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UNIT 7:

UNIT 7:. CHEMICAL BONDING. BONDS are forces which hold atoms and ions together in compounds form because this gives the “ lowest possible energy for the system ” ( lower P.E . means more stable bond thermodynamically) being broken absorbs (requires) energy

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UNIT 7:

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  1. UNIT 7: CHEMICAL BONDING

  2. BONDS • are forces which hold atoms and ions together in compounds • form because this gives the “lowest possible energy for the system” • (lower P.E. means more stable bond thermodynamically) • being brokenabsorbs (requires) energy • [endo, so DH = (+)] • being madereleases energy • [exo, so DH = (-)]

  3. OCTET RULE • atoms will lose or gain valence electrons, or share electrons so as to have a total of eight valence (outermost energy level) electrons (s2p6) in highest “n” value sublevels Ex: 3s23p6, 5s25p6

  4. BONDING TYPES • IONICbond is the • electrostatic attraction between • (+) cations and (-) anions • following the transfer of electrons • from metal atoms to non-metalatoms • called “ion-ion interactions • COVALENT bond is the • sharing of electron pairs between non-metal atoms • NON-POLARPOLAR • equalsharing of unequal sharing of • bonding electrons bonding electrons distorted cloud symmetrical cloud

  5. BOND TYPEdepends on the • ELECTRONEGATIVITY OF THE ATOMSin the bond: • ELECTRONEGATIVITY • is the attraction an atom has for another atom’s electrons in a bond • (or an atom’s ability to attract bonding electrons to itself, “greediness”) • is a relative value and has no units • values range from 0.7(Cs most active metal) • to 4.0(F most active non-metal) • generally increases leftright across a period, • decreases down a group

  6. Electronegativity • Electronegativity is the ability of atoms in a molecule to attract electrons to themselves. • On the periodic chart, electronegativity increases as you go… • …from left to right across a row. • …from the bottom to the top of a column.

  7. ELECTRONEGATIVITY “difference” indicates BOND TYPE EN DIFF 0------0.4-------------------------------------2.0-------------3.3 TYPE: non-polar -----> polar covalent -------> ionic covalent 0.5 1.9 Actually more a continuum than clear-cut boundaries: The greater the EN DIFF, the more the “ionic character” of the bond The lower the EN DIFF, the more the “non-polar character” of the bond

  8. Polar Covalent Bonds • When two atoms share electrons unequally, a bond dipole results. • The dipole moment, , produced by two equal but opposite charges separated by a distance, r, is calculated:  = Qr • It is measured in debyes (D).

  9. Polar Covalent Bonds • \ The greater the difference in electronegativity, the more polar is the bond.

  10. d+ d- d- Dipole moment is 1.90 Debye Polar molecule Dipole moment is zero Non-polar molecule

  11. Let’s go to handout

  12. Energetics of Ionic Bonding As we saw in the last chapter, it takes 496 kJ/mol to remove electrons from sodium.

  13. Energetics of Ionic Bonding We get 349 kJ/mol back by giving electrons to chlorine.

  14. Energetics of Ionic Bonding But these numbers don’t explain why the reaction of sodium metal and chlorine gas to form sodium chloride is so exothermic!

  15. Energetics of Ionic Bonding • There must be a third piece to the puzzle. • What is as yet unaccounted for is the electrostatic attraction between the newly formed sodium cation and chloride anion.

  16. Q1Q2 d Eel =  Lattice Energy • This third piece of the puzzle is the lattice energy: • The energy required to completely separate a mole of a solid ionic compound into its gaseous ions. • The energy associated with electrostatic interactions is governed by Coulomb’s law:

  17. IONIC COMPOUNDS • formed from metal cations& non-metal anions • the attractive force between these ions is expressed • as “LATTICE ENERGY” Def: energy released when gaseous ions form an ionic compound M+(g) + N-(g) ---> MN(s) + heat

  18. Attractive forces between ions described by COULOMB’S LAW: charges on the ions E = k Q1+ Q2- + - Lattice Energy: attractive force between ions r distance between ion centers in the lattice constant

  19. Lattice Energy • Lattice energy, then, increases with the charge on the ions. • It also increases with decreasing size of ions.

  20. Lattice Energy • Lattice energy, then, increases with the charge on the ions. • It also increases with decreasing size of ions.

