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This is a€?Energetics of Ion Formationa€?, section 7.3 from the book Principles of General Chemistry (v.
This content was accessible as of December 29, 2012, and it was downloaded then by Andy Schmitz in an effort to preserve the availability of this book. PDF copies of this book were generated using Prince, a great tool for making PDFs out of HTML and CSS. For more information on the source of this book, or why it is available for free, please see the project's home page. DonorsChoose.org helps people like you help teachers fund their classroom projects, from art supplies to books to calculators. To correlate ionization energies, electron affinities, and electronegativities with the chemistry of the elements. We have seen that when elements react, they often gain or lose enough electrons to achieve the valence electron configuration of the nearest noble gas. Because atoms do not spontaneously lose electrons, energy is required to remove an electron from an atom to form a cation. If an atom possesses more than one electron, the amount of energy needed to remove successive electrons increases steadily. The energy required to remove electrons from a filled core is prohibitively large under normal reaction conditions.
From their locations in the periodic table, predict which of these elements has the highest fourth ionization energy: B, C, or N. B Determine whether electrons are being removed from a filled or partially filled valence shell.
Because carbon and nitrogen have four and five valence electrons, respectively, their fourth ionization energies correspond to removing an electron from a partially filled valence shell. From their locations in the periodic table, predict which of these elements has the lowest second ionization energy: Sr, Rb, or Ar. The first ionization energies of the elements in the first six rows of the periodic table are plotted in Figure 7.10 "A Plot of Periodic Variation of First Ionization Energy with Atomic Number for the First Six Rows of the Periodic Table". The changes seen in the second (Li to Ne), fourth (K to Kr), fifth (Rb to Xe), and sixth (Cs to Rn) rows of the s and p blocks follow a pattern similar to the pattern described for the third row of the periodic table. Because of the first two trends, the elements that form positive ions most easily (have the lowest ionization energies) lie in the lower left corner of the periodic table, whereas those that are hardest to ionize lie in the upper right corner of the periodic table. Generally, I1 increases diagonally from the lower left of the periodic table to the upper right.
There is a decrease in ionization energy within a group (most easily seen here for groups 1 and 18). The darkness of the shading inside the cells of the table indicates the relative magnitudes of the ionization energies. Gallium (Ga), which is the first element following the first row of transition metals, has the following electron configuration: [Ar]4s23d104p1. As we noted, the first ionization energies of the transition metals and the lanthanides change very little across each row. Because their first, second, and third ionization energies change so little across a row, these elements have important horizontal similarities in chemical properties in addition to the expected vertical similarities. Use their locations in the periodic table to predict which element has the lowest first ionization energy: Ca, K, Mg, Na, Rb, or Sr. Use their locations in the periodic table to predict which element has the highest first ionization energy: As, Bi, Ge, Pb, Sb, or Sn.
Unlike ionization energies, which are always positive for a neutral atom because energy is required to remove an electron, electron affinities can be negative (energy is released when an electron is added), positive (energy must be added to the system to produce an anion), or zero (the process is energetically neutral).
There are many more exceptions to the trends across rows and down columns than with first ionization energies. In general, electron affinities of the main-group elements become less negative as we proceed down a column.
The electron affinities of elements B through F in the second row of the periodic table are less negative than those of the elements immediately below them in the third row. In general, electron affinities become more negative across a row and less negative down a column. As we have seen, the first electron affinity can be greater than or equal to zero or negative, depending on the electron configuration of the atom. Before analyzing the difference between first and second ionization energy, let us first discuss what ionization energy is.
The first ionization energy is defined as the energy absorbed by 1 mol of neutral gaseous atoms to remove the most loosely bound electron from atom to produce 1 mol of gaseous ions with +1 charge. The second ionization energy is defined as the energy absorbed by 1 mol of positively charged gaseous ions to produce 1 mol of gaseous ions with a +2 charge, by removing the loosely bound electron from the +1 ion.
First ionization energy (I1E): The energy required to remove the most loosely bound electron from 1 mol of gaseous atoms to produce 1 mol of gaseous ions with a positive charge (+1). Second ionization energy (I2E): The energy required to remove the most loosely bound electron from 1 mol of gaseous ions with a +1 charge to produce  mol of gaseous ions with +2 charge.


