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 otherA change = change of substance

4 Reactions Physical Molecules Rearranging, not changingEx. 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 matterG = gas L = liquid S = solid Aq = aqueous Some reactions can go back and forth Equilibrium Noted as

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

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

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

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

11 Classifying ReactionsCombustion 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 - SummaryCombination 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 ReactionColor Change Formation of a solid (precipitate) within a clear solution Evolution of a gas Evolution or absorption of heat

14 Balancing Simple EquationsC6H12O6 + O2 → CO2 + H2O The sides don’t balance _1_C6H12O6 + _6_O2 → _6_CO2 + _6_H2O

15 Balancing Simple EquationsOrder 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 limits the overall yield.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 Reaction REDOXOxidation – loss of electrons Fe+2 → Fe+3 + 1e- Reduction – gain of electrons Fe+3 + 1e- → Fe+2

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

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

23 Redox Reactions Rules: Number of atoms must balanceExchange of electrons must balance

24 Redox Reactions Half-Reaction Method SO3-2 + MnO4 → SO4-2 + Mn+2 AcidSteps: 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 - StepsBalance H2 atoms by adding H+ Note: H2O  H+ + OH+ H2O + SO3-2 → SO H+ 8H+ + MnO4 → Mn+2 + 4H2O Balance the charge by adding electrons -2 charge = 0 charge +2e- Must add e- +8-1 = -5e- Oxidation Reduction

26 Redox Reactions - StepsBalance 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 → 5SO H e- 10e- + 16H+ + 2MnO4 → 2Mn+2 + 8H2O

27 Redox Reactions - StepsRecombine; put H & H2O on ends 6H + 5SO MnO4- → 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, ReductionIdentifying Oxidation vs Reduction Identifying Oxidant vs Reductant Balancing reactions

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

31 Δ[A] or Δ[B] or Δ[C] or Δ[D]Chemical Kinetics Δ[A] or Δ[B] or Δ[C] or Δ[D] Δt Δt Δt Δt Negative Values Positive 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 StirringIncreases 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]mInclude anything that affects the rate

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

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

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

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 reactIf T1/2 = 2 hours T0 100M T2hrs 50M T M T M 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 enzyme2 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 EnzymesInhibitor 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 stateSubstrate 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 pressureHypotonic 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 Exp [NO] [Cl2]1 0.0125 0.0255 2.27x10-5 2 0.0510 4.55x10-5 3 0.0250 9.08x10-5 4 0.0500 0.1020

48 Rate of Reactions CH3N2CH3 → C2H6 + N2 Rate = K[CH3N2CH3]m Exp1 1.13x10-2 2.5x10-6 2 2.26x10-2 5.6x10-6

49 Rate of Reactions 2NO + O2 → 2NO2 Rate = K[NO]n[O2]m Exp [NO] [O2]1 0.0125 0.0253 0.0281 2 0.0250 0.112 3 0.0506 0.0561 4

50 2ClO2 + 2OH- → ClO3- + ClO2- + H2ORate of Reactions 2ClO2 + 2OH- → ClO3- + ClO2- + H2O Rate = K [ClO2]n [OH-]m Exp [ClO2] [OH-] Rate 1 0.060 0.030 0.0248 2 0.020 3 0.090 4 0.040

51 Thermochemistry Energy of a reaction 2 forms of energy → PhysicsPotential energy – stored energy Kinetic energy – energy of movement In terms of chemistry: Potential energy: energy stored in the bonds Kinetic energy: Vibrational movement of molecules Directional motion Breaking of bonds

52 Thermochemistry Work: Forming new bondsVibrational energy is perceived as temperature System vs. surroundings System: What we’re concerned with In most cases, the system is the chemical reaction being studied. Surroundings: Everything else

53 Thermochemistry SystemsOpen system – losing/exchanging energy and matter Closed system – losing/exchanging only energy Isolated system – energy and matter are confined.

54 Heat Heat is a measurement of energyTemperature is a relative experience Calorie: amount of energy required to raise 1 gram of water 1˚C

55 Energy Diagrams Understanding energy released from chemical reactionsPER ∆H PEP Exothermic Reaction

56 Energy Diagrams Understanding energy released from chemical reactionsPEP ∆H PER Endothermic Reaction

57 Measuring Exchange of EnergyBomb Calorimeter Closed system Tells total calories, not nutritional calories Not every reaction will work in B.C. Q=SMΔt Energy = Specific x Mass x ΔTemp Measuring Q via water QH2O = Qreaction Qreaction = - QH2O Q system = -Qsurroundings

58 Measuring Exchange of EnergyHess’s Law / Indirect Method Allows one to get calculations on reactions that can’t be done using the Bomb Calorimeter method.

