1. IIIIIIIV Ch. 9 – Chemical Names and Formulas I. Ion Formation Ionic Formulas Ionic Nomenclature

2. Ionic Bonds zWhen oppositely charged ions attract, electrostatic force that holds them together = ionic bond zCompounds containing ionic bonds = ionic compounds zElectrons are transferred from cations to anions zTypically bonds between metals and nonmetals, however polyatomic ions also involved.

3. A. Vocabulary zChemical Bond yattractive force between atoms or ions that binds them together as a unit ybonds form in order to… xfulfill octet rule xincrease stability

4. Octet Rule zAtoms will gain or lose electrons so that they have 8 electrons in their highest energy level zThe noble gases already have a full octet, 8 valence electrons, so are chemically stable.

5. zCation yPositively charged ion formed when an atom loses one or more valence e - yNumber of protons stays the same, but less electrons gives + charge A. Vocabulary Loses an e- Sodium ion

6. zAnions yNonmetals easily gain e - to form negative ions to get to 8 valence e - yName is changed to root + -ide A. Vocabulary Gains an e- Chloride ion

7. A. Vocabulary ION Polyatomic Ion Monatomic Ion 1 atom 2 or more atoms NO 3 - Na +

8. Monatomic Ions zIons formed from a single atom of an element zMost main group metals form one type of monatomic ion yGroup 1 metals: +1 ions (Na +, Li + …) yGroup 2 Metals: +2 ions (Mg 2+, Ca 2+ …) yGroup 13 Metals: +3 ions (Al 3+, Ga 3+ …)

9. Monatomic Ions zMain group nonmetals form one type of anion yGroup 15: -3 ions (N 3-, P 3- ) yGroup 16: -2 ions (O 2-, S 2- …) yGroup 17: -1 ions (F -, Cl - …)

10. Monatomic Ions zSome metals, mainly transition metals, can form more than one type of ion yIron: Fe 2+ or Fe 3+ yCopper: Cu + or Cu 2+ yLead: Pb 2+ or Pb 4+

11. D. Common Ions

12. B. Formula Unit

13. Chemical Formulas zChemical formulas yidentify the elements present in a compound using element symbols ythe number of atoms of each element is indicated with numbers in subscript Fe 2 O 3 2 Iron atoms 3 Oxygen atoms

14. Binary Ionic Compounds zBinary ionic compounds are ionic compounds formed between two elements. zAlways formed between a metal (positive ion) and nonmetal (negative ion) zMetal is written first in formula and then nonmetal

15. D. Ionic Formulas zWriting Ionic Formulas: zCalcium chloride zCa 2+ Cl 1- zcharges do not cancel, must criss-cross charges zRewrite as complete formula without charges zCaCl 2 2 1

16.  K + F   Mg 2+ N   Ba 2+ Cl   KF  Mg 3 N 2  BaCl 2 D. Ionic Formulas

17.  Ca 2+ O 2    Al 3+ S    Mg 2+ Br   D. Ionic Formulas Al 2 S 3 CaO MgBr 2

18. C. Lewis Structures zIonic – show transfer of electrons

19. G. Ionic Nomenclature Naming Binary Ionic Compounds zWrite names of both ions, cation (metal) first zChange ending of monatomic anions (nonmetal) to -ide zUse Roman numerals to show the ion’s charge if more than one is possible

20. G. Ionic Names with Type II Cations zYou must write the charge in parentheses using Roman numerals. To determine charge know that overall charge of compound = 0 zCr 2 O 3 zCrO O: 3 x -2 = - 6 Cr: 2 x ___= +6 +3 Chromium (III) oxide O: 1 x -2 = - 2 Cr: 1 x ___= +2 +2 Chromium (II) oxide Formula: Element: # atoms x charge = total charge

21. Name the following Compounds zCaCl 2 zAl 2 O 3 zNa 3 N zKI Calcium Chloride Aluminum Oxide sodium nitride Potassium Iodide

