1. Happy Birthday Tommy Lee Jones (1946) Read for Thursday Read for Thursday Chapter 5: Sections 1-5 Chapter 5: Sections 1-5 HOMEWORK – DUE Thursday 9/17/15 HOMEWORK – DUE Thursday 9/17/15 HW-WS 6 (Worksheet) (from course website) HW-WS 6 (Worksheet) (from course website) HOMEWORK – DUE TUESDAY 9/23/15 HOMEWORK – DUE TUESDAY 9/23/15 HW-BW 4.2 (Bookwork) CH 4 #’s 78-81 all, 83-89 odd, 90, 93, 95, 104-112 even, 116, 156 HW-BW 4.2 (Bookwork) CH 4 #’s 78-81 all, 83-89 odd, 90, 93, 95, 104-112 even, 116, 156 HW-WS 7 (Worksheet) (from course website) HW-WS 7 (Worksheet) (from course website) Lab this week Lab this week Wednesday/Thursday – EXP 5 Wednesday/Thursday – EXP 5 PRE LAB!!! PRE LAB!!! Lab next week Lab next week Monday/Tuesday – EXP 5 continued Monday/Tuesday – EXP 5 continued Wednesday/Thursday – EXP 6 Wednesday/Thursday – EXP 6 PRE LAB!!! PRE LAB!!!

2. Electrolytes Electrolyte: Any solution that conducts electricity In order to conduct electricity, a solution must contain “mobile charged particles”, i.e. ions that can move around

3. Strong electrolytes Strong acids are strong electrolytes Soluble salts (ionic compounds that do dissolve in water) are strong electrolytes Compounds that break COMPLETELY apart into ions in water HCl (aq) HBr (aq) HI (aq) HClO 3(aq) HClO 4(aq) HNO 3(aq) ANY other acid is a weak acid H 2 SO 4(aq) Types of Electrolytes

4. Strong electrolytes Soluble salts (ionic compounds that do dissolve in water) are strong electrolytes Compounds that break COMPLETELY apart into ions in water Types of Electrolytes

5. Strong electrolytes Strong acids are strong electrolytes Compounds that break COMPLETELY apart into ions in water Types of Electrolytes

6. H+H+ NO 3 – H+H+ H+H+ H+H+ H+H+ H+H+ H+H+ H+H+ H+H+ H+H+ H+H+ H+H+ KBr K+K+ Br – K+K+ K+K+ K+K+ K+K+ K+K+ K+K+ K+K+ K+K+ K+K+ K+K+ K+K+ HNO 3 TONS of ions formed Conducts electricity very well STRONG electrolyte TONS of ions formed Conducts electricity very well STRONG electrolyte KBr HNO 3

7. Weak electrolytes Weak acids are weak electrolytes Compounds remain ALMOST entirely as molecules in water, forming very VERY few ions Insoluble salts (ionic compounds that do NOT dissolve in water) are weak electrolytes Types of Electrolytes

8. Weak electrolytes Compounds remain ALMOST entirely as molecules in water, forming very VERY few ions Insoluble salts (ionic compounds that do NOT dissolve in water) are weak electrolytes Types of Electrolytes

9. Weak electrolytes Weak acids are weak electrolytes Compounds remain ALMOST entirely as molecules in water, forming very VERY few ions Types of Electrolytes

10. Ba 2+ SO 4 2– Very FEW ions formed Conducts electricity, but barely WEAK electrolyte Very FEW ions formed Conducts electricity, but barely WEAK electrolyte BaSO 4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 H2C2O4H2C2O4 C 2 O 4 2– H+H+ H+H+

11. Non electrolytes All molecular compounds are non electrolytes Compounds remain as molecules in water, forming absolutely NO IONS Types of Electrolytes

