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Chemistry - Chp 8 - Covalent Bonding - PowerPoint

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Chemistry - Chp 8 - Covalent Bonding - PowerPoint

  1. 1. Chapter 8“Covalent Bonding”
  2. 2. Section 8.1 Molecular CompoundsOBJECTIVES:–Distinguish between the melting points and boiling points of molecular compounds and ionic compounds.
  3. 3. Section 8.1 Molecular CompoundsOBJECTIVES:–Describe the information provided by a molecular formula.
  4. 4. Bonds are… Forces that hold groups of atoms together and make them function as a unit. Two types:1) Ionic bonds – transfer of electrons (gained or lost; makes formula unit)2) Covalent bonds – sharing of electrons. The resulting particle is called a “molecule”
  5. 5. Covalent BondsThe word covalent is acombination of the prefix co-(from Latin com, meaning “with”or “together”), and the verbvalere, meaning “to be strong”.Two electrons shared togetherhave the strength to hold twoatoms together in a bond.
  6. 6. Molecules Many elements found in nature are in the form of molecules:  a neutral group of atoms joined together by covalent bonds. For example, air contains oxygen molecules, consisting of two oxygen atoms joined covalently Called a “diatomic molecule” (O2)
  7. 7. How does H2 form? (diatomic hydrogen molecule)The nuclei repel each other,since they both have a positivecharge (like charges repel).+ + + +
  8. 8. How does H2 form?But, the nuclei are attracted tothe electronsThey share the electrons, andthis is called a “covalent bond”,and involves only NONMETALS! + +
  9. 9. Covalent bondsNonmetals hold on to their valenceelectrons.They can’t give away electrons to bond. – But still want noble gas configuration.Get it by sharing valence electrons witheach other = covalent bondingBy sharing, both atoms get to countthe electrons toward a noble gasconfiguration.
  10. 10. Covalent bondingFluorine has seven valenceelectrons (but would like to have 8) F
  11. 11. Covalent bondingFluorine has seven valenceelectronsA second atom also has seven F F
  12. 12. Covalent bondingFluorine has seven valenceelectronsA second atom also has sevenBy sharing electrons… F F
  13. 13. Covalent bondingFluorine has seven valenceelectronsA second atom also has sevenBy sharing electrons… F F
  14. 14. Covalent bondingFluorine has seven valenceelectronsA second atom also has sevenBy sharing electrons… F F
  15. 15. Covalent bondingFluorine has seven valenceelectronsA second atom also has sevenBy sharing electrons… F F
  16. 16. Covalent bondingFluorine has seven valenceelectronsA second atom also has sevenBy sharing electrons… F F
  17. 17. Covalent bondingFluorine has seven valence electronsA second atom also has sevenBy sharing electrons……both end with full orbitals F F
  18. 18. Covalent bondingFluorine has seven valence electronsA second atom also has sevenBy sharing electrons……both end with full orbitals F F 8 Valence electrons
  19. 19. Covalent bonding Fluorine has seven valence electrons A second atom also has seven By sharing electrons… …both end with full orbitals F F8 Valenceelectrons
  20. 20. Molecular CompoundsCompounds that are bondedcovalently (like in water, or carbondioxide) are called molecularcompoundsMolecular compounds tend to haverelatively lower melting and boilingpoints than ionic compounds – this isnot as strong a bond as ionic
  21. 21. Molecular CompoundsThus, molecular compounds tend tobe gases or liquids at roomtemperature – Ionic compounds were solidsA molecular compound has amolecular formula: – Shows how many atoms of each element a molecule contains
  22. 22. Molecular CompoundsThe formula for water is written asH2O – The subscript “2” behind hydrogen means there are 2 atoms of hydrogen; if there is only one atom, the subscript 1 is omittedMolecular formulas do not tell anyinformation about the structure (thearrangement of the various atoms).
  23. 23. - Page 215These are some of the 3. The ball and stick model isdifferent ways to represent the BEST, because it showsammonia: a 3-dimensional arrangement.1. The molecularformula showshow many atomsof each elementare present2. The structuralformula ALSOshows thearrangement ofthese atoms!
