1. UNIT 7: BONDING What previous knowledge will help us understand bonding? How can we describe energy involved in a chemical bond? How can we explain and draw ionic bonds? How can we explain and draw covalent bonds? What are metallic bonds and why are they good conductors? What is the difference between bond polarity and molecule polarity? How are molecules geometrically arranged? How does the VESPR theory influence the geometry of molecules? What are the different forces that hold molecules together? What are sigma and pi bonds?

2. AIM: Prior Knowledge   ENDOTHERMIC: chemical reaction that absorbs heat, producing products with more PE than the reactants   EXOTHERMIC: chemical reaction that produces heat with less PE than the reactants   POTENTIAL ENERGY: stored energy based on position or composition   Why do atoms become ions? To become stable   How do atoms become ions? Gain or loss of electrons   How do metals form ions? MELPS   How do nonmetals form ions? Opposite of MELPS

3. BONDING  Chemical bonds provide the glue that holds all compounds together  The electron structure of atoms helps explain many aspects of chemical bonding

4. ENERGY AND CHEMICAL BONDS  Chemical bonds are the forces that holds atoms together.  Energy is required to overcome these attractive forces and separate the atoms in a compound  Breaking a chemical bond is an endothermic process  ENDOTHERMIC: chemical reaction that absorbs energy producing products with more potential energy than the reactants Ex) N 2 + ENERGY  N + N

5. ENERGY AND CHEMICAL BONDS  Formation of a bond is an exothermic process. EXOTHERMIC: chemical reaction that releases energy producing products with less potential energy than the reactants Ex) N + N  N 2 + ENERGY

6. Bonding and stability  When bonds are formed their products are more stable  The compounds have smaller amount of potential energy POTENTIAL ENERGY: stored energy based on position or composition  The bonded elements of a compound are more stable than the individual atom or ions because the atoms have filled their valence electron shell

7. Types of bonds  There are three types of bonds: 1.Ionic 2.Covalent 3.Metallic  They differ in the types of elements involved.  Also how the valence electrons are handled

8. IONS  Atoms become ions so that they can become stable. Atoms become ions by either gaining or losing electron

9. IONS Metals form ions by losing electrons and becoming a positive ion with a smaller atomic radius. Positive ions are called cations. Nonmetals form ions by gaining electrons to become a negative ion with a larger atomic radius. Negative ions are also called anions

10. IONIC BONDING The bond that involves the transfer of one or more electrons from a metal atom to a nonmetal atom to form ions. The positive ion and negative ion they attract each other and create a bond.

11. IONIC BONDING There is a large electronegativity difference (E.D.) between a nonmetal and a metal. The nonmetal rips away the valence electron from the metal atom. Nonmetal becomes a negative ion or anion and the metal becomes a positive ion or cation

12. IONIC BONDING Examples: 1 - EX) when Na and Cl atoms come together  Na loses electrons and becomes Na +1 Cl gains electrons and becomes Cl -1 They attract to form ionic compound NaCl (aka: salts or electrolytes ) Ionic bonds have the highest polarity (most unequal type of bonding) and the most ionic character

13. CLUE FOR RECOGNIZING IONIC BONDS 1. Metal and nonmetal 2. Electronegativity difference in greater than 1.7 (approx)

14. Properties of Ionic Bonds 1. High melting and boiling point. 2. Good electrical conductor as a liquid or when dissolved in water 3. Not good electrical conductor as solid 4. Hard substances

15. DRAWING LEWIS DOT STRUCTURES FOR IONIC BONDS   Write the metal symbol with no dots in brackets   Place the charge at the top right of the bracket   Write the nonmetal symbol with 8 dots around it (except H!)   Draw brackets around the symbol and place the charge of the ion at the top right of the bracket

16. DRAWING LEWIS DOT STRUCTURES FOR IONIC BONDS   Example: Draw the Lew dot structure of the following elements Na and F

17. QUESTIONS:   What is an ionic bond?   What are the clues for recognizing an ionic bond?   What happened to electrons in an ionic bond?   We will be using dot structures in order to draw ionic bonds. Know the information from class about the roles of metals and nonmetals in forming ionic bonds; predict what the dot structure of a metal will look like.   Predict what the dot structure of a nonmetal will look like

18. PRACTICE

20. Bonding COVALENT/METALLIC

21. ionic Covalent molecular metallic Bonding m,nmall nm all m +,- transfershare M-SOME Polar covalent Non polar covalent

22. AIM - AIM - What is it about the structure of noble gases that leads to their stability?   Noble gases (group 18) are stable and undergo few chemical reactions – lack reactivity   ( Argon and Xenon combine with fluorine – rare)   Why? – All have 8 valence electrons except He with 2   Octet – configuration of 8 valence electrons represents max # of valence electrons an atom can have (except H and He – max 2)   Octet Rule – states atoms generally react by gaining, losing, or sharing electrons in order to achieve a complete octet of 8 valence electrons – noble gas configuration

