1. Molecular Bonding Unit 5

2. Covalent Bonds Sharing pairs of electrons Sharing pairs of electrons Covalent bonds are the inter-atomic attraction resulting from the sharing of electrons between atoms. Covalent bonds are the inter-atomic attraction resulting from the sharing of electrons between atoms. They result in ‘localized overlaps’ of orbitals of different atoms. They result in ‘localized overlaps’ of orbitals of different atoms. They also are the result of the attraction of electrons for the nucleus of other atoms. They also are the result of the attraction of electrons for the nucleus of other atoms. Typical of molecular substances. Typical of molecular substances.

3. Example of a Covalent Bond

4. Covalent Bonds Cont. Atoms bond together to form molecules Atoms bond together to form molecules –molecules are electrically neutral groups of atoms joined together by covalent bonds –strong attraction Molecules attracted to each other weakly form molecular compounds Molecules attracted to each other weakly form molecular compounds

5. Properties of Molecular Compounds Strong covalent bonds hold the atoms together within a molecule. Strong covalent bonds hold the atoms together within a molecule. The intermolecular forces that hold one molecule to another are much weaker. The intermolecular forces that hold one molecule to another are much weaker. Properties vary depending on the strength of the intermolecular forces. Properties vary depending on the strength of the intermolecular forces.

6. Lewis structures A Lewis structure is a representation of the valence e- in an atom, ion or molecule. A Lewis structure is a representation of the valence e- in an atom, ion or molecule. Element symbol represents nucleus and core e-. Element symbol represents nucleus and core e-. Dots represent valence e-. Dots represent valence e-.

8. Electron pairs In covalent compounds electrons are shared between atoms creating electron pairs. In covalent compounds electrons are shared between atoms creating electron pairs. Bonding pairs: e- that are shared between 2 atoms. Bonding pairs: e- that are shared between 2 atoms. Lone or unshared pairs: e- that are NOT involved in bonding.Lone or unshared pairs: e- that are NOT involved in bonding.

9. Covalent bonds Hydrogen follows the duet rule: sharing 2 electrons. Hydrogen follows the duet rule: sharing 2 electrons. Non-metals Carbon through Fluorine follow the octet rule: sharing 8 electrons. Non-metals Carbon through Fluorine follow the octet rule: sharing 8 electrons.

10. Writing Lewis Structures of Molecules 1. Determine the central atom (atom in the middle) - usually is the “single” atom - least electronegative element - H never in the middle; C always in the middle 2.Count the total number of valence e - (group #) - add ion charge for “-” - subtract ion charge for “+”

11. 3. Divide the total number of electrons by 2 - sharing involves 2 electrons 4. Attach the atoms together with one pair of electrons 5. All remaining e - = LONE PAIRS! - lone pairs are NOT involved in bonding - lone pairs are NOT involved in bonding

12. Writing Lewis Structures Cont. 6. Place lone pairs around non-central atoms to fulfill the “octet rule” - some elements may violate this octet rule – (H=2, Be=4, B=6) - some elements may violate this octet rule – (H=2, Be=4, B=6) 7. If more e - are still needed, create double or triple bonds around the central atom. or triple bonds around the central atom. - single = 1 pair of shared electrons (2 e - ) - single = 1 pair of shared electrons (2 e - ) - double = 2 pair of shared electrons (4 e - ) - double = 2 pair of shared electrons (4 e - ) - triple = 3 pair of shared electrons (6 e - ) - triple = 3 pair of shared electrons (6 e - )

13. Lewis structure water

14. Practice Give the Lewis structure for: Give the Lewis structure for: HCl HCl NH 3 NH 3 C 2 H 6 C 2 H 6 CO 2 CO 2

15. NAS When drawing Lewis structures, remember: terminal atoms, atoms that can only make one bond, must be on the outside. When drawing Lewis structures, remember: terminal atoms, atoms that can only make one bond, must be on the outside. N – A = S N – A = S #e- needed for – #e- available = #e- shared #e- needed for – #e- available = #e- shared octet or duet (valence e-) (bonds) octet or duet (valence e-) (bonds) S = the number of pairs of shared electrons S = the number of pairs of shared electrons S = the number of bonds (a dash may be used to S = the number of bonds (a dash may be used to 2 represent a pair of shared electrons) 2 represent a pair of shared electrons)

