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Regents Review

Regents Review. Unit 13 – Reduction & Oxidation Unit 14 – Organic Chemistry. Reduction – Oxidation (Redox). Chapters 20 + 21 Reference Table J. Oxidation Numbers (States). Positive, negative or neutral values assigned to an atom to keep track of the number of electrons lost or gained.

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Regents Review

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  1. Regents Review Unit 13 – Reduction & Oxidation Unit 14 – Organic Chemistry

  2. Reduction – Oxidation (Redox) Chapters 20 + 21 Reference Table J

  3. Oxidation Numbers (States) • Positive, negative or neutral values assigned to an atom to keep track of the number of electrons lost or gained. • Oxidation number of an element by itself is 0. • Oxidation number of a monatomic ion is equal to its charge • For any neutral compound, the sum of the oxidation numbers is 0 • For any polyatomic ion, the sum of the oxidation numbers must equal the charge of the ion

  4. Common Oxidation Numbers • Group 1 +1 • Group 2  +2 • Group 13  +3 • Group 16  -2 • Group 17  -1 • Exceptions to each above

  5. Redox Reactions • Reduction – Oxidation, or redox, involves the transfer of electrons • Reduction – gain of electrons • Oxidation – loss of electrons • LEO goes GER • Lose Electrons Oxidation • Gain Electrons Reduction

  6. Redox Reaction +2 0 0 -1 • Mg + Cl2 MgCl2 • Mg - lost electrons (oxidation) • Cl – gained electrons (reduction)

  7. Redox Reaction • Zn + CuSO4 Cu + ZnSO4 • One element loses electrons (oxidation) • One element gains electrons (reduction) • All other ions are spectators • Look for Single Replacement First

  8. Net Ionic Equation • Shows only the ions involved in the redox reaction, not spectator ions • Still shows conservation of mass and charge • Zn + CuSO4 Cu + ZnSO4 • Zn + Cu+2 Cu + Zn+2

  9. Half Reactions Only shows one element and how many electrons are gained or lost Zn + CuSO4 Cu + ZnSO4 Zn  Zn+2 + 2e- Oxidation Cu2+ + 2e- Cu Reduction

  10. Balancing Reactions • The number of electrons lost must equal the number of electrons gained • Example: • Zn + Na2SO4 2Na + ZnSO4 • Zn  Zn+2 + 2e- • 2(Na+ + e- Na)

  11. Spontaneous Reactions • More active element does not want to be alone • Table J • Metal being oxidized must be ABOVE metal being reduced for spontaneous reactions to occur • Nonmetal being reduced must be ABOVEnonmetal being oxidized for spontaneous reactions to occur

  12. Electrochemical Cells • any device that converts chemical energy into electrical energy or electrical energy into chemical energy • Two types • Voltaic (Chemical) • Electrolytic

  13. Electrochemical Cells • Electrode – conductor in an electrical circuit that carries electrons to or from a substance other than a metal. • Cathode – electrode where reduction takes place • Anode – electrode where oxidation takes place

  14. Electrochemical Cell Components • Salt Bridge • Allows for the passage of ions, not electrons • Switch • Device that opens(turns off) and closes(turns on) circuit

  15. Voltaic Cell • Flow of electrons is spontaneous • Chemical energy is converted to electrical energy • Cathode is positive, anode is negative • Examples: Batteries

  16. Electrolysis • Process in which electrical energy is converted to chemical energy • Example: • 2H2O  2H2 + O2

  17. Electrolytic Cells • Electrons are pushed by an outside power source • Electrical energy is converted to chemical energy • Cathode is negative, anode is positive • Examples: Electroplating, Electropolishing

  18. Voltaic Cell

  19. Electrolytic Cell

  20. Voltaic or Electrolytic?

  21. Voltaic or Electrolytic?

  22. Organic Chemistry Chapter 22-23 Reference Tables P, Q, R

  23. Organic Chemistry • Study of carbon-based compounds • Carbon has 4 valence electrons • Carbon always forms 4 covalent bonds • Carbon compounds form chains and/or rings

  24. Hydrocarbons • Molecule or compound composed of carbon and hydrogen only • Three Main Types • Alkanes, Alkenes, Alkynes • Homologous Series • Group of compounds with similar structure and function • Table Q (Alkanes, Alkenes, Alkynes)