Usually expelling the first electron from a ground state gaseous atom is easier than expelling the second electron from a positively charged ion. First ionization energy (I1E): First ionization energy values of atoms in every period show the same variation. See the license for more details, but that basically means you can share this book as long as you credit the author (but see below), don't make money from it, and do make it available to everyone else under the same terms. However, the publisher has asked for the customary Creative Commons attribution to the original publisher, authors, title, and book URI to be removed.
In this section, we develop a more quantitative approach to predicting such reactions by examining periodic trends in the energy changes that accompany ion formation. Larger values of I mean that the electron is more tightly bound to the atom and harder to remove.
Both Li+ and Be2+ have 1s2 closed-shell configurations, and much more energy is required to remove an electron from the 1s2 core than from the 2s valence orbital of the same element. Thus in the third row of the periodic table, the largest increase in ionization energy corresponds to removing the fourth electron from Al, the fifth electron from Si, and so fortha€”that is, removing an electron from an ion that has the valence electron configuration of the preceding noble gas. Predict which element has the highest fourth ionization energy, recognizing that the highest energy corresponds to the removal of electrons from a filled electron core.
The fourth ionization energy for boron, however, corresponds to removing an electron from the filled 1s2 subshell. This is because the valence electrons do not screen each other very well, allowing the effective nuclear charge to increase steadily across the row.
The electrons in aluminuma€™s filled 3s2 subshell are better at screening the 3p1 electron than they are at screening each other from the nuclear charge, so the s electrons penetrate closer to the nucleus than the p electron does.
They are presented numerically and graphically in Figure 7.11 "First Ionization Energies of the ". The transition metals are included in the fourth, fifth, and sixth rows, however, and the lanthanides are included in the sixth row. Although the principal quantum number n increases down a column, filled inner shells are effective at screening the valence electrons, so there is a relatively small increase in the effective nuclear charge. Consequently, ionization energies generally increase diagonally from lower left (Cs) to upper right (He). Its first ionization energy is significantly lower than that of the immediately preceding element, zinc, because the filled 3d10 subshell of gallium lies inside the 4p subshell, screening the single 4p electron from the nucleus. Differences in their second and third ionization energies are also rather small, in sharp contrast to the pattern seen with the s- and p-block elements. For example, all the first-row transition metals except scandium form stable compounds as M2+ ions, whereas the lanthanides primarily form compounds in which they exist as M3+ ions.
Based on trends in ionization energies across a row and down a column, identify the element with the lowest first ionization energy. Because we know that ionization energies increase from left to right in a row and from bottom to top of a column, we can predict that the element at the bottom left of the rectangle will have the lowest first ionization energy: Rb. This sign convention is consistent with our discussion of energy changes in Chapter 5 "Energy Changes in Chemical Reactions", where a negative value corresponded to the energy change for an exothermic process, which is one in which heat is released. Both figures show that the halogens, with their ns2np5 valence electron configuration, have the most negative electron affinities. Elements that do not form stable ions, such as the noble gases, are assigned an effective electron affinity that is greater than or equal to zero.
This is because as n increases, the extra electrons enter orbitals that are increasingly far from the nucleus. Apparently, the increased electrona€“electron repulsions experienced by electrons confined to the relatively small 2p orbitals overcome the increased electrona€“nucleus attraction at short nuclear distances. As you learned in Chapter 6 "The Structure of Atoms", the energy separation between the filled ns2 and the empty np subshells decreases with increasing n, so that formation of an anion from the heavier elements becomes energetically more favorable.
In contrast, the second electron affinity is always positive because the increased electrona€“electron repulsions in a dianion are far greater than the attraction of the nucleus for the extra electrons. In general, the ionization energy is referred to as the energy required to remove an electron from a gaseous atom or an ion. The magnitude of first ionization energy increases along a period in the periodic table and decreases along a group.
Therefore, the first ionization energy is less than the second ionization energy and the energy difference between first and second ionization energy is significantly large.
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First, the second electron is being removed from a positively charged species rather than a neutral one, so in accordance with Coulomba€™s law, more energy is required.
The chemical consequences are enormous: lithium (and all the alkali metals) forms compounds with the 1+ ion but not the 2+ or 3+ ions. This pattern explains why the chemistry of the elements normally involves only valence electrons.
