESS AND CHEMICAL EQUILIBRIUM • The formation of ammonia is a reversible process. we say that nitrogen and hydrogen react under appropriate conditions and give ammonia, then this statement is expressed in the form of an equation as • N2 (g) + 3 H2 (g) → 2 NH3 (g) • The reaction is taking place in forward direction and we, therefore, say it is a ‘forward reaction’.

  29. We can also write decomposition of NH3 as, 2NH3 (g) → N2 (g) + 3H2 (g) which is, in fact, a reverse of the forward reaction and takes place in opposite direction, and therefore, we call this reaction as ‘reverse reaction’. We represent forward and reverse reactions together as N2 (g) + 3H2 (g) 2 NH3 (g) • Such reactions never go to completion if performed in a close container.

  30. N2 (g) + 3 H2 (g)  2 NH3 (g) (Rate of formation of ammonia decreases with time) Rate of Reaction  Chemical Equilibrium Rate of decomposition of ammonia increases with time ) 2NH3 (g)  N2 (g) + 3H2( g ) O Time increases 

  31. For a reversible chemical reaction, an equilibrium state is attained when the rate at which a chemical reaction is proceeding in forward direction equals the rate at which the reverse reaction is proceeding. • At equilibrium, • Rate of forward reaction= Rate of reverse reaction • In case of ammonia, when reaction is performed at high pressure and temperature in a close container, at equilibrium, • Rate of formation of ammonia = Rate of decomposition of ammonia

  32. DYNAMIC EQUILIBRIUM • Since the rate of decomposition and rate of formation is the same at equilibrium, the concentrations of ammonia, hydrogen and nitrogen remain constant. • This equilibrium is dynamic in nature and is therefore, called dynamic equilibrium. • In a close container containing water, rate of evaporation and rate of condensation of water molecules become equal at equilibrium.

  33. CHARACTERISTIC FEATURES OF DYNAMIC EQUILIBRIUM (i) At equilibrium, macroscopic (observable) properties such as concentration, density, colour, etc., are constant under the given condition of temperature, pressure and initial amount of the substances. (ii) At equilibrium, microscopic (at molecular level) processes continue, but they are in balance. This means that no overall large scale (macroscopic) changes occur. Rate of forward process (reaction) is equal to the rate of reverse process, and as a consequence of this, no net change results.

  34. (iii) The equilibrium can be attained from either direction, beginning with only the materials on one side of the change. Such changes are described as reversible. (iv) Equilibrium can be achieved in a closed system. In a closed system, there is no loss or gain of matter to or from the surroundings. An open system may allow matter to escape or to enter. Therefore, in an open system, the equilibrium cannot be attained.

  35. EQUILIBRIUM CONSTANT • It has been experimentally found that at a particular temperature, when equilibrium is attained, the ratio between concentration of reactants and products becomes constant. For reaction • N2 (g) + 3 H2 (g) 2 NH3 (g) + 92 kJ • At equilibrium, Here, concentration of product (ammonia) occurs in numerator, and those of reactants (hydrogen and nitrogen) occur in denominator. Each concentration term, [NH3], or [N2], or [H2], is raised to a power equal to stoichiometric* coefficient in the balanced equation. K is called equilibrium constant.

  36. For a general reaction, • a A + b B c C + d D • at equilibrium, equilibrium constant can be • written as, • In the above expression, square bracket • denotes molar concentration, i.e. concentration • in mol/L. The expression • (1.2) is also known as Law of Chemical • Equilibrium. At a particular temperature, • equilibrium constant has a definite value. When • we express concentration in mol L-1, equilibrium • constant is denoted by Kc.

  37. Magnitude of equilibrium constant, K, indicates how far a chemical reaction can go. In other words, it is an indicator of the extent of a chemical reaction. • Larger the value of K, higher will be the equilibrium concentration of products on the right hand side of the reaction. Smaller value of K indicates lower concentration of the products at equilibrium. • For the reaction at 298 K, PCl3 (g) + Cl2 (g) _PCl5 (g) ; Kc = 1.9 • Low value of equilibrium constant shows that at equilibrium, the concentration of product (PCl5) will be low.

  38. . Similarly, for reaction at • 298 K, • N2 (g) + O2 (g) 2 NO (g) • K = c • Very small value of Kc implies that reactants • N2 and O2 will be dominant species in the system • at equilibrium.

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