16. Resonance When there is more than one Lewis structure for a molecule that differ only in the position of the electrons they are called resonance structures When there is more than one Lewis structure for a molecule that differ only in the position of the electrons they are called resonance structures –Lone Pairs and Multiple Bonds in different positions Resonance only occurs when there are double bonds and when the same atoms are attached to the central atom Resonance only occurs when there are double bonds and when the same atoms are attached to the central atom The actual molecule is a combination of all the resonance forms. The actual molecule is a combination of all the resonance forms. O S O

17. Exceptions to the Octet Rule H, Be, B (stable with 2, 4, and 6 e-, respectively) H, Be, B (stable with 2, 4, and 6 e-, respectively) Some molecules cannot be drawn with the Lewis structure rules, due to odd # of e-. i.e. NO & NO 2 (There is no way for N to get an octet) Some molecules cannot be drawn with the Lewis structure rules, due to odd # of e-. i.e. NO & NO 2 (There is no way for N to get an octet)

18. Some molecules are stable when the center atom has more than an octet. i.e. SF 6, PCl 5 sulfur has 12 electrons P has 10 electrons

19. Coordinate Covalent Bond A covalent bond in which one atom contributes both bonding electrons. A covalent bond in which one atom contributes both bonding electrons.

20. Polar Bonds: Electronegativity Measure of the ability of an atom to attract shared electrons Measure of the ability of an atom to attract shared electrons –Larger electronegativity means atom attracts more strongly –Values 0.7 to 4.0 Increases across period (left to right) on Periodic Table Increases across period (left to right) on Periodic Table Decreases down group (top to bottom) on Periodic Table Decreases down group (top to bottom) on Periodic Table

21. Larger difference in electronegativities means more polar bond Larger difference in electronegativities means more polar bond –negative end toward more electronegative atom Covalent bonding between unlike atoms results in unequal sharing of the electrons Covalent bonding between unlike atoms results in unequal sharing of the electrons –One end of the bond has larger electron density (more electronegative) than the other –Polar covalent – unequal sharing –Nonpolar covalent – equal sharing

22. Bond Polarity The result is bond polarity The result is bond polarity –The end with the larger electron density gets a partial negative charge –The end that is electron deficient gets a partial positive charge H F 

24. Predicting Molecular Geometry VSEPR Theory VSEPR Theory –Valence Shell Electron Pair Repulsion The shape around the central atom(s) can be predicted by assuming that the areas of electrons on the central atom will repel each other The shape around the central atom(s) can be predicted by assuming that the areas of electrons on the central atom will repel each other

25. Each Bond counts as 1 area of electrons Each Bond counts as 1 area of electrons –single, double or triple all count as 1 area Each Lone Pair counts a 1 area of electrons Each Lone Pair counts a 1 area of electrons –Even though lone pairs are not attached to other atoms, they do “occupy space” around the central atom –Lone pairs generally “push harder” than bonding electrons, affecting the bond angle

26. Shapes Straight Line Straight Line –molecule made up of only 2 atoms

27. Shapes - Linear –2 atoms on opposite sides of central atom, no lone pairs around CA –180° bond angles Trigonal Planar Trigonal Planar –3 atoms form a triangle around the central atom, no lone pairs around CA –Planar –120° bond angles 180° 120°

28. Tetrahedral Tetrahedral –4 surrounding atoms form a tetrahedron around the central atom, no lone pairs around the CA –109.5° bond angles 109.5°

29. Shapes Trigonal Pyramidal Trigonal Pyramidal –3 bonding areas and 1 lone pair around the CA –Bond angle = 107 0 V-shaped or Bent V-shaped or Bent –2 bonding areas and 2 lone pairs around the CA –bond angle = 104.5 0

30. Dipole Moment Bond polarity results in an unequal electron distribution, resulting in areas of partial positive and partial negative charge Bond polarity results in an unequal electron distribution, resulting in areas of partial positive and partial negative charge Any molecule that has a center of positive charge and a center of negative charge in different points is said to have a dipole moment Any molecule that has a center of positive charge and a center of negative charge in different points is said to have a dipole moment

31. If a molecule has more than one polar covalent bond, the areas of partial negative and positive charge for each bond will partially add to or cancel out each other If a molecule has more than one polar covalent bond, the areas of partial negative and positive charge for each bond will partially add to or cancel out each other The end result will be a molecule with one center of positive charge and one center of negative charge The end result will be a molecule with one center of positive charge and one center of negative charge The dipole moment affects the attractive forces between molecules and therefore the physical properties of the substance The dipole moment affects the attractive forces between molecules and therefore the physical properties of the substance

32. Charge distribution in the water molecule

33. Polarity of Molecules Molecule will be NONPOLAR if: Molecule will be NONPOLAR if: –the bonds are nonpolar (Br-Br, F-F) –there are no lone pairs around the central atom and all the atoms attached to the central atom are the same

