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COVALENT BONDING

COVALENT BONDING. When one nonmetal shares one or more electrons with an atom of another nonmetal so both atoms end up with eight valence electrons. COVALENT BOND FORMATION. There is a mutual attraction of different nuclei to the electron’s orbitals.

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COVALENT BONDING

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  1. COVALENT BONDING

  2. When one nonmetalshares one or more electrons with an atom of another nonmetal so both atoms end up with eight valence electrons COVALENT BOND FORMATION • There is a mutual attraction of different nuclei to the electron’s orbitals.

  3. Characteristics of covalent compounds: MOLECULES 1.Composed of non-metals that are sharing electrons. 2. Composed of molecules 3. There are two forces involved: Intramolecular – strong covalent bond with in the molecules Intermolecular – weak bondsbetweenthe molecules Van der Waals forces 4. Have low Melting Points (weak Van derWaals forces are breaking)

  4. Intramolecular O Intermolecular H H O Example: Water Strong covalent bonds between H & O, make up the water molecules • Every atom has full energy levels H H Weak attractive forces between different water molecules.

  5. 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 1. (London) Dispersion forces 2. Dipole Interactions

  6. (London) Dispersion Forces • Depend on the number of electrons • More electrons stronger forces • Bigger molecules more electrons • Fluorine is a gas • Bromine is a liquid • Iodine is a solid

  7. d+d- d+d- H F H F 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.

  8. d+d- d+d- d+d- d+d- d+d- d+d- d+d- Dipole Interactions d+d-

  9. Hydrogen bonding • Are the attractive forces caused by hydrogen bonded to F, O, or N. • F, O, and N are very electronegative so it is a very strong dipole. • The hydrogen partially share with the lone pair in the molecule next to it. • The strongest of the intermolecular forces.

  10. d- d+ O d+ H d+ d- H H O H d+ Hydrogen Bonding

  11. Diatomic Molecules (HONClBrIF) – Molecules composed of two atoms of the same element. ** All Diatomic molecules are purely covalent or non-polar covalent because: there is an equal sharing between the two atoms. EX: H2 1s1 1s1 H H H : H or H-H (180) Bond angle LINEAR SHAPE Both atoms strive to fill their 1s orbital so both hydrogen's attract the pair of bonding electrons equally Electrons spend time here; mutual attraction for the same electron pair. + : + Forms a NEW Molecular orbital * s-s bonding is the only bonding that is non-directional + +

  12. Fluorine has seven valence electrons • A second atom also has seven • By sharing electrons • Both end with full orbitals F F 8 Valence electrons 8 Valence electrons Linear shape (bond angle 180) Directional Bonding: Atoms approach at 2z Other possible bonding ( s-p, p-p, s-d or p-d)

  13. Single Covalent Bond • H2 and F2 are both examples of single covalent bonding. • A sharing of two valence electrons. • One pair of electrons shared between two atoms 6 unshared pairs How many unshared pairs of electrons does I2 have? I I (Unshared pairs are also called lone pairs or nonbonding pairs)

  14. Multiple Bonds • Sometimes atoms share more than one pair of valence electrons. • A double bond is when atoms share two pair of electrons. (4 electrons) • A triple bond is when atoms share three pair of electrons. (6 electrons)

  15. Double Covalent Bond : : :O::O: : O=O : or Ex: O2 Bond angle: 180 Shape: linear : : How many pairs of unshared electrons does O2 have? (4) :O=O: Triple Covalent Bond Ex: N2 :N:::N: Bond angle: 180 Shape: linear : N :N=N: How many pairs of unshared electrons does N2 have? (2)

  16. C O Coordinate Covalent Bond • When one atom donates both electrons in a covalent bond. • Carbon monoxide • CO

  17. Coordinate Covalent Bond • When one atom donates both electrons in a covalent bond. • Carbon monoxide • CO C O

  18. Coordinate Covalent Bond • When one atom donates both electrons in a covalent bond. • Carbon monoxide • CO :C=O: C O Triple Bond Oxygen donates a pair of electrons so both atoms now have 8 valence electrons.

  19. Ex: SO2 (Sulfur dioxide) : : : : : : . . . . . :S :S O: O: :O :O . . . . . .

  20. Ex: SO2 (Sulfur dioxide) : : : :O S: :O: : : Double bond

  21. Ex: SO2 (Sulfur dioxide) : : : : : : :O S=O: : :O S: :O: : : Double bond Sulfur donates a pair of electrons to oxygen so all 3 atoms have complete octets! Can this molecule be drawn another way and still be the same molecule? YES

  22. RESONANCE 2 or more valid electron dot formulas that can be written for a molecule. Ex: Ozone (O3) : : : . . . :O O: O: . . . Shares this pair of electrons

  23. RESONANCE : : • 2 or more valid electron dot formulas that can be written for a molecule. : : : :O=O O: : : :O::O:O: : or : : : :O:O::O: :

  24. Ex: SO2 (Sulfur dioxide) : : : :O S: :O: : : or : : : : : : : : O S O :

  25. Drawing Lewis Dot Structures • Draw a skeleton structure putting the first atom written in the center (except Hydrogen) • Add up all the valence electrons. • Count up the total number of electrons to make all atoms have a stable octet. • Subtract. • Divide by 2 • Tells you how many bonds - draw them. • Fill in the rest of the valence electrons to fill atoms up.

