PP 21: Acids & Bases
Drill: Rxn: 1 A + 2 B  1 C + 1 D
. (Kc = 2.0 x 10-12)
Calculate the equilibrium concentrations of each species when a solution is made with 1.0 M A & 1.0 M B
Acids & Bases
Properties of Acids: Sour taste, Change color of dyes, Conduct electricity in solution,
React with many metals, React with bases to form salts
Properties of Bases: Bitter taste, Feel slippery, Change color of dyes, Conduct electricity in solution,
React with acids to form salts
Types of acids & bases:
Arrhenius acids & bases
 Acids: release H+ or H3O+ in solution
 Bases: release OH- in solution
Bronsted/Lawry acids & bases:
 Acid: Proton donor

Base: Proton Acceptor
General
HA  H+ + AMOH  M+ + OH-
General
HA + H2O  H3O+ + AAcid base
CA
CB
NH3 + H2O  NH4+ + OHBase
acid
CA
CB
Specific
HCl  H+ + ClKOH  K+ + OHSpecific
HI + H2O  H3O+ + IAcid base
CA
CB
NH3 + H2O  NH4+ + OHBase
acid
CA
CB
Lewis Acids & Bases:


BF3 + NH3  H3N-BF3
Acid: Electron Acceptor
Base: Electron Donor
Common Names:
 H+
 H3O+
 H OH NH3
 NH4+
Hydrogen ion
Hydronium ion
Hydride ion
Hydroxide ion
Ammonia
Ammonium ion
Amphiprotism: Can act like an acid or a base or can donate or accept protons
Naming Acids:
 All acids are H-anion
o If the anion is:
 -ides 
 -ates 
 -ites 
hydro___ic acids
___ic acids
___ous acids
Naming bases:



Almost all bases are metal hydroxides
Name by normal method
Ammonia (NH3) as well as many amines are bases
Drill: Name each of the following: KOH
HBr
Al(OH)3
H2CO3
HClO4
Strong Acids or Bases:
 Strong acids or bases ionize 100 % in solution
 Weak acids or bases ionize <100 % in solution
Strong Acids:
 HClO 4
 H2SO4
 HNO3
 HCl
 HBr
 HI
Perchloric acid
Sulfuric acid
Nitric acid
Hydrochloric acid
Hydrobromic acid
Hydroiodic acid
Strong Bases
- Column IA meatl hydrocides
- Lower 3 of column II metal Hudrodrxides
Determining H+ or OH- concentrations in strong acids & bases:

0.10 M Strong Acids ionize 100%:
HA 
H+ + A0.10 – all 0.10 0.10

0.10 M Strong Bases ionize 100%:
MOH 
0.10 – all
M+ + OH0.10 0.10
Binary Acids: Acids containing only 2 elements
 HCl  hydrochloric acid
 HI  hydroiodic acid
Ternary Acids: Acids containing 3 elements
 H2SO4  sulfuric acid
 HNO3  nitric acid
Monoprotic Acids: Acids containing only 1 ionizable hydrogen
 HBr  hydrobromic acid
 HC2H3O2  acetic acid
Polyprotic Acids: Acids containing 2 or more ionizable hydrogen
 H2SO4  sulfuric acid
 H3PO4  phosphoric acid
Diprotic Acids: Acids containing 2 ionizable hydrogen
 H2SO4  sulfuric acid
 H2CO3  carbonic acid
Triprotic Acids: Acids containing 3 ionizable hydrogen
 H3PO4  phosphoric acid
 H3AsO4  arsenic acid
Monohydroxic Acids: Acids containing only 1 ionizable hydroxide
 KOH potassium hydroxide
 NaOH  sodium hydroxide
NH3
Neutralization Reaction: A reaction between an acid & a base making salt & H2O
· General Example:
HA(aq) + MOH(aq)  MA(aq) + H2O(l)
· General Example:
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
Drill: Identify the acid, the base, the conjugate acid, & the conjugate base in the following reaction:
General example:
HCO3- + H2O  H2CO3 + OH-
pH: The negative log of the hydrogen or hydronium ion concentration


pH = -log[H+]
pOH = -log[OH-]
Calculate the pH of each of the following:
 [H+] = 0.040 M
 [HCl] = 0.0025 M
 [HBr] = 0.080 M
Calculate the pOH of each of the following:
 [OH-] = 0.030 M
 [KOH] = 0.0025 M
 [NaOH] = 4.0 x 10-7 M
Standard Solution: A solution with known concentration
Titration: A method of determining the concentration of one solution by reacting it with a standard solution
Titration Fact:
 When titrating acids against bases, the end point of the titration is at the equivalence point
 No changes will be observed when titrating acids against bases; thus, one must use an
indicator to see changes
Equivalence Point: The point where the concentrations of the two solutions in the titration are equal
Acid/Base Equivalence Point: The point where the H+ concentration is equal to the OH- concentration
Indicator: An organic dye that changes color when the pH changes
Titration Formula for monoprotic solutions: MAVA = MBVB
Molarity (M): Moles of solute per liter of solution
Problem: Calculate the molarity of 25.0 mL HCl when it’s titrated to its equivalence point with
50.0 mL 0.200 M NaOH:
Dilution Formula: M1V1 = M2V2
Problems:
Calculate the mL of 16.0 M HNO3 it takes to make 4.0 L of 0.100 M HNO3

Calculate the mL of 12.5 M HCl required to make 2.5 L of 0.200 M HCl

Normality (N): Number of moles of hydrogen or hydroxide ions per liter of solution
Titration Formula for Acid/Base: NAVA = NBVB
Elliot’s Rule: #HMAVA = #OHMBVB
Problems:
 Calculate the molarity of 30.0 mL H2CO3 when it’s titrated to its equivalence point with
· 75.0 mL 0.200 M NaOH
 Calculate the molarity of 40.0 mL H3PO4 when it’s titrated to its equivalence point with
30.0 mL 0.20 M Ba(OH)2
 Calculate the volume of 0.250 M HCl needed to titrate 50.00 mL 0.200 M NaOH
to its equivalence point
 Calculate the molarity 25.0 mL H3PO4 that neutralizes 50.00 mL 0.200 M Ca(OH)2
to its equivalence point
Titration Curve: Strong acid vs strong base
Titration Curve: Strong acid vs strong base