1. Nomenclature

2. Chemical Nomenclature

3. Find the bonnet. Find the boot. 3

4. Nomenclature  A System of Naming Compounds  Compounds are two or more atoms of different elements bonded together. 4

5. IUPAC Nomenclature 5

6.  Geneva, Switzerland 1892  International Union of Pure and Applied Chemistry  IUPAC System of Nomenclature 6

7. Binary compounds  Binary Compounds  Compounds composed of two elements  Two types  Ionic  Covalent 7

8. “Perhaps one of you gentlemen would mind telling me just what is outside the window that you find so attractive..?” Image courtesy NearingZero.net 8

9. -- 1+ 2+3+1-2-3- Generally metals form cations and non-metals form anions.

10. Ions…  Electrons are arranged on levels or “shells”. Atoms are most stable with 8 electrons on their outermost shell. This is often referred to as the octet rule.  Number of electrons on the outer shell = Group number for elements in Groups I-VIII A (using the US convention labeling on periodic table) 10

11. The modern periodic table. US Convention IUPAC Convention 11

12. Criss-Cross Rule 12

13. Example: Aluminum Chloride 13 Step 1: Step 2: Step 3: 13 Step 4: AlCl 3 Criss-Cross Rule Al Cl 3+ 1- write out name with space write symbols & charge of elements criss-cross charges as subsrcipts combine as formula unit (“1” is never shown) Aluminum Chloride

14. Example: Aluminum Chloride 14 Step 1: AluminumChloride Step 2: Al 3+ Cl 1- Step 3: Al Cl 13 Step 4: AlCl 3 Criss-Cross Rule

15. Example: Aluminum Oxide 15 Step 1: Aluminum Oxide Step 2: Al 3+ O 2- Step 3: Al O 23 Step 4: Al 2 O 3 Criss-Cross Rule

16. Example: Magnesium Oxide 16 Step 1: Magnesium Oxide Step 2: Mg 2+ O 2- Step 3: Mg O 22 Step 4: Mg 2 O 2 Step 5: MgO Criss-Cross Rule

17. Chemical Nomenclature  System of rules for naming pure substances  Elements – element name used even if the substance is di- or polyatomic O 2 = oxygen S 8 = sulfur  Compounds - naming differs depending on whether a substance is held together primarily by ionic or covalent bonds. 17

18. Ions…  O Group 6 6 electrons on outer shell O will gain 2 e -.  Now #e - = 10 but #p + = 8 so the oxygen is an ion with the formula O 2-.  P Group 5 5 electrons on outer shell P will gain 3 e -  Now #e - = 18 but #p + = 15 so the phosphorus is an ion with the formula P 3- 18

19. Ions…  To achieve the octet, atoms with more than 4 electrons on the outer shell will gain enough electrons to reach 8.  Cl Group 7 7 electrons on outer shell Cl will gain 1 e -.  Now #e - = 18 but #p + = 17 so the chlorine is an ion with the formula Cl 1-. 19

20. Ions…  Na Group 1 1 electron on outer shell O will lose 1 e - to expose the complete inside shell.  Now #e - = 10 but #p + = 11 so the sodium is an ion with the formula Na 1+.  Mg Group 2 2 electrons on outer shell Mg will lose 2 e -  Now #e - = 10 but #p + = 12 so the magnesium is an ion with the formula Mg 2+ 20

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22. Naming binary ionic compounds The name of the cation is the same as the name of the metal. Many metal names end in -ium. The name of the anion (negatively charged ion) takes the root of the nonmetal name and adds the suffix -ide. Calcium and oxygen form calcium oxide. Aluminum and sulfur form aluminum sulfide. The name of the cation (positively charged ion) is written first, followed by that of the anion. 22

23. Ionic Compounds  Ionic Compounds  often a metal + nonmetal  anion (nonmetal), add “ide” to element name  Rules for naming Ionic Compounds All compound must be neutral UNLESS indicated otherwise. The cation is always named first and the anion second The cation takes the name of the element The anion is named by the root of the element name and adding “ide”. Chlorine becomes Chloride 23

24. Naming Ionic Binary Compounds 24

25. Naming Binary Compounds 25 FormulaName 1 BaO____________________ 2________________ sodium bromide 3 MgI 2 ____________________ 4 KCl____________________ 5________________ strontium fluoride 6________________ cesium fluoride barium oxide NaBr magnesium iodide potassium chloride SrF 2 CsF

26. Binary Ionic Compounds  Cations (positively charged ions) and anions (negatively charged ions) will associate with each other and form a neutral binary compound to reduce energy.  1 Na + 1 Cl - NaCl  1 Ca 2+ 1 O 2- CaO  3 K + 1 N 3- K 3 N  2 Al 3+ 3 S 2- Al 2 S 3 26

27. Examples 27 BaCl 2 barium chloride K2OK2O potassium oxide Mg(OH) 2 magnesium hydroxide KNO 3 potassium nitrate

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29. Lewis Dot structures 29

30. Keeping Track of Electrons  The electrons responsible for the chemical properties of atoms are those in the outer energy level.  Valence electrons - electrons in the outer energy level.  Core electrons -those in the energy levels below. 30

