Acid Properties
1. Three standard definitions (Arrhenius, BronstedLowry, and Lewis) but generally increase [H+] in a
solution
2. Tastes sour when dissolved in water
 do not taste to test chemicals
3. Reacts with metals to make H2(g)
 metals above H2 in activity series
 single replacement reaction
 Mgs) + H2SO4(aq) → MgSO4(aq) + H2(g)
Acid Properties
4. Turns litmus (an indicator) red
5. Reacts with bases to make salt and water in a
neutralization reaction.
 HBr(aq) + KOH(aq) → H2O(l) + KBr(aq)
6. Conducts electricity (electrolytes)
 strong acids conduct very well
 weak acids conduct a small amount
Naming Binary Acids
Think of HCl - “hydrochloric acid”
 Contain H and a nonmetal
 Always has “hydro-” prefix
 Root of other element’s name
 Ending “-ic”
 Examples: HI, H2S, HBr
Naming Ternary or Oxy- Acids
Think of HNO3 – “nitric acid”
 Contains H and a polyatomic anion
 Most have an oxygen containing polyatomic ion
 No prefix
 Name of polyatomic ion
 Ending “–ic” for “-ate” and
“–ous” for “-ite”
 Examples: HClO4, H3PO4, HNO2
Common Acids
 Sulfuric acid: H2SO4
 used in car batteries
 Nitric acid: HNO3
 used in explosives
 Phosphoric acid: H3PO4
 used as flavoring in soft drinks
 Hydrochloric acid: HCl
 stomach acid
 Acetic acid: HCH3COO or HC2H3O2
 vinegar contains it
Strengths of Acids
 Strong Acids (100% dissociation or ionization)
HCl, HBr, HI, HClO4, HNO3 , H2SO4
 Memorize this list!
 Weak Acids
 Only partially ionize
 Weak electrolytes
 If it’s not strong ….. it’s weak!
STRONG
WEAK
WEAK
Strengths of acids
 Depends on polarity and strength of the bond holding
the H to the rest of the molecule
 Strong acids ionize completely - good electrolytes
 Strength of Binary Acids
HF << HCl < HBr < HI
Decreasing bond strength …
more likely to ionize
 Strength of Oxyacids
HOCl < HOClO < HOClO2 < HOClO3
Increasing polarity … water
removes H+
The “-ic” acid is always stronger than the “-ous” acid!
Base Properties
1. [OH-] > [H+]
2. Tastes bitter
 most are caustic (cause burns) so do not taste to test
3. Feels slippery when dissolved in water
 do not touch to test
4. Turns litmus paper blue
5. React with acids to make salt and water in
neutralization reactions.
6. Conducts electricity
Common Bases
 Contains hydroxide ion:
 NaOH
 KOH
 Ca(OH)2
 Contains amine group (-NH2):
 NH3
 CH3NH2
Strengths of Bases
 Strong Bases (100% dissociation)
Alkali metal hydroxides and
Ca2+, Sr2+and Ba2+ hydroxides
 Strong electrolytes
 Weak Bases
 Weak electrolytes
 Examples: NH3, C6H5NH2
ACID-BASE DEFINITIONS
Arrhenius Acids and Bases
 Acid : Compound creates H+ in an aqueous solution
 Base : Compound creates OH- in an aqueous solution
Acid: HNO3 → H+ + NO3Base: NaOH → Na+ + OH Most specific/exclusive definition
 Created by Svante Arrhenius, Swedish
Acids with multiple Hydrogens
 Monoprotic acids
 Only have one acidic proton (a H+ that can be removed)
HNO3 → H+ + NO3-
 Polyprotic acids - diprotic, triprotic (not necessarily
stronger)
 Contains more than one acidic proton
H3PO4 → H+ + H2PO4H2PO4 - → H+ + HPO42HPO4 2- → H+ + PO43-
Bronsted-Lowry Acids and Bases
 Acid: Molecule or ion that is a H+ donor
 Base: Molecule or ion that is a H+ acceptor
HCl + H2O  H3O+ + Cl Less specific definition (encompasses the Arrhenius
definition)
Conjugate Pairs
Identifying the Conjugate Pairs
 An acid-base conjugate pair differs by 1 H+ ion.
 The acid in the pair has 1 more H+ ion.
 The base in the pair has 1 fewer H+ ion.
Base
Acid
Acid
Base
Identify the Acid-Base Pairs
HCl + OH-  Cl- + H2O
H2O + H2SO4  HSO4- + H3O+
Strength of
Conjugate Pairs
 If one of the pair is strong,
the other MUST be weak.
 The conjugate base of a strong acid must be weak base.
 The conjugate base of a weak acid must be a strong base.
NH3 + HNO3  NH4+ + NO3We know that NH3 is a weak base, therefore, NH4+ must be..
We know that HNO3 is a strong acid, therefore, NO3- must be..
Strength of Acids and Bases
NH3 + HNO3  NH4+ + NO3NH3 and NO3- are both weak bases
HNO3 and NH4+ are both strong acids
The K value will tell you which of the bases is stronger or
weaker, and which of the acids is stronger or weaker.
The Kb value of NH3 is 1.8 x 10-5. Any K value less than 1
indicates that the reactants are favored over products.
Practice Problem
 Given the equation:
HSO4- + H2O  OH- + H2SO4
 Identify the conjugate acid-base pairs.
 Determine the strongest acid and strongest base if Keq
value of this equation is 0.000004.
 Label each species as SA, SB, WA, WB.
Amphoterism
 A species that can act as an acid or a base is said to be
amphoteric or amphiprotic.
HCO3- + NH3  CO3-2 + NH4+
HCO3- + HCl  Cl- + H2CO3
 Could NH4+ ever be a Bronsted Lowry base? Explain.
 Could Cl- ever be a Bronsted Lowry acid? Explain.
Lewis Acids and Bases
(most general or broad definition)
 Acid: atom, ion, or molecule that accepts electron pair to
form covalent bond (electron acceptor)
 Base: atom, ion or molecule that donates and electron pair
to form covalent bond (electron donor)
 The most general and least traditional - used more in
biology
Most Popular Example
 Identify the Lewis acid and Lewis base
Electron
Acceptor
(acid)
Electron
Donor
(base)
 Identify the Lewis acid and Lewis base
Electron
Donor
(base)
Electron
Acceptor
(acid)
 These are usually formed from a transition metal
surrounded by ligands (polar molecules or negative
ions).
 As a "rule of thumb" you place twice the number of
ligands around an ion as the charge on the ion...
example: the dark blue Cu(NH3)42+ (ammonia is used
as a test for Cu2+ ions), and Ag(NH3)2+.
molecular
formula
Lewis
base/ligand
Lewis acid
donor
atom
coordination
number
Ag(NH3)2+
NH3
Ag+
N
2
[Zn(CN)4]2-
CN-
Zn2+
C
4
[Ni(CN)4]2-
CN-
Ni2+
C
4
[PtCl6] 2-
Cl-
Pt4+
Cl
6
[Ni(NH3)6]2+
NH3
Ni2+
N
6
Heme Group
 The Fe ion in hemoglobin is a
Lewis acid
 O2 and CO can act as Lewis
bases
Practice
Example 1
 Identify the Lewis acid and base.
Ag(NH3)+ (aq) + NH3 (aq)  Ag(NH3)2+
Electron
Acceptor
(acid)
Electron
Donor
(base)
(aq)
Practice
Example 2
 Excess hydrochloric acid is added to a solution of
cobalt(II) nitrate to produce a coordination complex.
Co2+ (aq) + 4 Cl- (aq)  [CoCl4]2-
(aq)
Neutralization Reactions
 When an Arrhenius acid and an Arrhenius base react, the
products are a salt and water
 A special type of double replacement reaction.
HCl + KOH → KCl + H2O
acid
base
salt
water
Identify the Products
 HNO3 + KOH 
 HCl + Mg(OH)2 
 H2SO4 + NaOH 
 HBr + Ba(OH)2 
Stoichiometric Problems
 How many moles of HNO3 are need to neutralize 0.86
moles of KOH?
 How many moles of HCl are needed to neutralize 3.5
moles of Mg(OH)2 ?
Neutralization Problems
MAVA = MBVB
 If polyprotic acid or polyhydroxic base
(x)MAVA = (y)MBVB
x = # of H+ in polyprotic acid (2 for H2SO4)
y= # OH- in polyhydroxic base [2 for Ca(OH)2]
Neutralization Problems
 If it takes 87 mL of an HCl solution to neutralize 0.67
moles of KOH what is the concentration of the HCl
solution?
 If it takes 58 mL of an H2SO4 solution to neutralize
0.34 moles of NaOH what is the concentration of the
H2SO4 solution?
Titration – Volumetric Analysis
 Using a certain volume of a solution
with a known concentration (titrant) to
a certain amount of solution being
analyzed (analyte)
 Uses neutralization of an acid or base
of unknown molarity by the opposite
(acid or base) of known molarity to
determine the unknown molarity
Titrations
 equivalence point- when there are equal moles of each
reactant present
 indicator- substance that changes color near
equivalence point to tell when neutralization has
occurred
 end point- when indicator changes color
Practice Problem
 To standardize a sodium hydroxide solution, a student
plans to titrate the solution with a strong monoprotic
acid, KHP (KHC8H4O4, potassium hydrogen
phthalate). She dissolves 1.30 g of KHP in water, adds
phenolphthalein, and titrates to the endpoint. (gramformula mass = 204 g/mol)
It took 41.2 mL of sodium hydroxide to titrate the KHP
solution.
Calculate the concentration of NaOH.
Example 4
 Write the balanced net ionic equation:
NaOH(aq) + KHP(aq) → NaKP(aq) + H2O(l)
Na+(aq) + OH-(aq) + KHP(aq) →
Na+(aq) + KP-(aq) + H2O(l)
OH-(aq) + KHP(aq) → H2O(l) + KP-(aq)
Example 4
 Find the moles of H+ in the KHP:
1molKHP
1molH 
1.30gKHP 

