CH 4 Reactions in Aqueous Solutions

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CH 4 Reactions in Aqueous Solutions Powered By Docstoc
					 Types of Equations Used to Describe Reactions
                  in Solution

   1.       Molecular: overall reaction stoichiometry- not actual
            forms
             MgSO4 (aq) + Na2CO3 (aq) --> MgCO3 (s) + Na2SO4 (aq)


   2. Complete Ionic: reactants and products that are strong
       electrolytes are represented as ions.
             Mg2+ + SO42- + 2Na+ + CO32- --> MgCO3 (s) + 2Na+ + CO32-


   3. Net Ionic: includes only those solution components
       undergoing a change. Spectator ions not included.
                           Mg2+ + CO32- --> MgCO3 (s)

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            Types of Chemical Reactions
Combination (Synthesis) reaction
          A + B  AB
Decomposition reactions AB  A + B
Displacement reactions
          AB + C  AC + B
Metathetical (change of position) reactions
          (double-replacement reactions)
          AB + CD  AD + CB
Combustion reactions  reactions with oxygen
          CxHy + nO2  xCO2 + (y/2) H2O

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Combination Reactions (Synthesis):               A+B→C

Metals + Oxygen:
      –
     Lithium + oxygen →
      –
     Magnesium + oxygen →
      –
     Gold + oxygen →
      –
     Platinum + oxygen →
                    Remember the diatomics
Metals with multiple charges:
  choose the one with higher charge; Cu+2 and not Cu+1

Nonmetals + Oxygen (Redox?)
      Excess carbon with oxygen →
      Limited amount of carbon with excess of oxygen →
      Phosphorus + excess oxygen →
      Phosphorus with limited amount of oxygen →
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Combination Reactions (Synthesis): A + B → C
Metals + nonmetals (Redox?)
            Cesium metal + iodine →
            Zinc + sulfur →
            Magnesium + nitrogen →
Metal Oxides (most are solid) + Water:
 (Redox?)
            Magnesium oxide + water →
            Lithium oxide + water →
            Aluminum oxide + water →
            Iron(III) oxide + water →
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Combination Reactions (Synthesis): A + B → C

 Nonmetal Oxides + Water : (Redox?)

             solid calcium oxide + water →
             solid lithium oxide + water →

 Can be Redox:
             2NO2(g) + H2O (l) → HNO3 (aq) + HNO2(aq)




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Combination Reactions (Synthesis): A + B → C

 Metal Oxides + Nonmetal Oxides (Redox?)
     calcium oxide + silicon dioxide →
     lithium oxide + tetra phosphorus deca oxide →


 Notes:
 The more electropositive (most metallic) element is always
   written first
        P4O10; CaO; H2O, CO2

                Check Periodic Table


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 Decomposition Reactions : C → A + B
              Reverse of combination (synthesis)

 Metallic oxides  metal + oxygen

 Nonmetallic oxides  nonmetal + oxygen

 Hydroxide  metal oxide + water

 Acid  nonmetallic oxide + water

            Which are Redox and which are not?
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Decomposition Reactions (Special Cases)
 Metal carbonates  metallic oxide + CO2

 Metal bicarbonates: metal oxide + CO2(g) + H2O (l)

 Metal sulfite  metallic oxide + SO2

 Metal chlorate  metal chloride + oxygen (O2)

 Binary compounds  elements

 Electrolysis of molten salts (ionic compounds)  elements


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 Decomposition Reactions : (Special Cases)
Decomposition of peroxides:
    peroxide  water + oxygen (O2)
Ammonium compounds  acid + ammonia;
               the acid may decompose

     (NH4)2CO3 (s)  2NH3(g) + CO2(g) + H2O(l)

       NH4NO2 (s)  N2(g) + 2H2O (l)

      NH4NO3(s)  N2O (g) + 2H2O(l)

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 Types of Equations Used to Describe Reactions
                  in Solution

   1.       Molecular: overall reaction stoichiometry- not actual
                    forms
             MgSO4 (aq) + Na2CO3 (aq) --> MgCO3 (s) + Na2SO4 (aq)


   2. Complete Ionic: reactants and products that are strong
                electrolytes are represented as ions.
             Mg2+ + SO42- + 2Na+ + CO32- --> MgCO3 (s) + 2Na+ + CO32-


   3. Net Ionic: includes only those solution components
       undergoing a change. Spectator ions not included.
                           Mg2+ + CO32- --> MgCO3 (s)

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            Writing Equations
Write a balanced molecular, ionic and net ionic
   equations for the following reactions:

1. Solution of silver nitrate was added to a
   solution of sodium chromate

2. A piece of solid zinc was placed in a solution
   of Copper(II) chloride
3.1
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Single Replacement or Displacement Reactions
                 A0 + B+C-  A+C- + B- ( metals)
                A0 + B+C-  B+A- + C0 (halogens)
                          All are Redox

                 Active metal replaces less active metal
                Active metal replaces H in water or acids
                Nonmetal replaces less active nonmetal
                  Activity series – used to predict Rx
              Standard Reduction Potential Chart and SHE

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             Single Replacement Reactions
        If a < reactive element is combined with a >
         reactive element in compound form → no Rx

1. Zinc metal reacts with copper (II) sulfate in water solution
        Molecular equation:
        Net Ionic equation :
                             Redox?

