Chapter 7. Chemical Bonds

1. Chapter 7.Chemical Bonds Chemical bonds are the attractive forces that hold atoms together in a complex unit.

2. Types of Bonds Ionic bonds are electrostatic forces between ions. Oppositely charged ions attract one another. Covalent bonds form when atoms share electrons. Bonds give shapes to molecules Atoms are not charged

3. Valence Electrons Valence electrons are the electrons in the outermost s sub-shell and electrons in any other unfilled subshell. Valence electrons are the least tightly held electrons in an atom, and they are the ones that "do chemistry". An element's position on the Periodic Table shows how many valence electrons it has.

4. Valence Electrons Representative elements (Groups 1A – 7A) have 1 to 7 electrons. We can represent them as dots around the element's symbol. These are called Lewis Symbols.

5. Lewis Symbols of 1st 20 Elements

6. Stable Electron Configurations Atoms are most stable when they have filled valence shells. For representative elements, each atom wants 8 electrons surrounding it. Transition metals are most stable with filled or half-filled outer s and d subshells. Atoms will lose, gain or share electrons to reach these stable configurations.

7. The Octet Rule Representative elements want 8 electrons in their valence shells. Noble gases have filled valence shells. Group 6A and 7A nonmetals gain elec- trons to fill their valence shells. Become anions Group 1A, 2A, and 3A metals lose elec- trons to fill their valence shells. Become cations

8. Predict Charges For ions of the following metals: Na, Ca, Al For ions of the following nonmetals: F, O, P

9. Ionic Bonds Ionic bonds are chemical bonds that result from the attraction of positive and negative ions for each other. Metal cation(s) and nonmetal anion(s) Higher charges, stronger bonds

10. Ionic Bonds Ionic compounds are chemical compounds characterized by ionic bonds between atoms. An ionic compound will have a formula such that charge neutrality is achieved. Na1+ + Cl1- NaCl Mg2+ + 2 Cl1- MgCl2 Al3+ + 3 Cl1- AlCl3

11. Structures of Ionic Compounds Ionic bonds are non-directional electrostatic forces Ions arrange themselves in an array (usually a crystal lattice) that Maximizes interactions between oppositely charged ions Minimizes interactions between ions of like charge

12. Structure of NaCl(2 dimensions)

13. Structure of NaCl(3 dimensions)

14. Other Crystal Structures

15. Details about Ionization

16. Ionization Energy

17. Electron Affinity

18. Covalent Bonds Covalent bonds are chemical bonds that result when two nuclei attract and share the same electrons. Covalent bonds form between atoms of nonmetals.

19. Covalent Bonds Covalent compounds are characterized by covalent bonds between atoms. A covalent compound will have a formula such that all the atoms share electrons in such a way that the octet rule is satisfied.

20. Lewis Structures Lewis Structures are groupings of Lewis symbols that show transfer of electrons in ionic compounds or sharing of electrons in covalent compounds.

21. Lewis Structure of H2

22. Bond Distance

23. Lewis Structures of Diatomic Molecules

24. Lewis Structure for Ammonia (NH3)

25. Multiple Covalent Bonds Two atoms can share more than two electrons.

26. Drawing Lewis Structures Determine the number of electrons A. Needed to give each atom an octet except hydrogen, which needs two B. Available (all valence electrons for all atoms in structure) C. Shared = Needed  Available

27. Drawing Lewis Structures Choose the central atom for the structure A. NEVER hydrogen B. CARBON if it's in the formula C. Choose atom furthest left and/or down on the Periodic Table

28. Drawing Lewis Structures Draw structure A. Write symbol of central atom B. Arrange other atoms around it C. Show shared electrons as dots between atoms D. Show other available electrons as dots around atoms

29. Drawing Lewis Structures Check structure A. Is the right atom in the center? B. Have the correct number of electrons been shared? C. Are all available electrons shown? D. Is the octet rule satisfied?

30. Drawing Lewis Structures Examples: H2O CO2 HCN SO3

31. Drawing Lewis Structures Resonance Structures are two or more Lewis Structures for the same species that differ only in the position of the electrons. There must be at least one double bond in the structure for there to be resonance.

32. Drawing Lewis Structures Polyatomic Ions are charged groups of atoms, held together by covalent bonds. They often appear in parentheses in chemical formulas, e.g. Al2(SO4)3. Draw their structures in the usual way, but account for the charge in the available electrons. Examples: PO43 NH41+ CO32

33. Drawing Lewis Structures Acids of Polyatomic Anions: Draw the structure of the anion, with hydrogen atoms bonded to the oxygen atoms. Examples: H3PO4 H2CO3

34. Molecular Geometry Lewis Structures give numbers and types of bonds in molecules and polyatomic ions. Lewis Structures do not convey information about the shapes of the molecules and ions.

35. Molecular Geometry VSEPR: Valence Shell Electron Pair Repulsion Theory. VSEPR theory is an explanation of the shapes of simple molecules that uses Lewis Struc-tures. It is based on the fact that electrons repel each other, and groups of electrons will get as far away from each other as possible in a molecule or polyatomic ion.

36. Molecular Geometry VSEPR electron groups are groups of valence electrons that are present in a localized region in a molecule. Each bond is a group (single or multiple) Each non-bonded pair is a group If the central atom follows the octet rule, there can be 2, 3, or 4 electron groups around it.

37. Molecular Geometry Electronic geometries are descriptions of the arrangement of electrons about the central atom in a molecule or ion. # of Electron Regions Electronic about Central Atom Geometry 2 Linear 3 Trigonal Planar 4 Tetrahedral

38. Molecular Geometry The Three Possible Electronic Geometries (where the electrons are)

39. Molecular Geometry What? That's not where you told me the atomic orbitals were! Nope! Now we're looking at molecules, with hybrid orbitals. Hybrid orbitals are combinations of atomic orbitals. These give rise to the shapes of molecules and polyatomic ions. A specific type of hybrid orbital is associated with each electronic geometry.

40. Linear Electronic Geometry, sp Hybridization

41. Trigonal Planar Electronic Geometry, sp2 Hybridization

42. Tetrahedral Electronic Geometry, sp3 Hybridization

43. Molecules with Linear Electronic Geometry, sp Hybridization Molecular Geometry is Linear CO2 (carbon dioxide) HCN (hydrogen cyanide) C2H2 (acetylene)

44. Molecules with Trigonal Planar Electronic Geometry, sp2 Hybridization Molecular Geometry SO3 (sulfur trioxide) Trigonal Planar SO2 (sulfur dioxide) Angular/Bent

45. Molecules with Tetrahedral Electronic Geometry, sp3 Hybridization Molecular Geometry CH4 (methane) Tetrahedral NH3 (ammonia) Trigonal Pyramidal H2O (water) Angular/Bent

46. Molecular Geometry Examples: CCl4 (carbon tetrachloride) H3O1+ (hydronium ion) C2H4 (ethylene)

47. Types of Bonds  (sigma) bonds result from end-to-end overlap of orbitals, electron density is on bond axis.

48. Types of Bonds • (pi) bonds result from parallel overlap of orbit-als; there is no electron density on bond axis.

49. Types of Bonds There can be more than one  bond between two atoms.

50. Types of Bonds Resonance is really a  bond distributed over several  bonds, as in the carbonate anion