  24. 24. Section 8.2The Nature of Covalent BondingOBJECTIVES: –Describe how electrons are shared to form covalent bonds, and identify exceptions to the octet rule.
  25. 25. Section 8.2The Nature of Covalent BondingOBJECTIVES: –Demonstrate how electron dot structures represent shared electrons.
  26. 26. Section 8.2The Nature of Covalent BondingOBJECTIVES: –Describe how atoms form double or triple covalent bonds.
  27. 27. Section 8.2The Nature of Covalent BondingOBJECTIVES: –Distinguish between a covalent bond and a coordinate covalent bond, and describe how the strength of a covalent bond is related to its bond dissociation energy.
  28. 28. Section 8.2The Nature of Covalent BondingOBJECTIVES: –Describe how oxygen atoms are bonded in ozone.
  29. 29. A Single Covalent Bond is...A sharing of two valence electrons.Only nonmetals and hydrogen.Different from an ionic bondbecause they actually formmolecules.Two specific atoms are joined.In an ionic solid, you can’t tell whichatom the electrons moved from or to
  30. 30. Sodium Chloride Crystal Lattice•Ionic compoundsorganize in acharacteristiccrystal lattice ofalternatingpositive andnegative ions,repeated over andover.
  31. 31. How to show the formation…It’s like a jigsaw puzzle.You put the pieces together to end upwith the right formula.Carbon is a special example - can itreally share 4 electrons: 1s22s22p2? – Yes, due to electron promotion!Another example: lets show how water isformed with covalent bonds, by using anelectron dot diagram
  32. 32. Water  Each hydrogen has 1 valenceH electron - Each hydrogen wants 1 more  The oxygen has 6 valence electronsO - The oxygen wants 2 more  They share to make each other complete
  33. 33. WaterPut the pieces togetherThe first hydrogen is happyThe oxygen still needs one more HO
  34. 34. WaterSo, a second hydrogen attachesEvery atom has full energy levels Note the two “unshared” pairs HO of electrons H
  35. 35. Examples:1. Conceptual Problem 8.1 on page 2202. Do PCl3
  36. 36. Multiple BondsSometimes atoms share more thanone pair of valence electrons.A double bond is when atoms sharetwo pairs of electrons (4 total)A triple bond is when atoms sharethree pairs of electrons (6 total)Table 8.1, p.222 - Know these 7elements as diatomic: Br2 I2 N2 Cl2 H2 O2 F2
  37. 37. Dot diagram for Carbon dioxide CO2 - Carbon is central atom ( more metallic ) C Carbon has 4 valence electrons Wants 4 more O Oxygen has 6 valence electrons Wants 2 more
  38. 38. Carbon dioxideAttaching 1 oxygen leaves theoxygen 1 short, and the carbon 3short CO
  39. 39. Carbon dioxideAttaching the second oxygenleaves both of the oxygen 1 short,and the carbon 2 short OC O
  40. 40. Carbon dioxideThe only solution is to share more O CO
  41. 41. Carbon dioxideThe only solution is to share more O CO
  42. 42. Carbon dioxideThe only solution is to share more O CO
  43. 43. Carbon dioxideThe only solution is to share more O C O
  44. 44. Carbon dioxideThe only solution is to share more O C O
  45. 45. Carbon dioxideThe only solution is to share more O C O
  46. 46. Carbon dioxideThe only solution is to share moreRequires two double bondsEach atom can count all theelectrons in the bond O C O
  47. 47. Carbon dioxideThe only solution is to share moreRequires two double bondsEach atom can count all the electrons inthe bond 8 valence electrons O C O
  48. 48. Carbon dioxideThe only solution is to share moreRequires two double bondsEach atom can count all the electrons inthe bond 8 valence electrons O C O
  49. 49. Carbon dioxideThe only solution is to share moreRequires two double bondsEach atom can count all the electrons inthe bond 8 valence electrons O C O
  50. 50. How to draw them? Use the handout guidelines:1) Add up all the valence electrons.2) Count up the total number of electrons to make all atoms happy.3) Subtract; then Divide by 24) Tells you how many bonds to draw5) Fill in the rest of the valence electrons to fill atoms up.