22. Name the following cont. zFeN zCu 2 S zZnCl 2 zV3P5zV3P5 zMnO 3 Iron (III) nitride Copper (I) sulfide Zinc chloride Vanadium (V) phosphide Manganese (VI) oxide

23. Write the formula for zStrontium nitride zLithium sulfide zGallium bromide zBarium oxide Sr 3 N 2 Li 2 S GaBr 3 BaO

24. Write the formulas zTin (IV) oxide zGold (III) Chloride zChromium (II) Nitride zMercury (I) sulfide SnO 2 AuCl 3 Cr 3 N 2 (Hg 2 ) 2 S

25. zCopper (II) bromide zTin (IV) oxide zManganese (II) chloride yCu 2+ + Br   ySn 4+ + O 2   yMn 2+ + Cl   F. Ionic Formulas with Type II Cations CuBr 2 Sn 2 O 4  SnO 2 MnCl 2

26. E. Polyatomic Ions zPolyatomic Ions yIons made of more than one atom yActs as an individual ion and its charge applies to the entire group of atoms yNEVER change the subscripts – add parentheses and subscripts outside, if necessary yListed on the back of your periodic table

27. zpotassium chlorate zmagnesium nitrate zammonium phosphate  K + ClO 3   Mg 2+ NO 3   NH 4 + PO 4   KClO 3  Mg(NO 3 ) 2  (NH 4 ) 3 PO 4 E. Ionic Formulas with PA Ions

28. zcalcium oxalate zaluminum perchlorate zstrontium phosphate  Ca 2+ C 2 O 4 2    Al 3+ ClO 4    Sr 2+ PO 4   E. Ionic Formulas with PA Ions CaC 2 O 4 Al(ClO 4 ) 3 Sr 3 (PO 4 ) 2

29. Nomenclature with PA zWhen naming compounds with polyatomic ions use the polyatomic ions name and follow all other rules for ionic compounds zCa 3 (PO 4 ) 2  calcium phosphate zFe(NO 3 ) 3  iron (III) nitrate

30. zCaBr zNa 2 CO 3 zNH 4 OH ycalcium bromide ysodium carbonate yammonium hydroxide G. Ionic Nomenclature

31. zCr 2 (SO 4 ) 3 zCu(NO 3 ) 2 zFeCl 3 yChromium (III) sulfate yCopper (II) nitrate yiron(III) chloride G. Ionic Names with Transition Metals

32. IIIIIIIV Ch. 9 – Chemical Names and Formulas II. Covalent Bond Formation Covalent Compound Names & Formulas

33. A. What is a covalent bond? zA chemical bond that results from the sharing of electrons zMolecule = two or more atoms that are held together by covalent bonds zMajority of covalent bonds form between nonmetals (CLOSE together on periodic table) H2OH2O

34. B. Examples: zWhich of the following are covalent compounds? yNaBr ySiO 2 yCO 2 yAlCl 3 yCH 4

35. C. Covalent Bonding Formation zDiatomic molecule ymolecule containing only two atoms zSome elements always exist this way because they are more stable than the individual atoms Cl 2

36. NOF Cl Br I H D. Diatomic Elements zThe Seven Diatomic Elements Br 2 I 2 N 2 Cl 2 H 2 O 2 F 2

37. E. Molecular Nomenclature zPrefix System (binary molecules) 1.Add prefixes to indicate # of atoms. Omit mono- prefix on first element. 2.Change the ending of the second element to -ide. 3.Second element ALWAYS gets a prefix.

38. PREFIX mono- di- tri- tetra- penta- hexa- hepta- octa- nona- deca- NUMBER 1 2 3 4 5 6 7 8 9 10 E. Molecular Nomenclature

39. zCCl 4 zN 2 O zSF 6 ycarbon tetrachloride ydinitrogen monoxide ysulfur hexafluoride E. Molecular Nomenclature

40. zarsenic trichloride zdinitrogen pentoxide ztetraphosphorus decoxide yAsCl 3 yN2O5yN2O5 yP 4 O 10 E. Molecular Nomenclature

41. Covalent Bonding – Common Names Diatomic Molecules E. Molecular Nomenclature NH 3 CH 4

42. Lewis Structures zThe Lewis Structures of Covalent compound represents elements in a compound, their valence electrons and how they are shared, and how the elements orient themselves around each other.

