1. “Chemical Names and Formulas” Mr. HUYNH

2. Naming Ions l OBJECTIVES: –Identify the charges on monatomic ions by using the periodic table, and name the ions.

3. Naming Ions l OBJECTIVES: –Define a polyatomic ion and write the names and formulas of the most common polyatomic ions.

4. Naming Ions l OBJECTIVES: –Identify the two common endings for the names of most polyatomic ions.

5. Atoms and Ions l Atoms are electrically neutral. –Because there is the same number of protons (+) and electrons (-). l Ions are atoms, or groups of atoms, with a charge (positive or negative) –They have different numbers of protons and electrons. l Only electrons can move, and ions are made by gaining or losing electrons.

6. An Anion is… l A negative ion. l Has gained electrons. l Nonmetals can gain electrons. l Charge is written as a superscript on the right. F 1- Has gained one electron (-ide is new ending = fluoride) O 2- Gained two electrons (oxide)

7. A Cation is… l A positive ion. l Formed by losing electrons. l More protons than electrons. l Metals can lose electrons K 1+ Has lost one electron (no name change for positive ions) Ca 2+ Has lost two electrons

8. Predicting Ionic Charges Group 1A: Lose 1 electron to form 1+ ions H1+H1+H1+H1+ Li 1+ Na 1+ K1+K1+K1+K1+ Rb 1+

9. Predicting Ionic Charges Group 2A: Loses 2 electrons to form 2+ ions Be 2+ Mg 2+ Ca 2+ Sr 2+ Ba 2+

10. Predicting Ionic Charges Group 3A: Loses 3 Loses 3 electrons to form 3+ ions B 3+ Al 3+ Ga 3+

11. Predicting Ionic Charges Group 4A: Do they lose 4 electrons or gain 4 electrons? Do they lose 4 electrons or gain 4 electrons? Neither! Group 4A elements rarely form ions (they tend to share)

12. Predicting Ionic Charges Group 5A: Gains 3 Gains 3 electrons to form 3- ions N 3- P 3- As 3- Nitride Phosphide Arsenide

13. Predicting Ionic Charges Group 6A: Gains 2 Gains 2 electrons to form 2- ions O 2- S 2- Se 2- Oxide Sulfide Selenide

14. Predicting Ionic Charges Group 7A: Gains 1 electron to form Gains 1 electron to form 1- ions F 1- Cl 1- Br 1- Fluoride Chloride Bromide I 1- Iodide

15. Predicting Ionic Charges Group 8A: Stable noble gases do not form ions! Stable noble gases do not form ions!

16. Predicting Ionic Charges Group B elements: Many transition elements Many transition elements have more than one possible oxidation state. have more than one possible oxidation state. Iron (II) = Fe 2+ Iron (III) = Fe 3+ Note the use of Roman numerals to show charges

17. Naming cations l Two methods can clarify when more than one charge is possible: 1)Stock system – uses roman numerals in parenthesis to indicate the numerical value 2)Classical method – uses root word with suffixes (-ous, -ic) Does not give true value

18. Naming cations l We will use the Stock system. l Cation - if the charge is always the same (like in the Group A metals) just write the name of the metal. l Transition metals can have more than one type of charge. –Indicate their charge as a roman numeral in parenthesis after the name of the metal.

19. Predicting Ionic Charges Some of the post-transition elements also Some of the post-transition elements also have more than one possible oxidation state. have more than one possible oxidation state. Tin (II) = Sn 2+ Lead (II) = Pb 2+ Tin (IV) = Sn 4+ Lead (IV) = Pb 4+

20. Predicting Ionic Charges Group B elements: Some transition elements Some transition elements have only one possible oxidation state, such as these three: have only one possible oxidation state, such as these three: Zinc = Zn 2+ Silver = Ag 1+ Cadmium = Cd 2+

21. Exceptions: l Some of the transition metals have only one ionic charge: –Do not need to use roman numerals for these: –Silver is always 1+ (Ag 1+ ) –Cadmium and Zinc are always 2+ (Cd 2+ and Zn 2+ )

22. Practice by naming these: l Na 1+ l Ca 2+ l Al 3+ l Fe 3+ l Fe 2+ l Pb 2+ l Li 1+

23. Write symbols for these: l Potassium ion l Magnesium ion l Copper (II) ion l Chromium (VI) ion l Barium ion l Mercury (II) ion

24. Naming Anions l Anions are always the same charge l Change the monatomic element ending to – ide l F 1- a Fluorine atom will become a Fluoride ion.

