Chapter 9 Ionic and Covalent Bonding

1. Chapter 9Ionic and Covalent Bonding

2. A chemical bond is a strong attractive force that exists between certain atoms in a substance. • There are three types of chemical bonds: Ionic bonds Covalent bonds Metallic bonds

3. An ionic bond is a chemical bond formed by the electrostatic attraction between positive and negative ions.

4. An ionic bond forms when one or more electrons are transferred from the valence shell of one atom to the valence shell of another atom. • Na ([Ne]3s1) + Cl ([Ne]3s23p5)  • Na+ ([Ne]) + Cl- ([Ne]3s23p6) • The atom that transferred the electron(s) becomes a cation. • The atom that gained the electron(s) becomes an anion.

5. A Lewis electron-dot symbol is a notation in which the electrons in the valence shell of an atom or ion are represented by dots placed around the chemical symbol of the element. • Note: Dots are placed one to a side, until all four sides are occupied.

6. Table 9.1 illustrates the Lewis electron-dot symbols for second- and third-period atoms.

7. [ ] 2- Ca O O • Represent the transfer of electrons in forming calcium oxide, CaO, from atoms. + + Ca2+

8. Let’s look next at the energy involved in forming ionic compounds. • The energy to remove an electron is the ionization energy. • The energy to add an electron is the electron affinity.

9. The combination of ionization energy and electron affinity is still endothermic; the process requires energy. • However, when the two ions bond, more than enough energy is released, making the overall process exothermic.

10. The lattice energy is the change in energy that occurs when an ionic solid is separated into gas-phase ions. • It is very difficult to measure lattice energy directly. It can be found, however, by using the energy changes for steps that give the same result.

11. For example, to find the lattice energy for NaCl, we can use the following steps.

12. The process of finding the lattice energy indirectly from other thermochemical reactions is called the Born–Haber cycle.

13. Ionic substances are typically high-melting solids. There are two factors that affect the strength of the ionic bond. They are given by Coulomb’s law: • The higher the ionic charge, the stronger the force; the smaller the ion, the stronger the force.

14. Based on this relationship, we can predict the relative melting points of NaCl and MgO. • The charge on the ions of MgO is double the charge on the ions of NaCl. Because the charge is double, the force will be four times stronger. • The size of Na+ is larger than that of Mg2+; the size of Cl- is larger than that of O2-. Because the distance between Mg2+ and O2- is smaller than the distance between Na+ and Cl-, the force between Mg2+ and O2- will be greater.

15. Based on the higher charge and the smaller distance for MgO, its melting point of MgO should be significantly higher than the melting point of NaCl. • The actual melting point of NaCl is 801°C; that for MgO is 2800°C.

16. When we examine the electron configuration of main-group ions, we find that each element gains or loses electrons to attain a noble-gas configuration.

17. [ ] - Cl • Give the electron configuration and the Lewis symbol for the chloride ion, Cl-. For chlorine, Cl, Z = 17, so the Cl- ion has 18 electrons. The electron configuration for Cl- is 1s2 3p6 2p6 3s2 2s2 The Lewis symbol for Cl- is

18. Group IIIA to VA metals often exhibit two different ionic charges: one that is equal to the group number and one that is 2 less than the group number. • The higher charge is due to the loss of both the s subshell electrons and the p subshells electron(s). The lower charge is due to the loss of only the p subshell electron(s). • For example, in Group IVA, tin and lead each form both +4 and +2 ions. In Group VA, bismuth forms +5 and +3 ions.

19. Polyatomic ions are atoms held together by covalent bonds as a group and that, as a group, have gained or lost one or more electron.

20. Transition metals form several ions. • The atoms generally lose the ns electrons before losing the (n– 1)d electrons. • As a result, one of the ions transition metals generally form is the +2 ion.

21. Give the electron configurations of Mn and Mn2+. Manganese, Z = 25, has 25 electrons;. Its electron configuration is 1s2 2s2 3p6 2p6 3s2 3d5 4s2 Mn2+ has 23 electrons.When ionized, Mn loses the 4s electrons first; the electron configuration for Mn2+ is 1s2 2p6 2s2 3s2 3d5 3p6

23. a. Fe2+: [Ar]3d44s2 • No. The 4s2 electrons would be lost before the 3d electrons. • b. N2-: [He]2s22p5 • No. Nitrogen will gain three electrons to fill the shell, forming N3-. • c. Zn2+: [Ar]3d10 • Yes! • d. Na2+: [He]2s22p5 • No. Sodium will lose only its one valence electron, forming Na+. • e. Ca2+: [Ne]3s23p6 • Yes!

24. Ionic radius is a measure of the size of the spherical region around the nucleus of an ion within which the electrons are most likely to be found. While ionic radius, like atomic radius, can be somewhat arbitrary, it can be measured in ionic compounds.

25. A cation is always smaller than its neutral atom. An anion is always larger than its neutral atom.

26. The term isoelectronic refers to different species having the same number and configuration of electrons. • For example, Ne, Na+, and F- are isoelectronic. • Ionic radius for an isoelectronic series decreases with increasing atomic number.

