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Chapter 9 Chemical Bonding I: Lewis Theory

Chapter 9 Chemical Bonding I: Lewis Theory

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Chapter 9 Chemical Bonding I: Lewis Theory

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  1. Chemistry: A Molecular Approach, 1st Ed.Nivaldo Tro Chapter 9Chemical Bonding I:Lewis Theory Roy Kennedy Massachusetts Bay Community College Wellesley Hills, MA 2008, Prentice Hall

  2. Bonding Theories • explain how and why atoms attach together • explain why some combinations of atoms are stable and others are not • why is water H2O, not HO or H3O • one of the simplest bonding theories was developed by G.N. Lewis and is called Lewis Theory • Lewis Theory emphasizes valence electrons to explain bonding • using Lewis Theory, we can draw models – called Lewis structures – that allow us to predict many properties of molecules • aka Electron Dot Structures • such as molecular shape, size, polarity Tro, Chemistry: A Molecular Approach

  3. Types of Bonds Tro, Chemistry: A Molecular Approach

  4. Types of Bonding

  5. Ionic Bonds • when metals bond to nonmetals, some electrons from the metal atoms are transferred to the nonmetal atoms • metals have low ionization energy, relatively easy to remove an electron from • nonmetals have high electron affinities, relatively good to add electrons to Tro, Chemistry: A Molecular Approach

  6. Covalent Bonds • nonmetals have relatively high ionization energies, so it is difficult to remove electrons from them • when nonmetals bond together, it is better for the atoms to share valence electrons Tro, Chemistry: A Molecular Approach

  7. Lewis Symbols of Atoms • aka electron dot symbols • use symbol of element to represent nucleus and inner electrons • use dots around the symbol to represent valence electrons • pair first two electrons for the s orbital • put one electron on each open side for p electrons • then pair rest of the p electrons Tro, Chemistry: A Molecular Approach

  8. Stable Electron ArrangementsAnd Ion Charge • Metals form cations by losing enough electrons to get the same electron configuration as the previous noble gas • Nonmetals form anions by gaining enough electrons to get the same electron configuration as the next noble gas • The noble gas electron configuration must be very stable Tro, Chemistry: A Molecular Approach

  9. Lewis Theory • the basis of Lewis Theory is that there are certain electron arrangements in the atom that are more stable • octet rule • bonding occurs so atoms attain a more stable electron configuration Tro, Chemistry: A Molecular Approach

  10. Melting an Ionic Solid Properties of Ionic Compounds • hard and brittle crystalline solids • all are solids at room temperature • melting points generally > 300C • the liquid state conducts electricity • the solid state does not conduct electricity • many are soluble in water • the solution conducts electricity well Tro, Chemistry: A Molecular Approach

  11. Li + Lewis Theory and Ionic Bonding • Lewis symbols can be used to represent the transfer of electrons from metal atom to nonmetal atom, resulting in ions that are attracted to each other and therefore bond + Tro, Chemistry: A Molecular Approach

  12. 2 Li + Predicting Ionic FormulasUsing Lewis Symbols • electrons are transferred until the metal loses all its valence electrons and the nonmetal has an octet • numbers of atoms are adjusted so the electron transfer comes out even Li2O Tro, Chemistry: A Molecular Approach

  13. Ionic BondingModel vs. Reality • ionic compounds have high melting points and boiling points • MP generally > 300°C • all ionic compounds are solids at room temperature • because the attractions between ions are strong, breaking down the crystal requires a lot of energy • the stronger the attraction (larger the lattice energy), the higher the melting point Tro, Chemistry: A Molecular Approach

  14. - - - - - - - - + + + + + + + + + + - - - - - - - - - - + + + + + + + + - - - - + + + + + - - - - - + + + + Ionic BondingModel vs. Reality • ionic solids are brittle and hard • the position of the ion in the crystal is critical to establishing maximum attractive forces – displacing the ions from their positions results in like charges close to each other and the repulsive forces take over Tro, Chemistry: A Molecular Approach

  15. Ionic BondingModel vs. Reality • ionic compounds conduct electricity in the liquid state or when dissolved in water, but not in the solid state • to conduct electricity, a material must have charged particles that are able to flow through the material • in the ionic solid, the charged particles are locked in position and cannot move around to conduct • in the liquid state, or when dissolved in water, the ions have the ability to move through the structure and therefore conduct electricity Tro, Chemistry: A Molecular Approach

  16. F F •• •• •• • • •• F F •• • • • • H H O •• •• •• •• •• •• •• •• •• F F •• H H •• O •• •• •• Single Covalent Bonds • two atoms share a pair of electrons • 2 electrons • one atom may have more than one single bond Tro, Chemistry: A Molecular Approach