23. Aim- How can we understand covalent (molecular) bonds?   Covalent bond – formed when two nuclei share electrons to achieve a stable arrangement of electrons – between all nonmetals   Diatomic – covalent bond formed between two nonmetals of the same element: Br 2, I 2, N 2, Cl 2 H 2, O 2, F 2,

24. PROPERTIES   Exist in gas, liquid, or solid state   Good insulators   Poor conductors   Low melting points   Many are soft substances

25. CONSTRUCTING LEWIS DOT STRUCTURE FOR SINGLE BINARY COVALENT MOLECULAR COMPOUNDS 1. 1. Determine valence electrons in total (add them up for each element in the compound 2. 2. Divide by 2 to determine the number of pairs of electrons in total for the compound 3. 3. Place first pair between the two elements (use a dash – to represent the shared pair) 4. 4. Place remaining pairs around each elements making sure not to violate the octet rule (Remember H can have a max of 2 electrons)

26. H2H2 Cl 2

27. HClHBr

28.  CONSTRUCTING LEWIS DOT STRUCTURES FOR SINGLE TERNANRY COVALENT MOLECULAR COMPOUNDS (more than two elements involved) 1. 1. Determine valence electrons in total (add them up for each element in the compound 2. 2. Divide by 2 to determine the number of pairs of electrons in total for the compound 3. 3. Determine the most electronegative element and place it in the middle 4. 4. Place the other elements around it 5. 5. Start placing pairs (as dash lines) between the central atom and the terminal atoms 6. 6. Place remaining around each elements making sure not to violate the octet rule (Remember H can have a max of 2 electrons)

29. NH 3 CH 4 CCl 4

30. CONSTRUCTING LEWIS DOT STRUCTURES FOR MULTIPLE COVALENT MOLECULAR COMPOUNDS (more than two elements involved) 1. 1. Determine valence electrons in total (add them up for each element in the compound 2. 2. Divide by 2 to determine the number of pairs of electrons in total for the compound 3. 3. Determine the most electronegative element and place it in the middle 4. 4. Place the other elements around it 5. 5. Start placing pairs (as dash lines) between the central atom and the terminal atoms 6. 6. Place remaining around each elements making sure not to violate the octet rule (Remember H can have a max of 2 electrons) 7. 7. If octet rule is not yet reached you can make additional pairs of electrons into double or triple bonds until octet rule is obeyed by all elements *can only be done with CNOPS

31. CO 2 O2O2 N2N2

32. Coordinate Covalent Bond

33. Sigma (  ) Bonds  Sigma bonds are characterized by  Head-to-head overlap.  Cylindrical symmetry of electron density about the internuclear axis.

34. Pi (  ) Bonds  Pi bonds are characterized by  Side-to-side overlap.  Electron density above and below the internuclear axis. © 2009, Prentice-Hall, Inc.

35. Sigma and Pi Bonds   Single bonds are always  bonds, because  overlap is greater, resulting in a stronger bond and more energy   Double bonds – contain one sigma and one pi bond   Triple bonds – contain one sigma and two pi bonds   Triple bonds are the shortest and strongest type of bond then double and single bonds are the longest and weakest type of bond

36. How can we use the 1.7 rule to predict bond polarity?  1.7 rule is applied primarily to binary compounds.  Determine the electronegativities of all atoms in the bond.  Take the difference between the bonded atoms  If the difference is: >1.7 then ionic < 1.7 but not “0”, then polar covalent < 1.7 but not “0”, then polar covalent =0 non polar covalent

38. Nonpolar covalent bond Polar covalent bond Ionic bond Covalent bonds

39. MOLECULE POLARITY: Nonpolar/polar shapes SNAP PAD  SNAP : S ymmetrical N onpolar A symmetrical P olar  PAD: P olar A symmetrical D ipole  OPEN: O dd P olar E ven N onpolar

40. DIPOLE MOMENT dipole moment --separation of the charge in a molecule; product of the size of the charge and the distance of separation align themselves with an electric field (figure b at right) align with each other as well in the absence of an electric field water—2 lone pairs establish a strong negative pole ammonia—has a lone pair which establishes a neg. pole note that the direction of the “arrow” indicating the dipole moment always points to the negative pole with the cross hatch on the arrow (looks sort of like we’re trying to make a + sign) is at the positive pole.

41. CompoundPolar or Nonpolar NH 3 H20H20

42. How can a molecule be both polar and non polar?  CX 4 tetrahedral shape. They are non polar.  The individual ligands or bonds are polar  Conclusion – nonpolar/polar.

43. What Determines the Shape of a Molecule?  Simply put, electron pairs, whether they be bonding or nonbonding, repel each other.  By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule. © 2009, Prentice-Hall, Inc.

