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Ammonia Reaction NH3 + H2O = NH4

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					Classification of Reactions

    Acid-Base Reactions

   Precipitation Reactions

Oxidation-Reduction Reactions
Electrolytes and Non
    Electrolytes

 • Strong Electrolytes
 • Weak Electrolytes
 • Non Electrolytes
Polarity of Water
Solvation of Ions
Because of Solvation by water, many
ionic compounds dissociate (ionize)
into their constituent ions when
dissolved in water



Compounds that “freely” ionize into independent ions
in aqueous solution are called electrolytes because
their aqueous solutions are capable of conducting an
electric current.
Some molecular compounds also dissociate
  into ions when they react with water.

                                    
  HCl(aq)  H (aq)  Cl (aq)
   Since the resulting solution is electrically
   conducting, the molecular substance is also
   classified as an electrolyte.
  Some molecular compounds dissolve
  but do not dissociate into ions.

C6 H12O6 (s) (glucose)  C6 H12O6 (aq)
                           H 2O


                         
C2 H 5OH(l ) (ethanol)   C2 H 5OH(aq)
                             H 2O



    These compounds are referred to as
    non-electrolytes since they dissolve in water to
    give a non-conducting solution.
   A weak electrolyte is an electrolyte that
   dissolves in water to give a relatively small
   percentage of ions.


           NH  (aq)  OH  (aq)
NH4OH(aq)    4



• Most soluble molecular compounds are either
   nonelectrolytes or weak electrolytes.
Acids and Bases
• The Arrhenius Concept

 – The Arrhenius concept defines acids as
   substances that produce hydrogen ions, H+,
   when dissolved in water.
 – An example is nitric acid, HNO3, a molecular
   substance that dissolves in water to give H+ and
   NO3-.

                                           
             
 HNO 3 (aq )  H (aq )  NO 3 (aq )
                 H 2O
• The Arrhenius Concept

 – The Arrhenius concept defines bases as
   substances that produce hydroxide ions, OH-,
   when dissolved in water.
 – An example is sodium hydroxide, NaOH, a
   substance that dissolves in water to give OH- and
   Na+.

                H2O
      NaOH(s)          Na+(aq) + OH-(aq)
• The Arrhenius Concept

   The molecular substance ammonia, NH3, is a
   base in the Arrhenius view, because it yields
   hydroxide ions when it reacts with water.
   Because the reaction of NH3 with water only
   reacts (at equilibrium) to a small extent (~5%) and
   produces relatively small amounts of OH-, it is
   described as a weak base (i.e. it is a weak
   electrolyte) water.

                       
                                 
NH 3 (aq )  H 2O(l )  NH 4 (aq )  OH  (aq )
• The Arrhenius Concept
   Acetic acid, HC2H3O2, is an acid in the Arrhenius
   view, because it dissociates to produce H+ ions
   when it reacts with water.
    However, HC2H3O2, is a weak acid because only a
    small amount (~4%) of the acetic acid molecules
    ionize to produce H+ ions. i.e. it is a weak
    electrolyte in water. Most of the acetic acid
    remains in its unionized molecular form.
                                        
                    
     HC2 H 3O2 (l )  H (aq)  C2 H 3O2 (aq)
• The Brønsted-Lowry Concept

– The Brønsted-Lowry concept of acids and bases
  involves the transfer of a proton (H+) from the acid
  to the base.
– In this view, acid-base reactions are proton-
  transfer reactions.

– Bronsted-Lowry Acid is Proton Donor
– Bronsted-Lowry Base is Proton Acceptor
          Bronsted-Lowry Acid

HNO3(aq) + H2O(I)  NO3-(aq) + H3O+(aq)




 Acid (1) + Base(1)  Base (2) + Acid (2)

HNO3 and H3O+ are acids (proton donors)
H2O and NO3- are bases (proton acceptors)
The hydronium ion, H3O+
• Which of the following is a strong
  electrolyte in aqueous solution?
• a. HF
• b. H2CO3
• c. Al(NO3)3
• d. HC2H3O2
• e. NH3
• Which of the following is a weak
  electrolyte in aqueous solution?
• a. lithium acetate
• b. ammonium carbonate
• c. acetic acid
• d. sulfuric acid
• e. calcium hydroxide
•   Which of the following is a strong acid?
•   a. ascorbic acid
•   b. sulfurous acid
•   c. lactic acid
•   d. acetic acid
•   e. hydrobromic acid
         Molecular and Ionic Equations

• A molecular equation is one in which the reactants
  and products are written as if they were molecules,
  even though they may actually exist in solution
  as ions.