2. zinc metal reacts with hydrochloric acid

3. aluminum metal reacts with sulfuric acid


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                 Single Replacement Reactions
Write formula and net ionic equations:

            sodium metal reacts with cold a water

            aluminum reacts with steam

            magnesium reacts with hot water

    Which metals will replace hydrogen from cold water?

    Which metals will replace hydrogen from hot water?

    Which metals will replace hydrogen from steam?

Activity series of metals:

http://www.chem.vt.edu/RVGS/ACT/notes/activity_series.html

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            Single Replacement Reactions:
                Halogen Displacement
Write molecular and net ionic equations:

    Chlorine gas reacts with aqueous solution
       with sodium bromide


Activity series:
                F2 > Cl2 > Br2 > I2

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            Double Replacement Reactions or
                Metathetical Reactions
                 A+B- + C+D-  A+ D- + C+B-


Reactions occur to completion when:
      Precipitate is produced
      Gas is produced
      Molecular substance such as H2O, CO2, NH3, SO2
             are produced
                   Redox or NonRedox ?
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            Double Replacement Reactions or
                Metathetical Reactions
Write the molecular
    complete ionic
           net ionic forms

Aqueous nickel (II) chloride reacts with aqueous sodium
    hydroxide

Aqueous sodium sulfide reacts with lead (II) nitrate

Aqueous potassium carbonate reacts with barium chloride

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            Double Replacement Reactions or
                Metathetical Reactions
Predict whether a reaction will occur in each of the following
  case. If so, write a net ionic equation for the reaction. If no
  reaction occurs, write NR after arrow.

    Al2(SO4)3 + NaOH 


      K2SO4(aq)+FeBr3(aq) 


    CdCl2(aq) + (NH4)2S(aq) 


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Double Replacement: Gas Formation
Common gases formed in DR Rx
  S2- + acid → H2S (g)+ salt
  CO32- + acid → CO2 (g)+ H2O + salt
  SO3- + acid → SO2 (g)+ H2O + salt
  NH4+ + OH- + Δ → NH3 g) + H2O + salt

  1. Sodium carbonate reacts with hydrochloric acid
  2. Ammonium chloride reacts with sodium hydroxide
  3. Magnesium nitride reacts with water
  4. Calcium sulfite reacts with hydrobromic acid
  5. Sodium chloride + sulfuric acid
  6. Sodium sulfide reacts with hydrochloric acid     20
            Selective Precipitation
Precipitation reactions allow us to target
  specific substances, and separate and recover
  them from a solution.

Example:
A solution contains Ca2+, Cu2+, and Pb2+. What
  anions can we add, and in what order , to
  separate and recover each cation?


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               Combustion

Write the products and balance the following
 combustion reaction:
     C6H12O6 (s) + O2 →
     C3H8O3 + O2 →
     CH3OH + O2



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            Acids and Bases: Arrhenius

Acid
      – Any substance that releases H+ ion in aqueous
        solution

Base
      – Any substance that releases OH- ion in aqueous
        solution


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    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




                           conjugate   conjugate.
                           base        acid

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       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
Acid-base reactions can be reversible:
reactants  products or products  reactants
                                 Conjugate acid:
                                 ____________
                                 Conjugate base:
                                 _________________


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Compounds that act as Brönsted Acid and
                Base
Write equations for the following reactions.
   Identify the acid, base, conjugated acid and
   conjugated base:
1.      HSO4-(aq) + H2O(l) →
2.      HSO4-(aq) + H2O(l) →
3.      H2O(l) + H2O(l) →
4.      HCO3-1 (aq) + H2O(l)



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              Important Acids and Bases
Strong Acids:
                               Strong Bases:
HCl             hydrochloric
                               NaOH            sodium
HBr             hydrobromic    hydroxide
HI              hydroiodic     KOH         potassium hydroxide
HNO3            nitric         Ca(OH)2         calcium
                               hydroxide
H2SO4           sulfuric
HClO4           perchloric
                               Weak Base:
Weak Acid:
                               NH3             ammonia
CH3CO2H         acetic

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              Ca(OH)2, Ba(OH)2 and Sr(OH)2
              Group IIA, heavy metals)




             Know the strong
              acids & bases!
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3.2
STRONG acids in water:
  100% of acid molecules form ions:
       HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq)



                                       H3O+ is
                                         hydronium ion




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 WEAK acids in water:
   ~5% or less of acid molecules form ions
            (acetic, H3PO4, H2CO3)




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  Polyprotic Acids: multiple acidic H atoms

               H2SO4  H+ + HSO4-
               HSO4-  H+ + SO42-

 Not all H’s are acidic:
        CH3CO2H




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            If H3PO4 reacts as an acid, which of the
                  following can it not make?