27. 34 Se 35 Br 38 Sr • Using the periodic table only, arrange the following ions in order of increasing ionic radius: Br-, Se2-, Sr2+. These ions are isoelectronic, so their size decreases with increasing atomic number: Se2- Br- < Sr2+ <

28. A covalent bond is a chemical bond formed by sharing a pair of electrons.

29. To consider how a covalent bond forms, we can monitor the energy of two isolated hydrogen atoms as they move closer together. • The energy decreases—first gradually, and then more steeply—to a minimum. As the atoms continue to move closer, it increases dramatically. • The distance between the atoms when energy is at a minimum is called the bond length. • This is illustrated on the following graph, from right to left.

31. As the hydrogen atoms move closer together, the electron of each atom is attracted to both its own nucleus and the nucleus of the second atom. The electron probability distribution illustrates this relationship.

32. A formula using dots to represent valence electrons is called a Lewis electron-dot formula. • An electron pair is represented by two dots. • A electron pair that is between two atoms is a bonding pair. It can also be represented by one line for each bonding pair. • Electron pairs that are not bonding are nonbonding, or lone pair electrons.

33. A coordinate covalent bond is formed when both electrons of the bond are donated by one atom. • The two electrons forming the bond with the hydrogen on the left were both donated by the nitrogen. Once shared, they are indistinguishable from the other N—H bonds.

34. In forming covalent bonds, atoms tend toward having a full eight electrons in their valence shell. This tendency is called the octet rule. • Hydrogen is an exception to the octet rule: it has two electrons in its valence shell (a duet).

35. A single bond is a covalent bond in which one pair of electrons is shared by two atoms. • A double bond is a covalent bond in which two pairs of electrons are shared by two atoms. • A triple bond is a covalent bond in which three pairs of electrons are shared by two atoms. • Double bonds form primarily with C, N, and O. Triple bonds form primarily with C and N.

36. A polar covalent bond (or polar bond) is a covalent bond in which the bonding electrons spend more time near one atom than near the other atom. • Electronegativity,X, is a measure of the ability of an atom in a molecule to draw bonding electrons to itself. Electronegativity is related to ionization energy and electron affinity.

37. Electronegativity increases from left to right and from bottom to top in the periodic table. F, O, N, and Cl have the highest electronegativity values.

38. The difference in electronegativity between the two atoms in a bond is a rough measure of bond polarity. • When the difference is very large, an ionic bond forms. When the difference is large, the bond is polar. When the difference is small, the bond is nonpolar.

39. Using electronegativities, arrange the following bonds in order by increasing polarity: C—N, Na—F, O—H. For C—N, the difference is 3.0 (N) – 2.5 (C) = 0.5. For Na—F, the difference is 4.0 (F) – 0.9 (Na) = 3.1. For O—H, the difference is 3.5 (O) – 2.1 (H) = 1.4. O—H < C—N < Na—F Bond polarities:

40. Writing Lewis Electron-Dot Formulas • Calculate the number of valence electrons. • Write the skeleton structure of the molecule or ion. • Distribute electrons to the atoms surrounding the central atom or atoms to satisfy the octet rule. • Distribute the remaining electrons as pairs to the central atom or atoms.

41. Write the electron dot formulas for the following: • a. OF2 • b. NF3 • c. NH2OH, hydroxylamine

42. Count the valence electrons in OF2: O 1(6) F 2(7) 20 valence electrons O is the central atom (it is less electronegative). Now, we distribute the remaining 16 electrons, beginning with the outer atoms. The last four electrons go on O.

43. Count the valence electrons in NF3: N 1(5) F 3(7) 26 valence electrons N is the central atom (it is less electronegative). Now, we distribute the remaining 20 electrons, beginning with the outer atoms. The last two electrons go on N.

44. Count the electrons in NH2OH: N 1(5) H 3(1) O 1(6) 14 valence electrons N is the central atom. Now, we distribute the remaining six electrons, beginning with the outer atoms. The last two electrons go on N.

45. Write electron-dot formulas for the following: • a. CO2 • b. HCN

46. Count the electrons in CO2: C 1(4) O 2(6) 16 valence electrons C is the central atom. Now, we distribute the remaining 12 electrons, beginning with the outer atoms. Carbon does not have an octet, so two of the lone pairs shift to become a bonding pair, forming double bonds.

47. Count the electrons in HCN: H 1(1) C 1(4) N 1(5) 10 valence electrons. C is the central atom. The remaining electrons go on N. Carbon does not have an octet, so two of the lone pairs shift to become a bonding pair, forming a triple bond.

48. Phosphorus pentachloride exists in solid state as the ionic compound [PCl4]+[PCl6]-; it exists in the gas phase as the PCl5 molecule. Write the Lewis formula of the PCl4+ ion.

49. Count the valence electrons in PCl4+: P 1(5) Cl 4(7) -1 32 P is the central atom. The remaining 24 nonbonding electrons are placed on Cl atoms. Add square brackets with the charge around the ion.

50. Delocalized bonding is a type of bonding in which a bonding pair of electrons is spread over a number of atoms rather than being localized between two atoms.