12. Non electrolytes C 6 H 12 O 6 NO ions formed Cannot conduct electricity NON electrolyte Types of Electrolytes

13. STRONG electrolytes soluble salts strong acids Compounds that break COMPLETELY apart into ions in water WEAK electrolytes insoluble salts weak acids Compounds remain ALMOST entirely as molecules in water, forming very VERY few ions NON electrolytes molecular compounds Compounds remain as molecules in water, forming absolutely NO IONS

16. Solubility Rules All ammonium and group 1 metal salts are SOLUBLE (no exceptions) All ammonium and group 1 metal salts are SOLUBLE (no exceptions) All nitrate, chlorate, perchlorate, and acetate salts are SOLUBLE (no exceptions) All nitrate, chlorate, perchlorate, and acetate salts are SOLUBLE (no exceptions) Most chloride, bromide, and iodide salts are SOLUBLE. EXCEPTIONS! Compound containing silver, mercury (I), and lead (II) Most chloride, bromide, and iodide salts are SOLUBLE. EXCEPTIONS! Compound containing silver, mercury (I), and lead (II) Most fluoride salts are SOLUBLE. EXCEPTIONS! Compounds containing magnesium, calcium, strontium, barium, and lead (II) Most fluoride salts are SOLUBLE. EXCEPTIONS! Compounds containing magnesium, calcium, strontium, barium, and lead (II) Most sulfate salts are SOLUBLE. EXCEPTIONS! Compounds containing strontium, barium, mercury (I), and lead (II) Most sulfate salts are SOLUBLE. EXCEPTIONS! Compounds containing strontium, barium, mercury (I), and lead (II) Most sulfide and hydroxide salts are INSOLUBLE. EXCEPTIONS! Compounds containing calcium, strontium, and barium Most sulfide and hydroxide salts are INSOLUBLE. EXCEPTIONS! Compounds containing calcium, strontium, and barium Most carbonate, phosphate, oxalate, and chromate salts are INSOLUBLE. Most carbonate, phosphate, oxalate, and chromate salts are INSOLUBLE. THESE ONLY APPLY TO IONIC COMPOUNDS!!!!!

17. Na 2 C 2 O 4(aq) + CaCl 2(aq)  2 NaCl (aq) + CaC 2 O 4(s) 3 H 2 SO 4(aq) + 2 Fe(OH) 3(s)  Fe 2 (SO 4 ) 3(aq) + 6 H 2 O (l) 2 LiCl (aq) + Hg 2 (C 2 H 3 O 2 ) 2(aq)  2 LiC 2 H 3 O 2(aq) + Hg 2 Cl 2(s) Solubility and Reactions 3 H 2 SO 4(aq) + 2 Fe(OH) 3(s)  Fe 2 (SO 4 ) 3(aq) + 6 H 2 O (l) 2 LiCl (aq) + Hg 2 (C 2 H 3 O 2 ) 2(aq)  2 LiC 2 H 3 O 2(aq) + Hg 2 Cl 2(s) S-E W-E S-EW-ES-EN-E S-E W-E

18. 2 HF (aq) + Ba(OH) 2(aq)  BaF 2(s) + 2 H 2 O (l) K 2 CO 3(aq) + 2 HClO 4(aq)  2 KClO 4(aq) + H 2 O (l) + CO 2(g) 2 NaNO 3(aq) + Mg(ClO 3 ) 2(aq)  2 NaClO 3(aq) + Mg(NO 3 ) 2(aq) Solubility and Reactions 2 HF (aq) + Ba(OH) 2(aq)  BaF 2(s) + 2 H 2 O (l) K 2 CO 3(aq) + 2 HClO 4(aq)  2 KClO 4(aq) + H 2 O (l) + CO 2(g) 2 NaNO 3(aq) + Mg(ClO 3 ) 2(aq)  2 NaClO 3(aq) + Mg(NO 3 ) 2(aq) W-ES-EW-EN-E S-E N-E S-E

