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CHEMISTRY 161 Chapter 9

CHEMISTRY 161 Chapter 9. Periodic Table of the Elements. ns 2 np x. ns 1. ns 2. chemical reactivity - valence electrons. THE OCTET RULE. atoms combine to form compounds in an attempt to obtain a stable noble gas electron configuration. ns 2 np 6. Iso electronic. A + B → AB.

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CHEMISTRY 161 Chapter 9

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  1. CHEMISTRY 161 Chapter 9

  2. Periodic Table of the Elements ns2npx ns1 ns2 chemical reactivity - valence electrons

  3. THE OCTET RULE atoms combine to form compounds in an attempt to obtain a stable noble gas electron configuration ns2np6 Iso electronic

  4. A + B → AB 1. ELECTRON FULLY TRANSFERED EXP I IONIC BONDING 2 Na(s) + Cl2(g)  2 NaCl(s) 2. ELECTRON SHARING COVALENT BONDING 2 H2(g) + O2(g)  2 H2O(l)

  5. LEWIS MODEL OF BONDING LEWIS DOT SYMBOL DOT represents one valence electron H. Gilbert Lewis (1875-1946)

  6. . . . . . . . . . . with the exception of He, the main group number represents number of ‘dots’ only valence electron are considered

  7. IONIC BONDING Na electron transfer Ne core implied in symbol 1s22s22p63s1 Lewis Symbol

  8. Na Ne core implied in symbol Cl 1s22s22p63s23p5 1s22s22p63s1 Lewis Symbol

  9. IONIC BONDING Cl Cl Na the formation of ionic bonds is represented in terms of Lewis symbols  Na+ 1s22s22p63s23p6 1s22s22p6 the loss or gain of electrons(dots) until both species have reached an octet of electrons

  10. Cl Cl [Ne] 3s23p6 represents one orbital (Pauli: 2 electrons)

  11. ions stack together in regular crystalline structures electrostatic interaction ionic solids typically 1. high melting and boiling points 2. brittle 3. form electrolyte solutions if they dissolve in water

  12. Li(s) + ½ F2(g) → LiF(s) enthalpy of formation lattice energy (up to few 1000 kJmol-1) Li+(g) + F-(g) → LiF(s) Hess’s Law Born-Haber Cycle

  13. Li(s) + ½ F2(g) → LiF(s) Li+(g) + F-(g) 5 ΔHoR= ΣΔHoi i=1 ΔHo4 ΔHo3 ΔHo5 Li(g) + F(g) ΔHo2 ΔHo1 ΔHoR LiF(s) Li(s) + ½ F2(g)

  14. Mg(s) + ½ O2(g) → MgO(s) Mg2+(g) + O2-(g) ΔHo6 7 ΔHo5 ΔHoR= ΣΔHoi i=1 Mg+(g) + O-(g) ΔHo7 ΔHo4 ΔHo3 Mg(g) + O(g) ΔHo1 ΔHo2 ΔHoR MgO(s) Mg(s) + ½ O2(g)

  15. COVALENT BONDING F F F F THE OCTET RULE sharing electrons (electron pair) electronic configuration of F is 1s22s22p5

  16. F F F F F F + non-bonding, or lone pair of electrons bonding pair of electrons

  17. H2is the simplest covalent molecule + H H H H + the bond length of H2 is the distance where the total energy of the molecule is minimum

  18. EXAMPLES NH3 HX H2O CH4 single bonds O2 CO2 C2H4 double bonds C2H2 N2 HCN triple bonds

  19. few 100 kJ/mol

  20. IONIC OR COVALENT electronegativity difference between two atoms involved in the bond

  21. ELECTRONEGATIVITY Li C N O F Na S Cl K Br Rb I Cs Ba Fr Ra is the tendency of an atom in a bond to attract shared electrons to itself F > O > N, Cl > Br > I, C, S …….. Na, Ba, Ra > K, Rb > Cs, Fr Electronegativity increases energies of the atomic orbital with the unpaired electron

  22. ELECTRONEGATIVITY Li C N O F Na S Cl P K Br Rb I Cs Ba Fr Ra F > O > N, Cl > Br > I, C, S …….. Na, Ba, Ra > K, Rb > Cs, Fr Electronegativity increases Se F is the most electronegative H has an electronegativity about the same a P

  23. IONIC VERSUS COVALENT BONDS bonds are neither completely ionic nor covalent (only in homonuclear molecules)

  24. IONIC VERSUS COVALENT BONDS compounds composed of elements with large difference in ELECTRONEGATIVITY significant ionic character in their bonding B has greater electronegativity A B

  25. IONIC VERSUS COVALENT BONDS B has a greater share A B

  26. H F H F HYDROGEN FLUORIDE Fluorine is more electronegative than hydrogen. + +

  27. H F H F HYDROGEN FLUORIDE Fluorine is more electronegative than hydrogen. + + d+ d– This is a polar covalent bond (dipole moment). The bond has a partly ionic and partly covalent nature.

  28. Microwave Spectroscopy molecules need a dipole moment

  29. Variation of ionic character with electronegativity.

  30. LEWIS SYMBOLS IONIC COMPOUDS COVALENT COMPOUNDS ELECTRONEGATIVITY

  31. H O H Lewis considers only valence electrons H2O bonding pair of electrons non-bonding, or lone pair of electrons single – double – triple

  32. LEWIS STRUCTURES 1. concept of resonances 2. exceptions to the octet rule

  33. O O O O O O 1. RESONANCES NO3- O: 1s22s22p4 N: 1s22s22p3 plus one extra electron for negative charge

  34. - + -

  35. O O O experiment shows all three bonds are the same 128 pm N bond angles 120 0 any one of the structures suggests one is different!

  36. O N O O 128 pm bond angles 120 0 modify the description by blending the structures blending of structures is called resonance

  37. O O O O O N N N O O O O RESONANCE use a double headed arrow between the structures electrons involved are said to be DELOCALIZED over the structure. blended structure is a RESONANCE HYBRID

  38. O O N O O O O O O N N N O O O O RESONANCE We use a double headed arrow between the structures..

  39. 2. Exceptions to the octet rule 1. more than 8 electrons around central atom 2. less than an octet around central atom 3. molecules with unpaired electrons

  40. F F F F F F S S F F F F F F 1. more than 8 electrons around central atom elements in rows 3 and following can exceed octet rule SF6 participation of d electrons

  41. Lewis structure for SF6 1s22s22p5 F has seven S has six 1s22s22p63s22p4 SF2SF4SF6 PF3PF5 NF3NF5 ClO4- SO42- I3-

  42. 2. less than an octet around central atom BeH2 AlF3 resonances BF3 NH3 (dative bond) Lewis base Lewis acids

  43. 3. molecules with unpaired electrons FREE RADICALS NO but not NO-

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