Molecular Compounds

1. Molecular Compounds Writing names and Formulas

2. Two Types of Compounds • Molecular compounds • Made of molecules. • Made by joining nonmetal atoms together into molecules.

3. Chemical Formulas • Shows the kind and number of atoms in the smallest piece of a substance. • Molecular formula- number and kinds of atoms in a molecule. • CO2 • C6H12O6

4. Molecular compounds • made of just nonmetals • smallest piece is a molecule • can’t be held together because of opposite charges. • can’t use charges to figure out how many of each atom

5. Easier • Ionic compounds use charges to determine how many of each. • Have to figure out charges. • Have to figure out numbers. • Molecular compounds name tells you the number of atoms. • Uses prefixes to tell you the number

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

7. Prefixes To write the Chemical Formulas write: Prefix name Prefix name -ide

8. Prefixes Exception 1: We don’t write mono- if there is only one of the first element. Example: CO2 Carbon Dioxide CO We do write mono- if the second element is only one. Carbon monoxide

9. Prefixes Exception 2: No double vowels when writing (oa oo) Example: C6O8 Name Hexacarbon octoxide NOT Hexacarbon octaoxide

10. Name These • N2O • NO2 • Cl2O7 • CBr4 • CO2 • BaCl2

11. Write formulas for these • diphosphorus pentoxide • tetraiodide nonoxide • sulfur hexaflouride • nitrogen trioxide • Carbon tetrahydride • phosphorus trifluoride • aluminum chloride

12. Acids Writing names and Formulas

13. Acids • Compounds that give off hydrogen ions when dissolved in water. • Must have H in them. • will always be some H next to an anion. • The anion determines the name.

14. Naming acids • If the anion attached to hydrogen is ends in -ide, put the prefix hydro- and change -ide to -ic acid • HCl - hydrogen ion and chloride ion • hydrochloric acid • H2S hydrogen ion and sulfide ion • hydrosulfuric acid

15. Naming Acids • If the anion has oxygen in it • it ends in -ate of -ite • change the suffix -ate to -ic acid • HNO3 Hydrogen and nitrate ions • Nitric acid • change the suffix -ite to -ous acid • HNO2 Hydrogen and nitrite ions • Nitrous acid

16. Name these • HF • H3P • H2SO4 • H2SO3 • HCN • H2CrO4

17. Writing Formulas • Hydrogen will always be first • name will tell you the anion • make the charges cancel out. • Starts with hydro- no oxygen, -ide • no hydro, -ate comes from -ic, -ite comes from -ous

18. Write formulas for these • hydroiodic acid • acetic acid • carbonic acid • phosphorous acid • hydrobromic acid

19. Covalent bonding

20. + + How does H2 form? • The nuclei repel

21. + + How does H2 form? • The nuclei repel • But they are attracted to electrons • They share the electrons

22. Covalent bonds • Nonmetals hold onto their valence electrons. • They can’t give away electrons to bond. • Still want noble gas configuration. • Get it by sharing valence electrons with each other. • By sharing both atoms get to count the electrons toward noble gas configuration.

23. F Covalent bonding • Fluorine has seven valence electrons

24. F F Covalent bonding • Fluorine has seven valence electrons • A second atom also has seven

25. F F Covalent bonding • Fluorine has seven valence electrons • A second atom also has seven • By sharing electrons

26. Covalent bonding • Fluorine has seven valence electrons • A second atom also has seven • By sharing electrons • Both end with full orbitals F F 8 Valence electrons

27. Covalent bonding • Fluorine has seven valence electrons • A second atom also has seven • By sharing electrons • Both end with full orbitals F F 8 Valence electrons

28. Single Covalent Bond • A sharing of two valence electrons. • Only nonmetals and Hydrogen. • Different from an ionic bond because they actually form molecules by sharing electrons. • Example of Single Covalent Bond • F2 F F

29. Double and Triple Covalent Bonds • If two atoms share two pairs of electrons or 4 ve- they are called double covalent bonds. • Example: O2 O O • If two atoms share three pairs of electrons or 6 ve- they are called triple covalent bonds. • Example: N2 N N

30. Structural Formulas • A structural formula is a molecular model that uses letters as symbols and bonds to show the relative position of atoms to one another

31. Drawing Lewis Structures • Determine the central atom, Always the one least in number in the chemical formula or the atom with the lowest electronegativity. • Draw the skeleton structure • Add up all the valence electrons for the molecule.

32. Drawing Lewis Structures • Divide the total number of available electrons by 2. This gives you the number of bonding pairs. • Place one bonding pair between the cental atom and everyother atom in the skeleton structure. • Subtract the number of bonding pairs used from the number of available electrons pairs. • Place the remaining pairs around the terminal atoms until each atoms fulfills the octet rule. • Any leftover electrons go on the central atom.

33. Examples N • NH3 • N - has 5 valence electrons • H - has 1 valence electrons • NH3 has 5+3(1) = 8 • 8/2= 4 available pairs H

34. Examples • Draw one bond between the N and each H. • Subtract the number of lines from your available pairs. 4 available pairs – 3 lines = one pair left. H H N H

35. Examples Examples • This leaves one available pair but Hydrogen only needs two electrons to be stable. • So the lone pair must go on the Nitrogen. H H N H

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

37. HCN • 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 • 10/2= 5 available pairs. • Draw one bond between the central atom and each terminal atom. H C N

38. HCN • Subtract 5 available pairs – 2 used pairs and you have 3 pairs left. • Carbon needs 2 more pairs and nitrogen needs 3 more pairs. H C N

39. HCN • When you don’t have enough pairs left, you have to make double or triple bonds. H C N

40. HCN H C N To make double or triple bonds you take pairs of electrons off the terminal atoms and share them with the central atom.

41. Polyatomic Ions • Polyatomic ions are charged, but with in the ion there is covalent bonding. • You Draw the Lewis structure the same as with other molecules, but what ever the charge number is, tells you how many e- to add or subtract from the total number of valence e- for the molecule.

42. Resonance • When more than one dot diagram with the same connections are possible. • NO2- • Which one is it? • Does it go back and forth. • It is a mixture of both, like a mule. • NO3-

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

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

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

46. VSEPR • Valence Shell Electron Pair Repulsion. • Predicts three dimensional geometry of molecules. • Name tells you the theory. • Valence shell - outside electrons. • Electron Pair repulsion - electron pairs try to get as far away as possible. • Can determine the angles of bonds.

47. VSEPR • Based on the number of pairs of valence electrons both bonded and unbonded. • Unbonded pair are called lone pair. • CH4 - draw the structural formula

48. VSEPR H • Single bonds fill all atoms. • There are 4 pairs of electrons pushing away. • The furthest they can get away is 109.5º. H C H H

49. 4 atoms bonded • Basic shape is tetrahedral. • A pyramid with a triangular base. • Same shape for everything with 4 pairs. H 109.5º C H H H

50. 3 bonded - 1 lone pair • Still basic tetrahedral but you can’t see the electron pair. • Shape is called trigonal pyramidal. N H N H H H <109.5º H H