Inorganic Chemistry
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Inorganic Chemistry

DisciplinaQuímica Inorgânica I3.874 materiais31.823 seguidores
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is related to the standard Gibbs free energy change. 
Equilibrium constants change with temperature in a way that depends on \u394H for the reaction. 
Reaction rates Reaction rates depend on the concentrations of reagents, and on a rate constant that itself 
depends on the energy barrier for the reaction. Reaction rates generally increase with rise in 
temperature. Catalysts provide alternative reaction pathways of lower energy. 
Related topics Inorganic reactions and synthesis (B6) 
Bond strengths (C8) 
Lattice energies (D6) 
Industrial chemistry: catalysts (J5) 
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reaction does depend on the pathway; this is the subject of chemical kinetics, and thermodynamic 
considerations alone cannot predict how fast a reaction will take place. 
Many known substances are thermodynamically stable, but others are only kinetically stable. 
For example, the hydrides B2H6 and SiH4 are thermodynamically unstable with respect to their 
elements, but in the absence of heat or a catalyst (and of atmospheric oxygen and moisture) the rate 
of decomposition is extremely slow. To assess why some substances are unknown, it is important to 
consider different possible routes to decomposition. For example, the unknown CaF(s) is probably 
thermodynamically stable with respect to the elements themselves, but certainly unstable 
(thermodynamically and kinetically) with respect to the reaction 
Thermodynamic and kinetic factors depend on temperature and other conditions. For example, CaF
(g) can be formed as a gas-phase molecule at high temperatures and low pressures. 
Enthalpy and Hess\u2019 Law 
The enthalpy change (\u394H) 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. The corrections are generally small, and enthalpy is 
commonly used as a measure of the energies involved in chemical reactions. Endothermic 
reactions (positive \u394H) are ones requiring a heat input, and exothermic reactions (negative \u394H) 
give a heat output. 
Hess\u2019 Law states that \u394H does not depend on the pathway taken between initial and final states, 
and is a consequence of the First Law of Thermodynamics, which asserts the conservation of total 
energy. Figure 1 shows a schematic thermodynamic cycle where the overall \u394H can be expressed as 
the sum of the values for individual steps: 
It is important that they need not represent any feasible mechanism for the reaction but can be any 
steps for which \u394H values are available from experiment or theory. Hess\u2019 Law is frequently used to 
estimate \u394H values that are not directly accessible, for example, in connection with lattice energy 
and bond energy calculations (see Topics D6 and C8). 
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), and 
Fig. 1. Schematic thermodynamic cycle illustrating the use of Hess\u2019 Law (see Equation 
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the temperature must be additionally specified, although 298 K is normally used. Corrections must 
be applied for any other conditions. The standard enthalpy of formation of any compound 
refers to formation from its elements, all in standard states. Tabulated values allow the standard 
enthalpy change \u394H\u398 in any reaction to be calculated from 
which follows from Hess\u2019 Law. By definition, is zero for any element in its stable (standard) 
Entropy and free energy 
Entropy (S) is a measure of molecular \u2018disorder\u2019, or more precisely \u2018the number of microscopic 
arrangements of energy possible in a macroscopic sample\u2019. Entropy increases with rise in 
temperature and depends strongly on the state. Entropy changes (\u394S) 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. To 
apply this to chemical reactions it is necessary to include entropy changes in the surroundings caused 
by heat input or output. Both internal and external changes are taken account of by defining the 
Gibbs free energy change (\u394G): for a reaction taking place at constant temperature (T, in kelvin) 
From the Second Law it can be shown that \u394G is always negative for a feasible reaction at constant 
temperature and pressure (and without any external driving force such as electrical energy) and is 
zero at equilibrium. 
As with enthalpies, \u394S and \u394G for reactions do not depend on the reaction pathway taken and so 
can be estimated from thermodynamic cycles like that of Fig. 1. They depend even more strongly 
than \u394H on concentration and pressure. Tabulated standard entropies may be used to estimate 
changes in a reaction from 
which is analogous to Equation 2 except that S\u398 values are not zero for elements. The direct analogy 
to Equation 2 may also be used to calculate \u394G\u398 for any reaction where the standard free energies of 
formation are known. 
Equilibrium constants 
For a general reaction such as 
the equilibrium constant is 
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where the terms [A], [B],\u2026strictly represent activities but are frequently approximated as 
concentrations or partial pressures. (This assumes ideal thermodynamic behavior and is a much 
better approximation for gases than in solution.) Pure liquids and solids are not included in an 
equilibrium constant as they are present in their standard state. A very large value (\u226b1) 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 (\u226a1) indicates very little tendency 
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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 (\u394G\u398) according to 
where R is the gas constant (=8.314 J K\u22121 mol\u22121) 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 \u394H\u398 and \u394S\u398 (see Equation 3). 
Equilibrium constants change with temperature in a way that depends on \u394H\u398 for the reaction. In 
accordance with Le Chatelier\u2019s principle, K increases 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, often according to a power 
law such as 
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 rate constants because only a small fraction of molecules have sufficient energy to react. 
This proportion may be increased by raising the temperature, and rate constants approximately 
follow the Arrhenius equation: 
Large activation energies arise in reactions where covalent