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Chapter 9 Ionic and Covalent Bonding. The shape of snowflakes results from bonding (and intermolecular) forces in H 2 O.

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Presentation on theme: "Chapter 9 Ionic and Covalent Bonding. The shape of snowflakes results from bonding (and intermolecular) forces in H 2 O."— Presentation transcript:

1 Chapter 9 Ionic and Covalent Bonding

2 The shape of snowflakes results from bonding (and intermolecular) forces in H 2 O.

3 Contents and Concepts Ionic Bonds Molten salts and aqueous solutions of salts are electrically conducting. This conductivity results from the motion of ions in the liquids. It suggests the possibility that ions exist in certain solids, held together by the attraction of ions of opposite charges. Describing Ionic Bonds Electron Configurations of Ions Ionic Radii

4 Covalent Bonds Not all bonds can be ionic. Hydrogen, H 2, is a clear example in which there is a strong bond between two like atoms. The bonding in the hydrogen molecule is covalent. A covalent bond forms between atoms by the sharing of a pair of electrons. 4. Describing Covalent Bonds 5. Polar Covalent Bonds; Electronegativity 6. Writing Lewis Electron-Dot Formulas 7. Delocalized Bonding: Resonance 8. Exceptions to the Octet Rule 9. Lewis Formulas 10. Bond Length and Bond Order

5 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


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

8 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]3s 1 ) + Cl ([Ne]3s 2 3p 5 )  Na + ([Ne]) + Cl - ([Ne]3s 2 3p 6 ) The atom that transferred the electron(s) becomes a cation. The atom that gained the electron(s) becomes an anion.

9 Figure 9.1: Sodium chloride crystals E.R. Degginger/

10 Molecular view of the electron transfer between Na and Cl

11 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.

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

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

14 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.


16 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.

17 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.

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

19 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.

20 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 Mg 2+ ; the size of Cl - is larger than that of O 2-. Because the distance between Mg 2+ and O 2- is smaller than the distance between Na + and Cl -, the force between Mg 2+ and O 2- will be greater.

21 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.

22 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.

23 Give the electron configuration and the Lewis symbol for the chloride ion, Cl -. Cl - ] [ For chlorine, Cl, Z = 17, so the Cl - ion has 18 electrons. The electron configuration for Cl - is 1s21s2 2s22s2 2p62p6 3s23s2 3p63p6 The Lewis symbol for Cl - is

24 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.

25 Figure 9.6: Common transition- metal cations in aqueous solution Photo by James Scherer. ©Houghton Mifflin Company. All rights reserved.

26 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.

27 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.

28 Give the electron configurations of Mn and Mn 2+. Manganese, Z = 25, has 25 electrons;. Its electron configuration is 1s21s2 2s22s2 2p62p6 3s23s2 3p63p6 4s24s2 3d53d5 Mn 2+ has 23 electrons. When ionized, Mn loses the 4s electrons first; the electron configuration for Mn 2+ is 1s21s2 2s22s2 2p62p6 3s23s2 3p63p6 3d53d5


30 a.Fe 2+ : [Ar]3d 4 4s 2 No. The 4s 2 electrons would be lost before the 3d electrons. b.N 2- : [He]2s 2 2p 5 No. Nitrogen will gain three electrons to fill the shell, forming N 3-. c.Zn 2+ : [Ar]3d 10 Yes! d.Na 2+ : [He]2s 2 2p 5 No. Sodium will lose only its one valence electron, forming Na +. e.Ca 2+ : [Ne]3s 2 3p 6 Yes!

31 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.

32 Figure 9.7: Determining the iodide ion radius in the lithium iodide (LiI) crystal

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

34 Figure 9.8: Comparison of atomic and ionic radii


36 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.

37 Using the periodic table only, arrange the following ions in order of increasing ionic radius: Br -, Se 2-, Sr 2+. 35 Br 34 Se 38 Sr These ions are isoelectronic, so their size decreases with increasing atomic number: Sr 2+ < Br - < Se 2-

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