25. Practice by naming these: l Cl 1- l N 3- l Br 1- l O 2- l Ga 3+

26. Write symbols for these: l Sulfide ion l Iodide ion l Phosphide ion l Strontium ion

27. Polyatomic ions are… l Groups of atoms that stay together and have an overall charge, and one name. l Usually end in –ate or -ite l Acetate: C 2 H 3 O 2 1- l Nitrate: NO 3 1- l Nitrite: NO 2 1- l Permanganate: MnO 4 1- l Hydroxide: OH 1- and Cyanide: CN 1- ?

28. l Sulfate: SO 4 2- l Sulfite: SO 3 2- l Carbonate: CO 3 2- l Chromate: CrO 4 2- l Dichromate: Cr 2 O 7 2- l Phosphate: PO 4 3- l Phosphite: PO 3 3- l Ammonium: NH 4 1+ If the polyatomic ion begins with H, then combine the word hydrogen with the other polyatomic ion present: H 1+ + CO 3 2- → HCO 3 1- hydrogen + carbonate → hydrogen carbonate ion (One of the few positive polyatomic ions) Polyatomic ions are…

29. Naming and Writing Formulas for Ionic Compounds l OBJECTIVES: –Apply the rules for naming and writing formulas for binary ionic compounds.

30. Naming and Writing Formulas for Ionic Compounds l OBJECTIVES: –Apply the rules for naming and writing formulas for compounds containing polyatomic ions.

31. Writing Ionic Compound Formulas Example: Barium nitrate (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! Ba 2+ NO 3 - 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss-cross method to balance subscripts. Not balanced! ( ) 2 Now balanced. = Ba(NO 3 ) 2

32. Writing Ionic Compound Formulas Example: Ammonium sulfate (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! NH 4 + SO 4 2- 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss-cross method to balance the subscripts. Not balanced! ( ) 2 Now balanced. = (NH 4 ) 2 SO 4

33. Writing Ionic Compound Formulas Example: Iron (III) chloride (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! Fe 3+ Cl - 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss-cross method to balance the subscripts. Not balanced! 3 Now balanced. = FeCl 3

34. Writing Ionic Compound Formulas Example: Aluminum sulfide (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! Al 3+ S 2- 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss-cross method to balance the subscripts. Not balanced! 23 Now balanced. = Al 2 S 3

35. Writing Ionic Compound Formulas Example: Magnesium carbonate (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! Mg 2+ CO 3 2- 2. Check to see if charges are balanced. They are balanced! = MgCO 3

36. Writing Ionic Compound Formulas Example: Zinc hydroxide (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! Zn 2+ OH - 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss-cross method to balance the subscripts. Not balanced! ( ) 2 Now balanced. = Zn(OH) 2

37. Writing Ionic Compound Formulas Example: Aluminum phosphate (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! Al 3+ PO 4 3- 2. Check to see if charges are balanced. They ARE balanced! = AlPO 4

38. Naming Ionic Compounds l 1. Name the cation first, then anion l 2. Monatomic cation = name of the element Ca 2+ = calcium ion l 3. Monatomic anion = root + -ide Cl  = chloride CaCl 2 = calcium chloride

39. Naming Ionic Compounds l some metals can form more than one charge (usually the transition metals) l use a Roman numeral in their name: PbCl 2 – use the anion to find the charge on the cation (chloride is always 1-) Pb 2+ is the lead (II) cation PbCl 2 = lead (II) chloride (Metals with multiple oxidation states)

40. Things to look for: 1) If cations have ( ), the number in parenthesis is their charge. 2) If anions end in -ide they are probably off the periodic table (Monoatomic) 3) If anion ends in -ate or –ite, then it is polyatomic

41. Practice by writing the formula or name as required… l Iron (II) Phosphate l Stannous Fluoride l Potassium Sulfide l Ammonium Chromate l MgSO 4 l FeCl 3

42. Naming and Writing Formulas for Molecular Compounds l OBJECTIVES: –Interpret the prefixes in the names of molecular compounds in terms of their chemical formulas.

43. Naming and Writing Formulas for Molecular Compounds l OBJECTIVES: –Apply the rules for naming and writing formulas for binary molecular compounds.