31. Keeping Track of Electrons  Atoms in the same column  Have the same outer electron configuration.  Have the same valence electrons.  Easily found by looking up the group number on the periodic table.  Group 2A - Be, Mg, Ca, etc.-  2 valence electrons 31

32. Lewis Structures 32 1) Count up total number of valence electrons 2) Connect all atoms with single bonds - “multiple” atoms usually on outside - “single” atoms usually in center; C always in center, H always on outside. 3) Complete octets on exterior atoms (not H, though) 4) Check - all atoms (except H) have an octet; if not, try multiple bonds - any extra electrons?Put on central atom

33. Lewis Structures 1) Write the element symbol. 2) Carbon is in the 4 th group, so it has 4 valence electrons. 3) Starting at the right, draw 4 electrons, or dots, counter- clockwise around the element symbol. 33

34. Lewis Structures 1) Check your work. 2) Using your periodic table, check that Carbon is in the 4 th group. 3) You should have 4 total electrons, or dots, drawn in for Carbon. 34

35. Covalent bonding FFluorine has seven valence electrons AA second F atom also has seven BBy sharing electrons BBoth end with full orbitals (stable octets) 35 FF 8 Valence electrons

36. Water Each hydrogen has 1 valence electron Each hydrogen wants 1 more The oxygen has 6 valence electrons The oxygen wants 2 more They share to make each other happy 36 H O

37. Water Put the pieces together The first hydrogen is happy The oxygen still wants one more 37 H O

38. Water  The second hydrogen attaches  Every atom has full energy levels  A pair of electrons is a single bond 38 HO H H HO

39. Carbon dioxide  CO 2 - Carbon is central atom  Carbon has 4 valence electrons  Wants 4 more  Oxygen has 6 valence electrons  Wants 2 more 39 OC

40. Carbon dioxide  Attaching 1 oxygen leaves the oxygen 1 short and the carbon 3 short 40 O C

41. Carbon dioxide l Attaching the second oxygen leaves both oxygen 1 short and the carbon 2 short O C O 41

42. l The only solution is to share more l Requires two double bonds l Each atom gets to count all the atoms in the bond 8 valence electrons Carbon dioxide O CO 42

43. Multiple Bonds  Double and triple bonds can form between atoms in order to fill the outer energy level  This occurs when two atoms share more than one pair of electrons

44. Lewis Structures try these elements on your own: a) H b) P c) Ca d) Ar e) Cl f) Al 44

45. Lewis Structures try these elements on your own: a) H b) P c) Ca d) Ar e) Cl f) Al 45

46. Lewis Structures try these elements on your own: a) H b) P c) Ca d) Ar e) Cl f) Al 46

47. Lewis Structures try these elements on your own: a) H b) P c) Ca d) Ar e) Cl f) Al 47

48. Lewis Structures try these elements on your own: a) H b) P c) Ca d) Ar e) Cl f) Al 48

49. Lewis Structures On your worksheet, try these elements on your own: a) H b) P c) Ca d) Ar e) Cl f) Al

50. Lewis Structures On your worksheet, try these elements on your own: a) H b) P c) Ca d) Ar e) Cl f) Al

51. Covalent Naming Nonmetals hold onto their valence electrons. They can’t give away electrons to bond. Get it by sharing valence electrons with each other. Different from an ionic bond because they actually form molecules. Called a binary molecular compound The first vowel is often dropped to avoid the combination of “ao” or “oo”. CO = carbon monoxide (monooxide

52. Greek Prefixes for Two Nonmetals Number Indicated Prefixes 1mono- 2di- 3tri- 4tetra- 5penta- 6hexa- 7hepta- 8octa- 9nona- 10deca-

53. Binary Molecular Compounds N 2 O dinitrogen monoxide N 2 O 3 dinitrogen trioxide N 2 O 5 dinitrogen pentoxide ICl iodine monochloride ICl 3 iodine trichloride SO 2 sulfur dioxide SO 3 sulfur trioxide

54. Binary Molecular Compounds 54 N 2 Odinitrogen monoxide N 2 O 3 dinitrogen trioxide N 2 O 5 dinitrogen pentoxide ICliodine monochloride ICl 3 iodine trichloride SO 2 sulfur dioxide SO 3 sulfur trioxide

55. Molecular Compound 55 Silicon dioxide, SiO 2, is a molecular compound. It is also a mineral called quartz (left). Quartz is found in nearly every type of rock. Most sand grains (center) are bits of quartz. Glass is made from sand. A compound containing atoms of two or more elements that are bonded together by sharing electrons.