   1. H4PO4+
   2. H2PO4-
   3. HPO42-
   4. PO43-

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     Reactions Involving Weak Bases

 HCl(aq) + NH3(aq)  NH4+(aq) + Cl-(aq)


Net-Ionic Equation:

          NH3(aq) + H+(aq)  NH4+(aq)

 Spectator Ion?

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   Acid-Base Reactions: Neutralization

The “driving force” is the formation of water.
    NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(liq)


Net ionic equation
      OH-(aq) + H+(aq) → H2O(liq)
    “Spectator Ions”? __________________________

  This applies to ALL reactions of STRONG acids and
                        bases.

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     Acid-Base Neutralization Rx
                 Polyprotic acids
             H2SO4 (SA)or H3PO4 (WA)
        H2SO4 : First H+ is ionized completely
                 H2SO4 → H+ + HSO4 -

1. If base is excess: all H+ form H2O
2. If equimolar acid + base: only 1 H+ ionizes
3. Acidic anhydrides (NMO) + Basic anhydrides (MO) :
   react with H2O before acid or base
            CH3CO2H(aq) + NaOH(aq) 

Choose the correct answer:
1. CH3CO2H2+(aq) + NaO(aq)
2. CH3CO2-(aq) + H2O(l) + Na+(aq)
3. CH4(g) + CO2(g) + H2O(l)

Complete Ionic equation:

    CH3COOH(aq) + OH-(aq) → CH3COO-(aq) + H2O(l)


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            HCN(aq) + NH3(aq) 

Answer?
 1. NH4+(aq) + CN-(aq)
 2. H2CN+(aq) + NH2-(aq)
 3. C2N2(s) + 3 H2(g)




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Hydrolysis Rx- Reverse Neutralization
      Salt + H2O → molecular species
 Formation of a weak acid and/or weak base
     NH4+ + Cl- + H2O → H+ + Cl- + NH4OH
    NH4Cl : salt from SA (HCl) + WB (NH3 )
   Forms acidic solution due to: NH4+> OH-
     *Salts of SA + WB → Acidic Solution
      *Salts of SA + WB → Basic Solution
     *Salts of SA + SB → Neutral Solution
   *Salts of WA + SB → ?? Check Ka and Kb
1. Aqueous potassium fluoride undergoes
   hydrolysis when placed in water.

2. Sodium chloride and water are mixed
    together.

3. Ammonium fluoride and water are mixed
   together.
            Oxidation-Reduction Reactions

Redox reactions:
  involve a transfer of electrons.

Assigning oxidation states to an element in a
  molecule:
           K2CrO4      LiSCN



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              LEO GER: Leo the Lion Says GER
                         OIL RIG

            Loss of Electrons is Oxidation
            Gain of Electrons is Reduction

              Oxidation Involves Loss
              Reduction Involves Gain


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            Determination of Oxidation States

                Fe2O3 + 2Al → Al2O3 + 2Fe

Iron (III) gains 3 electrons to become elemental
  iron.
Elemental aluminum lost 3 electrons to become
  the aluminum ion.
Write the half reactions:


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            N2 H4 + N2O4 → N2 + H2O
The combustion of hydrazine with dinitrogen
   tetroxide helps to keep the space shuttle in
   Earth Orbit.
Is it a Redox reaction? Explain.




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            Fe2O3 + 2Al → Al2O3 + 2Fe
Iron (III) ion gained electrons. It has been
  reduced.
The aluminum lost electrons. It has been
  oxidized.
The oxidizing agent is the species that is
  reduced (Iron (III)).
The reducing agent is the species that is
  oxidized (aluminum).


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        Rules for Assigning Oxidation States (OS)
  1. OS of an atom in an element is 0.
                            Na (s), O2 (g)
  2. OS of a monatomic ion is the same as its charge.
                      Na+ OS = +1, Cl- OS = -1
  3. In its covalent compounds with nonmetals, hydrogen is
        assigned an OS of +1.
                           HCl, NH3, H2O.
  4. Oxygen is assigned an OS of -2 in its covalent compounds.
                         CO, CO2, SO2, SO3
  The exception to this rules occurs in peroxides (compounds
        contains the O22- group), where each oxygen is
        assigned an OS of -1.
                                H2O2

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  5. In binary compounds the element with the greater
      attraction for the electrons in the bond is assigned a
      negative OS equal to its charge in its ionic compounds.

