2. COVALENT BONDING Chapter 16

3. AND THE SUBJECTS ARE… THE NAME IS BOND, COVALENT BOND SINGLES, DOUBLES & TRIPPPLES COORDINATE COVALENT BONDS RESONATE THIS! THERE’S ALWAYS AN EXCEPTION PRINCESS VSEPR HYBRIDS DISSOCIATIONS AND ATTRACTIONS

4. YOU KNOW THAT YOU MAY BE A COVALENT BOND WHEN… ELECTROSTATIC FORCES ARE WEAK  ELECTRONEGATIVITY IS SMALL THERE ARE NO METAL CATIONS IONS DON’T GET OR LOSE ELECTRONS

5. SINGLE COVALENT BOND HYDROGEN H· + ·H  H:H Hydrogen Hydrogen Hydrogen Atom Atom molecule Atom Atom molecule HYDROGEN ATOMS FORM SINGLE BONDS Shared electron pair

6. SINGLE COVALENT BOND DIATOMIC MOLECULES HYDROGENNITROGEN OXYGEN FLUORINE CHLORINEBROMINE IODINE  ELECTRONEGATIVITY = 0 SHARED ELECTRONS SHARED ELECTRONS

7. FOR CHLORINE, 1s 2s 2p 3s 3p 1s 2s 2p 3s 3p Cl ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑_ Cl ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ _↓ SINGLE COVALENT BOND

8. SINGLE COVALENT BOND DEFINITIONS EMPIRICAL FORMULA - LOWEST WHOLE NUMBER RATIO OF ELEMENTS IN A COMPOUND MOLECULAR FORMULA – SHOWS ACTUAL NUMBER AND KINDS OF ELEMENTS IN A COMPOUND STRUCTURAL FORMULA – SHOWS ARRANGEMENT OF ATOMS IN A COMPOUND

9. SINGLE COVALENT BOND

10. ELECTRON PAIRS REPRESENTED BY A LINE BETWEEN 2 ELEMENTS H - H OTHER ELECTRONS MUST BE SHOWN, INCLUDING UNSHARED PAIRS LONE PAIRS NON-BONDING PAIRS

11. SINGLE COVALENT BOND 2 CHLORINE ATOMS FORM A COVALENT BOND

12. WATER MOLECULE FORMED FROM 1 OXYGEN AND 2 HYDROGEN ATOMS USING COVALENT BONDS SINGLE COVALENT BOND

13. CARBON FORMS UP TO 4 BONDS C ↑↓ ↑↓ ↑_ ↑_ 1s 2s 2p 1s 2s 2p TO MAKE THE 4 TH BOND, A 2s ELECTRON MOVES TO A 2p ORBITAL C ↑↓ ↑_ ↑_ ↑_ ↑_ 1s 2s 2p 1s 2s 2p

14. DOUBLE AND TRIPLE COVALENT BONDS

15. CORDINATE COVALENT BONDS BOND FORMED WHEN BOTH ELECTRONS DONATED BY THE SAME ATOM

16. POLYATOMIC IONS POLYATOMIC IONS ARE COVALENTLY BONDED THE CHARGE ON THE ION SHOWS THE NUMBER OF ELECTRONS ADDED OR MISSING

17. RESONANCE RESONANCE STRUCTURES SHOW 2 OR MORE VALID DOT FORMULAS HAVING THE SAME NUMBER OF BONDS AND ELECTRONS

18. EXCEPTIONS to the OCTET RULE IN SOME MOLECULES, THE OCTET RULE CANNOT BE SATISFIED

19. EXCEPTIONS to the OCTET RULE DIAMAGNETIC – ALL ELECTRONS PAIRED PARAMAGNETIC - 1 OR MORE UNPAIRED ELECTRONS EXPANDED OCTET – MAY INCLUDE MORE THAN 8 ELECTRONS

20. MOLECULAR ORBITALS (MO) WHEN 2 ATOMIC ORBITALS OVERLAP, 2 MOLECULAR ORBITALS FORM 1 ORBITAL IS BONDING ENERGY LOWER THAN ATOMIC ORBITAL 1 ORBITAL IS ANTI-BONDING ENERGY HIGHER THAN ATOMIC ORBITAL ORBITAL BELONGS TO MOLECULE AS A WHOLE ORBITAL CONTAINS UP TO 2 ELECTRONS

21. MOLECULAR ORBITALS Energy Atomic Orbital Atomic Orbital Anti-bonding Molecular Orbital Bonding Molecular Orbital

22. MOLECULAR ORBITALS HYDROGEN TWO H 2 MOLECULE ELECTRONS OCCUPY A BONDING ORBITAL PROBABILITY OF FINDING ELECTRONS IN THE AREA BETWEEN 2 ATOMS IS HIGH BOND IS SYMETRICAL ALONG AXIS BETWEEN ATOMS CALLED A SIGMA (σ) BOND CALLED A SIGMA (σ) BOND

