CH 9: Ionic and Covalent Bonding

1. Vanessa N. Prasad-Permaul Valencia Community College CHM 1045 CH 9: Ionic and Covalent Bonding

2. Ionic Bonds • Ionic Bonds: a chemical bond formed by the electrostatic attraction between positive and negative ions • Bond forms when one or more electrons are transferred from the valence shell of one atom to the valence shell of the other. • Cation: atom losing electrons • Anion: atom gaining electrons 2

3. Ionic Bonds Na + Cl Na+ + Cl- [Ne]3s1 + [Ne]3s23p5 [Ne] + [Ne]3s23p6 3

4. Ionic Bonds • Lewis Electron-Dot Symbols: electrons in the valence shell of an atom or ion are represented by dots placed around the elemental symbol. Na + Cl Na+ + [ Cl ] - • Valence electrons are those electrons with the highest principal quantum number (n).

5. Ionic Bonds Lewis Electron-Dot Symbols for Atoms in the Periodic Table of Elements

6. Ionic Bonds EXAMPLE 9.1: Use Lewis Electron-Dot symbols to represent the transfer of electrons form magnesium to fluorine atoms to form ions with noble gas configuration. Mg + Fl MgFl2 F + Mg + F [ F ]- + Mg2+ + [ F ]-

7. Ionic Bonds EXERCISE 9.1: Use Lewis Electron-Dot symbols to represent the transfer of electrons form magnesium to oxygen atoms to form ions with noble gas configuration.

8. Electron Configurations of Ions • Common monatomic ions found in compounds of the • main-group elements fall into 3 categories: • Cations of Groups IA to IIIA having noble gas or • pseudo-noble-gas configurations ; the ion charges • equal group numbers. • Cations of Groups IIIA to VA having the ns2 electrons configurations; the ion charges equal the group • numbers minus two. Tl+, Sn2+, Pb2+, Bi3+. • 3. Anions of Groups VA to VIIA having noble gas configurations; the ion charges equal the group number minus 8.

9. Electron Configurations of Ions EXAMPLE 9.2: Write the electron configuration and Lewis symbol for N3-. N = [He]2s22p3 + 3e- [He]2s22p6 N = N N3- = [ N ]3-

10. Electron Configurations of Ions EXERCISE 9.2: Write the electron configuration and Lewis symbol for Ca2+ and S2-.

11. Electron Configurations of Ions EXERCISE 9.3: Write the electron configuration and Lewis symbol for Pband Pb2+.

12. Electron Configurations of Ions EXAMPLE 9.3: Write the electron configuration and Lewis symbol for Fe2+ and Fe3+. Fe = [Ar]3d64s2 = Fe Fe2+ = [Ar]3d6 = [ Fe ]2+ Fe3+ = [Ar]3d5 = [ Fe ]3+

13. Electron Configurations of Ions EXERCISE 9.4: Write the electron configuration and Lewis symbol for Mnand Mn2+.

14. Ionic Radii • Ionic Radius: A measure of the size of the spherical region around the nucleus of an ion within the electrons are most likely to be found. • If the atom loses an electron, the cation will be • smaller. • The electron-electron repulsion is initially less. • orbitals can shrink to increase the • attraction of the electrons for the nucleus.

15. Ionic Radii EXERCISE 9.5: Which has the larger radius. S or S2-. Explain.

16. Ionic Radii EXERCISE 9.6: Using only the Periodic Table, arrange the following ions in order of increasing ionic radius : Sr2 , Mg2+, Ca2+ .

17. Ionic Radii The Comparison of Atomic and Ionic Radii ATOMIC RADIUS DECREASES ACROSS A PERIOD INCREASES DOWN A GROUP

18. Ionic Radii • Isoelectronic: Different species having the same number and configuration of electrons. • Na+ < Mg2+ < Al3+ • Decrease in atomic radius • As the charge increases, the orbitals contract due to • the greater attractive forces of the nucleus  the • ionic radius decreases with increasing atomic • number. • In general , across a period the cations decrease in radius. • As the anions are reached, there is an abrupt increase in radius and then the radius decreases again.

19. Ionic Radii EXAMPLE 9.4: Arrange the following ions in order of decreasing ionic radius: F-, Mg2+, O2-. F- = 1s22s22p6 Mg2+ = 1s22s22p6 O2- = 1s22s22p6 All are isoelectronic, so as the nuclear charge increases, the ionic radius decreases. O2- , F- , Mg2+

20. Ionic Radii EXERCISE 9.7: Arrange the following ions in order of increasing ionic radius: Cl-, Ca2+, P3-.

21. Covalent Bonds • Covalent Bonds are formed by sharing at least one pair of electrons. • The attraction (nucleus/electrons) outweighs the repulsions (electron/electron & nucleus/nucleus)

22. Covalent Bonds Every covalent bond has a characteristic length that leads to maximum stability: Bond Length

23. Strength of Covalent Bonds Energy required to break a covalent bond in an isolated gaseous molecule is called the bond dissociation energy. Same amount of energy released when the bond forms.

