General Chemistry II

1. General Chemistry II Welcome!

2. Reactions • Nuclear • Dealing with subatomic particles (protons, neutrons) • Results in change in atom • If neutrons change = different isotope • If electrons change = different charge • If protons change = different element

3. Reactions • Chemical • Atom level • How atoms interact with each other • A change = change of substance

4. Reactions • Physical • Molecules • Rearranging, not changing • Ex. Water to ice

5. Reactions Reactants Products C6H12O6(s) + 6O2(g) 6 H2O(g) • Number in front tells you # molecules • No number = 1 • Tells ratio 1:6 • Subscript tells how many

6. Reactions • Often see state of matter • G = gas • L = liquid • S = solid • Aq = aqueous • Some reactions can go back and forth • Equilibrium • Noted as

7. Classifying Reactions • Combination reaction • 2 substances combine to a third substance • A+B→C • Na + Cl → NaCl • CaO(s) + SO2(g) → CaSO3(s)

8. Classifying Reactions • Decomposition Reaction • A single compound reacts to give 2 or more substances • C → A + B • 2HgO(s) → 2Hg(l) + O2(g) ∆

9. Classifying Reactions • Displacement Reaction • AKA: Single Replacement • A + BC → AC + B • Cu(s) + 2AgNO3(aq) → 2Ag(s) + Cu(NO3)2(aq)

10. Classifying Reactions • Metathesis • AKA: Double Replacement • AB + CD → AD + CB • 2Kl(aq) + Pb(NO3)(aq) → 2KNO3(aq) + PbI2(s)

11. Classifying Reactions • Combustion Reactions • Involve Oxygen • AB + O → AO + BO • C6H12O6 + 6O2 → 6CO2 + 6H2O • Can predict outcomes: • C → CO2 • H → H2O • S → SO2 • N → NO, NO2

12. Classifying Reactions - Summary • Combination A+B→C • Decomposition C → A + B • Single Replacement A + BC → AC + B • Double Replacement AB + CD → AD + CB • Combustion AB + O → AO + BO

13. Evidence of a Chemical Reaction • Color Change • Formation of a solid (precipitate) within a clear solution • Evolution of a gas • Evolution or absorption of heat

14. Balancing Simple Equations C6H12O6 + O2→ CO2 + H2O The sides don’t balance _1_C6H12O6 + _6_O2→ _6_CO2 + _6_H2O

15. Balancing Simple Equations • Order to balance equations: • Start with non-Oxygen, non-Hydrogen • Go to Hydrogen or Oxygen • Do anything in single atom form (elementary state) • Same atom in more than one compound on same side of reaction.

16. Collision Theory In order for a chemical reaction to occur, two or more compunds: • Must collide; • Must collide with the proper amount of energy; • Must collide in the proper orientation.

17. Stoichiometry • Find molecular weight of both reactant and product. • Determine ratio of reactant to product. • Ratio x known variable • Renders theoretical yield

18. Stoichiometry • In most cases, theoretical and actual yields are not the same. • Percent Yield = Actual/theoretical x 100 • Percent Error = 100 – percent yield

19. Limiting Reagent 3 eggs + 2 cups flour = 1 cake 15 eggs = 5 cakes 6 cups flour = 3 cakes 15 eggs + 6 cups flour = 3 cakes flour is limiting reagent Limiting reagent limits the overall yield.

20. Oxidation – Reduction ReactionREDOX • Oxidation – loss of electrons • Fe+2→ Fe+3 + 1e- • Reduction – gain of electrons • Fe+3 + 1e- → Fe+2

21. Indicators • Oxidation Indicator: • Increase O, decrease H • Reduction Indicator: • Decrease O, increase H

22. Redox Reactions • Oxidizing agent → Oxidant • Causes something to become oxidized • NH3 → NO + e- • Reduction agent → Reductant • -e- + O2 → H2O

23. Redox Reactions Rules: • Number of atoms must balance • Exchange of electrons must balance

24. Redox Reactions • Half-Reaction Method SO3-2 + MnO4→ SO4-2 + Mn+2 Acid • Steps: • Pull out like units; (this creates the half-reactions) • SO3-2 → SO4-2 • MnO4→ Mn+2 • Balance non-Oxygen, non-Hydrogen in half reac. • ___SO3-2 → ___SO4-2 • ___MnO4→ ___Mn+2

25. Redox Reactions - Steps • Balance H2 atoms by adding H+ • Note: H2O  H+ + OH+ • H2O + SO3-2 → SO4-2 + 2H+ • 8H+ + MnO4→ Mn+2 +4H2O • Balance the charge by adding electrons • H2O + SO3-2 → SO4-2 + 2H+ -2 charge -2+2 = 0 charge +2e- Must add e- • 8H+ + MnO4→ Mn+2 +4H2O +8-1 = +7 +2 -5e- Oxidation Reduction

