Chapter 9 Covalent Bonding
9.1 The Covalent Bond <ul><li>Covalent Bond  – A chemical bond that results from the sharing of valence electrons. </li></...
Covalent Bonds
Polar Covalent Bonds: Unevenly matched, but willing to share.
<ul><li>To describe the composition of a molecular compound we use a  molecular formula. </li></ul><ul><li>The molecular f...
<ul><li>The molecular formula for sucrose (table sugar) is C 12 H 22 O 11 . This indicates that one sucrose molecule conta...
<ul><li>Molecules are often shown using a  structural formula . </li></ul><ul><li>A  structural formula  specifies which a...
<ul><li>The principle that describes covalent bonding is the  octet rule . </li></ul><ul><ul><li>Atoms within a molecule w...
Single Covalent bonding <ul><li>Fluorine has seven valence electrons </li></ul><ul><li>A second F atom also has seven </li...
Covalent bonding <ul><li>Fluorine has seven valence electrons </li></ul><ul><li>A second atom also has seven </li></ul><ul...
Covalent bonding <ul><li>Fluorine has seven valence electrons </li></ul><ul><li>A second atom also has seven </li></ul><ul...
<ul><li>If more than two atoms join in a covalent bond the atoms with the fewest valence electrons will pair up with the a...
<ul><li>The Lewis structures that you have seen so far use a pair of dots to represent covalent bonds. </li></ul><ul><li>A...
<ul><li>Covalent compounds (molecules) tend to have low melting and boiling points </li></ul><ul><li>When an electron pair...
<ul><li>Bond length  depends on the sizes of the bonded atoms and the number of electron pairs they share. </li></ul><ul><...
9.2 Naming Molecules <ul><ul><li>Covalent compounds (Molecules) are named by adding prefixes to the element names. </li></...
<ul><ul><li>Remember to use the “ide” suffix. (Cl is chlor ide  not chlorine) </li></ul></ul><ul><ul><li>Exceptions to the...
Naming Binary Covalent Compounds:   Examples N 2 S 4 di nitrogen  tetra sulfide NI 3 nitrogen  tri iodide XeF 6 xenon  hex...
Naming Acids <ul><li>An  acid  is a molecular substance that dissolves in water to produce hydrogen ions (H+). </li></ul><...
Naming Acids <ul><li>The name of an acid usually results from the  name of the anion(-). </li></ul><ul><ul><li>For acids, ...
Naming Acids <ul><li>  Anion Corresponding Acid </li></ul><ul><li>F - , Fluoride  HF, hydrofluoric acid  </li></ul><ul><li...
9.3 Molecular Structures <ul><li>Structural formulas (Lewis Structures) use letter symbols and bonds to show the  position...
Water <ul><li>Each hydrogen has 1 valence electron </li></ul><ul><li>and wants 1 more </li></ul><ul><li>The oxygen has 6 v...
Water <ul><li>Put the pieces together </li></ul><ul><li>The first hydrogen is happy </li></ul><ul><li>The oxygen still wan...
Water <ul><li>The second hydrogen attaches </li></ul><ul><li>Every atom has full energy levels </li></ul>H H O
Multiple Bonds <ul><li>Sometimes atoms share more than one pair of valence electrons. </li></ul><ul><li>A double bond is w...
Carbon dioxide <ul><li>CO 2   -  Carbon is central atom   ( I have to tell you) </li></ul><ul><li>Carbon has  4 valence el...
Carbon dioxide <ul><li>Attaching 1 oxygen leaves the oxygen 1 short and the carbon 3 short </li></ul>C O
Carbon dioxide <ul><li>Attaching the second oxygen leaves both oxygen 1 short and the carbon 2 short </li></ul>C O O
Carbon dioxide <ul><li>The only solution is to share more </li></ul>C O O
Carbon dioxide <ul><li>The only solution is to share more </li></ul>C O O
Carbon dioxide <ul><li>The only solution is to share more </li></ul>C O O
Carbon dioxide <ul><li>The only solution is to share more </li></ul>C O O
Carbon dioxide <ul><li>The only solution is to share more </li></ul>C O O
Carbon dioxide <ul><li>The only solution is to share more </li></ul>O C O
Carbon dioxide <ul><li>The only solution is to share more </li></ul><ul><li>Requires two double bonds  </li></ul><ul><li>E...
Carbon dioxide <ul><li>The only solution is to share more </li></ul><ul><li>Requires two double bonds  </li></ul><ul><li>E...
Carbon dioxide <ul><li>The only solution is to share more </li></ul><ul><li>Requires two double bonds  </li></ul><ul><li>E...
Carbon dioxide <ul><li>The only solution is to share more </li></ul><ul><li>Requires two double bonds  </li></ul><ul><li>E...
How to draw them <ul><li>To figure out if you need multiple bonds </li></ul><ul><li>Add up all the valence electrons. </li...
Examples <ul><li>NH 3   </li></ul><ul><li>N - has 5 valence electrons wants 8 </li></ul><ul><li>H - has 1 valence electron...
Examples <ul><li>Draw in the bonds </li></ul><ul><li>All  8 electrons are accounted for </li></ul><ul><li>Everything is fu...