 0.00637mol H 
204gKHP 1molKHP
 Find the moles of NaOH used:

1molOH
1molNaOH

0.00637mol H 

 0.00637mol NaOH


1molH
1molOH
 Find molarity of NaOH used:
0.00637mol NaOH
NaOH 
 0.155 M NaOH
0.0412L
Example 5
A 0.3518 g mixture of CCl4 and HC7H5O2 was mixed
with water and titrated by 10.59 mL of 0.1546 M
NaOH.
Find the percent by mass of HC7H5O2 in the original
sample.
Example 5
 Write the net ionic equation
HC7H5O2(aq) + OH-(aq) →
H2O(l) + C7H5O2-(aq)
 Find moles of OH- used:
0.1546molN aOH 1molOH 
0.01059L 

 0.001637mo lOH 
1L
1molNaOH
Example 5
 Find moles of HC7H5O2 used:

1molHC 7 H 5 O 2
1molH

0.001637mo lOH 

 0.001637mo lHC 7 H 5 O 2


1molOH
1molH
 Find the grams of HC7H5O2:
122.2gHC 7 H 5 O 2
0.001637mo lHC 7 H 5 O 2 
 0.1999gHC 7 H 5 O 2
1molHC 7 H 5 O 2
 Find percent by mass:
0.1999gHC 7 H 5 O 2
%
 56.83%HC 7 H 5 O 2
0.3518gtotal