2. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Table of Contents Chapter 6 Chemical Bonding Section 1 Introduction to Chemical Bonding Section 2 Covalent Bonding and Molecular Compounds Section 3 Ionic Bonding and Ionic Compounds Section 4 Metallic Bonding Section 5 Molecular Geometry

3. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Objectives Define chemical bond. Explain why most atoms form chemical bonds. Describe ionic and covalent bonding. Explain why most chemical bonding is neither purely ionic nor purely covalent. Classify bonding type according to electronegativity differences. Section 1 Introduction to Chemical Bonding Chapter 6

4. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Chemical Bond An attractive force that holds atoms or ions together

5. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Types of Chemical Bonds Ionic -electrical attraction between positive and negative ions -transfer of electrons -usually a metal and a nonmetal Cation- positive ion Anion- negative ion

6. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Types of Chemical Bonds, cont’d Covalent -sharing of electrons between two atoms to make a bond -can be: nonpolar covalent (equal sharing) polar covalent (unequal sharing)

7. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Types of Chemical Bonds, cont’d Determining the Type of Bond -bonds are not 100% ionic or covalent -use the table of electronegativities (pg.161) 3.3 – 1.8 ionic bonds 1.7- 0.3 polar covalent 0.2- 0 nonpolar covalent

8. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Chemical Bonding, continued Sample Problem A Use electronegativity values listed on page 161, and Figure 2 in your book, on page 176, to classify bonding between sulfur, S, and the following elements: hydrogen, H; cesium, Cs; and chlorine, Cl. In each pair, which atom will be more negative? Section 1 Introduction to Chemical Bonding Chapter 6

9. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Objectives Define molecule and molecular formula. Explain the relationships among potential energy, distance between approaching atoms, bond length, and bond energy. State the octet rule. Chapter 6 Section 2 Covalent Bonding and Molecular Compounds

10. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Molecular Compounds A molecule is a neutral group of atoms that are held together by covalent bonds. A chemical compound whose simplest units are molecules is called a molecular compound. Section 2 Covalent Bonding and Molecular Compounds Chapter 6

11. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Visual Concepts Molecule Chapter 6

12. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Molecular Compounds One type of molecular compound is called a Diatomic molecule examples- O 2, this is composed of two atoms of oxygen others:

13. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Formation of a Covalent Bond Atoms have lower potential energy when they are bonded to other atoms The figure below shows potential energy changes during the formation of a hydrogen-hydrogen bond. Section 2 Covalent Bonding and Molecular Compounds Chapter 6

14. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Characteristics of the Covalent Bond The distance between two bonded atoms at their most stable energy state is the bond length. In forming bond, atoms release energy. The same amount of energy must be added to separate the bonded atoms. This is called Bond energy. Section 2 Covalent Bonding and Molecular Compounds Chapter 6

15. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Bond Energies & Bond Lengths for Single Bonds- Inverse relationship Section 2 Covalent Bonding and Molecular Compounds Chapter 6

16. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Characteristics of the Covalent Bond When two atoms form a covalent bond, their outer orbitals overlap. This achieves a noble-gas configuration.

17. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu The Octet Rule Bond formation follows the octet rule: Compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest energy level. Become like noble gases. Section 2 Covalent Bonding and Molecular Compounds Chapter 6

18. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Exceptions to the Octet Rule Atoms have more or less than 8 electron in outer energy level Hydrogen - only two electrons. Boron – only six electrons. Main-group elements in Periods 3-7 can form bonds with expanded valence, involving more than eight electrons. Section 2 Covalent Bonding and Molecular Compounds Chapter 6

19. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Electron-Dot Notation Bonding involves the outer s & p electrons (valence electrons). Using a dot notation shows the outer electrons each atom has for bonding.

20. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Electron-Dot Notation, continued Sample Problem B a.Write the electron-dot notation for hydrogen. b. Write the electron-dot notation for nitrogen. Section 2 Covalent Bonding and Molecular Compounds Chapter 6

21. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Objectives, continued Use the basic steps to write Lewis structures. Determine Lewis structures for molecules containing single bonds, multiple bonds, or both. Explain why scientists use resonance structures to represent some molecules. Section 2 Covalent Bonding and Molecular Compounds Chapter 6

22. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Lewis Structures Electron-dot notation can also be used to represent molecules. Section 2 Covalent Bonding and Molecular Compounds Chapter 6 The pair of dots between the two symbols represents the shared electron pair

23. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Lewis Structures The pair of dots between the two symbols represents the shared pair of a covalent bond. Section 2 Covalent Bonding and Molecular Compounds Chapter 6 In addition, each atom is surrounded by three pairs of electrons that are not shared in bonds. An unshared pair, also called a lone pair, is a pair of electrons that is not involved in bonding and that belongs exclusively to one atom.

24. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Lewis Structures The pair of dots representing a shared pair of electrons in a covalent bond is often replaced by a long dash. example: Section 2 Covalent Bonding and Molecular Compounds Chapter 6

25. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Rules for Drawing Lewis Structures 1.Determine the total valence electrons for all atoms. 2.Draw a skeleton. Central atom is C or least electronegative. Never H. 3.Give each atom an octet. (some exceptions) 4.Double or triple bond if you don’t have enough electrons. (COPSN) 5.Extra electrons? Put on the central atom.

26. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Lewis Structures, continued Examples: Draw the Lewis structure of a. iodomethane, CH 3 I. b. Ammonia, NH 3 c. Silane, SiH 4 d. Phosphorous trifluoride, PF 3 Section 2 Covalent Bonding and Molecular Compounds Chapter 6

27. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Multiple Covalent Bonds A double bond- two pairs of electrons are shared between two atoms. A triple bond-three pairs of electrons are shared between two atoms. In general, double bonds have greater bond energies and are shorter than single bonds. Triple bonds are even stronger and shorter than double bonds. Section 2 Covalent Bonding and Molecular Compounds Chapter 6

28. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Lewis Structures with Multiple Bonds Examples: a.Formaldehyde, CH 2 O b.Carbon dioxide, CO 2 c.Hydrogen cyanide, HCN d.Ethyne, C 2 H 2

29. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Polyatomic Ions A charged group of covalently bonded atoms is known as a polyatomic ion. Like other ions, polyatomic ions have a charge that results from either a shortage or excess of electrons. Section 3 Ionic Bonding and Ionic Compounds Chapter 6

30. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Some examples of Lewis structures of polyatomic ions nitrate NO 3 -1 ammonium, NH 4 +1 sulfate, SO 4 -2 Section 3 Ionic Bonding and Ionic Compounds Chapter 6 Polyatomic Ions, continued

31. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Objectives Compare a chemical formula for a molecular compounds with one for an ionic compound. Discuss the arrangements of ions in crystals. Define lattice energy and explain its significance. List and compare the distinctive properties of ionic and molecular compounds. Write the Lewis structure for a polyatomic ion given the identity of the atoms combined and other appropriate information. Section 3 Ionic Bonding and Ionic Compounds Chapter 6

32. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Ionic Compounds Most of the rocks and minerals that make up Earth’s crust consist of positive and negative ions held together by ionic bonding. example: table salt, NaCl, An ionic compound is composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal. Section 3 Ionic Bonding and Ionic Compounds Chapter 6

33. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Ionic Compounds Most ionic compounds exist as crystalline solids. In contrast to a molecular compound, an ionic compound is not composed of independent, neutral units that can be isolated. Section 3 Ionic Bonding and Ionic Compounds Chapter 6

34. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Ionic Compounds, continued The chemical formula of an ionic compound represents the simplest ratio of the compound’s ions. A formula unit is the simplest collection of atoms from which an ionic compound’s formula can be established. Section 3 Ionic Bonding and Ionic Compounds Chapter 6

35. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Ionic Vs. Covalent Bonding Section 3 Ionic Bonding and Ionic Compounds Chapter 6

36. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Formation of Ionic Compounds Section 3 Ionic Bonding and Ionic Compounds Chapter 6

37. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Formation of Ionic Compounds, continued In an ionic crystal, ions minimize their potential energy by combining in an orderly arrangement known as a crystal lattice. The combined attractive and repulsive forces within a crystal lattice determine: the distances between ions. the pattern of the ions’ arrangement in the crystal Section 3 Ionic Bonding and Ionic Compounds Chapter 6

38. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu NaCl and CsCl Crystal Lattices Section 3 Ionic Bonding and Ionic Compounds Chapter 6

39. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Objectives Describe the electron-sea model of metallic bonding, and explain why metals are good electrical conductors. Explain why metal surfaces are shiny. Explain why metals are malleable and ductile but ionic-crystalline compound are not. Section 4 Metallic Bonding Chapter 6

40. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Metallic Bonding Chemical bonding is different in metals than it is in ionic, molecular, or covalent-network compounds. The unique characteristics of metallic bonding gives metals their characteristic properties, listed below. electrical conductivity thermal conductivity malleability ductility shiny appearance Section 4 Metallic Bonding Chapter 6

41. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Properties of Substances with Metallic, Ionic, and Covalent Bonds Section 4 Metallic Bonding Chapter 6

42. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu The Metallic-Bond Model In a metal, the vacant orbitals in the atoms’ outer energy levels overlap. This overlapping of orbitals allows the outer electrons of the atoms to roam freely throughout the entire metal. The electrons are delocalized, which means that they do not belong to any one atom but move freely about the metal’s network of empty atomic orbitals. These mobile electrons form a sea of electrons around the metal atoms, which are packed together in a crystal lattice. Section 4 Metallic Bonding Chapter 6

43. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Metallic Bonding, continued Mobile electrons allow charge/heat to pass and allow the metal to be hammered because there is not a rigid structure. Section 4 Metallic Bonding Chapter 6

44. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Objectives Explain VSEPR theory. Predict the shapes of molecules or polyatomic ions using VSEPR theory. Explain the shapes of molecules or polyatomic ions using VSEPR theory. Section 5 Molecular Geometry Chapter 6

45. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Molecular Geometry The properties of molecules depend not only on the bonding of atoms but also on molecular geometry: the three-dimensional arrangement of a molecule’s atoms. Section 5 Molecular Geometry Chapter 6

46. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu VSEPR Theory As shown at right, diatomic molecules, like those of (a) hydrogen, H 2, and (b) hydrogen chloride, HCl, can only be linear because they consist of only two atoms. To predict the geometries of more-complicated molecules, one must consider the locations of all electron pairs surrounding the bonding atoms. This is the basis of VSEPR theory. Section 5 Molecular Geometry Chapter 6

47. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu VSEPR Theory VSEPR stands for “valence-shell electron-pair repulsion.” VSEPR theory states that repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible. Section 5 Molecular Geometry Chapter 6

48. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu VSEPR Theory Find the VSEPR formula A=central atom B=bonded atoms E=lone pairs of e- on the central atom Use the chart on pg.200 to determine the molecular shape Section 5 Molecular Geometry Chapter 6

49. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu VSEPR Theory, continued Sample Problems Use VSEPR theory to predict the molecular geometry of a. BCl 3 b. NH 3 c. CH 4 d. NO 3 -1 e. CO 2 Section 5 Molecular Geometry Chapter 6

50. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Unshared electron pairs repel other electron pairs more strongly than bonding pairs do. This is why the bond angles in ammonia and water are somewhat less than the 109.5° bond angles of a perfectly tetrahedral molecule. Section 5 Molecular Geometry Chapter 6 VSEPR Theory, continued

51. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Treat double and triple bonds the same way as single bonds. Treat polyatomic ions similarly to molecules. The next slide shows several more examples of molecular geometries determined by VSEPR theory. Section 5 Molecular Geometry Chapter 6 VSEPR Theory, continued

52. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu VSEPR and Molecular Geometry Section 5 Molecular Geometry Chapter 6

53. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu VSEPR and Molecular Geometry Section 5 Molecular Geometry Chapter 6

54. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Intermolecular Forces a.k.a. van der Waals Forces The forces of attraction between molecules are known as intermolecular forces. The boiling point is a good measure of the intermolecular forces between its molecules: the higher the boiling point, the stronger the forces between the molecules. Section 5 Molecular Geometry Chapter 6

55. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Intermolecular Forces, continued The strongest intermolecular forces exist between polar molecules. Identical atoms attached = nonpolar Different atoms attached = polar Polar bonds are the result of an asymmetrical distribution of electrons- causes a dipole Nonpolar bonds- equal distribution of electrons Section 5 Molecular Geometry Chapter 6

56. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Intermolecular Forces, continued A dipole is represented by an arrow with its head pointing toward the negative pole and a crossed tail at the positive pole. The dipole created by a hydrogen chloride molecule is indicated as follows: Section 5 Molecular Geometry Chapter 6

57. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Intermolecular Forces, continued Determine if each of the molecules is polar or nonpolar: CH 4 Cl 2 CCl 3 H Section 5 Molecular Geometry Chapter 6

58. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Intermolecular Forces a.k.a-van der Waals forces Three types 1. Hydrogen Bonding 2. Dipole-Dipole 3. London Dispersion Forces

59. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Hydrogen Bonding -Strongest -occurs between molecules that contain O-H, N-H or H-F atoms bonding. -example: H 2 O vs. H 2 S Boiling point 100 C-61 C

60. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Dipole-Dipole -2nd in strength - occurs in polar molecules - will see an increase in boiling point as the mass increases

61. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu London Dispersion Forces -3rd in strength -created from “momentary” dipoles that occur between two nonpolar molecules -also-increase in strength with increase in mass

62. Copyright © by Holt, Rinehart and Winston. All rights reserved. ResourcesChapter menu Intermolecular Forces Let’s try some: 1. Draw Lewis Structure 2. Determine if the molecule is polar or nonpolar 3. Pick based on polarity or if you see an HF, NH or OH you will know. 1. HCl 2. NH 3 3. CH 3 Cl 4. SiF 4