44. Valence Shell Electron Pair Repulsion Theory (VSEPR) “The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.” © 2009, Prentice-Hall, Inc.

45. THE VSPER MODEL AND MOLECULAR SHAPE   molecular geometry --the arrangement in space of the atoms bonded to a central atom not necessarily the same as the structural pair geometry lone pairs have a different repulsion since they are experiencing an attraction or “pull” from only one nucleus as opposed to two nuclei. They also take up more space around an atom as you can see on the left.   Each lone pair or bond pair repels all other lone pairs and bond pairs--they try to avoid each other making as wide an angle as possible.   - works well for elements of the s and p-blocks   - VSEPR does not apply to transition element compounds (exceptions)

46. MOLECULAR SHAPES -Linear -Bent (angular) -Pyramidal -Tetrahedral

47. Intermolecular Forces The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together. © 2009, Prentice-Hall, Inc.

48. Intermolecular Forces They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities. © 2009, Prentice-Hall, Inc.

49. Intermolecular Forces These intermolecular forces as a group are referred to as van der Waals forces. © 2009, Prentice-Hall, Inc.

50. Ion-Dipole Interactions  Ion-dipole interactions (a fourth type of force), are important in solutions of ions.  The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents. © 2009, Prentice-Hall, Inc.

51. van der Waals Forces  Dipole-dipole interactions  Hydrogen bonding  London dispersion forces © 2009, Prentice-Hall, Inc.

52. Dipole-Dipole Interactions  Molecules that have permanent dipoles are attracted to each other.  The positive end of one is attracted to the negative end of the other and vice-versa.  These forces are only important when the molecules are close to each other. © 2009, Prentice-Hall, Inc.

53. Dipole-Dipole Interactions The more polar the molecule, the higher is its boiling point. © 2009, Prentice-Hall, Inc.

54. London Dispersion Forces While the electrons in the 1s orbital of helium would repel each other (and, therefore, tend to stay far away from each other), it does happen that they occasionally wind up on the same side of the atom. © 2009, Prentice-Hall, Inc.

55. London Dispersion Forces At that instant, then, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side. © 2009, Prentice-Hall, Inc.

56. London Dispersion Forces Another helium nearby, then, would have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1. © 2009, Prentice-Hall, Inc.

57. London Dispersion Forces London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole. © 2009, Prentice-Hall, Inc.

58. London Dispersion Forces  These forces are present in all molecules, whether they are polar or nonpolar.  The tendency of an electron cloud to distort in this way is called polarizability. © 2009, Prentice-Hall, Inc.

59. Factors Affecting London Forces  The shape of the molecule affects the strength of dispersion forces: long, skinny molecules (like n-pentane tend to have stronger dispersion forces than short, fat ones (like neopentane).  This is due to the increased surface area in n-pentane. © 2009, Prentice-Hall, Inc.

60. Factors Affecting London Forces  The strength of dispersion forces tends to increase with increased molecular weight.  Larger atoms have larger electron clouds which are easier to polarize. © 2009, Prentice-Hall, Inc.

61. Which Will Have a Greater Effect? Dipole-Dipole Interactions or Dispersion Forces  If two molecules are of comparable size and shape, dipole-dipole interactions will likely the dominating force.  If one molecule is much larger than another, dispersion forces will likely determine its physical properties. © 2009, Prentice-Hall, Inc.

62. How Do We Explain This?  The nonpolar series (SnH 4 to CH 4 ) follow the expected trend.  The polar series follows the trend from H 2 Te through H 2 S, but water is quite an anomaly. © 2009, Prentice-Hall, Inc.

63. Hydrogen Bonding  The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong.  We call these interactions hydrogen bonds. © 2009, Prentice-Hall, Inc.

64. Hydrogen Bonding  Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine. © 2009, Prentice-Hall, Inc. Also, when hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed.

65. Summarizing Intermolecular Forces © 2009, Prentice-Hall, Inc.

66. Intermolecular Forces Affect Many Physical Properties The strength of the attractions between particles can greatly affect the properties of a substance or solution. © 2009, Prentice-Hall, Inc.

67. Single Bonds Single bonds are always  bonds, because  overlap is greater, resulting in a stronger bond and more energy © 2009, Prentice-Hall, Inc.

68. How can we understand and recognize metallic bonding?  Contains positively charged metals  Metals are arranged in a crystalline lattice structure immersed in a: sea of mobile electrons  Electrons are delocalized. This means no one atom owns any electrons they belong to the whole crystal. M-SOME Metals are good conductors because of mobile ions

70. Sea of Mobile Electrons

71. ionic Covalent molecular metallic Bonding m,nmall nmall m +,- Polar Covalent Non Polar Covalent