59 Measuring Exchange of EnergyStandard Heats of Formation See Table on A.18

60 Test 2 Review Rates of reaction & formation Rate Law ExpressionRelate rates and find an unknown Rate Law Expression Basic Rate Orders Rate orders related to enzymes Half Life Theoretical Models Collision Theory Catalysts Temperature

61 Test I Review Reaction mechanisms System vs. surroundingsOpen, closed, isolated systems Potential & Kinetic energies Exothermic vs Endothermic reactions Methods for measuring enthalpy First Law of Thermodynamics Energy Diagrams

62 Chemical Equilibrium 2A  4B + Energy  EquilibriumGoing both ways in a chemical reaction Rates of forward and back reactions are equal LaChatlier’s Principles ↑[A] As [A] reacts, [A]↓ [B] ↑ Eventually returns to equilibrium

63 Equilibrium 2A  4B + Heat [I] 4A  4B [Δ] 6A 4B [F] 5A  5BEquilibrium shifted to Right

64 Adding Energy Add energy, both EaR and EaP will increase in same proportion Back reaction benefits more Additional energy doesn’t help forward reaction because it is nearly completely activated.

65 4HCl(g) + O2(g)  2H2O(g) + 2Cl2Equilibrium Constant Kc = Equilibrium Constant 4HCl(g) + O2(g)  2H2O(g) + 2Cl2 Kc = [H2O]2 [Cl]2 [HCl]4 [O2]1 Concentrations of Products to power of molar ratio Concentrations of Reactants to power of molar ratio

66 Equilibrium Constant Kc can tell which side is favoredHigh Kc value means reactants favored Low Kc value means products favored Kw – dissociation constant for water 10-14 Ka– dissociation constant for an acid Q = Kc shift left Q = Kc at equilibrium Q < Kc shift right

67 Acids & Bases ArrheniusAcids release H+ (HCl → H++Cl-) Taste sour Blue litmus → red Base Bases release OH- (NaOH → Na++OH-) Taste bitter Red litmus →blue

68 Acids & Bases Bronsted - LowryProton donor (HCl + H2O  H3O + Cl-) Conjugate acid Conjugate base Base Proton acceptor

69 Acids & Bases Lewis Acid Electron pair acceptor (Na+ + Cl → NaCl) BaseElectron pair donor Nucleophile

70 pH pH [H+] [OH] pOH 14 10-14 100 13 10-13 10-1 1 12 10-12 10-2 2 ↓ 713 10-13 10-1 1 12 10-12 10-2 2 ↓ 7 10-7 1012 1013 1014 Base Neutral Acid

71 pH pH = -log [H+] pOH = -log [OH-] pKc = -log Kc pKw = -log KwpKa = -log Ka

72 pH Kw = [H+][OH-] = 10-14 pKw = pH + pOH = 14 Px = -log [x][x] = -log -px

73 pH pH [H+] [OH-] pOH A/B/N 3.8 x 10-5 4.88x10-11 5.8x10-8 8.88 3.225.85

74 Making Acid To make: [I] .025M → ___ H+M [F] 0 → .025M + .025M.025M of HCl [I] .025M → ___ H+M [F] 0 → .025M M pH = -log[H+] = log(.025) =

75 Weak Acid Quadratic Formula: X = -b - √b2 – 4ac / 2a +

76 Buffers Weak acid, combined with a conjugate base that produces a system that resists changes in pH. Any weak acid in equilibrium with a conjugate base is a buffer system.

77 Buffers 2 common buffers in the body Carbonic Acid SystemBuffer for blood Phosphoric Acid System Buffer for cell

78 Buffers Eat / Drink something acidic, equilibrium shifts to bring H+ down. CO2 transport in blood CO2 + H2O  H2CO3  H+ + HCO3- Some dissolve in blood, a little attached to RBC, most transported via Carbonic Acid System. O2 carried via hemoglobin via RBC

79 Buffers Bodies respond to change in pH Minor changes affect systemIf blood is out of pH balance, sign of something wrong. Hyperventilation

80 CO2 + H2O  H2CO3  H+ + HCO3- Buffering Blood BreathingLiver / Kidneys Digestive System

81 Entropy Δ G – change in free energy ΔH – EnthalpyAmount of energy available to do work Tells us if reaction is spontaneous or non-spontaneous Endogonic (+) Must add energy; non-spontaneous Exogonic (-) Losing energy; spontaneous ΔH – Enthalpy Non-spontaneous Exothermic – spontaneously give off energy

82 Entropy ΔS – Entropy – measure of disorderPositive entropy = spontaneous Negative entropy = non-spontaneous Energy required to create order again

83 Second Law of ThermodynamicsAll spontaneous process increase entropy of the universe, unless energy is put into it. Increasing entropy: Solid to liquid Liquid to gas Creating a solution Moles to moles

84 Third Law of ThermodynamicsThe entropy of a pure, perfect crystal at 0 Kelvin is zero.

85 Determining entropy ΔG = ΔH - T ΔS Conflicting info? Look at TempC6H12O6(s) + 6O2(g) → 6CO2(g) + 6H2O(g) ΔH = -500KJ ΔG = ΔH - T ΔS Conflicting info? Look at Temp Small T favors ΔH High T favors ΔS To determine ΔT Look at gases or Look at physical properties

86 Test 3 Review Chemical Equilibrium Strong acids/bases Weak acids/basesEx. 17.5, 17.6, 17.7 Buffers What are they? How do they operate? Determining pH pH in the body

87 Test 3 Review Spontaneous change Increasing entropy3 Laws of Thermodynamics