Examples <ul><li>HCN  C is central atom </li></ul><ul><li>N - has 5 valence electrons wants 8 </li></ul><ul><li>C - has 4 ...
HCN <ul><li>Put in single bonds </li></ul><ul><li>Need 2 more bonds </li></ul><ul><li>Must go between C and N </li></ul>N ...
HCN <ul><li>Put in single bonds </li></ul><ul><li>Need 2 more bonds </li></ul><ul><li>Must go between C and N </li></ul><u...
HCN <ul><li>Put in single bonds </li></ul><ul><li>Need 2 more bonds </li></ul><ul><li>Must go between C and N </li></ul><u...
Another way of indicating bonds <ul><li>Often use a line to indicate a bond </li></ul><ul><li>Called a structural formula ...
Structural Examples <ul><li>C has 8 electrons because each line is 2 electrons </li></ul><ul><li>Ditto for N </li></ul><ul...
Coordinate Covalent Bond <ul><li>A  Coordinate Covalent   Bond   is when one atom donates both electrons in a covalent bon...
Coordinate Covalent Bond <ul><li>When one atom donates both electrons in a covalent bond. </li></ul><ul><li>Carbon monoxid...
Coordinate Covalent Bond <ul><li>When one atom donates both electrons in a covalent bond. </li></ul><ul><li>Carbon monoxid...
Resonance <ul><li>Resonance  is a condition when more than one dot diagram with the same connections is possible. </li></u...
9.4 Molecular Shape (VSEPR) <ul><li>V alence  S hell  E lectron  P air  R epulsion.  </li></ul><ul><li>Predicts three dime...
VSEPR <ul><li>Based on the number of pairs of valence electrons both bonded and unbonded. </li></ul><ul><li>Unbonded pair ...
VSEPR <ul><li>Single bonds fill all atoms. </li></ul><ul><li>There are 4 pairs of electrons pushing away. </li></ul><ul><l...
4 atoms bonded <ul><li>Basic shape is tetrahedral. </li></ul><ul><li>A pyramid with a triangular base. </li></ul><ul><li>S...
3 bonded -  1 lone pair <ul><li>Still basic tetrahedral but you can’t see the electron pair. </li></ul><ul><li>Shape is ca...
2 bonded -  2 lone pair <ul><li>Still basic tetrahedral but you can’t see the 2 lone pair. </li></ul><ul><li>Shape is call...
3 atoms no lone pair <ul><li>The farthest you can the electron pair apart is 120º. </li></ul><ul><li>Shape is flat and cal...
2 atoms no lone pair <ul><li>With three atoms the farthest they can get apart is 180º. </li></ul><ul><li>Shape called line...
Molecular Orbitals <ul><li>The overlap of atomic orbitals from separate atoms makes molecular orbitals </li></ul><ul><li>E...
Sigma bonding orbitals  <ul><li>From s orbitals on separate atoms </li></ul>+ + s orbital s orbital + + +  + Sigma bonding...
Sigma bonding orbitals  <ul><li>From p orbitals on separate atoms </li></ul>p orbital p orbital Sigma bonding molecular or...
Pi bonding orbitals <ul><li>P orbitals on separate atoms </li></ul>    Pi bonding molecular orbital    
Sigma and pi bonds <ul><li>All single bonds are sigma bonds </li></ul><ul><li>A double bond is one sigma and one pi bond <...
Hybridization <ul><li>The mixing of several atomic orbitals to form the same number of  hybrid  orbitals. </li></ul><ul><l...
Hybridization <ul><li>When an “s” and “p” orbital come together they form a linear shaped molecule. </li></ul><ul><li>When...
9.5 Electronegativity and Bond Type Electro negativity Difference Bond Type <.4 Non-Polar Covalent 0.5 – 1.9 Polar Covalen...
Polar Bonds <ul><li>When the atoms in a bond are the same, the electrons are shared equally. </li></ul><ul><li>This is a  ...
Electronegativity <ul><li>A measure of how strongly the atoms attract electrons in a bond. </li></ul><ul><li>The bigger th...
How to show a bond is polar <ul><li>Isn’t a whole charge just a partial charge </li></ul><ul><li> means a partially pos...
Polar Molecules <ul><li>Polar molecules have  a partially positive end and a partially negative end </li></ul><ul><li>Requ...
Polar Molecules <ul><li>Symmetrical shapes are those without lone pair on central atom  </li></ul><ul><ul><li>Tetrahedral ...
Intermolecular Forces <ul><li>They are what make solid and liquid molecular compounds possible. </li></ul><ul><li>The weak...
Dispersion Force <ul><li>Depends only on the number of electrons in the molecule </li></ul><ul><li>Bigger molecules more e...
Dipole interactions <ul><li>Occur when polar molecules are attracted to each other. </li></ul><ul><li>Slightly stronger th...
Dipole interactions <ul><li>Occur when polar molecules are attracted to each other. </li></ul><ul><li>Slightly stronger th...
Properties of Molecular Compounds <ul><li>Made of nonmetals </li></ul><ul><li>Poor or nonconducting as solid, liquid or aq...
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Ch 9 Covalent Bonding

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Ch 9 Covalent Bonding

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