2. Unit 3: Chemical Bonding and Molecular Structure Cartoon courtesy of NearingZero.net

3. Bonds Forces that hold groups of atoms  Forces that hold groups of atoms together and make them function together and make them function as a unit – bonds form because it LOWERS the potential energy as a unit – bonds form because it LOWERS the potential energy Ionic bonds – transfer of electrons  Ionic bonds – transfer of electrons  Covalent bonds – sharing of electrons

4. Properties of Ionic Compounds Structure:Crystalline solids Melting point:Generally high Boiling Point:Generally high Electrical Conductivity: Excellent conductors, molten and aqueous Solubility in water: Electrons: Elements: Generally soluble transferred Metal + nonmetal

5. Properties of Covalent Compounds Electrons:shared Melting point:Generally low Boiling Point:Generally low Electrical Conductivity: Poor conductors, except aqueous acids Solubility in water: Elements: Two Types: Polar covalent are soluble Nonmetal + nonmetal Polar and nonpolar

6. Electronegativity The ability of an atom in a molecule to attract shared electrons to itself. The ability of an atom in a molecule to attract shared electrons to itself. Linus Pauling 1901 - 1994

7. Table of Electronegativities

8. Polar-Covalent bonds Nonpolar-Covalent bonds Covalent Bonds  Electrons are unequally shared  Electronegativity difference between.3 and 1.7  Electrons are equally shared  Electronegativity difference of 0 to 0.3

9. Determining Bond Type: BondEN DiffType?More EN? S to H2.5-2.1 Ca to Cl1.0-3.0 Mg to O1.2-3.5 F to F4.0-4.0 Na to Cl0.9-3.0 C to F2.5-4.0 Al to Cl1.5-3.0 B to H2.0-2.1 Si to O1.8-3.5

10. The Octet Rule – Ionic Compounds Ionic compounds tend to form so that each atom, by gaining or losing electrons, has an octet of electrons in its highest occupied energy level.

11. Ionic Bonding: The Formation of Sodium Chloride  Sodium has 1 valence electron Cl 1s 2 2s 2 2p 6 3s 2 3p 5 Na 1s 2 2s 2 2p 6 3s 1  Chlorine has 7 valence electrons  An electron transferred gives each an octet

12. Ionic Bonding: The Formation of Sodium Chloride Cl - 1s 2 2s 2 2p 6 3s 2 3p 6 Na + 1s 2 2s 2 2p 6 This transfer forms ions, each with an octet:

13. Ionic Bonding: The Formation of Sodium Chloride Cl - Na + The resulting ions come together due to electrostatic attraction (opposites attract): The net charge on the compound must equal zero

14. Examples of Ionic compounds All salts, which are composed of metals bonded to nonmetals, are ionic compounds and form ionic crystals. Examples: BaS MgCl 2 Na 2 O KI CaO LiF

15. Monatomic Cations Name H+H+ Hydrogen Li + Lithium Na + Sodium K+K+ Potassium Mg 2+ Magnesium Ca 2+ Calcium Ba 2+ Barium Al 3+ Aluminum

16. Monatomic Anions Name F-F- Fluoride Cl - Chloride Br - Bromide I-I- Iodide O 2- Oxide S 2- Sulfide N 3- Nitride P 3- Phosphide

17. Sodium Chloride Crystal Lattice Ionic compounds form solids at ordinary temperatures. Ionic compounds organize in a characteristic crystal lattice of alternating positive and negative ions.

18. Representation of Components in an Ionic Solid Lattice: A 3-dimensional system of points designating the centers of components (atoms, ions, or molecules) that make up the substance.

19. Ionic Bonds  Electrons are transferred  Electronegativity differences are generally greater than 1.7  The formation of ionic bonds is always exothermic!

20. Metallic Bonding  The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons  Vacant p and d orbitals in metal's outer energy levels overlap, and allow outer electrons to move freely  Electrons are “delocalized”

21. Properties of Metals  Metals are good conductors of heat and electricity  Metals are malleable  Metals are ductile  Metals have high tensile strength  Metals have luster

22. Bonding Forces – covalent compounds  Electron – electron repulsive forces repulsive forces  Proton – proton repulsive forces  Electron – proton attractive forces

23. Bond Length Diagram

24. Bond Energy It is the energy required to break a bond. (covalent) It gives us information about the strength of a bonding interaction. For ionic bonds, we measure Lattice Energy – the amount of energy it takes to form the bond

25. The Octet Rule – Covalent Compounds Covalent compounds tend to form so that each atom, by sharing electrons, has an octet of electrons in its highest occupied energy level. Diatomic Fluorine

26. Hydrogen Chloride by the Octet Rule

27. Formation of Water by the Octet Rule

28. Comments About the Octet Rule 2nd row elements C, N, O, F observe the octet rule. 2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive. 3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals. When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

29. Electron Dot Notation

30. Shows how valence electrons are arranged among atoms in a molecule. Reflects central idea that stability of a compound relates to noble gas electron configuration. Lewis Structures

31. C H H H Cl.. Completing a Lewis Structure - CH 3 Cl Add up available valence electrons: C = 4, H = (3)(1), Cl = 7 Total = 14 Join peripheral atoms to the central atom with electron pairs. Complete octets on atoms other than hydrogen with remaining electrons Make carbon the central atom..

32. Multiple Covalent Bonds: Double bonds Two pairs of shared electrons

33. Multiple Covalent Bonds: Triple bonds Three pairs of shared electrons

34. Bond Length and Bond Energy BondLength (pm)Energy (kJ/mol) C - C154346 C=C134612 CCCC120835 C - N147305 C=N132615 CNCN116887 C - O143358 C=O120799 COCO1131072 N - N145180 N=N125418 NNNN110942

35. Resonance Occurs when more than one valid Lewis structure can be written for a particular molecule. These are resonance structures. The actual structure is an average of the resonance structures.

36. Resonance in Ozone Neither structure is correct.

37. VSEPR Model The structure around a given atom is determined principally by minimizing electron pair repulsions. (Valence Shell Electron Pair Repulsion)

38. Predicting a VSEPR Structure Draw Lewis structure. Put pairs as far apart as possible. Put pairs as far apart as possible. Determine positions of atoms from the way electron pairs are shared. Determine positions of atoms from the way electron pairs are shared. Determine the name of molecular structure from positions of the atoms. Determine the name of molecular structure from positions of the atoms.

39. Table – VSEPR Structures

40. Polarity A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment. A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment.

41. Hydrogen Bonding Hydrogen bonding in Kevlar, a strong polymer used in bullet-proof vests. Bonding between hydrogen and more electronegative neighboring atoms such as oxygen and nitrogen

42. Hydrogen Bonding in Water

43. Hydrogen Bonding between Ammonia and Water

44. Dipole-Dipole Attractions Attraction between oppositely charged regions of neighboring molecules.

45. The water dipole

46. Relative magnitudes of forces The types of bonding forces vary in their strength as measured by average bond energy. Covalent bonds (400 kcal) Hydrogen bonding (12-16 kcal ) Dipole-dipole interactions (2-0.5 kcal) London forces (less than 1 kcal) Strongest Weakest