  26. Examples N • NH3 • N - has 5 valence electrons wants 8 • H - has 1 valence electrons wants 2 • NH3 has5+3(1) = 8 Valence • NH3 wants 8+3(2) = 14 Valence • (14-8)/2= 3 bonds • 4 atoms with 3 bonds H

  27. Examples • Draw in the bonds • 3 bonds = 6 electrons (8-6=2) • All 8 valence electrons are accounted for • Everything is full H H N H

  28. 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 • HCNhas 5+4+1 = 10 • HCNwants 8+8+2 = 18 • (18-10)/2= 4 bonds • 3 atoms with 4 bonds -will require multiple bonds - not to H

  29. HCN • Put in single bonds • Need 2 more bonds • Must go between C and N H C N

  30. HCN • Put in single bonds • Need 2 more bonds • Must go between C and N • Uses 8 electrons (2 more to add) 10 – 8 = 2 H C N

  31. 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 H C N

  32. 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 atoms may not be shared equally. • This is a polar covalent bond. • How do we measure how strong the atoms pull on electrons?

  33. 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 – 2.0 Covalent polar • >2.0 Ionic

  34. How to show a bond is polar • Isn’t a whole charge just a partial charge • d+ means a partially positive • d- means a partially negative (2.1)(3.0) difference = .9 (polar covalent) *The smaller the difference the more covalent the bond. • The Cl pulls harder on the electrons • The electrons spend more time near the Cl d+ d- H Cl

  35. Polar Molecules • Molecules with a positive and a 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.

  36. Is it polar? Bond (electronegativity) Bond Molecule • HF H=2.1 F=4.0 (1.9) v. polar polar • H2O H=2.1 O=3.5 (1.4) v. polar polar • NH3 H=2.1 N=3.0 (.9) polar polar • CCl4 C=2.5 Cl=3.0 (.5) polar non-polar • CO2 C=2.5 O=3.5 (1.0) v. polar non-polar

  37. MOLECULAR SHAPES OF COVALENT COMPOUNDS

  38. VSEPR THEORY ALENCE HELL VSEPR LECTRON AIR EPULSION

  39. What Vsepr means Since electrons do not like each other, because of their negative charges, they orient themselves as far apart as possible, from each other. Nonbonding electron pairs on the center atom strongly repel the bonding pairs, pushing the bonding pairs closer together This leads to molecules having specific shapes.

  40. HERE ARE THE RESULTING MOLECULAR SHAPES

  41. EXAMPLE: BeF2 Linear • Number of atoms = 3 • Number of Bonds = 2 • Number of Shared Pairs of Electrons = 2 • Bond Angle = 180° ** All molecules with 2 atoms are linear

  42. Bent #1 EXAMPLE: H2O • Number of atoms = 3 • Number of Bonds = 2 • Number of Shared Pairs of Electrons = 2 • Number of Unshared Pairs of Electrons = 2 • Bond Angle = 105°

  43. Bent #2 EXAMPLE: O3 • Number of atoms = 3 • Number of Bonds = 2 • Number of Shared Pairs of Electrons = 2 • Number of Unshared Pairs of Electrons = 1 • Bond Angle = 105°

  44. Trigonal Planar EXAMPLE: GaF3 • Number of atoms = 4 • Number of Bonds = 3 • Number of Shared Pairs of Electrons = 3 • Number of Unshared Pairs of Electrons = 0 • Bond Angle = 120°

  45. EXAMPLE: NH3 Pyramidal • Number of atoms = 4 • Number of Bonds = 3 • Number of Shared Pairs of Electrons = 3 • Number of Unshared Pairs of Electrons = 1 • Bond Angle = 107°

  46. EXAMPLE: CH4 Tetrahedral • Number of atoms = 5 • Number of Bonds = 4 • Number of Shared Pairs of Electrons = 4 • Number of Unshared Pairs of Electrons = 0 • Bond Angle = 109.5°

  47. COVALENT MOLECULES SUMMARY Use electronegativity chart Difference: 0-.4 non polar .5 – 1.7 polar covalent > 1.7 Ionic Depends on bond linear no 180 linear Non- polar (all terminal atoms same) No unshared pairs 180 bent Unshared pairs 105 polar Trigonal planar Non-polar (all terminal atoms same) 120 No unshared pairs pyramidal Unshared pairs 107 polar Non-polar (all terminal atoms same) No unshared pairs tetrahedral 109.5

  48. THE END

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