  25. Alkanes • All single C─C bonds H H H │ │ │ H─ C─ C─ C─ H │ │ │ H H H H H H H │ │ │ │ H─ C─ C─ C─ C─ H │ │ │ │ H H H H H H H H H │ │ │ │ │ H ─ C─ C─ C─ C─ C─ H │ │ │ │ │ H H H H H C3H8 C4H10 C5H12 General Formula CnH2n+2

  26. Alkenes • At least one double C=C bond H H H │ │ │ H─ C─ C= C─ H │ H H H H H │ │ │ │ H─ C─ C─ C= C─ H │ │ H H H H H H H │ │ │ │ │ H ─ C─ C─ C─ C= C─ H │ │ │ H H H C3H6 C4H8 C5H10 General Formula CnH2n

  27. Alkynes • At least one triple C≡C bonds H │ H─ C─ C≡ C─ H │ H H H │ │ H─ C─ C─ C≡ C─ H │ │ H H H H H │ │ │ H ─ C─ C─ C─ C≡ C─ H │ │ │ H H H C3H4 C4H6 C5H8 General Formula CnH2n-2

  28. Saturated vs. Unsaturated • Saturated Hydrocarbons • Contain only single Carbon to Carbon bonds (Alkanes) • Unsaturated Hydrocarbons • Contains at least one multiple (double, triple) Carbon to Carbon bond • Alkenes and Alkynes

  29. Naming Simple Hydrocarbons • Name is based on two parts • Number of carbons in longest continuous chain (Table P) • Type of bonds between carbons • Single –ane • Double –ene • Triple –yne

  30. Examples • Butane • Pentane H H H H │ │ │ │ H─ C─ C─ C─ C─ H │ │ │ │ H H H H H H H H H │ │ │ │ │ H ─ C─ C─ C─ C─ C─ H │ │ │ │ │ H H H H H

  31. Alkenes and Alkynes • Name includes location of multiple bond • Carbons numbered to give multiple bond the lowest possible number • 1-butyne 2-butyne H H │ │ H─ C─ C─ C≡ C─ H │ │ H H H H │ │ H─ C─ C≡ C ─ C─ H │ │ H H

  32. Condensed Structural Formula • Shows who is bonded to who, without the actual bonds H H H │ │ │ H─ C─ C─ C─ H │ │ │ H H H H H H │ │ │ H─ C─ C= C─ H │ H

  33. Branched Hydrocarbons • Alkyl Group • Hydrocarbon branch • Branch name ends with -yl • Location of branch is indicated by number • Number carbons to give lowest possible number • Multiple bond still takes priority in numbering

  34. Examples • 3-Methyl Pentane • 2-Methyl 1-Butene H H CH3 H H │ │ │ │ │ H ─ C─ C─ C─ C─ C─ H │ │ │ │ │ H H H H H H │ H─ C─ H │ H CH2 H │ │ │ H─ C─ C= C─ H │ H

  35. Functional Groups • Specific arrangement of atoms that give compounds a unique property • Table R • Examples • Hydroxyl Group, -OH • Carbonyl Group, O ǁ ─ C ─

  36. Halogen attached to a carbon Prefix indicates which halogen Table R # for which carbon halogen is attached Multiple bonds still take priority in numbering Halides H H Cl H │ │ │ │ H─ C─ C─ C─ C─ H │ │ │ │ H H H H Br H H │ │ │ H─ C─ C= C─ H │ H 2-Chlorobutane 3-Bromopropene CH3CH2CHClCH3 CH2BrCHCH2

  37. Hydroxyl group (-OH) attached to a carbon # for which alcohol group is attached Name ends in –ol Multiple bonds still take priority in numbering Alcohol H H H H │ │ │ │ HO─ C─ C─ C─ C─ H │ │ │ │ H H H H H H H H H H │ │ │ │ │ │ H ─ C─ C─ C─ C─ C─ C─ H │ │ │ │ │ │ H H H OH H H 3-Hexanol 1-Butanol CH3CH2CH2CH2OH CH3CH2CH2CHOHCH2CH3

  38. Carbonyl group at end of chain Name ends with –al Condensed structural formula ends with -CHO Aldehyde O ǁ ─ C ─ H H H H H O │ │ │ │ ǁ H ─ C─ C─ C─ C─ C─ H │ │ │ │ H H H H H H O │ │ ǁ H─ C─ C─ C─ H │ │ H H Propanal Pentanal CH3CH2CHO CH3CH2CH2CH2CHO