19. Step 1: Write a balanced chemical equation (molecular equation) A solution of chromium (III) sulfate is added to a solution of cesium phosphate. Give the net ionic equation for any reaction that happens. Cs 3 PO 4 Cr 2 (SO 4 ) 3(aq) + 2 Cs 3 PO 4(aq)  2 CrPO 4(s) + 3 Cs 2 SO 4(aq) Cr 2 (SO 4 ) 3 CrPO 4 Cr 2 (SO 4 ) 3(aq) + 2 Cs 3 PO 4(aq)  2 CrPO 4(s) + 3 Cs 2 SO 4(aq) Cs 2 SO 4  ++

20. Step 1: Write a balanced chemical equation (molecular equation) A solution of chromium (III) sulfate is added to a solution of cesium phosphate. Give the net ionic equation for any reaction that happens. Cr 2 (SO 4 ) 3(aq) + 2 Cs 3 PO 4(aq)  2 CrPO 4(s) + 3 Cs 2 SO 4(aq) Step 2: Break STRONG ELECTROLYTES ONLY into ions 2 Cr +3 (aq) + 3 SO 4 2- (aq) + 6 Cs + (aq) + 2 PO 4 -3 (aq)  2 CrPO 4(s) + 6 Cs + (aq) + 3 SO 4 2- (aq) 2 Cr +3 (aq) + 3 SO 4 2- (aq) 2 Cr +3 (aq) + 3 SO 4 2- (aq) + 6 Cs + (aq) + 2 PO 4 -3 (aq) 2 Cr +3 (aq) + 3 SO 4 2- (aq) + 6 Cs + (aq) + 2 PO 4 -3 (aq)  2 CrPO 4(s) SE WESE

21. Step 1: Write a balanced chemical equation (molecular equation) A solution of chromium (III) sulfate is added to a solution of cesium phosphate. Give the net ionic equation for any reaction that happens. Cr 2 (SO 4 ) 3(aq) + 2 Cs 3 PO 4(aq)  2 CrPO 4(s) + 3 Cs 2 SO 4(aq) Step 2: Break STRONG ELECTROLYTES ONLY into ions SE WESE 2 Cr +3 (aq) + 3 SO 4 2- (aq) + 6 Cs + (aq) + 2 PO 4 -3 (aq)  2 CrPO 4(s) + 6 Cs + (aq) + 3 SO 4 2- (aq) Step 3: Cross out all spectator ions 2 Cr +3 (aq) + 2 PO 4 -3 (aq)  2 CrPO 4(s) N.I.E

22. Step 1: Write a balanced chemical equation (molecular equation) A zinc acetate solution is combined with 3.50 M perchloric acid. Give the net ionic equation.  +HClO 4 HC 2 H 3 O 2 Zn(C 2 H 3 O 2 ) 2 2 HClO 4(aq) + Zn(C 2 H 3 O 2 ) 2(aq)  2 HC 2 H 3 O 2(aq) + Zn(ClO 4 ) 2(aq)

23. Step 1: Write a balanced chemical equation (molecular equation) A zinc acetate solution is combined with 3.50 M perchloric acid. Give the net ionic equation. Step 2: Break STRONG ELECTROLYTES ONLY into ions SE WESE 2 HClO 4(aq) + Zn(C 2 H 3 O 2 ) 2(aq)  2 HC 2 H 3 O 2(aq) + Zn(ClO 4 ) 2(aq) 2 H + (aq) + 2 ClO 4 - (aq) + Zn 2+ (aq) + 2 C 2 H 3 O 2 - (aq)  2 HC 2 H 3 O 2(aq) + Zn 2+ (aq) + 2 ClO 4 - (aq) 2 H + (aq) + 2 ClO 4 - (aq) + Zn 2+ (aq) + 2 C 2 H 3 O 2 - (aq)  2 HC 2 H 3 O 2(aq) 2 H + (aq) + 2 ClO 4 - (aq) + Zn 2+ (aq) + 2 C 2 H 3 O 2 - (aq) 2 H + (aq) + 2 ClO 4 - (aq)

