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•E X A M C L A R I F I C A T I O N / F R I D A Y
DEADLINE
•S E T C L I C K E R T O C H A N N E L 4 1
Bases
 Strong bases are hydroxide salts
Strong Bases
Weak Bases
LiOH, lithium
hydroxide
NH3, ammonia
NaOH, sodium
hydroxide
KOH,
potassium
hydroxide
 For now, only important weak
base is NH3

Ionization of weak acid produces a weak base
Ca(OH)2,
calcium
hydroxide
Ba(OH)2,
barium
hydroxide
If the oxalate anion (C2O4-) reacts in an acid-base reaction, which of the
following can’t it make?
1. H2C2O4
2. HC2O4-
53%
3. 2CO2
28%
19%
1
2
3
Acid Base Reactions
Acid Base Reactions
 Strong Acid + Strong Base
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
acid
base
“salt”
 What do we get if we mix:
HBr (aq) +
LiOH (aq) 
water
Acid Base Reactions
 Diprotic acids or bases
H2SO4(aq) + NaOH(aq) 
H2SO4(aq) + Ba(OH)2(aq) 
HCl(aq) + Ba(OH)2(aq) 
Acid Base Reactions
 Strong Acid + Weak Base
HCl(aq) + NH3(aq) 
 What do we get if we mix:
HNO3(aq) + NH3(aq) 
Acid Base Reactions
 Weak Acid + Strong Base (like strong acid+strong base)
HCN(aq) + NaOH(aq)  NaCN(aq) + H2O(l)
acid
base
“salt”
 What do we get if we mix:
HCOOH (aq) + KOH (aq) 
formic acid
water
Net Ionic Equations
1.
Write a balanced chemical equation
 Molecular equation
1.
Write out all the ions  Total ionic equation
2. Cancel out anything that appears on both sides
 Net ionic equation
Net Ionic Equations
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
What really happens:
H+(aq) + OH-(aq)  H2O(l)
Sodium ion and chloride ion are “spectator ions”
Reactions Involving a Weak Base
 Molecular equation:
HCl(aq) + NH3(aq)  NH4Cl(aq)
 Total ionic equation:
H+(aq) + Cl-(aq) + NH3(aq)  NH4+ + Cl-(aq)
 Net ionic equation:
H+(aq) + NH3(aq)  NH4+ (aq)
 What is the net ionic equation for:
HNO3(aq) + NH3(aq)  NH4NO3(aq)
CH3CO2H(aq) + NaOH(aq) 
CH3CO2H2+(aq) + NaO(aq)
2. CH3CO2-(aq) + H2O(l) + Na+(aq) 33%
3. CH4(g) + CO2(g) + H2O(l)
1.
1
33%
2
33%
3
20
HCN(aq) + NH3(aq) 
NH4+(aq) + CN-(aq)
2. H2CN+(aq) + NH2-(aq)
3. C2N2(s) + 3 H2(g)
1.
33%
1
33%
2
33%
3
20
The pH Scale
 Quantitative measure of solution acidity
 Remember solution concentration:


#
moles
solute
Molarity

1
L
solvent
[NaCl]=0.25M means 0.25 moles of NaCl are in 1L of solution
The pH Scale
 In pure water, some molecules ionize to form H3O+ and
OH-
H2O + H2O  OH– + H3O+
 In acidic and basic solutions, these concentrations are
not equal
acidic: [H3O+] > [OH–]
basic: [OH–] > [H3O+]
neutral: [H3O+] = [OH–]
The pH Scale
 pH scale= measure of [H3O+]
pH < 7.0 = acidic
pH > 7.0 = basic
pH = 7.0 = neutral
 Measure of H3O+ concentration
(moles per liter) in a solution
 As acidity increases, pH decreases
The pH Scale
 The pH scale is logarithmic
100 102
log(102) = 2
10 101
log(101) = 1
1
100
log(100) = 0
0.1 10–1 log(10–1) = –1
0.01 10–2 log(10–2) = –2
 pH = –log [H3O+]
 pH if [H3O+] = 10–5? 10–9?
Acidic or basic?
 pH if [H3O+] = 0.000057 M?
Finding [H3O+] from pH
 [H3O+] = 10-pH
or
[H3O+] = log-1 (-pH)
 Finding the inverse log (or log -1)of a number on your
calculator:
Enter the number, press the inverse (inv) or shift button, the press
the log button (it might be labeled 10x)
 What is [H3O+] if pH = 8.6?
pH: Quantitative Measure of Acidity
 Acidity is related to concentration of H+ (or H3O+)
 pH = -log[H3O+]
 [H3O+]=10-pH=log-1(-pH)