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CH 8: Bonding

CH 8: Bonding

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CH 8: Bonding

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  1. CH 8: Bonding General Concepts

  2. Chapter Outline – Part I • Types of chemical bonds (8.1) • Electronegativity and bond polarity (8.2/3 • Ions (8.4) • Energy changes when a binary ionic compound forms (8.5) • Ionic character of covalent bonds (8.6)

  3. Introduction to Bonding • Chemical bond – force that holds atoms together so that they function as a unit. • Consider 2 classes of bonds: • Ionic bonding • Covalent bonding

  4. Bond Types • Ionic bonds – attractive forces among oppositely charged ions • Forms when a metal loses electron(s) to a nonmetal. • Bond strength can be calculated using Coulomb’s law

  5. Ionic Bonds • Strength of the attraction between the ions can be calculated using Coulomb’s law. E = (2.31 x 10-19 J nm) (Q1Q2/r) • Q1 and Q2 are the charges on the ions. • r = distance between ion centers in nm

  6. Using Coulomb’s Law E = (2.31 x 10-19 J nm) (Q1Q2/r) • Sign on E??? • The more negative E, the stronger the attractive force between the ions.

  7. Using Coulomb’s Law E = (2.31 x 10-19 J nm) (Q1Q2/r) • Magnitude of E. • E is more negative when:

  8. Covalent Bonds • Covalent bond – bonded atoms share pairs of valence electrons • Covalent bonding results in formation of a molecule. • Covalent bonding occurs between nonmetals.

  9. Types of Covalent Bonds • “Pure” covalent bond – electrons are shared by like nonmetals • E.g. diatomic molecules • Results in equal sharing of the electrons • Aka – nonpolar covalent bond

  10. Types of Covalent Bonds • Polar covalent bond – unequal sharing of electrons by the bonded atoms • bond between different nonmetals each with its own ability to attract the shared electrons

  11. Polar Covalent Bonds • Showing bond polarity: • Consider the HF molecule. • See board and/or page 290. • Experimental determination of bond polarity, page 289

  12. Bond Polarity • To predict bond polarity…consider the electronegativity (EN) of the bonded atoms. • EN – the ability of an atom in a molecule to attract shared electrons.

  13. EN Values • The higher the EN the greater the atom’s ability to attract shared electrons. • EN values and the periodic table • EN ________ down a group. • EN ________ across a period. • See “back” of the periodic table.

  14. EN and Bond Polarity • As the difference in EN between bonded atoms increases so does the polarity of the bond. • Can also say that the ionic character of the bond is increasing. • See table 8.1 on page 289.

  15. Bond Polarity and Dipoles • Polar molecules have a preferred orientations when placed in an electric field. • Said to have a dipole moment. • Dipole moment – molecule has a center of positive charge and a center of negative charge

  16. Bond Polarity and Dipoles • Not all molecule with polar bonds have dipole moments! • Bond polarities cancel each other in molecules with symmetrical dipoles. • Molecules with equal, opposing dipoles. • See page 291and 8.2 on page 292 • Dog walking example!

  17. Compound Formation • Atoms gain, lose, or share enough electrons to achieve the same stable electron configuration as a noble gas • Nonmetals share electrons • Form molecules with covalent bonds • Representative metals lose electrons to nonmetals in ionic compounds • Ions are isoelectronic to noble gases

  18. 8.4 Ion Formation • Binary ionic compounds • The metal loses electron(s) to a nonmetal • Focus on represenative metals • The atoms lose/gain enough electrons to obtain a noble gas electron configuration.

  19. Cations • Group IA metals form ions with a _____ charge. • Na atom • Na ion • Isoelectronic to: ____________________

  20. Anions • Group VIA elements form ions with a ______ charge. • Sulfur atom • Sulfur ion (called _________________) • Isoelectronic to: ____________________

  21. Ionic Compound • Consider the compound formed between sodium and sulfur. • Each sodium atom loses 1 electron. • Each sulfur atom needs 2 electrons. • Formula for compound:

  22. Ion Size • Cations are smaller than their parent atom. • Atoms lose their valence shell when the ion forms. • Na 1s22s22p63s1 ____ protons • Na+ 1s22s22p6 ____ protons

  23. Ion Size • Anions are larger than their parent atom. • Atoms add electrons to their valence shell when the ion forms – proton # remains the same. • F 1s22s22p5 ____ protons • F1- 1s22s22p6 ____ protons

  24. Ion Size • Isoelectronic ions decrease in size as the number of protons increases. • Example: ions with 10 electrons 10 e O2- F1- Na1+ Mg2+ Al3+ # p 8 9 11 12 13

  25. Isoelectronic Ions 10 e O2- F1- Na1+ Mg2+ Al3+ # p 8 9 11 12 13 Radius* 140 136 95 65 50 • * picometers • The diagram on page 296 should make sense.