24. A zinc acetate solution is combined with 3.50 M perchloric acid. Give the net ionic equation. Step 3: Cross out all spectator ions N.I.E 2 H + (aq) + 2 C 2 H 3 O 2 - (aq)  2 HC 2 H 3 O 2(aq) Step 1: Write a balanced chemical equation (molecular equation) Step 2: Break STRONG ELECTROLYTES ONLY into ions SE WESE 2 HClO 4(aq) + Zn(C 2 H 3 O 2 ) 2(aq)  2 HC 2 H 3 O 2(aq) + Zn(ClO 4 ) 2(aq) 2 H + (aq) + 2 ClO 4 - (aq) + Zn 2+ (aq) + 2 C 2 H 3 O 2 - (aq)  2 HC 2 H 3 O 2(aq) + Zn 2+ (aq) + 2 ClO 4 - (aq)

25. Net Ionic Equations molecular equation Break ALL strong electrolytes into ions. Leave non and weak electrolytes alone: ionic equation Start with a balanced chemical equation (including subscripts): 2 HF (aq) + Ba(OH) 2(aq)  BaF 2(s) + 2 H 2 O (l) W.E.S.E.W.E.N.E. 2 HF (aq) 2 HF (aq) + Ba 2+ (aq) + 2OH - (aq) 2 HF (aq) + Ba 2+ (aq) + 2OH - (aq)  BaF 2(s) 2 HF (aq) + Ba 2+ (aq) + 2OH - (aq)  BaF 2(s) + 2 H 2 O (l)

26. Net Ionic Equations Break ALL strong electrolytes into ions. Leave non and weak electrolytes alone: Cross out every thing that is EXACTLY the same on both sides: net ionic equation 2 HF (aq) + Ba 2+ (aq) + 2OH - (aq)  BaF 2(s) + 2 H 2 O (l) ionic equation

27. Net Ionic Equations molecular equation (M.E.) Break ALL strong electrolytes into ions. Leave non and weak electrolytes alone: ionic equation (I.E.) Start with a balanced chemical equation (including subscripts): NaNO 3(aq) + LiCl (aq)  NaCl (aq) + LiNO 3(aq) S.E. Na + (aq) + NO 3 - (aq) Na + (aq) + NO 3 - (aq) + Li + (aq) + Cl - (aq)  Na + (aq) + Cl - (aq) Na + (aq) + NO 3 - (aq) + Li + (aq) + Cl - (aq) Na + (aq) + NO 3 - (aq) + Li + (aq) + Cl - (aq)  Na + (aq) + Cl - (aq) + Li + (aq) + NO 3 - (aq)

28. Net Ionic Equations Break ALL strong electrolytes into ions. Leave non and weak electrolytes alone: ionic equation (I.E.) Na + (aq) + NO 3 - (aq) + Li + (aq) + Cl - (aq)  Na + (aq) + Cl - (aq) + Li + (aq) + NO 3 - (aq) Cross out every thing that is EXACTLY the same on both sides: no reaction (N.R.) Na + (aq) + NO 3 - (aq) + Li + (aq) + Cl - (aq)  Na + (aq) + Cl - (aq) + Li + (aq) + NO 3 - (aq)

29. Na 2 C 2 O 4(aq) + CaCl 2(aq)  2 NaCl (aq) + CaC 2 O 4(s) Net Ionic Equations Start with a balanced chemical equation: Na 2 C 2 O 4(aq) + CaCl 2(aq)  2 NaCl (aq) + CaC 2 O 4(s) M.E. Break ALL strong electrolytes into ions. Leave non and weak electrolytes alone: I.E. S.E. W.E. 2Na + (aq) + C 2 O 4 -2 (aq) + Ca 2+ (aq) + 2Cl - (aq)  2Na + (aq) + 2Cl - (aq) + CaC 2 O 4(s) 2Na + (aq) + C 2 O 4 -2 (aq) 2Na + (aq) + C 2 O 4 -2 (aq) + Ca 2+ (aq) + 2Cl - (aq) 2Na + (aq) + C 2 O 4 -2 (aq) + Ca 2+ (aq) + 2Cl - (aq)  2Na + (aq) + 2Cl - (aq)

