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Chapter 5 Types of Chemical Reactions

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Chapter 5 Types of Chemical Reactions Powered By Docstoc
					  Chapter 4: Types of Chemical
           Reactions

Goals:
• To be able to predict chemical reactivity.
• To know how to synthesize specific
  compounds.
            Types of Reactions
•   Acid-Base: proton-transfer
•   Oxidation-Reduction: electron-transfer
•   Precipitation: formation of insoluble salts
•   Gas Forming
•   Organic:
    – Substitution
    – Addition
    – Elimination
  Reactions in Aqueous Solution

Unless mentioned, all reactions studied this
 and next week occur in aqueous solution.
                   Electrolytes

Strong Electrolytes: solute breaks apart to give ions in solution.
       NaCl  Na+ + Cl-


Weak Electrolytes: solute partially breaks apart to give ions.
       CH3CO2H  CH3CO2- + H+              happens less than 5%


Nonelectrolytes: no ions formed.
       CH3CH2OH
      Brønsted-Lowery Acid-Base
              Definitions

• An acid is a substance that donates a proton (H+)
  to a base
• A base is a substance that accepts a proton (H+)
  from an acid
     Brønsted-Lowery Definitions
• acid: donates a proton (H+) to a base
• base: accepts a proton (H+) from an acid
• Acid-base reactions can be reversible:
  reactants  products or
  products  reactants
     Brønsted-Lowery Definitions
• An acid is a substance that donates a proton (H+)
  to a base
• A base is a substance that accepts a proton (H+)
  from an acid
• Acid-base reactions can be reversible:
  reactants  products or
  products  reactants
          Important Acids and Bases
Strong Acids:
                               Strong Bases:
HCl             hydrochloric
                               LiOH            lithium hydroxide
HBr             hydrobromic
                               NaOH            sodium hydroxide
HI              hydroiodic
                               KOH       potassium hydroxide
HNO3            nitric
                               Ca(OH)2         calcium hydroxide
H2SO4           sulfuric
                               Ba(OH)2         barium hydroxide
HClO4           perchloric
                               Weak Base:
Weak Acid:
                               NH3             ammonia
CH3CO2H         acetic
Any other acids are WEAK
STRONG acids in water: 100% of acid molecules form
  ions:
        HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq)




                                       H3O+ is
                                        hydronium ion
WEAK acids in water, ~5% or less of acid molecules
 form ions (acetic, H3PO4, H2CO3)
          Polyprotic Acids
       multiple acidic H atoms

H2SO4  H+ + HSO4-
HSO4-  H+ + SO42-



Not all H’s are acidic:
        CH3CO2H
If H3PO4 reacts as an acid, which
of the following can it not make?

•   1. H4PO4+
•   2. H2PO4-
•   3. HPO42-
•   4. PO43-
       2-
If C2O4 reacts in an acid-base
reaction, which of the following
can it not make?
• 1. H2C2O4
• 2. HC2O4-
• 3. 2 CO2
            Acid-Base Reactions

Strong Acid + Strong Base

  HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
  acid       base       “salt”   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)  NH4Cl(aq)
           Acid-Base Reactions

Weak Acid + Strong Base

  HCN(aq) + NaOH(aq)  NaCN(aq) + H2O(l)
  acid      base       “salt”     water
          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 weak bases
HCl(aq) + NH3(aq)  NH4+(aq) + Cl-(aq)



Net-Ionic Equation:

NH3(aq) + H+(aq)  NH4+(aq)
CH3CO2H(aq) + NaOH(aq) 
• 1. CH3CO2H2+(aq) + NaO(aq)
• 2. CH3CO2-(aq) + H2O(l) + Na+(aq)
• 3. CH4(g) + CO2(g) + H2O(l)
HCN(aq) + NH3(aq) 
• 1. NH4+(aq) + CN-(aq)
• 2. H2CN+(aq) + NH2-(aq)
• 3. C2N2(s) + 3 H2(g)
Solution Concentration: Molarity
• Molarity = moles solute per liter of solution

• 0.30 mol NH3 dissolved in 0.500 L

  Concentration =
• Written like: [NH3] = 0.60 M
         pH Scale
• In pure water, a few 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–]
             pH Scale
• Measure how much H3O+ is in a
  solution using pH
• pH < 7.0 = acidic
• pH > 7.0 = basic
• pH = 7.0 = neutral
• Measure of H3O+ and OH–
  concentration (moles per liter) in a
  solution
• As acidity increases, pH decreases
           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 [H3    O +]     from pH
[H3O+] = 10-pH

What is [H3O+] if pH = 8.9?
     pH: Quantitative Measure of
              Acidity

• Acidity is related to concentration of H+
  (or H3O+)

• pH = -log[H3O+]

				
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