  39. Ketones • Carbonyl group not on end of chain • Number indicates which carbon the oxygen is attached to • Name ends with –one • Condensed structural formula has -CO- in it O ǁ ─ C ─ H O H H H │ ǁ │ │ │ H ─ C─ C─ C─ C─ C─ H │ │ │ │ H H H H H O H │ ǁ │ H─ C─ C─ C─ H │ │ H H Propanone 2-Pentanone CH3COCH3 CH3COCH2CH2CH3

  40. Amine • Nitrogen attached to a carbon • Number indicates which carbon the nitrogen is attached to • Name ends in –amine • Multiple bonds still take priority in numbering H H H H │ │ │ │ H2N─ C─ C─ C─ C─ H │ │ │ │ H H H H H H H H H H │ │ │ │ │ │ H ─ C─ C─ C─ C─ C─ C─ H │ │ │ │ │ │ H H H NH2 H H 3-Hexanamine 1-Butanamine CH3CH2CH2CH2NH2 CH3CH2CH2CHNH2CH2CH3

  41. Amide • Carbonyl group with an amine group attached to it • Must be on an end • Name ends in -amide O H ǁ ǀ ─ C ─ N ─ H H H H H O │ │ │ │ ǁ H ─ C─ C─ C─ C─ C─ NH2 │ │ │ │ H H H H H H O │ │ ǁ H─ C─ C─ C─ NH2 │ │ H H Propanamide Pentanamide CH3CH2CONH2 CH3CH2CH2CH2CONH2

  42. O ǁ ─ C ─ O ─ H Organic Acid • Carbonyl group with a hydroxyl group attached to it • Must be on an end • Name ends in –oic acid • Condensed structural formula ends with -COOH • Hydroxyl H is the acidic H H H H O │ │ │ ǁ H ─ C─ C─ C─ C─ OH │ │ │ H H H H O │ ǁ H ─ C─ C─ OH │ H Ethanoic Acid Butanoic Acid CH3COOH CH3CH2CH2COOH

  43. Ether • Single oxygen between 2 carbon chains • Name each carbon chain H H H H │ │ │ │ H─ C─ C─O ─ C─ C─ H │ │ │ │ H H H H H H H H H H │ │ │ │ │ │ H ─ C─ C─ C─ C─ O─ C─ C─ H │ │ │ │ │ │ H H H H H H Diethyl Ether Butyl Ethyl Ether CH3CH2OCH2CH3 CH3CH2CH2CH2OCH2CH3

  44. O ǁ ─ C ─ O ─ Ester • Carbonyl group with single oxygen between carbon chains • Name in two parts • 1st Branch off oxygen first as alkyl group • 2nd Chain containing Carbonyl group • Ending in –oate H O H │ ǁ │ H ─ C─ O ─ C─ C─ H │ │ H H H H H O H H │ │ │ ǁ │ │ H ─ C─ C─ C─ C─ O─ C─ C─ H │ │ │ │ │ H H H H H Methyl Ethanoate Ethyl Butanoate CH3COOCH3 CH3CH2CH2COOCH2CH3

  45. Organic Reactions • Saponification • Fermentation • Combustion • Addition • Substitution • Polymerization (2 types) • Esterification

  46. Organic Reactions • Saponification • Production of Soap • Fermentation • Production of ethanol and CO2 from sugar • C6H12O6 2C2H5OH + 2CO2

  47. Combustion • Complete Combustion of a Hydrocarbon • CxHy + O2 CO2 + H2O • Example: • C3H8 + 5O2 3CO2 + 4H2O

  48. Addition • Addition of a halogen onto an alkene or alkyne Br Br │ │ H─ C≡ C─ H + Br─Br  H─ C= C─ H Br Br Br Br │ │ │ │ H─ C= C─ H + Cl─Cl  H─ C─ C─ H │ │ Cl Cl

  49. Substitution • Substitution of one halogen in place of a hydrogen on an alkane H H H H │ │ │ │ H─ C─ C─ H + Cl─Cl  H─ C─ C─ H + H ─ Cl │ ││ │ H H H Cl H H Br H │ │ │ │ H─ C─ C─ H + Br─Br  H─ C─ C─ H + H ─ Br │ ││ │ H Cl H Cl

  50. Polymerization • Connecting of smaller pieces into a long repeating chain • Plastics, starches, nylon • Two types: • Addition • Condensation -(C2H4)n-

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