Ca(OH )2 (aq)  Na2CO3 (aq)  CaCO3 (s)  2NaOH(aq)


     Ca(OH)2, Na2CO3, and NaOH are all soluble
     compounds but CaCO3 is not.
 • An ionic equation represents strong electrolytes as
   separate independent ions because they are fully
   ionized in solution.


Ca(OH )2 (aq)  Na2CO3 (aq)  CaCO3 (s)  2NaOH(aq)
Molecular Equation
                                              2
Ca 2 (aq )  2OH  (aq )  2Na  (aq )  CO 3 (aq ) 
  Ionic Equation      CaCO 3 (s )  2Na  (aq )  2OH  (aq )


     2                2
  Ca (aq)  CO3 (aq)  CaCO ( s)
                           3
 Net Ionic Equation
Acid-Base Neutralization Reactions
    Strong Acid + Strong Base
     Acid + Base = Salt + Water

     HCl + NaOH → NaCl + H2O

    HNO3 + KOH → KNO3 + H2O

            H+ + OH- → H2O

The net ionic equation is the same for all
  strong acid – strong base reactions
  Acid-Base Neutralization Reactions
       Weak Acid + Strong Base

        HF + NaOH → NaF + H2O

    HF + Na+ + OH - → Na+ + F- + H2O

            HF + OH- → F- + H2O

For weak electrolytes, the predominant (i.e.
non-ionized) species in solution is written
in the net-ionic equation
   Acid-Base Neutralization Reactions
        Strong Acid + Weak Base

          HBr + NH3 → NH4+ + Br-

        H+ + Br- + NH3 → NH4+ + Br-

               H+ + NH3 → NH4+

For weak electrolytes, the predominant
(unionized) species in solution is written
in the net-ionic equation
     Reaction of Bicarbonate With an Acid

    NaHCO3 + HCl → NaCl + H2O + CO2
    Molecular Equation



Na+ HCO3- + H+ Cl- → Na+ Cl- + H2O + CO2
 Ionic Equation



     HCO3- + H+ → + H2O + CO2
    Net Ionic Equation Equation
The net ionic equation for the reaction of nitrous acid with lithium hydroxide

•   a. HNO3(aq) + LiOH(aq) → LiNO3(aq) + H2O(l).

•   b. HNO3(aq) + LiOH(aq) → Li+(aq) + NO3–(aq) + H2O(l).

•   c. HNO2(aq) + OH–(aq) → NO2 –(aq) + H2O(l).

•   d. H+(aq) + NO2 –(aq) + Li+(aq) + OH–(aq) → Li+(aq) + NO2 –(aq) + H2O(l).

•   e. H+(aq) + OH–(aq) → H2O(l).
The net ionic equation for the reaction of weak acid, phosphoric acid,
with the strong base, potassium hydroxide, is

a.       H3PO4(aq) + 3 KOH(aq)  K3PO4(aq) + 3 H2O(l)

b.       H3PO4(aq) + 3 KOH(aq)  3 K+(aq) + PO43-(aq) + 3 H2O(l)

c.       H+(aq) + OH-(aq)  H2O(l)

d.       H3PO4(aq) + 3 OH-(aq)  PO43-(aq) + 3 H2O(l)
      Insolubility and Precipitation
Reaction of Lead(II) Nitrate and Potassium Iodide Solution

          Pb(NO3)2 + 2KI → PbI(s) + 2KNO3
          Soluble Soluble Insoluble Soluble
Pb(NO3)2 + 2KI → PbI(s) + 2KNO3
Soluble Soluble Insoluble Soluble

       Pb(NO3)2 + 2 KI 

  Pb2+ + 2 NO3- + 2 K+ + 2I-  ?




                      Pb2+ + 2I-  PbI2(s)
                      net ionic equation
Silver nitrate + sodium chromate → silver chromate + sodium nitrate


2 AgNO3(aq) + Na2CrO4(aq)  Ag2CrO4(s) + 2NaNO3(aq)

2 Ag+ NO3- + 2 Na+ + CrO42-  Ag2CrO4(s) + 2Na+ NO3-

                                                       Spectator ions

            2 Ag+ + CrO42-  Ag2CrO4(s)

                net ionic equation
Essential Solubility Rules for Ionic Compounds in Water

1.   All Group 1A and ammonium compounds (M) are soluble
     M = Li+, Na+, K+, Rb+, NH4+.
     i.e. all MnXm salts are soluble, where X = any anion

2.   All nitrate and acetate salts are soluble.
      NO3-, C2H3O2-
     i.e. All metal nitrates and metal acetates, ammonium
     nitrate and ammonium acetate are soluble
• All the following compounds are insoluble
  in water EXCEPT

•   a. MgCO3
•   b. CuS
•   c. BaSO4.
•   d. Fe(NO3)2.
•   e. Ca3(PO4)2
Which of the following is the correct net ionic equation for the
   reaction that occurs when aqueous solutions of Pb(ClO3)2 and
   Na2SO4 are mixed?