  26. 8.5 Energy in Binary Ionic Compounds • Lattice energy – change in energy when separated gaseous ions form an ionic solid. M+(g) + X-(g) MX(s) LE < 0

  27. Lattice Energy • LE = k (Q1Q2)/r • K is the proportionality constant • Q1 and Q2 are the charges on the ions • r is the ionic radius

  28. Lattice Energy • LE becomes more exothermic as the ion charges increase and the ion radius decreases. • Small highly charged ions have more exothermic LE • See board for examples.

  29. Formation of ionic compounds. • Consider energy changes associated with formation of a binary ionic compound. • 5 step process, page 297/298 • Most common series of steps is shown on the next slide.

  30. Formation of ionic compounds. • Sublime the metal. • Ionize the gaseous metal atoms. • ionization energy(ies) • Dissociate the nonmetal (if diatomic). • Bond energy • Ionize the gaseous nonmetal atoms. • Electron affinity • Form the solid from the gaseous ions • LE

  31. Born Haber Practice… • NaF • MgF2 • IE – page 272; EA – page 275 • Bond energies – page 306 • Sublimation: Na 109 kJ; Mg 147 kJ • Lattice energy: NaF -923kJ; MgF2 2913

  32. 8.6 Partial Ionic Character • When atoms with different EN bond the result is either a polar covalent or an ionic bond. • There’s evidence that some level of electron sharing occurs in all bonds. • Even in what we consider as ionic bonds.

  33. 8.6 Partial Ionic Character • Classify a bond as ionic if it conducts electricity when melted. • Essentially all compounds with metals meet this criteria. • These compounds generally have more than 50% ionic character.

  34. 8.8 Covalent Bond Energies • Strength of a given bond depends upon the compound. • Not all C-H bonds are of the same energy! • See page 305. • Bond energies given in tables are averages based on experimental data.

  35. Bond Energies • Consider the bond energies on page 306. • Compare the bond energies and bond length associated with single, double, and triple bonds between a given pair of atoms.

  36. 8.8 Using Bond Energies • The DH for a reaction can be estimated from bond energies. DH = energy needed to break bonds of reactants – energy released when product bonds form

  37. 8.8 Using Bond Energies • Estimate the DH for.

  38. CH 8 Part II: Bonding Models • Introduction to models (8.7) • Localized Electron (LE) Bonding Model (8.9) • Lewis Structure (8.10) • Resonance (8.12) • Exceptions to the Octet Rule(8.11) • VSEPR Theory (8.13) • Key pages: 326/27

  39. 8.7 Models • Read “Fundamental Properties of Models” on page 350.

  40. 8.9 LE Bonding Model • Localized electron bonding model • Assumes a molecule is made of atoms bound together by sharing pairs of electrons using the orbitals of the bonding atoms.

  41. 8.9 LE Bonding Model • Localized electron bonding model • Shared electrons are pictures to be localized in the space between the atoms • Called bonding pairs • Non-bonding valence electrons are pictured to be localized on the parent atom. • Called lone pairs • Consider HCl

  42. 8.10 Lewis Structures • Lewis structures show the arrangement of the valence electrons in molecules (and ions). • Representative atoms will have the same number of valence electrons as one of the noble gases • 2 electrons to be like H • 8 electrons to be like all other noble gases

  43. Lewis Structures • Lewis structures illustrate LE bonding model. • Show the bonding electrons and the lone pairs. • Lewis structures can be used to predict the 3D geometry of a molecule. • Requires application of VSEPR Theory • More to come on this…..

  44. 1st Goal: To Write Lewis Structures • Sum the valence electrons. • Use a pair of electrons to form a bond between each of the bonded atoms. • Put the atom that needs the most electrons in the center when the molecule contains more than 2 atoms. • Arrange the remaining electrons to satisfy the duet rule for H and the octet rule for elements in the 2nd row of elements.

  45. Writing Lewis Structures • Practice! • H2O • O2 • HCN • NO31- • PH3

  46. 8.12 Resonance • More than one valid Lewis structure can often be drawn for molecules with multiple bonds (double, triple..) • Consider NO21- • 2 valid Lewis structures can be drawn.

  47. Resonance Structures • Lewis structure just drawn indicate 2 types of bonds in NO21- -- single bond and a double bond • However….the data shows that both bonds in NO21- are of the same energy and bond length • Both bonds are stronger and shorter than a single bond, but not as strong or short as a double bond!

  48. Exceptions to the Octet Rule • Less than an octet. • Be and B • More than an octet • 3rd period elements and “up” • Odd number of electrons

  49. Exceptions to the Octet Rule • Less than an octet. • Be - satisfied/stable with 4 electrons • B - satisfied/stable with 6 electrons

  50. Exceptions to the Octet Rule 2. More than an octet • Atoms in the 3rd period and “up*” can use their unfilled d orbitals to accommodate more than 8 electrons • Commonly see 10 electrons and 12 electrons around the central atom. • “Up” refers to periods 4, 5,6,…