Inorganic Chemistry
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Inorganic Chemistry

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state will reduce the potential. For example, 
cyanide (CN\u2212) complexes much more strongly with Mn3+ than with Mn2+, and at unit activity 
reduces the Mn3+/Mn2+ potential from its standard value of +1.5 V to +0.22 V. Conversely, the 
potential increases if the lower oxidation state is more strongly complexed. 
Diagrammatic representations 
A Latimer diagram shows the standard electrode potentials associated with the different oxidation 
states of an element, as illustrated in Fig. 1 for manganese. Potentials not given explicitly can be 
calculated using Equation 1 and taking careful account of the number of electrons involved. Thus the 
free energy change for the Mn3+/Mn reduction is the sum of those for Mn3+/Mn2+ and Mn2+/Mn. 
From Equation 1 therefore 
Fig. 1. Latimer diagram for Mn at pH=0. 
In a Frost or oxidation state diagram (see Fig. 2) each oxidation state (n) is assigned a volt 
equivalent equal to n times its value with respect to the element. The potential in volts 
between any two oxidation states is equal to the slope of the line between the points in this diagram. 
Steep positive slopes show strong oxidizing agents, steep negative slopes strong reducing agents. 
Frost diagrams are convenient for displaying the comparative redox properties of elements (see 
Topics F9 and H3). 
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Fig. 2. Frost diagram for Mn at pH=0 (solid line) and pH=14 (dashed line). 
Frost diagrams also provide a visual guide to when disproportionation of a species is expected. For 
example, in Fig. 2 the Mn3+ state at pH=0 is found above the line formed by joining Mn2+ with 
MnO2. It follows that the Mn
3+/Mn2+ potential is more positive than MnO2/Mn
3+, and 
disproportionation is predicted: 
The equilibrium constant of this reaction can be calculated by noting that it is made up from the half 
reactions for MnO2/Mn
3+ and Mn3+/Mn2+ each with n=1, and has from 
Fig. 1, giving K=2×109. The states MnV and MnVI are similarly unstable to disproportionation at 
pH=0, whereas at pH=14, also shown in Fig. 2, only MnV will disproportionate. 
Latimer and Frost diagrams display the same information but in a different way. When 
interpreting electrode potential data, either in numerical or graphical form, it is important to 
remember that a single potential in isolation has no meaning, 
Kinetic limitations 
Electrode potentials are thermodynamic quantities and show nothing about how fast a redox 
reaction can take place (see Topic B3). Simple electron transfer reactions (as in Mn3+/Mn2+) are 
expected to be rapid, but redox reactions where covalent bonds are made or broken may be much 
slower (see Topics F9 and H7). For example, the potential is well above that for the 
oxidation of water (see O2/H2O in Table 1), but the predicted reaction happens very slowly and 
aqueous permanganate is commonly used as an oxidizing agent (although it should always be 
standardized before use in volumetric analysis). 
Kinetic problems can also affect redox reactions at electrodes when covalent substances are 
involved. For example, a practical hydrogen electrode uses specially prepared platinum with a high 
surface area to act as a catalyst for the dissociation of dihydrogen into atoms (see Topic J5). On other 
metals a high overpotential may be experienced, as a cell potential considerably larger than the 
equilibrium value is necessary for a reaction to occur at an appreciable rate. 
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Section F\u2014 
Chemistry of nonmetals 
Covalent chemistry 
Nonmetallic elements include hydrogen and the upper right-hand portion of the p block (see Topic 
B2, Fig. 1). Covalent bonding is characteristic of the elements, and of the compounds they form with 
other nonmetals. The bonding possibilities depend on the electron configurations of the atoms (see 
Topics A4 and C1). Hydrogen (Topic F2) is unique and normally can form only one covalent bond. 
Boron (Topic F3) is also unusual as compounds such as BF3 have an incomplete octet. Electron 
deficiency leads to the formation of many unusual compounds, especially hydrides (see also Topic 
The increasing number of valence electrons between groups 14 and 18 has two possible 
consequences. In simple molecules obeying the octet rule the valency falls with group number (e.g. 
in CH4, NH3, H2O and HF, and in related compounds where H is replaced by a halogen or an 
organic radical). On the other hand, if the number of valence electrons involved in bonding is not 
limited, then a wider range of valencies becomes possible from group 15 onwards. This is most 
easily achieved in combination with the highly electronegative elements O and F, and the resulting 
compounds are best classified by the oxidation state of the atom concerned (see Topic B4). Thus the 
maximum possible oxidation state increases from +5 in group 15 to +8 in group 18. The +5 state is 
found in all periods (e.g. PF5) but higher oxidation states in later groups require octet 
expansion and occur only from period 3 onwards (e.g. SF6 and in group 18 only xenon can 
do this, e.g. XeO4). 
Key Notes 
Hydrogen and boron stand out in their chemistry. In the other elements, valence states depend on 
the electron configuration and on the possibility of octet expansion which occurs in period 3 
onwards. Multiple bonds are common in period 2, but are often replaced by polymerized 
structures with heavier elements. 
Simple anionic chemistry is limited to oxygen and the halogens, although polyanions and 
polycations can be formed by many elements. 
Many halides and oxides are Lewis acids; compounds with lone-pairs are Lewis bases. Brønsted 
acidity is possible in hydrides and oxoacids. Halide complexes can also be formed by ion 
The oxidizing power of elements and their oxides increases with group number. Vertical trends 
show an alternation in the stability of the highest oxidation state. 
Electronegativity and bond type (B1) Chemical periodicity (B2) 
Electron pair bonds (C1) 
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Octet expansion or hypervalence is often attributed to the involvement of d orbitals in the same 
principal quantum shell (e.g. 3d in period 3; see Topics A3 and A4). Thus six octahedrally directed 
bonds as in SF6 could be formed with sp
3d2 hybrid orbitals (see Topic C6). In a similar way the 
multiple bonding normally drawn in species such as (1) is often described as d\u3c0-p\u3c0 bonding. 
These models certainly overestimate the contribution of d orbitals. It is always possible to draw 
valence structures with no octet expansion provided that nonzero formal charges are allowed. For 
example, the orthonitrate ion is drawn without double bonds (2), and could be similarly 
represented. One of many equivalent valence structures for SF6 where sulfur has only eight valence-
shell electrons is shown in 3. Three-center four-electron bonding models express similar ideas (see 
Topic C6). Such models are also oversimplified. It is generally believed that d orbitals do play some 
role in octet expansion, but that two other factors are at least as important: the larger size of elements 
in lower periods, which allows higher coordination numbers, and their lower electronegativity, 
which accommodates positive formal charge more easily. 
Another very important distinction between period 2 elements and others is the ready formation of 
multiple bonds by C,