  51. 51. Example NH3, which is ammoniaN N – central atom; has 5 valence electrons, wants 8 H - has 1 (x3) valence electrons, wants 2 (x3)H NH3 has 5+3 = 8 NH3 wants 8+6 = 14 (14-8)/2= 3 bonds 4 atoms with 3 bonds
  52. 52. ExamplesDraw in the bonds; start with singlesAll 8 electrons are accounted forEverything is full – done with this one. H H NH
  53. 53. Example: HCNHCN: C is central atomN - has 5 valence electrons, wants 8C - has 4 valence electrons, wants 8H - has 1 valence electron, wants 2HCN has 5+4+1 = 10HCN wants 8+8+2 = 18(18-10)/2= 4 bonds3 atoms with 4 bonds – this will requiremultiple bonds - not to H however
  54. 54. HCNPut single bond between each atomNeed to add 2 more bondsMust go between C and N (Hydrogen is full) HC N
  55. 55. HCNPut in single bondsNeeds 2 more bondsMust go between C and N, not the HUses 8 electrons – need 2 more toequal the 10 it has HC N
  56. 56. HCNPut in single bondsNeed 2 more bondsMust go between C and NUses 8 electrons - 2 more to addMust go on the N to fill its octet HC N
  57. 57. Another way of indicating bondsOften use a line to indicate a bondCalled a structural formulaEach line is 2 valence electronsHOH H O H =
  58. 58. Other Structural ExamplesH C NH C OH
  59. 59. A Coordinate Covalent Bond... When one atom donates both electrons in a covalent bond. Carbon monoxide (CO) is a good example:Both the carbon COand oxygen giveanother singleelectron to share
  60. 60. Coordinate Covalent Bond When one atom donates both electrons in a covalent bond. Carbon monoxide (CO) is a good example: Oxygen gives both ofThis carbon theseelectron C O electrons,moves to since it hasmake a pair no morewith the other singles tosingle. share.
  61. 61. Coordinate Covalent Bond When one atom donates both electrons in a covalent bond. Carbon monoxide (CO)Thecoordinatecovalent bondis shown withan arrow as:C O C O
  62. 62. Coordinate covalent bondMost polyatomic cations andanions contain covalent andcoordinate covalent bondsTable 8.2, p.224Sample Problem 8.2, p.225The ammonium ion (NH41+) can beshown as another example
  63. 63. Bond Dissociation Energies...The total energy required to breakthe bond between 2 covalentlybonded atomsHigh dissociation energy usuallymeans the chemical is relativelyunreactive, because it takesa lot of energy to break it down.
  64. 64. Resonance is...When more than one valid dotdiagram is possible.Consider the two ways to draw ozone(O3)Which one is it? Does it go back andforth?It is a hybrid of both, like a mule; andshown by a double-headed arrowfound in double-bond structures!
  65. 65. Resonance in Ozone Note the different location of the double bondNeither structure is correct, it isactually a hybrid of the two. To showit, draw all varieties possible, and jointhem with a double-headed arrow.
  66. 66. ResonanceOccurs when more than one valid Lewisstructure can be written for a particularmolecule (due to position of double bond)•These are resonance structures of benzene.•The actual structure is an average (or hybrid)of these structures.
  67. 67. Polyatomic ions – note the different positions of the double bond.Resonancein acarbonateion (CO32-):Resonancein anacetate ion(C2H3O21-):
  68. 68. The 3 Exceptions to Octet rule For some molecules, it is impossible to satisfy the octet rule #1. usually when there is an odd number of valence electrons – NO2 has 17 valence electrons, because the N has 5, and each O contributes 6. Note “N” page 228 It is impossible to satisfy octet rule, yet the stable molecule does exist
  69. 69. Exceptions to Octet rule• Another exception: Boron • Page 228 shows boron trifluoride, and note that one of the fluorides might be able to make a coordinate covalent bond to fulfill the boron • #2 -But fluorine has a high electronegativity (it is greedy), so this coordinate bond does not form• #3 -Top page 229 examples exist because they are in period 3 or beyond
  70. 70. Section 8.3 Bonding TheoriesOBJECTIVES:–Describe the relationship between atomic and molecular orbitals.