24. Electron-Dot Structures • The electron-dot structures provide a simple, but useful, way of representing chemical reactions. • Ionic: • Covalent:

25. Covalent Bonds Coordinate covalent bond: a bond formed when both electrons of the bond are donated by one atom A + B A B A + B A B H + H+ + NH3 H N H H

26. Covalent Bonds OCTET RULE • Group 1A tends to lose their ns1 valence shell electron to adopt a noble gas electron configuration. • Group 2A lose both ns2 • Group 3A lose all three ns2 np1 • Group 7A Gains one electron to attain noble gas configuration • Group 8A inert, rarely lose or gain electrons

27. Covalent Bonds • Single Bonds: • Double Bonds: • Triple Bonds:

28. Covalent Bonds

29. Covalent Bonds ELECTRONEGATIVITY: a measure of the ability of an atom in a molecule to draw bonding electrons to itself. • Bond polarity is due to electronegativity differences between atoms. • Pauling Electronegativity: is expressed on a scale where F = 4.0

30. Covalent Bonds Pauling Electronegativities

31. Covalent Bonds EXAMPLE 9.5: Use electronegativity values to arrange the following bonds in order of increasing polarity: P H, H O, C Cl. P H = 0.0 H O = 1.4 C Cl = 0.5 P H , C Cl , H O

32. Covalent Bonds EXERCISE 9.8: Use electronegativity values to arrange the following bonds in order of increasing polarity: C O, C S, H Br.

33. Drawing Lewis-Dot Structures Rule 1: Count the total valence electrons. Rule 2: Draw the structure using single bonds. Rule 3: Distribute the remaining electron pairs around the peripheral atoms. Rule 4: Put remaining pairs on central atom. Rule 5: Share lone pairs between bonded atoms to create multiple bonds.

34. Drawing Lewis-Dot Structures

35. Lewis-Dot Structures • NH2F Amino Fluoride: In this molecule, nitrogen is the central atom. • Rule 1: Number of electrons = 5 + (2 x 1) + 7 = 14 = 7 pairs Rule 2 Rule 3 Rule 4

36. Lewis-Dot Structures EXAMPLE 9.6: Sulfur dichloride SCl2, is a red fuming liquid used in the manufacture insecticides. Write the Lewis formula for the molecule. S = 6 Cl = 7 each for a total of 20 electrons :Cl : S : Cl: :Cl S Cl: ¨ ¨ ¨ ¨ ¨ ¨ ¨ ¨ ¨ ¨ ¨ ¨

37. Lewis-Dot Structures EXERCISE 9.9: Dichlorodifluoromethane CCl2F2, is a gas used as a refrigerant and aerosol propellant. Write the Lewis formula for the molecule.

38. Lewis-Dot Structures EXAMPLE 9.7: Carbonyl chloride or phosgene, COCl2, is a highly toxic gas used as a starting material for the preparation of polyurethane plastics. What is the electron dot structure of this compound? C = 4, O = 6, Cl =7 each for a total of 24 electrons : Cl : C : Cl : :O: : Cl C Cl : :O: ¨ ¨ ¨ ¨ ¨ ¨ ¨ ¨ ¨ ¨

39. Lewis-Dot Structures EXERCISE 9.10: Write the electron-dot structure of carbon dioxide.

40. Lewis-Dot Structures EXAMPLE 9.8: Obtain the electron-dot formula for the BF4- ion. B = 3, F = 7 each (7 x 4) = 28 for a total 31 electrons. It is an ion with one more electron so a total of 32 electrons. : F : : F : B : F : : F : ¨ ¨ : F : : F B F : : F : ¨ ¨ ¨ ¨ ¨ ¨ ¨ ¨ ¨ ¨ ¨ ¨

41. Lewis-Dot Structures EXERCISE 9.11: Write the electron-dot structure of: A. the hydronium ion, H3O+ B. The chlorite ion, ClO2-

42. Resonance Structures • When multiple structures can be drawn, the actual structure is an average of all possibilities. • The average is called a resonance hybrid. A straight double-headed arrow indicates resonance. O O O O O O

43. Lewis-Dot Structures EXAMPLE 9.9: Describe the electron structure of the carbonate ion CO32-, in terms of electron-dot formulas. C = 4, O = 3 x 6 = 18 for a total of 22 electrons, but it has gained two electrons so there is a total of 24 electrons. 2- 2- 2- : O : : O : : O : C CC : O : : O : : O : O : : O : O : ¨ ¨ ¨ ¨ ¨ ¨ ¨ ¨

44. Lewis-Dot Structures EXERCISE 9.12: Describe the bonding in NO3- using resonance formulas.

45. Formal Charge • Formal Charge: Determines the best resonance structure. • We determine formal charge and estimate the more accurate representation. Formal Charge = valence e- - # of e- in a bond - (# of lone-pair e-) 2

46. Formal Charge Cl = 7 – (2/2) – 6 = 0 O = 6 – (4/2) – 4 = 0 C = 4 – (8/2) – 0 = 0 ¨ ¨ :Cl C Cl: :O: ¨ ¨ ¨ ¨ Cl = 7 – (4/2) – 4 = +1 O = 6 – (2/2) – 6 = -1 C = 4 – (8/2) – 0 = 0 :Cl C Cl :O: + ¨ ¨ ¨ - ¨ ¨ + Cl C Cl: :O: Cl = 7 – (4/2) – 4 = +1 O = 6 – (2/2) – 6 = -1 C = 4 – (8/2) – 0 = 0 ¨ ¨ ¨ -

47. Formal Charge EXAMPLE 9.11: Write the Lewis formula that best describes the charge distribution in the sulfuric acid molecule, H2SO4, according to the rules of formal charge. :O: H O S O H H :O: ¨ ¨ ¨ O ¨ ¨ +2 ¨ ¨ ¨ ¨ O S O H ¨ ¨ ¨ ¨ O ¨

48. Exercise 9.15: Write the Lewis formula that best describes the phosphoric acid molecule, H3PO4.

49. Resonance Structures • How is the double bond formed in O3? • The correct answer is that both are correct,but neither is correct by itself.

50. Example 1: Which of the following is correct? • Energy is absorbed to form a bond • Energy is released when a bond is formed