34. Molecule will be POLAR if: Molecule will be POLAR if: –the central atom has lone pairs –there are no lone pairs around the central atom and all the atoms attached to the central atom are NOT the same

35. Intermolecular forces From weak to strong: From weak to strong: - Dispersion - Dipole-dipole - Hydrogen bonding - Ion-dipole (attraction between ions and dipole molecules) - Ionic Dispersion, dipole-dipole, and hydrogen bonding are Van der Waals forces

36. Intermolecular Forces Hydrogen Bonding – extreme dipole bonding involving hydrogen and a very electronegative element (FON) Hydrogen Bonding – extreme dipole bonding involving hydrogen and a very electronegative element (FON) Examples: Examples: –H2O–H2O–H2O–H2O –NH 3 Properties – universal solvent (H 2 O), unique properties (H 2 O) Properties – universal solvent (H 2 O), unique properties (H 2 O)

37. Hydrogen Bonding in H 2 O H-bonding is especially strong in water because the O—H bond is very polar There are 2 lone pairs on the O atom Accounts for many of water’s unique properties.

38. Intermolecular Forces Dipole-Dipole – interactions between 2 polar bonds or molecules Dipole-Dipole – interactions between 2 polar bonds or molecules Examples: Examples: –sugar and H 2 O –acids Properties – produce acids, dissolve molecular (organic) solids in H 2 O Properties – produce acids, dissolve molecular (organic) solids in H 2 O

39. Intermolecular Forces Dispersion (aka London Dispersion, Induced Dipole) – interaction that is proportional to the number of e - and proportional to the size of the e - cloud Dispersion (aka London Dispersion, Induced Dipole) – interaction that is proportional to the number of e - and proportional to the size of the e - cloud –Results from motion of electrons Examples: non-polar molecules Examples: non-polar molecules Properties – help explain the states of matter and the states of the halogens Properties – help explain the states of matter and the states of the halogens

41. Hybridization refers to a mixture or a blending refers to a mixture or a blending Biology – refers to genetic material Biology – refers to genetic material Chemistry – refers to blending of orbitals Chemistry – refers to blending of orbitals Remember, orbitals can only predict an area in space where an e - may be located. Remember, orbitals can only predict an area in space where an e - may be located. Sometimes blending orbitals can produce a lower, more stable bonding opportunity. Sometimes blending orbitals can produce a lower, more stable bonding opportunity. Orbital hybridization occurs through e - promotion in orbitals that have similar energies (i.e. same energy level). Orbital hybridization occurs through e - promotion in orbitals that have similar energies (i.e. same energy level).

42. Hybridization Cont. Hybridization occurs WITHIN the atom to enhance bonding possibilities. Hybridization occurs WITHIN the atom to enhance bonding possibilities. Do not confuse this concept with orbital overlap (bonding). Do not confuse this concept with orbital overlap (bonding). Hybridization is a concept used to explain observed phenomenon about bonding that can’t be explained by dot structures. Hybridization is a concept used to explain observed phenomenon about bonding that can’t be explained by dot structures. EXAMPLES – draw box diagrams for Be, B, and C (use noble gas core). EXAMPLES – draw box diagrams for Be, B, and C (use noble gas core).

46. How do I know if my central atom is hybridized? If your central atom is B, Be, C, Si, or Al then it is hybridized. If your central atom is B, Be, C, Si, or Al then it is hybridized. If your molecule has multiple bonds in it then it is hybridized. If your molecule has multiple bonds in it then it is hybridized. –Double bonds – sp 2 hybridized –Triple bonds – sp hybridized

48. Sigma & Pi Bonds Sigma Bond (σ) – combination of orbitals that is symmetrical around the axis connecting the two nuclei. Sigma Bond (σ) – combination of orbitals that is symmetrical around the axis connecting the two nuclei. Pi Bond – parallel overlap of p orbitals causing bonding electrons to be found above & below the bond axis. Pi Bond – parallel overlap of p orbitals causing bonding electrons to be found above & below the bond axis. Sigma bonds: 2 “s” orbitals overlapping 2 “s” orbitals overlapping 1 “s” and 1 “p” orbital overlapping 1 “s” and 1 “p” orbital overlapping 2 “p” orbitals overlapping (same axes; end-to-end) 2 “p” orbitals overlapping (same axes; end-to-end) Pi bonds: 2 “p” orbitals overlapping (parallel axes; side-by-side) 2 “p” orbitals overlapping (parallel axes; side-by-side)