23. MOLECULAR ORBITALS Energy Atomic Orbital Atomic Orbital   Anti-bonding Molecular Orbital Bonding Molecular Orbital  H2H2

24. MOLECULAR ORBITALS (MO)

25. MOLECULAR ORBITALS HELIUM 2 MOLECULES OF He CANNOT BE FORMED 2 ELECTRONS IN BONDING ORBITAL 2 ELECTRONS IN ANTI-BONDING ORBITAL REPULSIVE FORCES IN ANTI-BONDING ORBITAL NEGATES ATTRACTIVE FORCES IN BONDING ORBITAL

26. MOLECULAR ORBITALS Energy Atomic Orbital Atomic Orbital  Anti-bonding Molecular Orbital Bonding Molecular Orbital  He

27. MOLECULAR ORBITALS “p” ORBITALS CAN OVERLAP OVERLAP END-TO-END – SIGMA (σ) BOND OVERLAP SIDE-BY-SIDE – PI (  BOND

28. MOLECULAR ORBITALS

30. VALENCE SHELL ELECTRON PAIR REPULSION THEORY VSEPR ELECTRON PAIRS REPEL DUE TO THE SAME CHARGE MOLECULAR SHAPE ADJUSTS TO SET FURTHEST DISTANCE BETWEEN PAIRS CREATES 3 DIMENSIONAL STRUCTURES

31. MOLECULES in 3-D METHANE (CH 4 ) CARBON in CENTER, with H in the CORNERS to form a TETRAHEDRON H-C-H ANGLES ARE 109.5°

32. MOLECULES in 3-D MOLECULAR GEOMETRY

33. MOLECULES in 3-D UNSHARED ELECTRON PAIRS IMPACT SHAPE UNSHARED ELECTRONS HELD CLOSER TO NUCLEUS UNSHARED ELECTRONS ALSO MORE STRONGLY REPEL OTHER ATOMS  ANGLE SHIFTS TO 107°

34. MOLECULES in 3-D DIHYDROGEN OXIDE (H 2 O) MOLECULE BECOMES BENT PLANAR WITH 2 PAIRS UNSHARED ELECTRONS UNSHARED PAIRS REPEL BONDING PAIRS, COMPRESSING ANGLE H-O-H ANGLE IS 105°

35. MOLECULES in 3-D CARBON DIOXIDE (CO 2 ) NO UNSHARED ELECTRON PAIRS IN CARBON DIOXIDE MOLECULE IS LINEAR ANGLE IS 180 o

36. MOLECULES in 3-D MOLECULES WITH NO UNPAIRED ELECTRONS, BUT WITH DOUBLE BOND MOLECULE IS TRIAGONAL PLANAR SHAPE

37. HYBRID ORBITALS SEVERAL ATOMIC ORBITALS CAN MIX TO FORM HYBRID ORBITALS ORBITALS EACH ATOM HAS SEPARATE s AND p ORBITALS

38. HYBRID ORBITALS SOME ATOMIC ORBITALS MERGE, FORMING 4 sp 3 HYBRID ORBITALS ORBITALS EACH ORBITAL BENDS TOWARD CORNER OF TETRAHEDRON

39. HYBRID ORBITALS METHANE (CH 4 ) s ORBITALS OF HYDROGEN OVERLAP CARBON sp 3 ORBITAL ORBITAL FORMS σ BONDS

40. HYBRID ORBITALS DOUBLE BONDS ATOMIC ORBITALS IN DOUBLE BONDS MERGE TO FORM 3 sp 2 HYBRID ORBITALS & 1 p ORBITAL REMAINS ORBITALS BEND TOWARD CORNER OF TRIANGLE, PERPENDICULAR TO REMAINING p ORBITAL

41. HYBRID ORBITALS ETHYLENE (C 2 H 4 ) HYDROGEN ORBITALS OVERLAP WITH CARBON sp 2 ORBITALS ORBITALS 1 sp 2 ORBITAL FROM EACH CARBON BONDS REMAINING CARBON p ORBITALS FORM A  BOND

42. HYBRID ORBITALS TRIPLE BONDS ATOMIC ORBITALS MERGE TO FORM 2 sp HYBRID ORBITALS, AND 2 p ORBITALS REMAIN ORBITALS ORBITAL IS LINEAR AND PERPENDICULAR TO REMAINING p ORBITALS

