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CHEM 1133 Study Guide - Midterm Guide: Buffer Solution, Rice Chart, Conjugate Acid

Course Code
CHEM 1133
Ruth Heisler
Study Guide

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Chemistry 2 Final Exam Review Exam 2
Chapter 18: Acid-Base Equilibria
Three acid-base reactions:
o Acid + water H3O+ + conj. base (Ka)
o Base + water OH- + conj. acid (Kb)
o Acid + base salt + water
Equilibrium: [reactants] and products] doesn’t change
o Rate forward rxn = rate reverse rxn
o G = 0, reach minimum in free energy
o Equilibrium constant K= 
 can also use partial pressures
o Solving equilibrium problems use K (or Q) and ICE tables
Manipulations of K (and Q):
o If the rxn is reversed: K’ = 1/K
o If the rxn is multiplied by a coefficient n: K’ = Kn
o If two rxns are added together: K’ = K1 x K2
o If two rxns are subtracted: K’ = K2 / K1
Arrhenius acids and bases: aqueous solutions of acids or bases conduct electricity ions present in solution
o Acids: ionize in water to product H+ ions (protons)
o Bases: have OH in formula and ionize in water to produce OH- hydroxide ions
Strong and weak acids:
o A strong acid dissociates completely into ions in water
A dilute solution of a strong acid contains no HA molecules
o A weak acid dissociates slightly to form ions in water
A dilute solution of a weak acid contains has most HA molecules undissociated
o Strong acids: HCl, HBr, HI, HNO3, HClO4, H2SO4
Strong and weak bases:
o Strong bases completely dissociate in water
Group 1 and 2 hydroxides
o Weak bases only slightly dissociate in water
Most weak bases are amine bases (contain nitrogen)
Ka acid dissociation reaction equilibrium constant
o Reactions of acids always have this form: HA + H2O A- + H3O+
o 
o Stronger acid = larger Ka value
Bronsted-Lowry acid-base definition:
o An acid is a proton donor, any species that donates an H+ ion
An acid must contain H in its formula
o A base is a proton acceptor, any species that accepts an H+ ion
A base must contain a lone pair of electrons to bond H+
o An acid-base rxn is a proton-transfer process
Autoionization of water: H2O + H2O OH- + H3O+
o Kw = [H3O+] [OH-] = 1x10-14
Ka x Kb = Kw
At 25°C: [H3O+] =[OH-] = 1x10-7
pH: the [H3O+] in aq solution commonly covers a wide range
o pH = -log[H3O+]
o As pH increases, acidity decreases
pOH: pH + pOH = 14
o pKw = 14
pKa = -logKa
o A low pKa corresponds to a high Ka
Conjugate acid-bases:
o Conjugate acid-base pairs differ by one proton (H+)
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Chemistry 2 Final Exam Review Exam 2
o Generally: the stronger the acid, the weaker its conjugate base
Conjugate bases of weak acids are weak bases
Conjugates bases of strong acids are not considered bases
Weak acids: weak acids and bases do NOT dissociate completely so need to use Ka with ICE tables to calculate concentrations
Percent ionization: % ionization of a weak acid is the ratio of ionized acid concentration to the initial acid concentration, time
o   
 
o The % ionization of a weak acid decreases with increasing concentration of the acid
o Adding water decreases all concentrations but increases the ration of [A-]/[HA]0
Using pH to determine Ka
o Make ICE table
o Use pH to calculate [H3O+] and plug that into equilibrium on ICE table
o Calculate Ka
Polyproptic acids: contain more than one acidic proton
o Only lose one proton at a time
o Remaining protons are harder to remove
o Ka1 is always larger than Ka2 as it becomes a weaker acid and becomes harder to remove successive H+
o To find the pH: use Ka1, Ka2 is so small it is negligible
Predicting the direction of equilibrium
o Weaker acid + weaker base stronger acid + stronger base
o In acid-base rxns, equilibrium favors the weaker acid
Acidity increases as:
o Bond strength decreases (down periodic table)
o Electronegativity increases (across periodic table)
o Number of O atoms increases
Salt solutions
o Cation strong base + anion strong acid neutral solution
o Cation weak base + anion strong acid acidic solution
o Cation strong base + anion weak acid basic solution
o Cation strong base + anion polyproptic acid pH depends on Ka or Kb
Anions that come from polyproptic acids are amphiprotic can act as an acid or base, if Kb > Ka solution will
be basic, or vice versa
Lewis Acid-Base definition:
o Lewis acid: any species that accepts an electron pair to form a bond
Molecules that contain a polar multiple bond often function as Lewis acids
A metal cation acts as a Lewis acid when it dissolves in water to form a hydrated ion (positive ion)
o Lewis base: any species that donates an electron pair to form a bond
Negative ion
o Lewis definition views an acid-base rxn as the donation and acceptance of an electron pair to form a covalent bond
o A buffer works through the common-ion effect
If a compound is added with an ion that is common to both solutions, the equilibrium will shift away from it
Acid-base buffers
o An acid-base buffer is a solution that lessens the impact of pH from the addition of acid or base
o Usually consists of a conjugate acid-base pair where both species are present in appreciable quantities in solution
o Therefore, it is a solution of a weak acid and its conjugate base, or vice versa
o How they work:
Their relative changes in the amounts of the buffer components are small
Their concentration ratio changes by very little
pH changes by very little
Quantitative calculations of buffers
o When there is both acid and conjugate base in the equilibrium line of an ICE table use the Henderson-Hasselbalch
o    
Buffer capacity: ability of a buffer to maintain a constant pH
o Buffer capacity is higher in solutions containing: higher concentrations of the buffer components
o Equimolar ([HA]=[A-]) amount of the buffer components
Buffer range: an acid/base pair can buffer over a pH range of pKa +/- 1
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