44. Molecular compounds are… l made of just nonmetals l smallest piece is a molecule l can’t be held together by opposite charge attraction l can’t use charges to figure out how many of each atom (there are no charges present)

45. Molecular compounds are easier! l Ionic compounds use charges to determine how many of each. –You have to figure out charges. –May need to criss-cross numbers. l Molecular compounds: the name tells you the number of atoms. –Uses prefixes to tell you the exact number of each element present!

46. Prefixes l 1 = mono- l 2 = di- l 3 = tri- l 4 = tetra- l 5 = penta- l 6 = hexa- l 7 = hepta- l 8 = octa-

47. Prefixes l 9 = nona- l 10 = deca- l To write the name, write two words: l One exception is we don’t write mono if there is only one of the first element. l Normally, we do not have double vowels when writing names (oa oo) PrefixnamePrefixname-ide

48. Practice by naming these: lN2OlN2O l NO 2 l Cl 2 O 7 l CBr 4 l CO 2 l BaCl 2 (This one will not use prefixes, since it is an ionic compound!) = dinitrogen monoxide (also called nitrous oxide or laughing gas) = nitrogen dioxide = dichlorine heptoxide = carbon tetrabromide = carbon dioxide

49. Write formulas for these: l diphosphorus pentoxide l tetraiodine nonoxide l sulfur hexafluoride l nitrogen trioxide l carbon tetrahydride l phosphorus trifluoride l aluminum chloride (Ionic compound)

50. Balancing Equations l Law of Conservation of Matter: –In a chemical reaction, matter can be neither created nor destroyed. –In a chemical reaction, the amount of reactants equal the amount of products. –

51. Balancing Equations l Paraphrase: l Law of Conservation of Atoms: –The number of atoms of each type of element must be the same on each side of the equation.

52. Balancing Equations l Hydrogen and oxygen are diatomic elements. l Their subscripts cannot be changed. l The subscripts on water cannot be changed. Hydrogen + oxygen water H 2 + O 2 H 2 O

53. Balancing Equation l Count the atoms on each side. –Reactant side: 2 atoms H and 2 atoms O –Product side: 2 atoms H and 1 atom O H 2 + O 2 H 2 O

54. Balancing Equations l H 2 + O 2 H 2 O l If the subscripts cannot be altered, how can the atoms be made equal? l Adjust the number of molecules by changing the coefficients.

55. Balancing Equations l Reactants: 2 atoms of H and 2 atoms of O l Products: 4 atoms of H and 2 atoms of O l H is no longer balanced! H 2 + O 2 2H 2 O

56. Balancing Equations l Reactant side: 4 atoms of H and 2 atoms of O l Product side: 4 atoms of H and 2 atoms of O l It’s Balanced! 2H 2 + O 2 2H 2 O

57. Balancing Equations l Count atoms. –Reactants: 2 atoms N and 2 atoms H –Products: 1 atom N and 3 atoms of NH 3 N 2 + H 2 NH 3 Nitrogen + hydrogen ammonia

58. Balancing Equations l Nothing is balanced. l Balance the nitrogen first by placing a coefficient of 2 in front of the NH 3. N 2 + H 2 2NH 3

59. Balancing Equations l Hydrogen is not balanced. l Place a 3 in front of H 2. l Reactant side: 2 atoms N, 6 atoms H l Product side: 2 atoms N, 6 atoms H N 2 + 3H 2 2NH 3

60. Balancing Equations l Count atoms. l Reactants: Ca – 3 atoms, P – 2 atoms, O – 8 atoms; H – atoms, S – 1 atom, O – 4 atoms Ca 3 (PO 4 ) 2 + H 2 SO 4 CaSO 4 + H 3 PO 4

61. Balancing Equations l Side note on Ca 3 (PO 4 ) 2 l The subscript after the phosphate indicates two phosphate groups. l This means two PO 4 3- groups with two P and eight O atoms.