The valence electrons are therefore attracted more strongly to the nucleus, so atomic sizes decrease and ionization energies increase.


The first ionization energies of the transition metals are somewhat similar to one another, as are those of the lanthanides. Experiments have revealed something of even greater interest: the second and third electrons that are removed when gallium is ionized come from the 4s2 orbital, not the 3d10 subshell. As the d orbitals are filled, the effective nuclear charge causes the 3d orbitals to be slightly lower in energy than the 4s orbitals. In general, electron affinities become more negative as we go across a row of the periodic table. Atoms with the largest radii, which have the lowest ionization energies (affinity for their own valence electrons), also have the lowest affinity for an added electron.
Since electrons are attracted to the positive nucleus, the energy needs to be supplied for this process.
Second, removing the first electron reduces the repulsive forces among the remaining electrons, so the attraction of the remaining electrons to the nucleus is stronger.
Similarly, beryllium (and all the alkaline earth metals) forms compounds with the 2+ ion but not the 3+ or 4+ ions.
These effects represent two sides of the same coin: stronger electrostatic interactions between the electrons and the nucleus further increase the energy required to remove the electrons.
This makes the S atom slightly less stable than would otherwise be expected, as is true of all the group 16 elements. Ionization energies increase from left to right across each row, with discrepancies occurring at ns2np1 (group 13), ns2np4 (group 16), and ns2(nA a?’A 1)d10 (group 12) electron configurations. Valence electrons that are farther from the nucleus are less tightly bound, making them easier to remove, which causes ionization energies to decrease. The chemistry of gallium is dominated by the resulting Ga3+ ion, with its [Ar]3d10 electron configuration. The [Ar]3d2 electron configuration of Ti2+ tells us that the 4s electrons of titanium are lost before the 3d electrons; this is confirmed by experiment.
This pattern corresponds to the increased effective nuclear charge felt by the valence electrons across a row, which leads to increased electrostatic attraction between the added electron and the nucleus (a more negative electron affinity).
Consequently, the elements of the third row (nA =A 3) have the most negative electron affinities. The energy required to remove electrons from a filled core is prohibitively large and simply cannot be achieved in normal chemical reactions. This and similar electron configurations are particularly stable and are often encountered in the heavier p-block elements.
A similar pattern is seen with the lanthanides, producing cations with an (nA a?’A 2)fn valence electron configuration. Farther down a column, the attraction for an added electron decreases because the electron is entering an orbital more distant from the nucleus. They are sometimes referred to as pseudo noble gas configurationsThe (na?’1)d10 and similar electron configurations that are particularly stable and are often encountered in the heavier p-block elements.. Some of the alkaline earths (group 2), the elements of group 12, and all the noble gases (group 18) have effective electron affinities that are greater than or equal to zero, while the electron affinities for the elements of group 15 are usually less negative than those for the group 14 elements. Electrona€“electron repulsions also decrease because the valence electrons occupy a greater volume of space. In fact, for elements that exhibit these configurations, no chemical compounds are known in which electrons are removed from the (nA a?’A 1)d10 filled subshell. These effects tend to cancel one another, so the changes in electron affinity within a family are much smaller than the changes in ionization energy.
In other words; first ionization energy relates to neutral gaseous atoms and second ionization energy relates to the gaseous ions with a (+1) charge. Both the alkaline earth metals and the noble gases have valence electron shells with filled subshells (ns2 and ns2np6, respectively).
The magnitude of the ionization energy varies depending on the charge of the nucleus, the distance of the electron form of the nucleus and the number of electrons between the nucleus and the outer shell electrons. In each case, the added electron must enter a higher-energy orbital, requiring an input of energy.
All the group 15 elements have an ns2np3 valence electron configuration, in which each of the three p orbitals has a single electron, in accordance with Hunda€™s rule; hence the added electron must enter an already occupied p orbital.
The resulting electrona€“electron repulsions destabilize the anion, causing the electron affinities of these elements to be less negative than we would otherwise expect. In the case of nitrogen, the 2p orbital is quite small, and the electrona€“electron repulsions are so strong that nitrogen has approximately zero affinity for an extra electron. In the heavier elements, however, the effect is relatively small because they have larger valence p orbitals.



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