- Online Text
- 1. Introduction
- 2. Valence Electron Patterns and Lewis Structures
- 3. Ionic Bonds
- 4. Covalent Bonds and the Octet Rule
- 5. Polarity and Basic Lewis Structures
- 6. Advanced Lewis Structures
- 7. VSEPR Theory
- 8. Hybrid Orbitals
- 9. Intermolecular Forces
- 10. Physical Properties of Molecules
- 11. Conclusion
- 12. Further Reading
- Unit Guide (PDF)
Section 3: Ionic Bonds
Most of the elements in the periodic table are quite reactive, and this reactivity relates directly to the valence electrons of the elements. They have incomplete valence shells and thus are able to react with other substances to try to change the numbers of electrons in that outermost shell. As atoms react, they begin to form compounds, and the atoms of compounds are held together by bonds. A bond is an arrangement of a pair of atoms in such a way that those atoms want to stay associated with each other rather than separate. In all cases, the bonds between atoms—regardless of the type of bond—result from the electrons on the atoms that have come together.
Figure 5-4. Coulombic Forces Hold Ionic Bonds Together
By rubbing this balloon on her head, some of the electrons from this young lady's hair get transferred to the balloon, making it negatively charged and her hair positively charged. Since opposite charges attract, her hair stays connected to the balloon even as it floats away.
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Ionic bonds are the simplest sort of bond. For ionic bonds to happen, atoms must have opposite charges. Because of the charges, they attract each other and so stay paired up due to Coulombic forces. Coulombic forces happen when negatively-charged things are attracted to positively-charged things, or when like charges repel each other (Figure 5-4). How did these ions form? Electrons jumped from the valence shell of one atom to another. The result is one atom with extra electrons creates a negatively charged anion, and one atom with a shortfall of electrons creates a positively charged cation.
Figure 5-5. Formation of an Ionic Bond
The electron configuration of sodium (Na, 1s22s22p63s1) and chlorine (Cl, 1s22s22p63s23p5) makes it possible for them to form an ionic bond. After the transfer of an electron from the sodium to the chlorine, the compound sodium chloride (NaCl), common table salt, forms. Note that it is made up of Na+ (1s22s22p6) and Cl- (1s22s22p63s23p6), each of which has 8 valence electrons between its outermost s and p orbitals. For each of these, the atoms and ions are written as Lewis dot formulas.
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Ionic bonds are common between atoms with large differences in electronegativity. Recall from the last unit that electronegativity is a measure of an atom's "love" for electrons, so if one atom has a high electronegativity it would like to gain electrons, and vice versa. This means that ionic compounds usually occur between atoms on opposite sides of the periodic table, such as sodium and chloride, which form sodium chloride, NaCl, or calcium and oxygen, which form calcium oxide, CaO. In both cases, the ionic compounds form from one non-metallic element and one main group metal.
To understand this more deeply, we need to return to this magic number eight that comes from the eight possible valence electrons that an atom can have. There is increased stability if an atom is able to completely fill its valence shell. Thus, there is always motivation for elements that are close to eight electrons to gain more in order to reach eight, and these elements are the ones with high electronegativities. On the other hand, some elements have only one or two electrons in their valence shell, and in that case it is easier for these elements to give up these electrons and have an empty shell. These elements are usually metals and have very low electronegativities. This actually still means that these elements have a full valence shell after losing electrons, it is just one level below the one that gave up those electrons.
Let's return to sodium chloride (NaCl) which forms if sodium (Na) and chlorine (Cl) come together. If an electron departs from sodium and joins the chlorine, both resulting ions will have filled valence shells (Figure 5-5). The chlorine will have gained the one electron it needed to make a complete third shell. The sodium, by losing the one electron in its third shell will now have the second shell as its valence shell, and this shell is already complete with eight electrons.
Sodium, then, has a positive charge because it is short one electron but the number of protons in the nucleus hasn't changed. Chlorine, meanwhile, is negative because it now has an extra electron, one more than the number of protons in the nucleus. The opposite charges on these atoms then hold them together in an ionic bond.