Ch 9 Covalent Bonding

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  • 1. Chapter 9 Covalent Bonding
  • 2. 9.1 The Covalent Bond
    • Covalent Bond – A chemical bond that results from the sharing of valence electrons.
      • Between nonmetallic elements of similar electronegativity.
    • A group of atoms united by a covalent bond is called a molecule .
    • A substance made up of molecules is called a molecular substance .
    • Molecules may have as few as two atoms. (CO), (O 2 )
  • 3. Covalent Bonds
  • 4. Polar Covalent Bonds: Unevenly matched, but willing to share.
  • 5.
    • To describe the composition of a molecular compound we use a molecular formula.
    • The molecular formula tells you how many atoms are in a single molecule of the compound.
      • The molecular formula for oxygen is O 2 , which tells you that this molecule contains 2 atoms of oxygen.
  • 6.
    • The molecular formula for sucrose (table sugar) is C 12 H 22 O 11 . This indicates that one sucrose molecule contains 12 carbon atoms, 22 hydrogen atoms and 11 oxygen atoms.
    • The empirical formula shows the ratio of atoms in the molecule. You get it by dividing all subscripts by the smallest subscript and rounding. C 6 H 12 O 6 becomes CH 2 O.
  • 7.
    • Molecules are often shown using a structural formula .
    • A structural formula specifies which atoms are bonded to each other in a molecule.
      • One type is the Lewis structure.
      • The Lewis structure uses dots as before and has a circle drawn around the dots that represent valence electrons that are joined in a covalent bond.
  • 8.
    • The principle that describes covalent bonding is the octet rule .
      • Atoms within a molecule will share electrons so as to acquire a full set (8) of valence electrons.
    • If each atom has 7 valence electrons, then they will share one so that they both will have 8. The sharing of that electron is a single covalent bond .
  • 9. Single Covalent bonding
    • Fluorine has seven valence electrons
    • A second F atom also has seven
    • By sharing electrons
    • Both end with full orbitals
    F F
  • 10. Covalent bonding
    • Fluorine has seven valence electrons
    • A second atom also has seven
    • By sharing electrons
    • Both end with full orbitals
    F F 8 Valence electrons
  • 11. Covalent bonding
    • Fluorine has seven valence electrons
    • A second atom also has seven
    • By sharing electrons
    • Both end with full orbitals
    F F 8 Valence electrons
  • 12.
    • If more than two atoms join in a covalent bond the atoms with the fewest valence electrons will pair up with the atom that has the most up to a total of 8.
    • If a pair of electrons is not shared, they are called an unshared pair .
    • If two atoms share two pairs of electrons, it is called a double covalent bond.
    • If two atoms share three pairs of electrons, it is called a triple covalent bond.
  • 13.
    • The Lewis structures that you have seen so far use a pair of dots to represent covalent bonds.
    • Another kind of Lewis structure uses dashes to represent covalent bonds.
      • A single dash represents a single covalent bond.
      • A double dash represents a double covalent bond
      • A triple dash represents a triple covalent bond.
  • 14.
    • Covalent compounds (molecules) tend to have low melting and boiling points
    • When an electron pair is shared by the direct overlap of bonding orbitals, a sigma bond results. Single bonds are sigma bonds.
    • The overlap of parallel orbitals forms a pi bond . Multiple bonds involve both sigma and pi bonds.
  • 15.
    • Bond length depends on the sizes of the bonded atoms and the number of electron pairs they share.
    • Bond dissociation energy is the energy needed to break a covalent bond.
      • Double bonds have larger bond dissociation energies than single,
      • triple even larger
    • Bond length and bond dissociation energy are directly related.
  • 16. 9.2 Naming Molecules
      • Covalent compounds (Molecules) are named by adding prefixes to the element names.
        • Covalent’ (in this context) means both elements are nonmetals .
      • The names of molecular compounds will include prefixes that indicate the number of atoms in the molecule. (CO 2 would be called Carbon Di oxide, CCl 4 is called Carbon Tetra chloride)
      • The following prefixes are used:
        • mono-1 di-2
        • tri-3 tetra-4
        • penta-5 hexa-6
        • hepta-7 octa-8
        • nona-9 deca-10
    • Note : When a prefix ending in ‘o’ or ‘a’ is added to ‘oxide’, the final vowel in the prefix is dropped.
  • 17.
      • Remember to use the “ide” suffix. (Cl is chlor ide not chlorine)
      • Exceptions to the rule :
        • The prefix “mono” is usually not written with the first word of the compounds name.