                         HF, NH3, H2S, HI

  6. The sum of the oxidation states must be zero for an
     electrically neutral compound and must be equal to the
     overall charge for an ionic species.
                            NH4+, CO32-




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            Which Atoms Undergo Redox?
        2H2 (g) + O2 (g) → 2H2O (g)

        Zn (s) + Cu2+(aq) → Zn2+ (aq) + Cu(s)

        2AgCl (s) + H2 (g) → 2H+ (aq) + 2Ag(s) + 2Cl- (aq)

        2MnO4- (aq) + 16H+ (aq) + 5C2O42- (aq) →
        2Mn2+(aq) + 10 CO2 (g) + 8 H2O (l)


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   Methods for Balancing Redox Reactions

1. Oxidation states method
   CdS + I2 + HCl → CdCl2 + HI + S
           (1, 1, 2, 1, 2, 1)
   Cl2 + Ca(OH)2 → CaCl2 + Ca(ClO3) + H2O
           ( 6, 6, 5, 1, 6)

2. Half reaction method


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   Balancing a Redox Equation by the Oxidation
                 States Method
   1. Assign the oxidation states of all atoms.
   2. Decide which element is oxidized and determine the
      increase in oxidation state.
   3. Decide which element is reduced and determine the
      decrease in oxidation state.
   4. Choose coefficients for the species containing the atom
      oxidized and the atom reduced such that the total
      increase in oxidation state equals the total decrease in
      oxidation state.
   5. Balance the remainder of the equation by inspection.



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  The Half-Reaction or Ion-electron Method for
  Balancing Redox Reactions in Acidic Solutions
    1. Split Rx into Half Reactions
       2 parts of a REDOX RX- pick 1 to start then repeat
               oxidation - RA
               reduction – OA
     2. Balance # of atoms first– not O or H
     3. Balance O by + H2O to side deficient in O
     4. + H+ to balance H
     5. Mass balance achieved

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6. Balance charges with e-
7. Be sure # e- lost = #e- gained
8. Cancel common terms on opposite side of →
9. Σ the half Rx
10. Check to make sure the Rx balances

Balance the equation in acid solution using the half
   reaction method
      Cu(s) + HNO3 (aq) --> Cu2+ (aq) + NO(g)
 The Half-Reaction or Ion-electron Method for
 Balancing Redox Reactions in Basic Solutions

1.   Follow the same procedure as in acidic solutions
2.   +OH- to both sides to cancel out the H+
3.   Make H2O from the OH- and H+
4.   Cancel out H2O that is common to both sides
5.   Be sure #e-lost = #e- gained
6.   Cancel common terms on opposite side of →
7.   Σ the half Rx
8.   Check to make sure the Rx balances
            Balancing by Half Reaction Method
    Balance in acidic environment:
                 Cr2O72- (aq) + NO (g) → Cr +3 (aq) + NO3- (aq)
               Cr2O72-(aq) + 2NO (g) + 6H+ (aq) → 2Cr3+ (aq) + 2NO3- (aq) + 3H2O (l)



  Balance in basic environment:
               Cr2O72- (aq) + NO (g) → Cr +3 (aq) + NO3- (aq)

            Cr2O72-(aq) + 2NO (g) + 3 H2O --> 2Cr3+ (aq) + 2NO3- (aq) + 6OH- (aq)




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    Predicting Types of Redox Rx
Simple Redox
  Hydrogen displacement      Metal displacement
  Halogen displacement       Combustion
                 Decomposition
Oxoanions
  Table of Common Oxidizing and Reducing Agents
             Atypical Redox Rx
1. Hydrogen + MO (hot) → M + HOH

2. MS + O2 → MO + SO2

3. Cl2 (g) + NaOH(dilute) → NaClO + NaCl + HOH

4. Cu + H2SO4 (conc) → CuSO4 + SO2 + HOH

5. Cu + HNO3 (dilute) → CuNO3 + NO + HOH

6. Cu + HNO3 (conc) → CuNO3 + NO2 + HOH
            Disproportionation Reactions

    Simultaneous oxidation and reduction of one
                      species.
          3NO2 + H2O → 2H+ + 2NO3- + NO

             2H2O2(aq) → 2H2O(l) + O2(g)

Cl2(g) + 2OH-(aq) → ClO-(aq) + Cl-(aq) + H2O(l)

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