56. Chemical Bonding 56 …atoms share electrons to get a full valence shell Covalent Bonds C 2s 2 2p 2 (4 v.e – ) F 2s 2 2p 5 (7 v.e – ) both need 8 valence electrons for a full outer shell (octet rule) 1s 2

57. 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. 57

58. Covalent Bonding  Bonds between atoms are formed through the sharing of electrons  Covalent bonds form between two non-metal atoms through sharing of pairs of electrons  Atoms have a “desire” to have their outer energy levels filled (Octet Rule)  Covalent bonding can be represented with Lewis Dot Diagrams 58

59. Lewis Dot Diagrams  Lewis Dot Diagrams show the sharing of electrons between atoms and where the bonds form  atoms share electrons to fill their outer energy levels (8 electrons in their outer shell)  The exception is hydrogen (2 electrons in its outer shell) 59

60. Lewis Dot Diagrams for Hydrogen and Chlorine Gas The first row shows the atoms before they are bonded The second row shows the sharing of electrons to fill the outer energy level The third row has circles around the electrons to show those that belong to each atom. Where the circles overlap represents a covalent bond 60

61. Transition metal Ionic compounds  Transition metal ionic compounds  indicate charge on metal with Roman numerals FeCl 2 2 Cl - -2 so Fe is +2 iron(II) chloride FeCl 3 3 Cl - -3 so Fe is +3 iron(III) chloride Cr 2 S 3 3 S -2 -6 so Cr is +3 (6/2)chromium(III) sulfide

62. Common Monatomic Cations and Anions

63. Common Type II Cations

64. Ternary Compounds Ternary compounds are those containing three different elements. (NaNO 3, NH 4 Cl, etc.). The naming of ternary compounds involves the memorization of several positive and negative polyatomic ions, (two or more atoms per ion), and adding these names to the element with which they combine. i.e., Sodium ion, Na 1+ added to the nitrate ion, NO 3 1-, to give the compound, NaNO 3, sodium nitrate. Binary rules for indicating the oxidation number of metals and for indicating the numbers of atoms present are followed. The polyatomic ions are listed in a separate handout.

65. phosphate sulfate carbonate chlorate nitrate Polyatomic Ions - Memorize phosphATE sulfATE carbonATE chlorATE nitrATE PO 4 3- …………… SO 4 2- …………… CO 3 2- ………….. ClO 3 1- ………….. NO 3 1- ………..…. Eight “-ATE’s” Exceptions: ammonium hydroxide cyanide NH 4 1+ …………… OH 1- …………… CN 1- …………..

66. Ternary Compounds NaNO 2 KClO 3 Ca 3 (PO 4 ) 2 Fe(OH) 3 NaHCO 3 sodium nitrite potassium chlorate calcium phosphate iron (III) hydroxide sodium bicarbonate sodium hydrogen carbonate

67. Common Polyatomic Ions Names of Common Polyatomic Ions Ion Name Ion Name NH 4 1+ ammoniumCO 3 2- carbonate NO 2 1- nitriteHCO 3 1- hydrogen carbonate NO 3 1- nitrate (“bicarbonate” is a widely SO 3 2- sulfite used common name) SO 4 2- sulfateClO 1- hypochlorite HSO 4 1- hydrogen sulfateClO 2 1- chlorite (“bisulfate” is a widelyClO 3 1- chlorate used common name)ClO 4 1- perchlorate OH 1- hydroxideC 2 H 3 O 2 2- acetate CN 1- cyanideMnO 4 1- permanganate PO 4 3- phosphateCr 2 O 7 2- dichromate HPO 4 2- hydrogen phosphateCrO 4 2- chromate H 2 PO 4 1- dihydrogen phosphateO 2 2- peroxide Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 100 Print Version

68. Common Polyatomic Ions

69. Rules for Parentheses Parentheses are used only when the following two condition are met: 1.There is a radical (polyatomic ion) present and… 2.There are two or more of that radical in the formula. Examples: NaNO 3 NO 3 1- is a radical, but there is only one of it. Co(NO 3 ) 2 NO 3 1- is a radical and there are two of them (NH 4 ) 2 SO 4 NH 4 1+ is a radical and there are two of them; SO 4 2- is a radical but there is only one of it. Co(OH) 2 OH 1- is a radical and there are two of it. Al 2 (CO 3 ) 3 CO 3 2- is a radical and there are three of them. NaOHOH 1- is a radical but there is only one of it.

70. Vocabulary CHEMICAL FORMULA molecular formula unit IONICCOVALENT CO 2 NaCl

71. Vocabulary COMPOUND ternary compound binary compound 2 elements more than 2 elements NaNO 3 NaCl

72. Vocabulary ION polyatomic Ion monatomic Ion 1 atom 2 or more atoms NO 3 - Na + Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

73. Naming Simple Chemical Compounds Ionic (metal and nonmetal)Covalent (2 nonmetals) Metal Forms only one positive ion Forms more than one positive ion Nonmetal Use the name of element Use element name followed by a Roman numeral to show the charge First nonmetal Second nonmetal Before element name use a prefix to match subscript Use a prefix before element name and end with ide Single Negative Ion Polyatomic Ion Use the name of the element, but end with ide Use the name of polyatomic ion (ate or Ite)