30. Net Ionic Equations Break ALL strong electrolytes into ions. Leave non and weak electrolytes alone: I.E. Cross out every thing that is EXACTLY the same on both sides: 2Na + (aq) + C 2 O 4 -2 (aq) + Ca 2+ (aq) + 2Cl - (aq)  2Na + (aq) + 2Cl - (aq) + CaC 2 O 4(s) C 2 O 4 -2 (aq) + Ca 2+ (aq)  CaC 2 O 4(s) N.I.E. 2Na + (aq) + C 2 O 4 -2 (aq) + Ca 2+ (aq) + 2Cl - (aq)  2Na + (aq) + 2Cl - (aq) + CaC 2 O 4(s)

31. Net Ionic Equations Start with a balanced chemical equation (including subscripts): M.E. Break ALL strong electrolytes into ions. Leave non and weak electrolytes alone: I.E. S.E.W.E.S.E.W.E. Fe 2 (SO 4 ) 3(aq) + 3 PbCl 2(s)  2 FeCl 3(aq) + 3 PbSO 4(s) 2Fe +3 (aq) + 3SO 4 -2 (aq) + 3PbCl 2(s)  2Fe +3 (aq) + 6Cl - (aq) 2Fe +3 (aq) + 3SO 4 -2 (aq) + 3PbCl 2(s) 2Fe +3 (aq) + 3SO 4 -2 (aq) 2Fe +3 (aq) + 3SO 4 -2 (aq) + 3PbCl 2(s)  2Fe +3 (aq) + 6Cl - (aq) + 3PbSO 4(s)

32. 3SO 4 -2 (aq) + 3PbCl 2(s)  2Cl - (aq) + 3PbSO 4(s) Net Ionic Equations Break ALL strong electrolytes into ions. Leave non and weak electrolytes alone: I.E. Cross out every thing that is EXACTLY the same on both sides: N.I.E. 2Fe +3 (aq) + 3SO 4 -2 (aq) + 3PbCl 2(s)  2Fe +3 (aq) + 6Cl - (aq) + 3PbSO 4(s) 3SO 4 -2 (aq) + 3PbCl 2(s)  6Cl - (aq) + 3PbSO 4(s)

33. Net Ionic Equations M.E. Break ALL strong electrolytes into ions. Leave non and weak electrolytes alone: I.E. Start with a balanced chemical equation (including subscripts): 2 HC 2 H 3 O 2(aq) + Ba(OH) 2(aq)  Ba(C 2 H 3 O 2 ) 2(aq) + 2 H 2 O (l) W.E.S.E. N.E. 2 HC 2 H 3 O 2(aq) 2 HC 2 H 3 O 2(aq) + Ba 2+ (aq) + 2 OH - (aq)  Ba 2+ (aq) + 2 C 2 H 3 O 2 - (aq) 2 HC 2 H 3 O 2(aq) + Ba 2+ (aq) + 2 OH - (aq) 2 HC 2 H 3 O 2(aq) + Ba 2+ (aq) + 2 OH - (aq)  Ba 2+ (aq) + 2 C 2 H 3 O 2 - (aq) + 2 H 2 O (l)

