1. Molecular Compounds Unit 7

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. 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 How can we distinguish between an ionic and a molecular compound? How can we distinguish between an ionic and a molecular compound? –Ionic = contains a metal –Molecular = composed of 2 or more nonmetals

4. 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. Molecular Compounds can exist in all 3 states: Molecular Compounds can exist in all 3 states: –Solids – sugar, ice, aspirin –Liquids – water, alcohols –Gases – O 2, CO 2, and N 2 O (laughing gas)

5. Properties Cont. Properties vary depending on the strength of the intermolecular forces. Properties vary depending on the strength of the intermolecular forces. Usually have a lower MP (than ionic, metallic) Usually have a lower MP (than ionic, metallic) Generally soft Generally soft Non-conductors (in any state & aqueous) Non-conductors (in any state & aqueous) Very important to organic chemistry and pharmaceutics. Very important to organic chemistry and pharmaceutics.

6. 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 Larger difference in electronegativities means more polar bond Larger difference in electronegativities means more polar bond –negative end toward more electronegative atom

7. Bond Polarity 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 than the other –Polar covalent – unequal sharing –Nonpolar covalent – equal sharing 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 

8. Naming Molecular Compounds A molecular compound is a compound that is made up of 2 or more nonmetals. A molecular compound is a compound that is made up of 2 or more nonmetals. When naming these compounds simply use the appropriate prefix before the name of the element. When naming these compounds simply use the appropriate prefix before the name of the element. The most electronegative element is written second and ends in –ide. The most electronegative element is written second and ends in –ide. Do not write the prefix mono for the first element, but you should for the second. Do not write the prefix mono for the first element, but you should for the second.

9. PrefixesSubscriptPrefix 1mono- 2di- 3tri- 4tetra- 5penta- 6hexa- 7hepta- 8octa- 9nona- 10deca-

10. Practice CO 2 CO 2 –Carbon dioxide CO CO –Carbon monoxide PCl 5 PCl 5 –Phosphorus pentachloride N 2 O 5 N 2 O 5 –Dinitrogen pentoxide NOTE – Be sure to drop the last vowel of the prefix if there would be any a-o, o-o, or a-a combinations (pentoxide, not pentaoxide). NOTE – Be sure to drop the last vowel of the prefix if there would be any a-o, o-o, or a-a combinations (pentoxide, not pentaoxide).

11. More Practice Cl 2 O Cl 2 O –Dichlorine monoxide Tetraiodine nonoxide Tetraiodine nonoxide –I4O9–I4O9–I4O9–I4O9 SF 6 SF 6 –Sulfur hexafluoride

12. Naming Arrhenius Acids Arrhenius acids are compounds which lose H + ion in H 2 O. Arrhenius acids are compounds which lose H + ion in H 2 O. The general form for an acid is “HA” where H is hydrogen and A is either a monoatomic or polyatomic anion. The general form for an acid is “HA” where H is hydrogen and A is either a monoatomic or polyatomic anion. Here are the rules for naming acids: Here are the rules for naming acids:

13. Naming Arrhenius Acids 1. If the anion part normally ends in –ide (binary acid), then the acid name begins with the prefix hydro and ends with –ic. Ex. HCl is hydrochloric acid 2. If the anion part ends in –ate (polyatomic) then NO hydro is used and the ending is – ic. Ex. HNO 3 is nitric acid (notice – no hydro). 3. If the anion part normally ends in –ite no hydro is used and the ending is –ous. Ex. HNO 2 is nitrous acid

14. Naming Arrhenius Acids 4. To write formulas for acids just use the number of H’s equal to the negative charge of the anion (since each H is +1). Ex. Carbonic acid – no hydro is used so the anion must be polyatomic. The acid name ends in –ic so the anion must end in –ate, i.e. carbonate. Since carbonate is CO 3 2- two H’s are necessary and the formula is H 2 CO 3.

