chapter 2 lecture 1 kinetics/thermo & acids/bases i.review of simple kinetics and thermodynamics...
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III. Kinetics A.Rate Laws a)Describe how fast a reaction occurs and how we can effect that speed 1.For simple organic reactions, we can directly write the rate law based on the stoichiometry of the reactants 1. NO 2 + NO 2 NO 3 + NO rate = k[NO 2 ] 2 2. NO 3 + CO NO 2 + CO 2 rate = k[NO 3 ][CO] c)Other examples: A + A + B productsrate = k[A] 2 [B] A + B + C productsrate = k[A][B][C] rate = k[A][B] A + BC k = a constant unique to each reaction [A], [B] = concentration of reactants (M)TRANSCRIPT
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Chapter 2 Lecture 1 Kinetics/Thermo & Acids/Bases
I. Review of Simple Kinetics and ThermodynamicsA. Definitions
1) Thermodynamics = changes in energy during a process or reaction. Determines extent of completion of the reaction or process
2) Kinetics = rate of a process or reaction. Determines how fast the reaction or process occurs.
B. Equilibria1) Equilibrium = state of a system in which the concentrations of reactants
and products are no longer changing.
2) Equilibrium Constanta) If K is large, reaction goes forwardb) If K is small, reaction goes in reverse
D C B A K [A][B][C][D]K
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II. Equilibrium ConstantsA. The Law of Mass Action
1. This is an empirical law discovered in 18642. Every reaction has a constant associated with it telling us where the
equilibrium position is.
jA + kB lC + mD
3. K = Equilibrium Constant = tells us where the equilibrium position isa) K > 1 tells us the equilibrium lies to the rightb) K < 1 tells us the equilibrium lies to the left
4. If we know the concentrations, we can find K from its equation
5. K is written without units, even in cases where there are units left not cancelled. This is correct for nonideal behavior of molecules.
6. Sample Ex. 13.1 Write K for: 4NH3 + 7O2 4NO2 + 6H2O
7. Don’t include solvents, pure liquids or pure solids in the K equation
kj
ml
[B][A][D][C]K K
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III. KineticsA. Rate Laws
a) Describe how fast a reaction occurs and how we can effect that speed
1. For simple organic reactions, we can directly write the rate law based on the stoichiometry of the reactants
1. NO2 + NO2 NO3 + NO rate = k[NO2]2
2. NO3 + CO NO2 + CO2 rate = k[NO3][CO]
c) Other examples:A + A + B products rate = k[A]2[B]A + B + C productsrate = k[A][B][C]
rate = k[A][B]
A + B C k = a constant unique to each reaction[A], [B] = concentration of reactants (M)
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IV. Bronsted-Lowery Model of Acids and Basesa) Acid is an H+ donorb) Base is an H+ acceptorc) HCl + H2O H3O+ + Cl-
1) General Acid Equation
HA + H2O H3O+ + A-
a) Conjugate base = what is left after H+ leaves acidb) Conjugate acid = base + H+ c) Conjugate acid-base pair are related by loss/gain of H+
d) Competition for H+ by A- and H2O; strongest base wins
acid base hydronium ion
H O
HH Cl H O
H
H Cl+ +
acid baseconjugateacid
conjugatebase
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2) Ka = acid dissociation constant
3) Sample Exercise: Write simple Ionizations for:HCl, HC2H3O2, NH4
+, C6H5NH3+, Al(H2O)6
3+
4) Bronsted-Lowery theory allows for non-aqueous solutionsNH3 + HCl NH4Cl
[HA]]][A[H
[HA]]][AO[HK 3
a
H N
H
H
H Cl H N
H
H
H
Cl+ +
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B. Acid Strength1) Acid strength describes the equilibrium position of the ionization reaction
HA + H2O H3O+ + A-
2) Strong Acid = equilibrium lies far to the right (Large Ka)
a) Almost all HA has ionized to H+ and A- ([H+] = [HA]0)
b) A strong acid has a weak conjugate basei. To ionize fully, the conjugate base must have low proton affinityii. The conjugate base must be weaker that water
3) Weak Acid = equilibrium lies far to the left (Small Ka)
a) Almost all HA remains unionized ([H+] << [HA]0)
b) A weak acid has a strong conjugate basec) The conjugate base is much stronger than water
Strong acid Weak Acid
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C. Water as an Acid and Base1) An amphoteric substance can behave as an acid or a base (water)2) Autoionization of water (reaction with itself)
H2O + H2O H3O+ + OH-
3) Ionization constant for water = KW = [H3O+][OH-] = [H+][OH-]
a) For any water solution at 25 oC, [OH-] x [H+] = KW = 1 x 10-14
b) Neutral solutions (pure water) have [OH-] = [H+] = 1 x 10-7 c) Acidic solutions: [H+] > [OH-] d) Basic solutions: [OH-] > [H+]
e) Sample Ex. Calculate [OH-] or [H+] for the following:i. [OH-] = 1 x 10-5 Mii. [OH-] = 1 x 10-7 Miii. [H+] = 10 M
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D. pH Scale1) pH = -log[H+] (simplifies working with small numbers)
2) If [H+] = 1.0 x 10-7, pH = -log(1 x 10-7) = -(-7.00) = 7.00
3) pOH = -log[OH-] pKa = -logKa
4) pH changes by 1 unit for every power of 10 change in [H+]
a) pH = 3 [H+] = 10 times the [H+] at pH = 4a) pH decreases as [H+] increases (pH = 2 more acidic than pH = 3)
E. Meaning of pKa
HA + H2O H3O+ + A-
The lower the pKa, the stronger the acid
5 )10 x -log(1 -logKpK 10 x 1HA
AHK 5-aa
5a
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B. Predicting Acid/Base Strength1) Size of A-: HI > HBr > HCl > HF
a) F- is small, more concentrated charge, holds on to H+
b) I- is large, less concentrated charge, gives up H+
2) Electronegativity of A-: HF > H2O > NH3 > CH4
3) Resonance Forms of A-
C. Lewis Acids and Bases1) Lewis Acid = electron pair acceptor = Electrophile2) Lewis Base = electron pair donor = Nucleophile3) Some covalently bonded molecules can be considered Lewis Acid/Base
pairs
4) Dissociation of a Lewis Acid/Base Pair (Mechanisms)
NH3 BCl3+ H3N BCl3
C CH3
CH3
Cl
CH3
C CH3
CH3
CH3
+Cl- +
OH2C CH3
CH3
H2O+
CH3
H++ C CH3
CH3
HO
CH3