a.   ClO3-(aq) + Na+(aq) → NaClO3(s)
b.   Pb2+(aq) + SO42-(aq) → PbSO4(s)
c.   Pb(ClO3)2(aq) + Na2SO4(aq) → PbSO4(s) + NaClO3(aq)
d. Pb2+(aq) + 2ClO3-(aq) + 2Na+(aq) + SO42-(aq) →
   PbSO4(s) + 2Na+(aq) + SO42-(aq)
How many mL of 0.250 M NaOH solution will react exactly with 35.0 mL of 0.175 M H2SO4
according to the equation?

      H2SO4(aq) + 2 NaOH(aq) → 2 H2O(l) + Na2SO4(aq)
     How many mL of 0.250 M NaOH solution will react exactly with 35.0 mL of 0.175 M H2SO4
     according to the equation?

           H2SO4(aq) + 2 NaOH(aq) → 2 H2O(l) + Na2SO4(aq)


a.   14.2 mL          mol of H2SO4 reacting = 35.0 mL x 0.175 M (i.e. mol = V x M)
b.   23.6 mL                                  1000
c.   32.5 mL
d.   49.0 mL          mol of NaOH reacting = 2 (35.0 mL x 0.175 M)     = VmL x 0.250
                                                 1000                    1000

                      Hence, V = 49.0 mL



                2 x Vacid x Macid =     Vbase x Mbase

               2 x mol acid reacting = mol base reacting
•   Which of the following is a strong acid?
•   a. oxalic acid, H2C2O4
•   b. phosphoric acid, H3PO4
•   c. acetic acid, HC2H3O2
•   e. nitric acid, HNO3
• All the following compounds are insoluble
  in water?

•   a. Fe(NO3)2.
•   b. (NH4)2SO3
•   c. NaC2H3O2
•   d. MgCO3.
•   e. Li2S.
Oxidation-Reduction Reactions

       Redox Reactions

    Electron-Transfer Reactions
    Oxidation of Mg by O2 to form MgO


            2 Mg + O2   → 2 MgO

 Mg - 2e- → Mg2+ (loss of electrons, oxidation)

           the Mg is oxidized to Mg2+

O2 + 4e- → 2 O2- (gain of electrons, reduction)

         each O atom is reduced to O2-
    Oxidation of Mg by O2 to form MgO




 2 Mg - 4e- → 2 Mg2+ Oxidation Half-Reaction

  O2 + 4e- → 2 O2- Reduction Half Reaction
_________________________________________

           2 Mg + O2       → 2 MgO
                       .
 Oxidation-Reduction Reactions
Oxidizing Agent is a species that oxidizes
another species, and is itself reduced.
Reducing Agent is a species that reduces
another species, and is itself oxidized

        Loss of e-, oxidation


reducing agent

         oxidizing agent

                 Gain of e-, reduction
Reaction of Fe with Cu2+(aq) to yield Fe2+(aq) and Cu(s)

              Fe + Cu2+    → Fe2+ + Cu
Oxidation-Reduction

Cu2+ is the Oxidizing Agent. It is the species
that oxidizes the Fe, and is itself reduced.
Fe is the Reducing Agent. It is the species
that reduces the Cu2+, and is itself oxidized

       Loss of 2 e-, oxidation

 reducing agent

          oxidizing agent


                  Gain of 2 e-, reduction
Oxidation-Reduction
Oxidizing Agent
The species in an oxidation-reduction
reaction that oxidizes another species, and is
itself reduced (to a lower oxidation number).
It is the species that accepts the electrons.
Reducing Agent
The species in an oxidation-reduction
reaction that reduces another species, and is
itself oxidized (to a higher oxidation number).
It is the species that provides the electrons.
Oxidation-Reduction

Oxidation
Loss of electrons
Increase in oxidation number

Reducion
Gain of electrons
Decrease in oxidation number
An atom in its elemental form has an oxidation
number of 0.
Mg, Cl2, O2, N2, Fe

Fe → Fe2+ + 2e-
On losing electrons, the iron is oxidized to a higher
oxidation state (+2). Its oxidation number increases
from 0 to +2

Cl2 + 2e- → 2 Cl-
On gaining electrons, each Cl atom is reduced to a
lower oxidation state (-1). The oxidation number of Cl
decreases from 0 to -1
• The oxidation number of chromium in
  sodium chromite, NaCrO2, is

•   a. –2.
•   b. –1.
•   c. +1.
•   d. +2.
•   e. +3.
In a reaction in which SO32- is converted into S2O42-,
 the sulfur
    a.      is oxidized
    b.      is reduced
    c.      is the reducing agent
    d.      does not undergo a redox change
Which are Oxidation-Reduction reactions?

