AS.030.103 Midterm: Chem 2 Midterm 1 Notes (Equilibrium Constant, Acids and Bases, Le Chatelier, Buffer Solutions, Titrations, Precipitation Reaction, Solubility, Electrochemistry, Redox Reactions, Rate Law, Reaction Half Life)
Lecture 1:
Equilibrium Constant (K):
aA + bB
!
cC + dD
K =
"#$%"&$'
"($)"*$+
*products over reactants!
DGo =
,-. /012345
Acid-Base Concepts:
Types:
• Arrhenius: acid = substance yields H+, base = substance yields OH- [H+ + OH-
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H2O]
o Does not explain how substances like NH3 dissolve in water to form basic
solutions (does not contain OH)
o Does not explain acid-base reactions that do not take place in aqueous
solution
• Brønsted-Lowry: acid = proton (H+) donor, base = proton acceptor [HCl(aq) +
NH3(aq)
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NH4+(aq) + Cl-(aq)]
• Lewis: acid = electron pair acceptor, base = electron pair donor [A + :B
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A:B]
o H+ accepts electron pair from bases, such as OH-
o Need not involve H+
o
• Arrhenius is special case of Brønsted-Lowry, Brønsted-Lowry is special case of
Lewis
Lecture 2: ***FIND LE CHATELIER PRINCIPLE/LAST SEM. CHEAT SHEET
Autoionization of Water:
H2O(l) + H2O(l)
!
H3O+(aq) + OH-(aq)
Kw = [H3O+][OH-] = 1.0 x 10-14 at 25oC
Concentration of H3O+ and OH- in pure H2O:
[H3O+] = [OH-] = y
y2 = 1 x 10-14
y = 1.0 x 10-7
Strong Acids and Bases:
Strong acid dissociates completely in water
Leveling effect: both reactions go almost completely to the right, so there is little to no
difference in acid “strength” for HCl and HClO4 (difference not apparent in water
as a solvent)
[H3O+] for a 0.1M solution of any strong acid = 0.1M
[OH-] for a 0.1M solution of any strong acid:
[H3O+][OH-] = Kw = 1.0x10-14
[OH-] = Kw/[H3O+] = 1.0x10-14/0.10 = 1.0x10-13
*amount of H3O+ from autoionization of water is neglected; too small when strong acid
present
Weak acids and bases:
Weak acid (ie acetic acid):
Weak base:
The pH function:
pH < 7: acidic
pH = 7: neutral
pH > 7: basic
[H3O+] = 10-pH
[H+] = 10-pH
pH of 10-12 M HCl = 7 *amount of HCl is less than amount of water by magnitude
of 5 (basically nothing compared to the water)
Buffer Solutions:
Function: to resist changes in the pH of a solution
HCl + H2O vs. HCl + (H2PO4- + HPO42-) <- buffer solution
Buffers made from weak acid + conj base (ie HOAc + OAc-, H2PO4- + HPO42-, NH4+ + NH3)
Henderson-Hasselbalch Equation:
pH largely determined by pKa, but is then adjusted by the ratio of weak acid and
conjugate base
*This is only valid for weak acids!!!
[H3O+] depends on the ratio of weak acid and conj base
For buffer to work, [HA] and [A-] must be NEARLY EQUAL and LARGE compared
to any acid or base added to solution
Addition of acid/base consumes a little HA or A-, but the ratio of [HA]/[A-]
remains the same
Buffer is MOST effective when:
acid/base ratio = 1
0.1 < [base]:[acid] < 10
[acid] and [base] are large
Titration:
Titrant: solution of unknown concentration
Titrant slowly added to solution of known concentration from burette until reaction is
complete
Indicator: chemical that changes color when pH changes
Equivalence point: moles of H3O+ = moles of OH-
AKA neutralization reaction
H+ + OH-
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H2O
Lecture 3:
Acid-Base Titration:
Indicator color changes at equivalence point, or titration’s endpoint (enough base/acid
has been added to neutralize all of the acid/base)
***KNOW THE TITRATION CURVES/GRAPHS***
Strong acid titrated with a strong base:
Phenolphthalein: indicator example
pH at halfway point: [HA] = [A-]
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[H3O+] = Ka
6
pH = pKa
Lecture 4:
Precipitation Reaction:
Add cation to anion, forms immediate precipitate that is INSOLUBLE
Predicting Precipitation:
Calculate initial concentrations
Calculate initial reaction quotient (Q)
Check if Q > Ksp, Q = Ksp, or Q < Ksp:
