First ionization energy alkali metals

Experiments and calculations both indicate that electron transfer from potassium to water is spontaneous and rapid, whereas electron transfer from silver to water does not occur. In redox terms, potassium oxidizes easily, but silver resists oxidation. Because oxidation involves the loss of electrons, these differences in reactivity of silver and potassium can be traced to how easily each metal loses electrons to become an aqueous cation. One obvious factor is their first ionization energies, which show that it takes much more energy to remove an electron from silver than from potassium 731 kJ/mol for Ag and 419 kJ/mol for K. The other alkali metals with low first ionization energies, Na, Rb, Cs, and Fr, all react violently with water. 

The first ionization energy (f) of element X is relatively low when compared to I2 and I3. This means that X is probably a member of the Group I alkali metals. Thus, the formation of X2+ and X3+ would be difficult to achieve. Therefore, the formula is most likely to be XC1. 

TABLE 6.4 Properties of Alkali Metals Melting Boiling Point (°C) Point (°C) Density (g/cm3) First Ionization Energy (kj/mol) Abundance on Earth (%) Atomic Radius (pm) Ionic (M+) Radius (pm)... 

The alkaline earth metals undergo the same kinds of redox reactions that the alkali metals do, but they lose two electrons rather than one to yield dipositive ions, M2+. Because their first ionization energy is larger than that of alkali metals (Figure 6.3), the group 2A metals tend to be somewhat less reactive than alkali metals. The general reactivity trend is Ba > Sr > Ca > Mg > Be. 

The radii of the metals increase with increasing atomic number and their atomic sizes are the largest in their respective periods. Such features lead to relatively small first ionization energy (/1) for the atoms. Thus the alkali metals are highly reactive and form M+ ions in the vast majority of their compounds. The very high second ionization energy (I2) prohibits formation of the M2+ ions. Even though the electron affinities (T) indicate only mild exothermicity, M- ions can be produced for all the alkali metals (except Li) under carefully controlled conditions. 

The highest first ionization energy, about 2400 kJ mol-1, is for He. As a group, the noble gases have the highest ionization energies and the alkali metals have the lowest. 

The first ionization energy of H atom (1311kJmoU ) is very high in comparison with that of many other elements, such as alkali metals. Removal of the Is electron leaves a bare proton, which, having a radius of only about 1.5 x 10 pm, never exists in the condensed phase. However, when bonded to other species it is well known in solutions and in solids (e.g. H3O+, NH4+, etc.). The proton affinity of water and the solvation energy of the H+ ion in water have been estimated... 

TABLE 12.7 First Ionization Energies for the Alkali Metals and Noble Gases... 

Some important properties of the first five alkali metals are shown in Table 12.9. The data in Table 12.9 show that when we move down the group, the first ionization energy decreases and the atomic radius increases. This agrees with the general trends discussed in Section 12.15. 

For reactions of this type, the relative reducing powers of the alkali metals can be predicted from the first ionization energies listed in Table 12.9. Since it is much easier to remove an electron from a cesium atom than from a lithium atom, cesium should be the better reducing agent. The expected trend in reducing ability is... 

FIGURE 3.4 First ionization energy plotted versus atomic number shows periodic behavior. Symbols for the noble gases are shown in red those for alkali metals are shown in blue. 

Arrange the members of each of the following sets of elements in order of increasing first ionization energies (a) the alkali metals (b) the halogens (c) the elements in the second period (d) Br, F, B, Ga, Cs, and H. 

Compare the alkali metals with the alkaline earth metals with respect to (a) atomic radii, (b) densities, (c) first ionization energies, and (d) second ionization energies. Explain the comparisons. 

FIGURE 8.11 Variation of the first ionization energy with atomic number. Note that the noble gases have high ionization energies, whereas the alkali metals and alkaline earth metals have low ionization energies. 

The group 2A elements (the alkaline earth metals) have higher first ionization energies than the alkali metals do. The alkaline earth metals have two valence electrons (the outermost electron confignration is ns ). Because these two s electrons do not shield each other well, the effective nuclear charge for an alkaline earth metal atom is larger than that for the preceding alkali metal. Most alkaline earth compounds contain dipositive ions (Mg +, Ca, Sr, Ba +). The Be ion is isoelectronic with Li and with He, Mg is isoelectronic with Na and with Ne, and so on. 

Figure 8.11 Periodicity of first ionization energy (IE,). A plot of IEi vs. atomic number for the elements in Periods 1 through 6 shows a periodic pattern the lowest values occur for the alkali metals (brown) and the highest for the noble gases (purple). This is the inverse of the trend in atomic size (see Figure 8.10).

Since ionization energies decrease going down a colunrn in the periodic table, francium should have the lowest first ionization energy of all the alkali metals. As a result, Fr should be the most reactive of all the Group 1A elements toward water and oxygen. The reaction with oxygen would probably be similar to that of K, Rb, or Cs. 

Whereas the first Ionization energies the energy to remove one electron) of the alkali metals are relatively low, the second Ionization energies are very high. 

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