P.A. Cox - Inorganic chemistry (793955), страница 9
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For example, CaF(g) can be formedas a gas-phase molecule at high temperatures and low pressures.Enthalpy and Hess’ LawThe enthalpy change (ΔH) in a reaction is equal to the heat input under conditions of constant temperature and pressure.It is not exactly equal to the total energy change, as work may be done by expansion against the external pressure. Thecorrections are generally small, and enthalpy is commonly used as a measure of the energies involved in chemicalreactions. Endothermic reactions (positive ΔH) are ones requiring a heat input, and exothermic reactions(negative ΔH) give a heat output.Hess’ Law states that ΔH does not depend on the pathway taken between initial and final states, and is aconsequence of the First Law of Thermodynamics, which asserts the conservation of total energy.
Figure 1 shows aschematic thermodynamic cycle where the overall ΔH can be expressed as the sum of the values for individual steps:(1)It is important that they need not represent any feasible mechanism for the reaction but can be any steps for which ΔH valuesare available from experiment or theory. Hess’ Law is frequently used to estimate ΔH values that are not directlyaccessible, for example, in connection with lattice energy and bond energy calculations (see Topics D6 and C8).Fig. 1. Schematic thermodynamic cycle illustrating the use of Hess’ Law (see Equation 1).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 (1 bar), andB2—STABILITY AND REACTIVITY35the temperature must be additionally specified, although 298 K is normally used. Corrections must be applied for anyother conditions. The standard enthalpy of formationof any compound refers to formation from itselements, all in standard states. Tabulated values allow the standard enthalpy change ΔHΘ in any reaction to becalculated from(2)which follows from Hess’ Law. By definition,is zero for any element in its stable (standard) state.Entropy and free energyEntropy (S) is a measure of molecular ‘disorder’, or more precisely ‘the number of microscopic arrangements ofenergy possible in a macroscopic sample’.
Entropy increases with rise in temperature and depends strongly on the state.Entropy changes (ΔS) are invariably positive for reactions that generate gas molecules. The Second Law ofThermodynamics asserts that the total entropy always increases in a spontaneous process, and reaches a maximumvalue at equilibrium. To apply this to chemical reactions it is necessary to include entropy changes in the surroundingscaused by heat input or output. Both internal and external changes are taken account of by defining the Gibbs freeenergy change (ΔG): for a reaction taking place at constant temperature (T, in kelvin)(3)From the Second Law it can be shown that ΔG is always negative for a feasible reaction at constant temperature andpressure (and without any external driving force such as electrical energy) and is zero at equilibrium.As with enthalpies, ΔS and ΔG for reactions do not depend on the reaction pathway taken and so can be estimatedfrom thermodynamic cycles like that of Fig.
1. They depend even more strongly than ΔH on concentration andpressure. Tabulated standard entropies may be used to estimate changes in a reaction fromwhich is analogous to Equation 2 except that SΘ values are not zero for elements. The direct analogy to Equation 2 mayalso be used to calculate ΔGΘ for any reaction where the standard free energies of formationare known.Equilibrium constantsFor a general reaction such asthe equilibrium constant iswhere the terms [A], [B],…strictly represent activities but are frequently approximated as concentrations or partialpressures. (This assumes ideal thermodynamic behavior and is a much better approximation for gases than insolution.) Pure liquids and solids are not included in an equilibrium constant as they are present in their standard state.A very large value (≫ 1) of K indicates a strong thermodynamic tendency to react, so that very little of the reactants (A36SECTION B—INTRODUCTION TO INORGANIC SUBSTANCESand B) will remain at equilibrium.
Conversely, a very small value (<<1) indicates very little tendency to react: in thiscase 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 (ΔGΘ) according to(4)where R is the gas constant (=8.314 J K−1 mol−1) and T the absolute temperature (in K). Thus equilibrium constants canbe estimated from tabulated values ofand trends may often be interpreted in terms of changes in ΔHΘ and ΔSΘ(see Equation 3).Equilibrium constants change with temperature in a way that depends on ΔHΘ for the reaction. In accordance withLe Chatelier’s principle, K increases with rise in temperature for an endothermic reaction, and decreases for anexothermic one.Reaction ratesThe rate of reaction generally depends on the concentration of reactants, often according to a power law such aswhere k is the rate constant and n and m are the orders of reaction with respect to reactants A and B.
Orders ofreaction depend on the mechanism and are not necessarily equal to the stoichiometric coefficients a and b. The rateconstant depends on the mechanism and especially on the energy barrier or activation energy associated with thereaction pathway. High activation energies (Ea) give low rate constants because only a small fraction of molecules havesufficient energy to react. This proportion may be increased by raising the temperature, and rate constantsapproximately follow the Arrhenius equation:Large activation energies arise in reactions where covalent bonds must be broken before new ones are formed, or whereatoms must move through solids.
Reactions involving free radicals, or ions in solution, often have small (sometimeszero) activation energies.Reactions may be accelerated by the presence of a catalyst, which acts by providing an alternative pathway withlower activation energy. A true catalyst by definition can be recovered unchanged after the reaction, and so does notalter the thermodynamics or the position of equilibrium (see Topic J5).Section B—Introduction to inorganic substancesB4OXIDATION AND REDUCTIONKey NotesDefinitionsOxidation statesBalancing redoxreactionsExtraction of theelementsRelated topicsOxidation means combination with a more electronegative elementor the removal of electrons. Reduction means combination with aless electronegative element or the addition of electrons.
A completeredox reaction involves both processes.Oxidation states of atoms in a compound are calculated by assigningelectrons in a bond to the more electronegative element. In simpleionic compounds they are the same as the ionic charges. In any redoxreaction the oxidation states of some elements change.In complete redox reactions the overall changes in oxidation statemust balance. When reactions involve ions in water it is convenientto split the overall reaction into two half reactions.
To balance theseit may also be necessary to provide water and H+ or OH−.Redox reactions are used in the extraction of nearly all elements fromnaturally occurring compounds. Carbon is used to reduce some metaloxides, but many elements require stronger reducing agents, or theuse of electrolysis.Inorganic reactions andElectrode potentials (E5)synthesis (B6)DefinitionsOxidation originally meant ‘combination with oxygen’ and reduction ‘removal of oxygen’. These definitions have beengreatly expanded. Oxidation implies combination with a more electronegative element, the removal of a lesselectronegative one, or simply the removal of electrons.
Reduction is the reverse of oxidation and in generalimplies addition of electrons. In any reaction where one species is oxidized, another must be reduced: the termredox reaction is used to express this.Two examples are: the reaction of zinc in aqueous acid,(1)where zinc metal is oxidized to Zn2+, and hydrogen reduced from H+ to H2; and the reduction of zinc oxide by carbon,(2)38SECTION B—INTRODUCTION TO INORGANIC SUBSTANCESwhere zinc is reduced from ZnO to the metal, elemental carbon is oxidized to CO, and oxygen, combined with a lesselectronegative element on both sides, is not oxidized or reduced.A strong oxidizing agent is a substance capable of oxidizing many others, and is thus itself easily reduced;conversely, a strong reducing agent is itself easily oxidized; these terms usually imply thermodynamic reactiontendency although kinetics may also be important (see Section B3).
Atmospheric dioxygen is a good oxidizing agent, butmany substances (e.g. organic compounds) are kinetically stable in air. Strong reducing agents include electropositivemetals, especially those of group 1 (see Section G2).Oxidation statesThe oxidation state (or oxidation number) is a number applied to each atom in a compound in such as way as tokeep track of changes occurring in redox reactions.
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