          • CO 2 is not called mono carbon di oxide, just carbon dioxide.
        • Some compounds have common names. (H 2 O is water or dihydrogen monoxide)
  • 18. Naming Binary Covalent Compounds: Examples N 2 S 4 di nitrogen tetra sulfide NI 3 nitrogen tri iodide XeF 6 xenon hexa fluoride CCl 4 carbon tetra chloride P 2 O 5 di phosphorus pent oxide SO 3 sulfur tri oxide 1 mono 2 di 3 tri 4 tetra 5 penta 6 hexa 7 heptaa 8 octa 9 nona 10 deca * Second element in ‘ide’ from * Drop –a & -o before ‘oxide’
  • 19. Naming Acids
    • An acid is a molecular substance that dissolves in water to produce hydrogen ions (H+).
      • Acids behave like an ionic compound because when they dissolve in water, they create cations and anions.
      • The cation is always H+ and the anion depends on the particular acid.
        • Example : When the acid HCl dissolves in water it forms H+ cations and Cl- anions.
        • HNO 3 acid will form H+ and NO 3 - ions.
  • 20. Naming Acids
    • The name of an acid usually results from the name of the anion(-).
      • For acids, where the anion ends with the suffix “-ide” (ex.Chloride), mainly binary acids, the name of the acid begins with the prefix hydro.
      • The acids name also includes the root name of the anion(-) and the word acid.
      • In addition you need to change the suffix of the anion from “ide” to “ic”.
        • By these rules HCl is named hydrochloric acid.
    • Other acids are named without the prefix “hydro”. The acid HNO 3 - derives its name from NO 3 - which is nitrate. HNO 3 is named Nitric Acid.
    • The following are names and formulas of common acids:
  • 21. Naming Acids
    • Anion Corresponding Acid
    • F - , Fluoride HF, hydrofluoric acid
    • Cl - , Chloride HCl, hydrochloric acid
    • Br - , Bromine HBr, hydrobromic acid
    • I - , iodide HI, hydroiodic acid
    • S 2 - , sulfide H 2 S, hydrosulfuric acid
    • NO 3 - , nitrate HNO 3 , nitric acid
    • CO 3 -2 , carbonate H 2 CO 3 , carbonic acid
    • SO 4 -2 , sulfate H 2 SO 4 , sulfuric acid
    • PO 4 -3 , phoshate H 3 PO 4 , Phosphoric acid
    • C 2 H 3 O 2 - , acetate HC 2 H 3 O 2 , acetic acid
  • 22. 9.3 Molecular Structures
    • Structural formulas (Lewis Structures) use letter symbols and bonds to show the position of atoms.
    • Drawing Lewis Structures:
  • 23. Water
    • Each hydrogen has 1 valence electron
    • and wants 1 more
    • The oxygen has 6 valence electrons
    • and wants 2 more
    • They share to make each other “happy”
    H O
  • 24. Water
    • Put the pieces together
    • The first hydrogen is happy
    • The oxygen still wants one more
    H O
  • 25. Water
    • The second hydrogen attaches
    • Every atom has full energy levels
    H H O
  • 26. Multiple Bonds
    • Sometimes atoms share more than one pair of valence electrons.
    • A double bond is when atoms share two pair (4) of electrons.
    • A triple bond is when atoms share three pair (6) of electrons.
  • 27. Carbon dioxide
    • CO 2 - Carbon is central atom ( I have to tell you)
    • Carbon has 4 valence electrons
    • Wants 4 more
    • Oxygen has 6 valence electrons
    • Wants 2 more
    C O
  • 28. Carbon dioxide
    • Attaching 1 oxygen leaves the oxygen 1 short and the carbon 3 short
    C O
  • 29. Carbon dioxide
    • Attaching the second oxygen leaves both oxygen 1 short and the carbon 2 short
    C O O
  • 30. Carbon dioxide
    • The only solution is to share more
    C O O
  • 31. Carbon dioxide
    • The only solution is to share more
    C O O
  • 32. Carbon dioxide
    • The only solution is to share more
    C O O
  • 33. Carbon dioxide
    • The only solution is to share more
    C O O
  • 34. Carbon dioxide
    • The only solution is to share more
    C O O
  • 35. Carbon dioxide
    • The only solution is to share more
    O C O
  • 36. Carbon dioxide
    • The only solution is to share more
    • Requires two double bonds
    • Each atom gets to count all the atoms in the bond
    O C O
  • 37. Carbon dioxide
    • The only solution is to share more
    • Requires two double bonds
    • Each atom gets to count all the atoms in the bond
    O C O 8 valence electrons
  • 38. Carbon dioxide
    • The only solution is to share more
    • Requires two double bonds
    • Each atom gets to count all the atoms in the bond
    O C O 8 valence electrons
  • 39. Carbon dioxide
    • The only solution is to share more
    • Requires two double bonds
    • Each atom gets to count all the atoms in the bond
    O C O 8 valence electrons
  • 40. How to draw them
    • To figure out if you need multiple bonds
    • Add up all the valence electrons.