34. HC 2 H 3 O 2(aq) + OH - (aq)  C 2 H 3 O 2 - (aq) + H 2 O (l) Net Ionic Equations Break ALL strong electrolytes into ions. Leave non and weak electrolytes alone: I.E. Cross out every thing that is EXACTLY the same on both sides: 2 HC 2 H 3 O 2(aq) + Ba 2+ (aq) + 2OH - (aq)  Ba 2+ (aq) + 2C 2 H 3 O 2 - (aq) + 2 H 2 O (l) N.I.E. 2 HC 2 H 3 O 2(aq) + 2OH - (aq)  2C 2 H 3 O 2 - (aq) + 2 H 2 O (l) 2 HC 2 H 3 O 2(aq) + Ba 2+ (aq) + 2 OH - (aq)  Ba 2+ (aq) + 2 C 2 H 3 O 2 - (aq) + 2 H 2 O (l)

35. Net Ionic Equations M.E. Break ALL strong electrolytes into ions. Leave non and weak electrolytes alone: I.E. Start with a balanced chemical equation (including subscripts): H 2 SO 4(aq) + 2 NaOH (aq)  Na 2 SO 4(aq) + 2 H 2 O (l) S.E. N.E. 2H + (aq) + SO 4 -2 (aq) + 2Na + (aq) + 2OH - (aq)  2Na + (aq) + SO 4 -2 (aq) 2H + (aq) + SO 4 -2 (aq) 2H + (aq) + SO 4 -2 (aq) + 2Na + (aq) + 2OH - (aq) 2H + (aq) + SO 4 -2 (aq) + 2Na + (aq) + 2OH - (aq)  2Na + (aq) + SO 4 -2 (aq) + 2H 2 O (l)

36. Net Ionic Equations Break ALL strong electrolytes into ions. Leave non and weak electrolytes alone: I.E. H + (aq) + OH - (aq)  H 2 O (l) 2 H + (aq) + 2 OH - (aq)  2 H 2 O (l) 2H + (aq) + SO 4 -2 (aq) + 2Na + (aq) + 2OH - (aq)  2Na + (aq) + SO 4 -2 (aq) + 2H 2 O (l) Cross out every thing that is EXACTLY the same on both sides: N.I.E. 2H + (aq) + SO 4 -2 (aq) + 2Na + (aq) + 2OH - (aq)  2Na + (aq) + SO 4 -2 (aq) + 2H 2 O (l)

37. Oxidation Numbers/States

38. Monatomic ions Some charges can be predicted based on group number Some charges can be predicted based on group number 2+ Be A Bit of Review

39. Variable Oxidation State metals Charge can NOT be predicted based on periodic table Charge can NOT be predicted based on periodic table ANYTHING IN HERE CAN BE A VOS METAL(ALMOST)

40. A Bit of Review Variable Oxidation State metals Charge can NOT be predicted based on periodic table Charge can NOT be predicted based on periodic table Can assume multiple ion charges Can assume multiple ion charges non VOS metal – potassium ion = K + VOS metal – lead ion = Pb 2+ or Pb 4+ The charge of VOS metals MUST be indicated in the name with the use of roman numerals The charge of VOS metals MUST be indicated in the name with the use of roman numerals Sn = tin () ionSn = tin () ion Sn 2+ = tin (II) ionSn 4+ = tin (IV) ion

41. (III) CO 3 2- Co 2 (CO 3 ) 3 -6 cobaltcarbonate CO 3 2- Co CO 3 2- Total cation charges + total anion charges = 0 2x+= 0 x = 3 = oxidation state of Co Co CO 3 2-

42. (IV) CrO 4 2- Pb(CrO 4 ) 2 -4 leadchromate(IV) CrO 4 2- Pb CrO 4 2- Total cation charges + total anion charges = 0 1x+= 0 x = 4 = oxidation state of Pb

43. O 2- MnO 4 - -8 O 2- Mn O 2- Total atom “A” charges + total atom “B” charges = charge 1x+= x = 7 = oxidation state of Mn O 2- Total cation charges + total anion charges = 0

44. O 2- Cr 2 O 7 -2 -14 O 2- Cr O 2- Total atom “A” charges + total atom “B” charges = charge 2x+= x = 6 = oxidation state of Cr O 2- Cr -2