  71. 71. Section 8.3 Bonding TheoriesOBJECTIVES:–Describe how VSEPR theory helps predict the shapes of molecules.
  72. 72. Molecular Orbitals are...The model for covalent bondingassumes the orbitals are those ofthe individual atoms = atomic orbitalOrbitals that apply to the overallmolecule, due to atomic orbitaloverlap are the molecular orbitals – A bonding orbital is a molecular orbital that can be occupied by two electrons of a covalent bond
  73. 73. Molecular Orbitals - definitions Sigma bond- when two atomic orbitals combine to form the molecular orbital that is symmetrical along the axis connecting the nuclei Pi bond- the bonding electrons are likely to be found above and below the bond axis (weaker than sigma) Note pictures on the next slide
  74. 74. - Pages 230 and 231Sigma bond is symmetrical along theaxis between the two nuclei. Pi bond is above and below the bond axis, and is weaker than sigma
  75. 75. VSEPR: stands for...Valence Shell Electron Pair RepulsionPredicts the three dimensional shape ofmolecules.The name tells you the theory:– Valence shell = outside electrons.– Electron Pair repulsion = electron pairs try to get as far away as possible from each other.Can determine the angles of bonds.
  76. 76. VSEPRBased on the number of pairs ofvalence electrons, both bonded andunbonded.Unbonded pair also called lone pair.CH4 - draw the structural formulaHas 4 + 4(1) = 8wants 8 + 4(2) = 16(16-8)/2 = 4 bonds
  77. 77. VSEPR for methane (a gas): Single bonds fill all atoms. H There are 4 pairs ofH C H electrons pushing away. H The furthest theyThis 2-dimensional drawing doesnot show a true representation of can get away is the chemical arrangement. 109.5º
  78. 78. 4 atoms bondedBasic shape istetrahedral.A pyramid with a H 109.5ºtriangular base.Same shape foreverything with C4 pairs. H H H
  79. 79. Other angles, pages 232 - 233 Ammonia (NH3) = 107o Water (H2O) = 105o Carbon dioxide (CO2) = 180 o Note the shapes of these that are pictured on the next slide
  80. 80. - Page 232Methane hasan angle of109.5o, calledtetrahedralAmmonia hasan angle of107o, calledpyramidalNote the unshared pair that is repulsion for other electrons.
  81. 81. Section 8.4Polar Bonds and MoleculesOBJECTIVES:–Describe how electronegativity values determine the distribution of charge in a polar molecule.
  82. 82. Section 8.4Polar Bonds and MoleculesOBJECTIVES:–Describe what happens to polar molecules when they are placed between oppositely charged metal plates.
  83. 83. Section 8.4Polar Bonds and MoleculesOBJECTIVES:–Evaluate the strength of intermolecular attractions compared with the strength of ionic and covalent bonds.
  84. 84. Section 8.4Polar Bonds and MoleculesOBJECTIVES:–Identify the reason why network solids have high melting points.
  85. 85. Bond PolarityCovalent bonding means sharedelectrons– but, do they share equally?Electrons are pulled, as in a tug-of-war, between the atoms nuclei– In equal sharing (such as diatomic molecules), the bond that results is called a nonpolar covalent bond
  86. 86. Bond PolarityWhen two different atoms bondcovalently, there is an unequalsharing – the more electronegative atom will have a stronger attraction, and will acquire a slightly negative charge – called a polar covalent bond, or simply polar bond.
  87. 87. Electronegativity? The ability of an atom in a molecule to attract shared electrons to itself.Linus Pauling 1901 - 1994
  88. 88. Table of Electronegativities
  89. 89. Bond PolarityRefer to Table 6.2, p. 177 (or handout)Consider HClH = electronegativity of 2.1Cl = electronegativity of 3.0– the bond is polar– the chlorine acquires a slight negative charge, and the hydrogen a slight positive charge
  90. 90. Bond PolarityOnly partial charges, much lessthan a true 1+ or 1- as in ionic bondWritten as: δ+ δ− H Clthe positive and minus signs (withthe lower case delta: δ + and δ − )denote partial charges.