Identify the redox changes occuring

Identify the oxidizing agent

2 Al + 3 S2 → 2 Al2S3

N2 + 3 H2 → 2 NH3

2 NO + O2 → 2 NO2

2 HgO → 2 Hg + O2

CaCO3 → CaO + CO2
The balanced oxidation half-reaction in the reaction
  Cu2+(aq) + Fe(s) → Cu(s) + Fe2+(aq)
  is?
• a. Cu2+(aq) + 2e– →Cu(s).
• b. Fe2+(aq) + 2e– →Fe(s).
• c. Fe(s) → Fe2+(aq) + 2e–
• d. Cu(s) + 2e– →Cu(s).
• e. Cu(s) → Cu2+(aq) + 2e–
    Balancing Redox Equations
     Oxidation Number Method




2 MnO4- + Br- + 2 H+ → 2 MnO2 + BrO3- + H2O
        Balancing Redox Equations

Half-Reaction Method

 MnO4- + Br- → MnO2 + BrO3-

 2 (MnO4- + 4 H+ + 3 e- → MnO2 + 2 H2O ) Reduction

  Br- + 3 H2O → BrO3- + 6 H+ + 3 H2O + 6 e- Oxidation
  _____________________________________
 2 MnO4- + Br- + 2 H + → 2 MnO2 + BrO3- + H2O
 ______________________________________
     After balancing the redox equation below, the ratio of the
     coefficients x : y is


     x Cr2O72- (aq) + y Fe2+(aq) → Cr3+(aq) + Fe3+(aq) + H2O(l)


a.    1: 2
b.    1:3
c.    2:5
d.    1:6
Highest Oxidation State = Group Number (IA – VIIA)

Gp 1A 1+ is only oxidation number Li+, Na+, K+, Rb+,Cs+

Gp 2A. 2+ is only oxidation n umber, e.g. Mg2+, Ca2+

Gp 3A. 3+ is only oxidation number, e.g. Al3+, BF3 Al2O3

Gp 4A. 4+ is highest oxidation number. CO2, CO32-
Gp 5A. 5+ is highest oxidation number. N2O5, NO3-, HNO3

Gp 6A. 6+ is highest oxidation number. SF6, SO3, H2SO4

Gp 7A. 7+ is highest oxidation number. Cl2O7, HClO4
Displacing one metal by another.

Activity Series of Metals and H2

Most active metal is at the top of the series (strongest
reducing agent)

The least active metals (weakest reducing agents) are at the
bottom.

Metals below H2 cannot displace H2 from any source
Displacing one Metal by Another.

Activity Series of Metals and H2

Fe(s) + Cu2+(aq) → Fe2+(aq) + Cu(s)
Cu(s) + 2Ag+ (aq) → Cu2+(aq) + Ag(s)

2Na(s) + 2 H+(aq) → 2 Na+(aq) + H2(g)

Mg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s)

Mg > Fe > Cu > Ag Reducing Strength

 Most active metal is at the top of the series (strongest reducing agent)
• The oxidation state of nitrogen shown is
  correct for all the following species
  EXCEPT

•   a. N2H4 (–2).
•   b. NH2OH (–1).
•   c. N2O (+1).
•   d. HN3 (–1).
•   e. HNO2 (+3).
Which of the following equations describes an
oxidation process?
      a.     MnO4 + 8 H+ + 5e  Mn2+ + 4 H2O
      b.     Cl2 + 2 e  2 Cl
      c.     O2 + 4 H+ + 4 e  H 2O
      d.     2 Cr3+ + 7 H2O      Cr2O72 + 14 H+ + 6 e
Identify the false statement describing events in a
redox reaction:

a.      the oxidation number of the reducing agent increases.

b.      the reducing reagent provides the electrons (loses
        electrons).

c.      the reducing agent is more likely to be a metallic
        element than a non-metallic element.

d.      the oxidizing agent undergoes oxidation
           Chapter 4. Operational Skills
• Identifying strong and weak acids and bases
• Acid-base neutralization reactions
• Writing net Ionic equations (strong acid-strong base, weak
  acid-strong base, strong acid-weak base).
• Solubility rules (nitrates, acetates, Gp I, NH4+ salts soluble)
• Predicting insolubility, writing net ionic equations.
• Oxidation-Reduction
• Identifying oxidation and reduction changes and oxidizing
  and reducing agents
• Assigning oxidation numbers
• Balancing oxidation-reduction reactions (oxidation number
  method)
Recommended end-of-chapter problems

Chapter 4

6, 14, 20, 29, 32, 47, 66, 74, 79, 113, 132

				
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