15. Naming Arrhenius Acids NOTE: For a couple of elements only part, or even none, of the element ending is dropped before adding the acid ending. Ex. H 2 SO 4 is not sulfic acid it is sulfuric acid. H 3 PO 4 is phosphoric acid, no phosphic acid.

16. 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. Calculate N-A=S Finding “N” : The amount of electrons needed is eight for every atom, except Finding “N” : The amount of electrons needed is eight for every atom, except (H=2, Be=4, B=6). Finding “A” : This is the number of valence electrons they ALREADY HAVE. Example: H=1, C=4, F= 7 Note: Some compounds can be ions as well, if your compound is charged then the amount of electrons it “already has” must be adjusted to match. Just like you did when you wrote their electron configurations! Example: - add electron for “-” - subtract electron for “+” Writing Lewis Structures of Molecules N – A = S Needed to be “happy” Already have Shared Pairs

17. Writing Lewis Structures of Molecules 3. Divide the result by 2 (because covalent bonds are shared pairs). This is the number of bonds your compound will contain. 4. Write your structure, starting with your central atom. Connect each atom to the central atom with at least 1 bond, then add doubles and triples as required. Meaning/summary: - single = 1 pair of shared electrons (2 e - ) - double = 2 pair of shared electrons (4 e - ) - triple = 3 pair of shared electrons (6 e - ) 5. Finish your masterpiece by adding nonbonding pairs of electrons until your atoms are all filled. Work from the outside towards the central atom. Remember some due not hold eight. (H=2, Be= 4, B=6) 6. Post on refrigerator and admire.

18. Writing Lewis Structures of Molecules - Method #2 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 “+” 3.Divide the total number of electrons by 2 - sharing involves 2 electrons 4.Attach the atoms together with one pair of electrons

19. Writing Lewis Structures Method #2 Cont. 5. All remaining e - = LONE PAIRS! - lone pairs are NOT involved in bonding - lone pairs are NOT involved in bonding 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 - )

20. 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

21. 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.

22. 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 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

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

24. 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 Tetrahedral Tetrahedral –4 surrounding atoms form a tetrahedron around the central atom, no lone pairs around the CA –109.5° bond angles 180° 120° 109.5°

25. 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

26. 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 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 effects the attractive forces between molecules and therefore the physical properties of the substance The dipole moment effects the attractive forces between molecules and therefore the physical properties of the substance

27. 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 Molecule will be POLAR if: Molecule will be POLAR if: –the central atom has lone pairs

28. 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. 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) 2 “p” orbitals overlapping (parallel axes; side-by-side) 2 “p” orbitals overlapping (parallel axes; side-by-side)

31. 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).

32. 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).

36. How do I know if my CA is hybridized? If your CA is B, Be, C, Si, or Al then it is hybridized. If your CA 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

38. 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) Example: Example: –H 2 O and NaCl –NH 3 Properties – universal solvent (H 2 O), unique properties (H 2 O) Properties – universal solvent (H 2 O), unique properties (H 2 O)

39. 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

40. 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: Examples: 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

42. Bond Dissociation Energy energy required to break a single bond energy required to break a single bond EX  H-H + 435 kJ  H ● + H ● EX  H-H + 435 kJ  H ● + H ● –435 kJ of energy/mole of H 2 Shorter bond = Stronger bond = higher energy Shorter bond = Stronger bond = higher energy Calculate the amount of energy needed to dissociate 1 mole of C 2 H 6. Calculate the amount of energy needed to dissociate 1 mole of C 2 H 6.

43. Bondlength (pm) and bond energy (kJ/mol) BondLengthEnergyBondLengthEnergy H--H74435H--C109393 C--C154348H--N101391 N--N145170H--O96366 O--O148145H--F92568 F--F142158H--Cl127432 Cl-Cl199243H--Br141366 Br-Br228193H--I161298 I--I267151 C--C154348 C--C154348C=C134614 C--N147308 CCCC 120839 C--O143360 O--O148145 C=O121736O=O121498 C--F135488N--N145170 C--Cl177330 NNNN 110945