  91. 91. Bond Polarity Can also be shown: H Cl – the arrow points to the more electronegative atom. Table 8.3, p.238 shows how the electronegativity can also indicate the type of bond that tends to form
  92. 92. Polar moleculesSample Problem 8.3, p.239A polar bond tends to make theentire molecule “polar”– areas of “difference”HCl has polar bonds, thus is a polarmolecule.– A molecule that has two poles is called dipole, like HCl
  93. 93. Polar moleculesThe effect of polar bonds on thepolarity of the entire molecule dependson the molecule shape– carbon dioxide has two polar bonds, and is linear = nonpolar molecule!
  94. 94. Polar moleculesThe effect of polar bonds on thepolarity of the entire molecule dependson the molecule shape– water has two polar bonds and a bent shape; the highly electronegative oxygen pulls the e- away from H = very polar!
  95. 95. Polar moleculesWhen polar molecules areplaced between oppositelycharged plates, they tend tobecome oriented with respectto the positive and negativeplates.Figure 8.24, page 239
  96. 96. Attractions between molecules and liquid They are what make solid molecular compounds possible. The weakest are called van der Waal’s forces - there are two kinds:#1. Dispersion forces weakest of all, caused by motion of e- increases as # e- increases halogens start as gases; bromine is liquid; iodine is solid – all in Group 7A
  97. 97. #2. Dipole interactionsOccurs when polar molecules areattracted to each other.2. Dipole interaction happens inwater– Figure 8.25, page 240– positive region of one molecule attracts the negative region of another molecule.
  98. 98. #2. Dipole interactionsOccur when polar molecules areattracted to each other.Slightly stronger than dispersion forces.Opposites attract, but not completelyhooked like in ionic solids. δ+ δ− δ+ δ− H F H F
  99. 99. #2. Dipole Interactions δ+ δ+ − δ δ− + δ− δ δ− δ+ δ+ δ− δ− δ+ δ+ δ + δ− δ−0
  100. 100. #3. Hydrogen bonding …is the attractive force caused by hydrogen bonded to N, O, F, or Cl N, O, F, and Cl are very electronegative, so this is a very strong dipole. And, the hydrogen shares with the lone pair in the molecule next to it. This is the strongest of the intermolecular forces.1
  101. 101. Order of Intermolecular attraction strengths 1) Dispersion forces are the weakest 2) A little stronger are the dipole interactions 3) The strongest is the hydrogen bonding 4) All of these are weaker than ionic bonds2
  102. 102. #3. Hydrogen bonding defined: When a hydrogen atom is: a) covalently bonded to a highly electronegative atom, AND b) is also weakly bonded to an unshared electron pair of a nearby highly electronegative atom. – The hydrogen is left very electron deficient (it only had 1 to start with!) thus it shares with something nearby – Hydrogen is also the ONLY element with no shielding for its nucleus when involved in a covalent bond!3
  103. 103. Hydrogen Bonding (Shown in water) δ- H O δ+ δ+ δ+ δ- H H O δ+ This hydrogen is bonded H covalently to: 1) the highly negative oxygen, and 2) a nearby unshared pair.4
  104. 104. Hydrogen bonding allows H2O to be a liquid at room conditions. H H H O O H O H H O H H H H H H O O O H H5
  105. 105. Attractions and properties Why are some chemicals gases, some liquids, some solids? – Depends on the type of bonding! – Table 8.4, page 244 Network solids – solids in which all the atoms are covalently bonded to each other6
  106. 106. Attractions and properties Figure 8.28, page 243 Network solids melt at very high temperatures, or not at all (decomposes) – Diamond does not really melt, but vaporizes to a gas at 3500 oC and beyond – SiC, used in grinding, has a melting point of about 2700 oC7
  107. 107. Covalent Network Compounds Some covalently bonded substances DO NOT form discrete molecules. Diamond, a network of Graphite, a network of covalently bonded carbon covalently bonded carbon atoms atoms8
  108. 108. 9

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