  17. •• •• • • • • O O •• •• •• •• O •• •• •• •• O Double Covalent Bond • two atoms sharing two pairs of electrons • 4 electrons Tro, Chemistry: A Molecular Approach

  18. •• •• • • • • N N • • N N •• •• •• •• •• Triple Covalent Bond • two atoms sharing 3 pairs of electrons • 6 electrons Tro, Chemistry: A Molecular Approach

  19. Covalent BondingModel vs. Reality • molecular compounds have low melting points and boiling points • MP generally < 300°C • molecular compounds are found in all 3 states at room temperature • melting and boiling involve breaking the attractions between the molecules, but not the bonds between the atoms • the covalent bonds are strong • the attractions between the molecules are generally weak • the polarity of the covalent bonds influences the strength of the intermolecular attractions Tro, Chemistry: A Molecular Approach

  20. Molecular BondingModel vs. Reality • some molecular solids are brittle and hard, but many are soft and waxy • the kind and strength of the intermolecular attractions varies based on many factors • the covalent bonds are not broken, however, the polarity of the bonds has influence on these attractive forces Tro, Chemistry: A Molecular Approach

  21. Molecular BondingModel vs. Reality • molecular compounds do not conduct electricity in the liquid state • molecular acids conduct electricity when dissolved in water, but not in the solid state • in molecular solids, there are no charged particles around to allow the material to conduct • when dissolved in water, molecular acids are ionized, and have the ability to move through the structure and therefore conduct electricity Tro, Chemistry: A Molecular Approach

  22. Bond Polarity • covalent bonding between unlike atoms results in unequal sharing of the electrons • one atom pulls the electrons in the bond closer to its side • one end of the bond has larger electron density than the other • the result is a polar covalent bond • bond polarity • the end with the larger electron density gets a partial negative charge • the end that is electron deficient gets a partial positive charge Tro, Chemistry: A Molecular Approach

  23. Electronegativity • measure of the pull an atom has on bonding electrons • increases across period (left to right) and • decreases down group (top to bottom) • fluorine is the most electronegative element • francium is the least electronegative element • the larger the difference in electronegativity, the more polar the bond • negative end toward more electronegative atom Tro, Chemistry: A Molecular Approach

  24. Electronegativity Scale Tro, Chemistry: A Molecular Approach

  25. Percent Ionic Character 4% 51% 0 0.4 2.0 4.0 Electronegativity Difference Electronegativity and Bond Polarity • If difference in electronegativity between bonded atoms is 0, the bond is pure covalent • equal sharing • If difference in electronegativity between bonded atoms is 0.1 to 0.4, the bond is nonpolar covalent • If difference in electronegativity between bonded atoms 0.5 to 1.9, the bond is polar covalent • If difference in electronegativity between bonded atoms larger than or equal to 2.0, the bond is ionic “100%”

  26. ENCl = 3.0 3.0 - 3.0 = 0 Pure Covalent ENCl = 3.0 ENH = 2.1 3.0 – 2.1 = 0.9 Polar Covalent ENCl = 3.0 ENNa = 1.0 3.0 – 0.9 = 2.1 Ionic Bond Polarity Tro, Chemistry: A Molecular Approach

  27. Bond Dipole Moments • the dipole moment is a quantitative way of describing the polarity of a bond • a dipole is a material with positively and negatively charged ends • measured • dipole moment, m, is a measure of bond polarity • it is directly proportional to the size of the partial charges and directly proportional to the distance between them • m = (q)(r) • not Coulomb’s Law • measured in Debyes, D • the percent ionic character is the percentage of a bond’s measured dipole moment to what it would be if full ions Tro, Chemistry: A Molecular Approach

  28. Example 9.3(c) - Determine whether an N-O bond is ionic, covalent, or polar covalent. • Determine the electronegativity of each element N = 3.0; O = 3.5 • Subtract the electronegativities, large minus small (3.5) - (3.0) = 0.5 • If the difference is 2.0 or larger, then the bond is ionic; otherwise it’s covalent difference (0.5) is less than 2.0, therefore covalent • If the difference is 0.5 to 1.9, then the bond is polar covalent; otherwise it’s covalent difference (0.5) is 0.5 to 1.9, therefore polar covalent Tro, Chemistry: A Molecular Approach

  29. Lewis Structures of Molecules • shows pattern of valence electron distribution in the molecule • useful for understanding the bonding in many compounds • allows us to predict shapes of molecules • allows us to predict properties of molecules and how they will interact together Tro, Chemistry: A Molecular Approach

