Lecture 8 - kau

Lecture 8 - kau

Lecture 8 Stability and reactivity We tend to say that substances are stable or unstable, reactive or unreactive but these terms are relative and may depend on many .factors Thermodynamic and kinetic factors can also

.be important Stability and reactivity can be controlled by thermodynamic factors (depending only on the initial and final states and not on the reaction pathway) or kinetic ones (very .dependent on the reaction pathway)

Both factors depend on the conditions, and on the possibility of different routes to .decomposition or reaction Enthalpy and Hess Law The enthalpy change (H) in a reaction is equal to the heat input under conditions of constant

.temperature and pressure Enthalpy is commonly used as a measure of the .energies involved in chemical reactions Endothermic reactions (positive H) are ones requiring a heat input, and exothermic .reactions (negative H) H) give a heat output

Hess Law states that H does not depend on the reaction pathway( taken between initial and final states), and is a consequence of the First .Law of Thermodynamics H can be expressed as the sum of the values for :individual steps +.. H = H1+ H2

Enthalpy change does depend on conditions of temperature, pressure and concentration of the initial and final states, and it is important to specify these. Standard states are defined as pure substances at standard pressure . (1bar), and temperature(298K)

Schematic thermodynamic cycle illustrating the use of Hess Law The standard enthalpy of formation of any compound refers to formation from its elements, all in standard states. By definition,it is zero for any

.element in its stable(standard ) state enthalpy change H in any reaction to be calculated from Entropy and free energy Entropy (S) is a measure of moleculardisorder, or more precisely the number of microscopic arrangements of energy possible

.in a macroscopic sample Entropy increases with rise in temperature and depends strongly on the state Entropy changes (S) are invariably positive for .reactions that generate gas molecules The Second Law of Thermodynamics asserts that the total entropy always increases in a spontaneous process, and reaches a maximum

.value at equilibrium Both internal and external changes are taken account of by defining the Gibbs free energy :change (G) for a reaction taking place at constant temperature (T, in kelvin) Gibbs free energy change (G)

Gibbs free energy change (G) From the Second Law it can be shown that G is always negative for a feasible reaction at constant temperature and pressure and is .zero at equilibrium S and G for reactions do not depend on the G for reactions do not depend on the

reaction pathway. They depend even more strongly than H on concentration and .pressure Tabulated standard entropies may be used to estimate changes in a reaction from where S values are not zero for elements

Equilibrium constants For a general reaction such as cC + dD aA +bB the equilibrium constant is K =[A]a[B]b/[C]c[D]d

where the terms [A], [B] as concentrations or partial pressures. (This assumes ideal thermodynamic behavior and is a much better approximation for gases than in solution.) A very large value (1) of 1) of K indicates a strong thermodynamic tendency to react, so that very

little of the reactants (A and B) will remain at equilibrium. Conversely, a very small value (1) indicates very little tendency1) indicates very little tendency to react: in this case the reverse reaction (C and D going to .A and B) will be very favorable For any reaction K may be related to the standard Gibbs

free energy change (H) G) according to R T lnK) = - H) G( where R is the gas constant (=8.314 J K1 mol1) and T the absolute temperature (in K). Thus equilibrium constants can be estimated from tabulated values of and trends may often be .interpreted in terms of changes in H and G for reactions do not depend on the S

Important notes Equilibrium constants change withtemperature in a way that depends on H for the reaction In accordance with Le Chateliers principle, Kincreases with rise in temperature for an endothermic reaction, and decreases for an .exothermic one

Reaction rates The rate of reaction generally depends on the concentration .of reactants Rate =k[A]n[B]m where k is the rate constant and n and m are the orders of.reaction with respect to reactants A and B Orders of reaction depend on the mechanism and are not.necessarily equal to the stoichiometric coefficients a and b

The rate constant depends on the mechanism and especially on the energy barrier or activation energy associated with .the reaction pathway High activation energies (Ea) give low rateconstants because only a small fraction of .molecules have sufficient energy to react This proportion may be increased by raisingthe temperature, and rate constants .approximately follow the Arrhenius equation

Arrhenius equation K =Ae-Ea/RT Large activation energies arise in reactionswhere covalent bonds must be broken before new ones are formed, or where atoms must . move through solids

Reactions involving free radicals, or ions insolution, often have small (sometimes zero) .activation energies Catalyst Reactions may be accelerated by the presence of a catalyst, which acts by providing an alternative pathway with lower activation .energy

A true catalyst by definition can be recoveredunchanged after the reaction, and so does not alter the thermodynamics or the position of equilibrium

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