1. Chapter 6: Ionic Bonds and Some Main-Group Chemistry
2/23/2019
Copyright © 2008 Pearson Prentice Hall, Inc.

2. Ionic Bonding Occurs in ionic compound
Results from transferring electron
Created a strong attraction among the closely pack compound

3. Formation of ions
Ions of main-group elements (groups 1A – 7A) lose or gain electrons to achieve a noble gas configuration.
Atoms or ions with same electron configuration are isoelectronic.
The Aufbau principle applies to the formation of ionic compound:
Cations are formed through the loss of the highest energy electron(s) (a.k.a valence electrons) in an atom.
How will the electron configuration of Na and Al change when they become anions?
Na:
Al:
Anions are formed by adding electron(s) to the lowest-energy unoccupied orbital.
How will the electron configuration of O and Cl change when they become anions? O: Cl

4. Electron Configurations of Ions

5. Ions and their configuration
Although the main group elements form easily predicted ions, the transition metals don’t follow the same pattern.
Transition metals form cations of various oxidation states that are not isoelectronic with the noble gases.
Transition metals react with nonmetals to form cations by first losing their valence-shell s electrons then losing one or more d electrons.
Atoms
Ions
Fe:
[Ar] 4s2 3d6
- 2 e-
- 3 e-
Fe2+:
Fe3+:
[Ar] 3d6
[Ar] 3d5
Mn:
Mn2+

6. Chapter 6: Ionic Bonds and Some Main-Group Chemistry
Ionic Radii
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Atomic radii increases down a column of the periodic table because successively larger valence-shell orbitals are occupied, and radii decrease from left to right across a row of the periodic table because the Zeff increases across the row.
The cation has higher Zeff than the respective neutral element. This causes the remaining electrons in the ion to be attracted more strongly to the nucleus, reducing the size of the ion
Effective nuclear charge.
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7. Ionic Radii
The expansion that occurs when nonmetals gain electron to yield anions can’t account for the change of the quantum numbers of the valence shell.
Thus, the expansion is due entirely to the decrease in Zeff and the increase in electron-electron repulsions that occurs when an extra electron added.

8. Example Predicting Relative Sizes of Ions
Which of the following spheres represents a K+ ion, which a Ca2+ ion, and which a Cl− ion?

9. Chapter 6: Ionic Bonds and Some Main-Group Chemistry
Ionization Energy
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Ionization Energy (Ei): The amount of energy necessary to remove the highest-energy electron from an isolated neutral atom in the gaseous state.
Increasing Ei
Students sometimes mistakenly equate ionization energy to the amount of energy required to form an ion.
Decreasing Ei
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10. First Ionization Energy
For each period, the lowest value Ei is for the alkali metal and the highest value for the noble gas.
As with the periodicity in atomic size, this trend can be attributed to the increase in Zeff as you go across period. The greater the effective positive charge that attract an electron, the more energy that will be required to remove it.
As you descend a group, the electron that will be lost is a higher “n” shell. The higher “n” orbital is higher energy and extends farther away from the positive charge in the nucleus, making it easier to remove electron

11. Chapter 6: Ionic Bonds and Some Main-Group Chemistry
Ionization Energy
2/23/2019
Boron has a lower Ei due to a smaller Zeff
(shielding by the 2s electrons)
Copyright © 2008 Pearson Prentice Hall, Inc.

12. Oxygen has a lower Ei, since the first electron is
removed from a filled orbital.
Chapter 6, Unnumbered Figure 2, Page 201

13. Higher Ionization Energies
M+ + e–
M + energy
M2+ + e–
M+ + energy
M3+ + e–
M2+ + energy
Successive ionization energies for a given element always increases (e.g Ei1 < Ei2 < Ei3). This is because an electron is removed from a neutral element (M) for Ei1, a monocation (M+), for Ei2, a dication (M2+), for Ei3, etc. The successive increase in positive charge causes a greater attraction for the electron and the larger ionization energy
First ionization energy, second ionization energy, third ionization energy, etc.

14. Higher Ionization of Energy
Larger jumps in successive energies are also found for other elements, as is indicated by the zigzag line. This evidence for the shell structure of electron energies predicted by quantum mechanics, and the special stability associate with noble electron configuration

15. Chapter 6: Ionic Bonds and Some Main-Group Chemistry
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Electron Affinity
Electron Affinity (Eea): The energy released when a neutral atom gains an electron to form an anion.
Books use different conventions. The one this book uses assigns a negative sign to electron affinity.
Metals tend to have low Ei and low Eea.
They tend to lose one or more electrons.
Nonmetals tend to have high Ei and high Eea.
They tend to gain one or more electrons.
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16. Electron Affinity
Large negative Eea’s are found for the halogen (F, Cl, Br, I) because each of these elements has both a high Zeff and room in its valence shell for an additional electron. Halide ions, halogens with negative charge, have a noble-gas electron configurations with filled s and p subshells and the attraction between the additional electron and the atomic nucleus is high.
Near – zero values of electron affinity indicate that element doesn’t want to pick up an electron and won’t form a stable anion. Among the elements with near-zero values are the noble gas. These elements already have a have a stable electron configuration with no orbital available in the valence; placing an additional electron into the next higher energy shell is unfavorable.

17. Comparing Electron Affinities of Different Elements
Why does nitrogen have a less favorable (more positive) Eea than its neighbors on either side, C and O?

18. Ionic Bonds and the Formation of Ionic Solids
Chapter 6: Ionic Bonds and Some Main-Group Chemistry
Ionic Bonds and the Formation of Ionic Solids
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1s2 2s2 2p63s1
1s2 2s2 2p6 3s23p5 Na + Cl
Na1+
Cl1-
1s2 2s2 2p6
1s2 2s2 2p6 3s2 3p6
A transfer of electrons from the metal to the nonmetal.
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