  30. C B N O F Lewis Structures • use common bonding patterns • C = 4 bonds & 0 lone pairs, N = 3 bonds & 1 lone pair, O= 2 bonds & 2 lone pairs, H and halogen = 1 bond, Be = 2 bonds & 0 lone pairs, B = 3 bonds & 0 lone pairs • often Lewis structures with line bonds have the lone pairs left off • their presence is assumed from common bonding patterns • structures which result in bonding patterns different from common have formal charges Tro, Chemistry: A Molecular Approach

  31. Writing Lewis Structures of Molecules HNO3 • Write skeletal structure • H always terminal • in oxyacid, H outside attached to O’s • make least electronegative atom central • N is central • Count valence electrons • sum the valence electrons for each atom • add 1 electron for each − charge • subtract 1 electron for each + charge N = 5 H = 1 O3 = 3∙6 = 18 Total = 24 e- Tro, Chemistry: A Molecular Approach

  32. Writing Lewis Structures of Molecules HNO3 • Attach central atom to the surrounding atoms with pairs of electrons and subtract from the total Electrons Start 24 Used 8 Left 16 Tro, Chemistry: A Molecular Approach

  33. Writing Lewis Structures of Molecules HNO3 • Complete octets, outside-in • H is already complete with 2 • 1 bond and re-count electrons N = 5 H = 1 O3 = 3∙6 = 18 Total = 24 e- Electrons Start 24 Used 8 Left 16 Electrons Start 16 Used 16 Left 0 Tro, Chemistry: A Molecular Approach

  34. Writing Lewis Structures of Molecules HNO3 • If all octets complete, give extra electrons to central atom. • elements with d orbitals can have more than 8 electrons • Period 3 and below • If central atom does not have octet, bring in electrons from outside atoms to share • follow common bonding patterns if possible Tro, Chemistry: A Molecular Approach

  35. CO2 SeOF2 NO2-1 H3PO4 SO3-2 P2H4 Practice - Lewis Structures Tro, Chemistry: A Molecular Approach

  36. CO2 SeOF2 NO2-1 H3PO4 SO3-2 P2H4 : : :O::C::O: Practice - Lewis Structures 16 e- 32 e- 26 e- 26 e- 18 e- 14 e- Tro, Chemistry: A Molecular Approach

  37. O S O •• •• •• • • • • • • • • •• •• Formal Charge • during bonding, atoms may wind up with more or less electrons in order to fulfill octets - this results in atoms having a formal charge FC = valence e- - nonbonding e- - ½ bonding e- left O FC = 6 - 4 - ½ (4) = 0 S FC = 6 - 2 - ½ (6) = +1 right O FC = 6 - 6 - ½ (2) = -1 • sum of all the formal charges in a molecule = 0 • in an ion, total equals the charge Tro, Chemistry: A Molecular Approach

  38. Metallic Bonds • low ionization energy of metals allows them to lose electrons easily • the simplest theory of metallic bonding involves the metals atoms releasing their valence electrons to be shared by all to atoms/ions in the metal • bonding results from attraction of cation for the delocalized electrons Tro, Chemistry: A Molecular Approach

  39. Metallic Bonding Tro, Chemistry: A Molecular Approach

  40. Metallic BondingModel vs. Reality • metallic solids conduct electricity • because the free electrons are mobile, it allows the electrons to move through the metallic crystal and conduct electricity • as temperature increases, electrical conductivity decreases • heating causes the metal ions to vibrate faster, making it harder for electrons to make their way through the crystal Tro, Chemistry: A Molecular Approach

  41. Metallic BondingModel vs. Reality • metallic solids conduct heat • the movement of the small, light electrons through the solid can transfer kinetic energy quicker than larger particles • metallic solids reflect light • the mobile electrons on the surface absorb the outside light and then emit it at the same frequency Tro, Chemistry: A Molecular Approach

  42. Metallic BondingModel vs. Reality • metallic solids are malleable and ductile • because the free electrons are mobile, the direction of the attractive force between the metal cation and free electrons is adjustable • this allows the position of the metal cation islands to move around in the sea of electrons without breaking the attractions and the crystal structure Tro, Chemistry: A Molecular Approach

  43. Metallic BondingModel vs. Reality • metals generally have high melting points and boiling points • all but Hg are solids at room temperature • the attractions of the metal cations for the free electrons is strong and hard to overcome • melting points generally increase to right across period • the charge on the metal cation increases across the period, causing stronger attractions • melting points generally decrease down column • the cations get larger down the column, resulting in a larger distance from the nucleus to the free electrons Tro, Chemistry: A Molecular Approach