43. A. Drawing Lewis Structures 1.Count ALL Valence electrons on all atoms in the molecule. zFor an anion ion, add one electron for each negative charge. zFor a cation, subtract one electron for each positive charge.

44. A. Drawing Lewis Structures 2. The atom with the least amount is central atom and place the other atoms around the central atom. zHydrogen is never central atom zDraw a line connecting the peripheral atoms to the central atom zEach line represent 2 electrons zCheck for octet

45. A. Drawing Lewis Structures 3.Place pairs of valence electrons around each peripheral atom, except hydrogen, until octet is reached. 4.If any electrons remain place around the central atom until octet is reached. 5.If central atom still does not have an octet, use a lone pair of electrons on a neighboring atom to form a multiple bond to the central atom.

46. Examples zCF 4 zCO 2 z HCN zClO - zNH 4 +

47. B. Drawing Lewis Diagrams  CF 4 1 C × 4e - = 4e - 4 F × 7e - = 28e - 32e - F F C F F 16 pairs of e - - 4 pairs of e - 12 pairs of e - + 2

48. A. Drawing Lewis Diagrams  CO 2 1 C × 4e - = 4e - 2 O × 6e - = 12e - 16e - O C O 6 pairs of e - + 2 8 pairs of e - -2 pairs of e -

49. B. Polyatomic Ions  To find total # of valence e - : yAdd 1e - for each negative charge ySubtract 1e - for each positive charge  Place brackets around the ion and label the charge

50. B. Polyatomic Ions  ClO 4 - 1 Cl × 7e - = 7e - 4 O × 6e - = 24e - 31e - O O Cl O O + 1e - 32e - 16 e - pairs - 4 e - pairs 12 e - pairs 2 = +

51.  NH 4 + 1 N × 5e - = 5e - 4 H × 1e - = 4e - 9e - H H N H H - 1e - 8e - 4 pairs of e - -4 pairs of e - 0 pairs of e - B. Polyatomic Ions 2 = +

52. C. Resonance Structures  Molecules that can’t be correctly represented by a single Lewis diagram  Actual structure is an average of all the possibilities  Show all possible structures separated by double-headed arrows

53. C. Resonance Structures O O S O O O S O O O S O  SO 3

54. A. Octet Rule  Remember… yMost atoms form bonds in order to have 8 valence electrons

55. yHydrogen  2 valence e - yGroups 1,2,3 get 2,4,6 valence e - yExpanded octet  more than 8 valence e - (e.g. S, P, Xe)  Exceptions: D. Octet Rule F B F F H O HN O Very unstable!! F F S F F

56. E. Drawing Lewis Diagrams  BeCl 2 1 Be × 2e - = 2e - 2 Cl × 7e - = 14e - 16e - Cl Be Cl 8 pairs of e - -2 pairs of e - 6 pairs of e - 2 +

57. Expanded Octet zSome atoms may have more than 8 electrons around them. Follow the rules for normal Lewis Diagrams zNonmetals and metalloids in rows 3 or higher can have expanded octets

58. Examples zSF 6 zPCl 5 zXeF 4 zICl 3

59. E. Drawing Lewis Diagrams  SF 6 1S× 6e - = 6e - 6F× 7e - = 42e - 48e - F F S F F 24 pairs of e - - 6 pairs of e - 18 pairs of e - 2 +

60. Octet Deficient Elements zBeryllium and boron and do not need an octet zBeryllium (Be) covalently bonds and only needs 4 electrons around it zBoron (B) needs 6 around it

61. zBCl 3 zBeF 2