45. Rules for Assigning Oxidation Numbers 1) The sum of the oxidation numbers will always equal the particle’s charge 2(Co) + 3(-2) = 0 Co 2 (CO 3 ) 3 MnO 4 - 1(Mn) + -8 = -1 Cr 2 O 7 -2 2(Cr) + -14 = -2 1(C) + 3(-2) = -2 CO 3 2- Co = +3 C = +4 Mn = +7 Cr = +6

46. Rules for Assigning Oxidation Numbers 1) The sum of the oxidation numbers will always equal the particle’s charge 2) The oxidation number for a neutral atom is always zero 2 K (s) + 2 H 2 O (l)  2 KOH (aq) + H 2(g) Oxidation states are zero

47. Rules for Assigning Oxidation Numbers 1) The sum of the oxidation numbers will always equal the particle’s charge 2) The oxidation number for a neutral atom is always zero 3) Oxidation numbers for non-VOS metals depend on their group Li = +1Ba = +2Al = +3

48. Rules for Assigning Oxidation Numbers 1) The sum of the oxidation numbers will always equal the particle’s charge 2) The oxidation number for a neutral atom is always zero 3) Oxidation numbers for non-VOS metals depend on their group 4) Oxidation numbers for VOS metals are found based on anion 5) Oxidation numbers for nonmetals are typically found based on their group P = -3Se = -2Cl = -1

49. Rules for Assigning Oxidation Numbers 5) Oxidation numbers for nonmetals are typically found based on their group typically?!?! Fluorine = -1 always! Hydrogen = +1 unless paired with a metal Oxygen = -2 unless peroxide NaH (Na = +1, H = -1) BaO 2 (Ba = +2, O = -1) Other nonmetals: the element closest to fluorine on the PT gets to keep its “usual” O.S. OF 2 (F = -1, O = +2) ClO 3 -1 SCl 4 O.S. oxygen = -2 O.S. chlorine = +5O.S. sulfur = +4 O.S. chlorine = -1 P2S5P2S5 O.S. phosphorus = +5 O.S. sulfur = -2

50. O 2- Fe 3 O 4 -8 O 2- Fe O 2- Total atom “A” charges + total atom “B” charges = charge 3x+= x = 8/3 = oxidation state of Fe O 2- 0 Fe HUH??

51. O 2- Fe 3 O 4 -8 Fe O 2- Total atom “A” charges + total atom “B” charges = charge 3x+= x = 8/3 = oxidation state of Fe O 2- 0 Fe HUH??

52. Fe 3 O 4 Total atom “A” charges + total atom “B” charges = charge x = 8/3 = oxidation state of Fe Fe 3 O 4 = FeO + Fe 2 O 3 O.S. Fe = +2 O.S. Fe = +3 Average = 8/3

53. Practice!! Assign oxidation numbers to all atoms in each of the following: a) H 2 CO b) S 2 O 3 2- c) NH 4 + d) NO 3 - e) Br 2 f) Ca 3 (PO 4 ) 2 g) 2 K (s) + 2 H 2 O (l)  2 KOH (aq) + H 2(g)

54. Oxidation and Reduction Oxidation is the process that occurs when Oxidation is the process that occurs when the oxidation number increases (gets more positive) the oxidation number increases (gets more positive) an element loses electrons an element loses electrons a compound adds bonds to oxygen a compound adds bonds to oxygen a compound loses bonds to hydrogen a compound loses bonds to hydrogen a half-reaction has electrons as products a half-reaction has electrons as products Reduction is the process that occurs when Reduction is the process that occurs when the oxidation number of an element decreases (gets more negative) the oxidation number of an element decreases (gets more negative) an element gains electrons an element gains electrons a compound loses bonds to oxygen a compound loses bonds to oxygen a compound gains bonds to hydrogen a compound gains bonds to hydrogen a half-reaction has electrons as reactants a half-reaction has electrons as reactants