  21. Which compound has a) more ionic bond character? ______ CaO the one with the greater E.N. difference CaO b) a stronger ionic bond? ______ the one with the greater lattice energy 1) CaO or NaCl MPt: 801oC MPt: 2613oC = -1 = -4 Ca2+ O2- Na+ Cl- larger chargesgreater lattice energy so stronger bond (the more heat released, the lower the PE) E.N Diff: 3.5-1.0 = 2.5 E.N Diff: 3.0-0.9 = 2.1

  22. 2) CaO or MgO MPt: 2825oC 2+ EN Diff: 3.5-1.2 = 2.3 2- 2- 2+ E = -4 r E = -4 r So MgO has stronger ionic bond But CaO has more ionic bond character • CaCl2 or K2S both ions smaller E = -2 r E = -2 r EN Diff: 2.5-0.8 = 1.7 EN Diff: 3.0-1.0 = 2.0 So CaCl2 has stronger ionic bond and more ionic bond character

  23. The OVERALL energy changein “formation of an ionic solid” must be calculated in steps: (then add energies together) 1. Enthalpy of sublimationM(s) -----> M(g) 2. Ionization EnergyM(g) -----> M+(g) to form the cation 3. Dissociation EnergyN2(g) -----> 2N(g) for diatomic (also called Bond Energy) 4. Electron AffinityN(g) + e- ----> N-(g) to form the anion 5. Lattice EnergyM+(g)+ N-(g)---> MN(s) when gaseous ions come together

  24. Energetics of Ionic Bonding By accounting for all three energies (ionization energy, electron affinity, and lattice energy), we can get a good idea of the energetics involved in such a process.

  25. Use Table of Bond Energies!! Chemical reactions involve bond-breakingand bond-making. Each chemical bond has an ___________________ in kJ/mol. BOND ENERGY is __________________________________ Always ______ (_____) BOND LENGTH is __________________________________ Like leggos!! See Handout with Table “average” bond energy the energy required to break a bond (+) endo the distance between 2 nuclei connected by a bond BOND ENERGY is __________________________________

  26. - Bond Order 1 2 3 atoms pulled closer # e-prs in a bond

  27. - shortens 1) As # of shared pairs increases, the bond length ___________ 2) The _______the bond energy, the ________ the bond. greater stronger “more stable”

  28. To calculate the DH of the reaction(enthalpy change) using bond energy values: DHrxn= SD (bonds broken) - SD (bonds made) energy absorbed energy released energy/mol Thus if net energy change is (+) meaning “more energy absorbed than released” then DH= (+)indicating an overall endothermic process • initial minus final “D” means bond energy [Note: The only time we use “initial minus final”, rather than “final minus initial”!]

  29. Average Bond Enthalpies • Table 8.4 lists the average bond enthalpies for many different types of bonds. • Average bond enthalpies are positive, because bond breaking is an endothermic process.

  30. Enthalpies of Reaction So, H = [D(C—H) + D(Cl—Cl)] − [D(C—Cl) + D(H—Cl)] = [(413 kJ) + (242 kJ)] − [(328 kJ) + (431 kJ)] = (655 kJ) − (759 kJ) = −104 kJ

  31. making bonds breaking bonds 436 kJ/mol 239 kJ/mol 427 kJ/mol Moles cancel!! 2 431 kJ 242 kJ 436 kJ [ 671 kJ ] - [854kJ ] -184kJ/ 2mol HCl produced We can get to -92kJ/mol f

  32. 1 2 3 1 2 3 4 5 6 • 391 kJ/mol • 941 kJ/mol • 436 kJ/mol • 6 391 kJ • 3 436 kJ • 941 kJ • [ 941 + 1308 ] - [2346 ] • -97kJ/ 2mol NH3produced • -49kJ/mol • -46 • f

  33. Group A # gives # val e-s 5 7 5 1 NH3 PCl3 5 + 21 = 26 total 5 + 3 = 8 total

  34. lost e- gained e-

  35. 2 3

  36. now spread bonds equidistant from each other trigonal planar (polyatomic ion)

  37. Day 5 • FORMAL CHARGE • is a hypothetical charge on an atom in a molecule or ion • helps to determine the best possible • Lewis structure for a molecule or ion • is the difference between the total number of valence electrons of a particular atom and the number of electrons involved in bonds and/or lone pairs • The sum of all the formal charges for a molecule/ion is equal to the charge on that molecule/ion.

  38. Rules: • To count electrons in formal charges: • lonepairs = 2e- • (“unshared” or “non-bonding pairs”) • 2) singlebonds = 1e- • (“shared” or “bonding pairs”) • 3) doublebonds = 2e- • 4) triple bonds = 3e- • Now: Go to Overhead

  39. For 2 “non-equivalent”Lewis Structures, choose the one with: formal charges closest to zero, and the (-) formal charge is on the most electronegative atom

  40. Day 6 Electronic Geometry: geom. of ELECTRON DOMAINS around an atom (used to find hybridization around a particular atom) Molecular Geometry: geom. of ATOMS in the molecule VSEPR Theory meansValence Shell Electron-Pair Repulsion Theory 1) all electrons “paired” 2) all atoms have stable octet, H has duet 3) pairs are spread equidistant from each other around atom used to determine 3-dimensional geometry of molecules and ions

  41. KEY IDEA: bonds (shared pairs) and lone pairs (unshared pairs) arrange themselves so that repulsion is “minimized” • (they are as far apart as they can get!) • 2 ways electrons are positioned around an atom in a molecule • or ion: • 1) in bonds • 2) in lone pairs • called “electron domains”

  42. Molecular Shapes • The shape of a molecule plays an important role in its reactivity. • By noting the number of bonding and nonbonding electron pairs, we can easily predict the shape of the molecule.

  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.

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