62. Balancing Equations l Ca 3 (PO 4 ) 2 + H 2 SO 4 CaSO 4 + H 3 PO 4 l Count atoms in the product. l Ca atoms – 1, S atom – 1, O atoms – 4; H atoms – 3, P atom – 1, O atoms - 4

63. Balancing Equations l In this equation, the ion groups do not break up. l Instead of counting individual atoms, ion groups may be counted. l Ca 3 (PO 4 ) 2 + H 2 SO 4 CaSO 4 + H 3 PO 4

64. Balancing Equations l Ca 3 (PO 4 ) 2 + H 2 SO 4 CaSO 4 + H 3 PO 4 l Reactants: Ca 2+ – 3, PO 4 3- - 2, H + – 2, SO 4 2+ - 1 l Products: Ca 2+ - 1, SO 4 2- - 1, H + - 3, PO 4 3- - 1

65. Balancing Equations l Balance the metal first by placing a coefficient of 3 in front of CaSO 4. l Products: Ca – 3 atoms, SO 4 2- - 3 groups l Ca 3 (PO 4 ) 2 + H 2 SO 4 3CaSO 4 + H 3 PO 4

66. Balancing Equations l Three sulfate groups are needed on the reactant side so place a coefficient of 3 in front of H 2 SO 4. l 3H 2 SO 4 gives 6 H + and 3 SO 4 2-. l Neither phosphate nor calcium is balanced. l Ca 3 (PO 4 ) 2 + 3H 2 SO 4 3CaSO 4 + H 3 PO 4

67. Balancing Equations l A coefficient of 2 placed in front of H 3 PO 4 which balances both hydrogen and phosphate. l Ca 3 (PO 4 ) 2 + 3H 2 SO 4 3CaSO 4 + 2H 3 PO 4

68. Balancing Equations l The sulfate group breaks up. Each atom must be counted individually. Ugh! l Reactants: Cu – 1, H – 2, S – 1, O – 4 l Products: Cu – 1, S – 1, O - 4, H – 2, O – 1, S – 1, O - 2 Cu + H 2 SO 4 CuSO 4 + H 2 O + SO 2

69. Balancing Equations l Sulfur is not balanced. l Place a two in front of sulfuric acid. l Count atoms: 2 H 2 SO 4 H – 4, S – 2, O - 8 l Cu + 2H 2 SO 4 l CuSO 4 + H 2 O + SO 2

70. Balancing Equations l Hydrogen needs to be balanced so place a 2 in front of the H 2 O. l Count the number of atoms. l Cu + 2H 2 SO 4 l CuSO 4 + 2H 2 O + SO 2

71. Balancing Equations l Reactants: Cu – 1, H – 4, S – 2, O – 8 l Products: Cu – 1, S – 1, O – 4, H – 4, O – 2, S – 1, O – 2 = Cu – 1, S – 2, H – 4, O – 8 l It’s balanced! l Cu + 2H 2 SO 4 l CuSO 4 + 2H 2 O + SO 2

72. Balancing Equations l Balancing hints: –Balance the metals first. –Balance the ion groups next. –Balance the other atoms. –Save the non ion group oxygen and hydrogen until the end.

73. Balancing Equations l This method of balancing equations is the inspection method. l The method is trial and error. l Practice.

74. Writing and Naming l Write the corresponding formula equation and then balance the equation. Nickel + hydrochloric acid Nickel(II) chloride + hydrogen

75. Writing and Naming l Write each formula independently. l Ignore the rest of the equation. l Balance the equation after writing the formulas. Ni + HCl NiCl 2 + H 2 Ni + 2HCl NiCl 2 + H 2

76. Writing and Naming l Remember the diatomic elements: H 2, N 2, O 2, F 2, Cl 2, Br 2, and I 2.

77. Writing and Naming l Balance the formula equation. l Write the word equation. Cu + H 2 SO 4 CuSO 4 + H 2 O + SO 2

78. Writing and Naming l Cu + 2H 2 SO 4 l CuSO 4 + 2H 2 O + SO 2 l Write the names: l Cu by itself is just copper. Copper(I) or copper(II) would be incorrect. l H 2 SO 4 should be named as an acid. l Sulfuric acid

79. Writing and Naming l CuSO 4 has a SO 4 2- group so Cu must be 2+. Some metals must have Roman Numerals. Copper(II) sulfate l H 2 O is known as water. l SO 2 is a nonmetal compound. Its name is either sulfur dioxide or sulfur(IV) oxide.

80. Writing and Naming l Copper + sulfuric acid  l Copper(II) sulfate + water + sulfur dioxide Cu + 2H 2 SO 4 CuSO 4 + 2H 2 O + SO 2