    • Count up the total number of electrons to make all atoms happy.
    • Subtract.
    • Divide by 2
    • Tells you how many bonds - draw them.
    • Fill in the rest of the valence electrons to fill atoms up.
  • 41. Examples
    • NH 3
    • N - has 5 valence electrons wants 8
    • H - has 1 valence electrons wants 2
    • NH 3 has 5+3(1) = 8
    • NH 3 wants 8+3(2) = 14
    • (14-8)/2= 3 bonds
    • 4 atoms with 3 bonds
    N H
  • 42. Examples
    • Draw in the bonds
    • All 8 electrons are accounted for
    • Everything is full
    N H H H
  • 43. Examples
    • HCN C is central atom
    • N - has 5 valence electrons wants 8
    • C - has 4 valence electrons wants 8
    • H - has 1 valence electrons wants 2
    • HCN has 5+4+1 = 10
    • HCN wants 8+8+2 = 18
    • (18-10)/2= 4 bonds
    • 3 atoms with 4 bonds -will require multiple bonds - not to H
  • 44. HCN
    • Put in single bonds
    • Need 2 more bonds
    • Must go between C and N
    N H C
  • 45. HCN
    • Put in single bonds
    • Need 2 more bonds
    • Must go between C and N
    • Uses 8 electrons - 2 more to add
    N H C
  • 46. HCN
    • Put in single bonds
    • Need 2 more bonds
    • Must go between C and N
    • Uses 8 electrons - 2 more to add
    • Must go on N to fill octet
    N H C
  • 47. Another way of indicating bonds
    • Often use a line to indicate a bond
    • Called a structural formula
    • Each line is 2 valence electrons
    H H O = H H O
  • 48. Structural Examples
    • C has 8 electrons because each line is 2 electrons
    • Ditto for N
    • Ditto for C here
    • Ditto for O
    H C N C O H H
  • 49. Coordinate Covalent Bond
    • A Coordinate Covalent Bond is when one atom donates both electrons in a covalent bond.
    • Carbon monoxide
    • CO
    O C
  • 50. Coordinate Covalent Bond
    • When one atom donates both electrons in a covalent bond.
    • Carbon monoxide
    • CO
    O C
  • 51. Coordinate Covalent Bond
    • When one atom donates both electrons in a covalent bond.
    • Carbon monoxide
    • CO
    O C O C
  • 52. Resonance
    • Resonance is a condition when more than one dot diagram with the same connections is possible.
    • Resonance structures differ only in the position of the electron pairs, never the atom positions.
  • 53. 9.4 Molecular Shape (VSEPR)
    • V alence S hell E lectron P air R epulsion.
    • Predicts three dimensional geometry of molecules.
    • The name tells you the theory.
    • Valence shell - outside electrons.
    • Electron Pair Repulsion - electron pairs try to get as far away as possible.
    • Can determine the angles of bonds.
    • And the shape of molecules
  • 54. VSEPR
    • Based on the number of pairs of valence electrons both bonded and unbonded.
    • Unbonded pair are called lone pair .
    • CH 4 - draw the structural formula
  • 55. VSEPR
    • Single bonds fill all atoms.
    • There are 4 pairs of electrons pushing away.
    • The furthest they can get away is 109.5º.
    C H H H H
  • 56. 4 atoms bonded
    • Basic shape is tetrahedral.
    • A pyramid with a triangular base.
    • Same basic shape for everything with 4 pairs.
    C H H H H 109.5º
  • 57. 3 bonded - 1 lone pair
    • Still basic tetrahedral but you can’t see the electron pair.
    • Shape is called trigonal pyramidal.
    N H H H N H H H <109.5º
  • 58. 2 bonded - 2 lone pair
    • Still basic tetrahedral but you can’t see the 2 lone pair.
    • Shape is called bent.
    O H H O H H <109.5º
  • 59. 3 atoms no lone pair
    • The farthest you can the electron pair apart is 120º.
    • Shape is flat and called trigonal planar.
    • Will require 1 double bond
    C H H O C H H O 120º
  • 60. 2 atoms no lone pair
    • With three atoms the farthest they can get apart is 180º.