26. Redox Reactions - Steps • Balance exchange of electrons between half reactions (2e- and 5 e-) • 1/2 = 5/10 – multiply top by 5 • +1/5 = 2/10 – multiply bottom by 2 5H2O + 5SO3-2 → 5SO4-2 + 10H+ + 10 e- 10e- + 16H+ + 2MnO4→ 2Mn+2 +8H2O

27. Redox Reactions - Steps • Recombine; put H & H2O on ends 6H + 5SO3-2 + 2MnO4- → 5SO4-2+2Mn+2+3H2O • If in base condition, convert to base • Recheck atoms and charges.

28. Test I Review • Differences between physical, chemical and nuclear reaction • Different parts of a chemical equation • Classical classification of reactions • Balance chemical equations • Stoichiometry • Limiting Reagent • % yield • % error

29. Test I Review • Redox Reactions • Definitions of Oxidation, Reduction • Identifying Oxidation vs Reduction • Identifying Oxidant vs Reductant • Balancing reactions

30. Chemical Kinetics • Rate of reaction • How fast the concentration of a reactant or product changes with time A + 2B → C + 3D Rate of Reaction Rate of Formation

31. Chemical Kinetics Δ[A] or Δ[B] or Δ[C] or Δ[D] Δt Δt Δt Δt Negative ValuesPositive Values

32. Chemical Kinetics Reminder: In order for a chemical reaction to occur, two or more compunds: • Must collide; • Must collide with the proper amount of energy; • Must collide in the proper orientation.

33. Chemical Kinetics • What will affect the rate? • Adding heat • Increases collisions • Stirring • Increases collisions & surface area • Increasing surface area • Adding a catalyst • Works on Ea & orientation • Add more reactants • Remove products as they form

34. Catalysts • Any chemical that will reduce Ea without itself becoming part of the chemical reaction. • Catalysts do not direct the reaction; just speed it up.

35. Rate Law Expression Rate = K[A]n[B]m • Include anything that affects the rate

36. Basic Rate Orders 0˚ - Concentration does not affect the rate Rate = K[A]˚ Rate = K Rate [A]

37. Basic Rate Orders 1˚ - Concentration is proportional Rate = K[A]1 Rate [A]

38. Basic Rate Orders 2˚ - Increased concentration = rate x 4 Rate = K[A]2 Rate [A]

39. Basic Rate Orders Inverse 1˚ (-1˚) Increased rate = decreased concentration Rate = K[A]-1 or K/[A] Rate [A]

40. Half Life • T1/2 • How long it takes half of the reactant to react • If T1/2 = 2 hours • T0 100M • T2hrs 50M • T4 25M • T6 12.5M etc…

41. Enzymes • Enzymes are catalysts in the body • Enzymes have 2 parts: • Apoenzyme – protein portion • Cofactor – some mineral; often cofactor is a catalyst • May have co-enzyme instead of cofactor • A protein surrounds the enzyme creating specificity in catalyzing reactions. • Holoenzyme = apoenzyme + cofactor

42. Enzymes Substrate • Generic for any chemical acted upon by an enzyme • 2 criteria for substrate • Proper shape to fit into active site • Lock & Key Model • Induced Fit Model / Hand in Glove • Proper Chemical Interaction

43. Enzymes 2 major categories of enzymes • Nonallosteric Enzymes • Inhibitor site • Allows inhibitor to connect • Not competing for active site, so non-competitive inhibitor. • Changes shape of active site, prevents substrate from attaching. • Adding substrate has no effect; only change by removing inhibitor.

44. Enzymes Allosteric Enzymes • Sequential Model • Starting at inactive state • Substrate attaches to active site • Active site doesn’t have proper shape; takes longer for substrate to attach • Once it attaches, shape of all other active sites change and become active • Converts all to active state.

45. Glycogen • Storage form of glucose in animals ↑ Osmotic pressure Hypotonic H2O 1000 glucose 1 glycogen

46. Enzymes Non-Allosteric Allosteric Rate Rate 0˚ 0˚ 1˚ 1˚ 2˚ [A] [A]

47. Rate of Reactions 2NO + Cl2 → 2NOCl Rate = K[NO]n[Cl2]m

48. Rate of Reactions CH3N2CH3→ C2H6 + N2 Rate = K[CH3N2CH3]m

49. Rate of Reactions 2NO + O2 → 2NO2 Rate = K[NO]n[O2]m

50. Rate of Reactions 2ClO2 + 2OH-→ ClO3- + ClO2- + H2O Rate = K [ClO2]n [OH-]m