43. HYBRID ORBITALS ACETYLENE (C 2 H 2 ) HYDROGEN s ORBITALS OVERLAP WITH CARBON sp 2 ORBITALS, AND 1CARBON sp 2 ORBITAL EACH BOND ORBITAL REMAINING CARBON p ORBITALS FORM  BOND

44. POLAR BONDS BOND POLARITY ELECTRONS SHARED IN COVALENT BOND BUT NOT NECESSARILY EQUALLY NON-POLAR BONDS ELECTRONS SHARED EQUALLY POLAR BOND ELECTRONS PULLED STRONGLY TO ATOM WITH HIGHER ELECTRONEGATIVITY

45. POLAR BONDS HYDROGEN IODIDE (HI) HI IODINE (EN = 2.5) HAS A HIGHER ELECTRONEGATIVITY THAN HYDROGEN (EN = 2.1) IODINE (EN = 2.5) HAS A HIGHER ELECTRONEGATIVITY THAN HYDROGEN (EN = 2.1) ELECTRONS DRAWN TOWARD IODINE NUCLEUS REPRESENTED BY: δ + δ - H – I orH - I HYDROGEN DEVELOPS PARTIAL POSITIVE CHARGE, WHILE IODINE COUNTERS WITH A PARTIAL NEGATIVE CHARGE

46. POLAR BONDS WATERWATER AND CHLOROMETHANE ARE EXAMPLES OF POLAR MOLECULES WATER

47. POLAR BONDS TYPE OF BOND IS DETERMINED BY THE ELECTRONEGATIVITY OF THE BONDING ELEMENTS Electronegativity Differences and Bond Type Electronegativity Difference Type of Bond Example 0.0-0.4 Non-polar covalent H-H (0.0) 0.4-1.0 Moderately polar covalent H-Cl (0.9) 1.0-2.0 Very polar covalent H-F (1.9) > 2.0 Ionic Na + Cl - (2.1)

48. POLAR MOLECULES POLAR MOLECULE HAS 1 OR MORE POLAR BONDS CHARGED REGION IS A ‘POLE’ MOLECULE HAVING 2 POLES IS A ‘DIPOLE’ WHEN POLAR MOLECULES ARE PLACED IN AN ELECTRIC FIELD, THEY BECOME ORIENTED WITH RESPECT TO THE FIELD

49. POLAR MOLECULES IN AN ELECTRIC FIELD Negative Plate Positive Plate

50. EFFECT OF POLAR BONDS ON THE MOLECULE’S POLARITY DEPENDS ON SHAPE OF THE MOLECULE AND POLAR BOND ORIENTATION WATER IS A POLAR MOLECULE, BUT CARBON DIOXIDE IS NOT POLAR MOLECULES : O HH : : : : : O = C = O

51. BOND DISSOCIATION ENERGY ENERGY NEEDED TO BREAK A COVALENT BOND

52. INTERMOLECULAR ATTRACTIONS INTERMOLECULAR FORCES DETERMINE COMPOUND’S PHYSICAL STATE VERY WEAK – COMPOUND IS A GAS SOMEWHAT STRONGER – COMPOUND IS A LIQUID MUCH STRONGER – COMPOUND IS A SOLID MOLECULAR COMPOUND MELTING AND BOILING POINTS ARE LOWER THAN IONIC COMPOUNDS MOLECULAR COMPOUNDS ARE INSOLUBLE IN WATER MOLECULAR COMPOUNDS DO NOT CONDUCT ELECTRICITY IN THEIR LIQUID STATE

53. Van der Waals FORCES DISPERSION FORCES CAUSED BY ELECTRON MOTION INCREASES AS ELECTRON NUMBER INCREASES DIPOLE INTERACTIONS ELECTROSTATIC INTERACTION BETWEEN OPPOSITE CHARGED REGIONS OF POLAR MOLECULES

54. HYDROGEN BONDING CONDITIONS FOR HYDROGEN BONDING COVALENT HYDROGEN BOND WITH HIGHLY ELECTRONEGATIVE ELEMENT HYDROGEN WEAKLY BONDED TO UNSHARED ELECTRONDS IN ANOTHER MOLECULE HYDROGEN BONDING OCCURS BECAUSE THERE IS NO SHIELDING EFFECT AROUND HYDROGEN NUCLEUS

55. IONIC & COVALENT COMPOUND CHARACTERISTICS Characteristics Ionic Compounds Covalent Compounds Representative unit Formula Unit Molecule Bond Formation Transfer of one or more electrons between atoms Sharing of electron pairs between atoms Type of elements Metallic and non-metallic Non-metallic Physical State Solid Solid, liquid or gas Melting Point High (usually above 300  C) Low (usually below 300  C) Solubility in water Usually high Usually low Electrical conductivity of (aq) Good conductor Poor to non-conducting

56. COVALENT BONDING THE END