    • Shape called linear.
    • Will require 2 double bonds or one triple bond
    C O O 180º
  • 61. Molecular Orbitals
    • The overlap of atomic orbitals from separate atoms makes molecular orbitals
    • Each molecular orbital has room for two electrons
    • Two types of MO
      • Sigma ( σ ) between atoms
      • Pi ( π ) above and below atoms
  • 62. Sigma bonding orbitals
    • From s orbitals on separate atoms
    + + s orbital s orbital + + + + Sigma bonding molecular orbital
  • 63. Sigma bonding orbitals
    • From p orbitals on separate atoms
    p orbital p orbital Sigma bonding molecular orbital      
  • 64. Pi bonding orbitals
    • P orbitals on separate atoms
        Pi bonding molecular orbital    
  • 65. Sigma and pi bonds
    • All single bonds are sigma bonds
    • A double bond is one sigma and one pi bond
    • A triple bond is one sigma and two pi bonds.
  • 66. Hybridization
    • The mixing of several atomic orbitals to form the same number of hybrid orbitals.
    • When an atom approaches another atom to form a bond, the orbitals of its electrons may be changed.
    • The changing of the orbitals forms what is called hybrid orbitals. This is a cross between all of the orbitals involved (s,p,d,f).
  • 67. Hybridization
    • When an “s” and “p” orbital come together they form a linear shaped molecule.
    • When an “s” and “p2” orbital come together they form a trigonal planer molecule.
    • When an “s” and “p3” orbital come together they form tetrahedral, pyramidal or bent molecules.
    • Atoms with higher orbital numbers form longer bonds. Multiple bonds are shorter than single bonds because bonds with more electrons attract the nuclei of the bonding atoms more strongly.
  • 68. 9.5 Electronegativity and Bond Type Electro negativity Difference Bond Type <.4 Non-Polar Covalent 0.5 – 1.9 Polar Covalent >2.0 Ionic
  • 69. Polar Bonds
    • When the atoms in a bond are the same, the electrons are shared equally.
    • This is a nonpolar covalent bond .
    • When two different atoms are connected, the electrons may not be shared equally.
    • This is a polar covalent bond .
  • 70. Electronegativity
    • A measure of how strongly the atoms attract electrons in a bond.
    • The bigger the electronegativity difference the more polar the bond.
    • 0.0 - 0.4 Covalent nonpolar
    • 0.5 - 1.0 Covalent moderately polar
    • 1.0 -2.0 Covalent polar
    • >2.0 Ionic
  • 71. How to show a bond is polar
    • Isn’t a whole charge just a partial charge
    •  means a partially positive
    •  means a partially negative
    • The Cl pulls harder on the electrons
    • The electrons spend more time near the Cl
    H Cl  
  • 72. Polar Molecules
    • Polar molecules have a partially positive end and a partially negative end
    • Requires two things to be true:
      • The molecule must contain polar bonds.This can be determined from differences in electronegativity.
      • Symmetry can not cancel out the effects of the polar bonds.
    • Must determine geometry first.
  • 73. Polar Molecules
    • Symmetrical shapes are those without lone pair on central atom
      • Tetrahedral
      • Trigonal planar
      • Linear
    • Will be nonpolar if all the atoms are the same
    • Shapes with lone pair on central atom are not symmetrical
    • Can be polar even with the same atom
  • 74. Intermolecular Forces
    • They are what make solid and liquid molecular compounds possible.
    • The weakest are called van der Waal’s forces - there are two kinds
      • Dispersion forces
      • Dipole Interactions
  • 75. Dispersion Force
    • Depends only on the number of electrons in the molecule
    • Bigger molecules more electrons
    • More electrons stronger forces
        • F 2 is a gas
        • Br 2 is a liquid
        • I 2 is a solid
  • 76. Dipole interactions
    • Occur when polar molecules are attracted to each other.
    • Slightly stronger than dispersion forces.
    • Opposites attract but not completely hooked like in ionic solids.
  • 77. Dipole interactions
    • Occur when polar molecules are attracted to each other.
    • Slightly stronger than dispersion forces.
    • Opposites attract but not completely hooked like in ionic solids.
    H F     H F    
  • 78. Properties of Molecular Compounds
    • Made of nonmetals
    • Poor or nonconducting as solid, liquid or aqueous solution
    • Low melting point
    • Two kinds of crystals
      • Molecular solids held together by IMF
      • Network solids- held together by bonds
      • One big molecule (diamond, graphite)