Chapter 5 Reactions Between Ions in Aqueous Solutions

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					  Chapter 5: Reactions Between
   Ions in Aqueous Solutions
• A solution is a homogeneous mixture in
  which the two or more components mix
  freely
• The solvent is taken as the component
  present in the largest amount
• A solute is any substance dissolved in the
  solvent
Formation of a solution of
iodine molecules in ethyl
alcohol. Ethyl alcohol is the
solvent and iodine the
solute.

Solutions have variable
composition. They may be
characterized using a solute-
to-solvent ratio called the
concentration.
• For example, the percentage concentration
  is the number of grams of solute per 100 g
  of solution
• The relative amounts of solute and solvent
  are often given without specifying the
  actual quantities
                   The dilute solution on the left
                   has less solute per unit volume
                   than the (more) concentrated
                   solution on the right.

                   Concentrated and dilute are
                   relative terms.
• There is usually a limit to the amount of
  solute that can dissolve in a given amount
  of solvent
  – For example, 36.0 g NaCl is able to dissolve in
    100 g of water at 20°C
• A solution is said to be saturated when no
  more solute can be dissolved at the current
  temperature
• The solubility of a solute is the number of
  grams of solute that can dissolve in 100
  grams of solvent at a given temperature
• Solubilities of some common substances
                                         Solubility
 Substance               Formula       (g/100 g water)
 Sodium chloride         NaCl          35.7 at 0°C
                                       39.1 at 100°C
 Sodium hydroxide NaOH                 42 at 0°C
                                       347 at 100°C
 Calcium carbonate CaCO3               0.0015 at 25°C

A solution containing less solute is called unsaturated
because it is able to dissolve more solute.
• Solubility usually increases with
  temperature
• Supersaturated solutions contain more
  solute than required for saturation at a given
  temperature
• They can be formed, for example, by
  careful cooling of saturated solutions
• Supersaturated solutions are unstable and
  often result in the formation of a
  precipitate
• A precipitate is the solid substance that
  separates from solution
• Precipitates can also form from reactions
• Reactions that produce a precipitate are
  called precipitation reactions
• Many ionic compounds dissolve in water
• Solutes that produce ions in solution are
  called electrolytes because their solutions
  can conduct electricity
• An ionic compounds dissociates as it
  dissolves in water
                      Ions separate from the solid
                      and become hydrated or
                      surrounded by water
                      molecules.

                      The ions move freely and the
                      solution is able to conduct
                      electricity.


Ionic compounds that dissolve completely are
strong electrolytes
• Most solutions of molecular compounds do
  not conduct electricity and are called
  nonelectrolytes
                                The molecules of a
                                nonelectrolyte separate
                                but stay intact. The
                                solution is nonconducting
                                because no ions are
                                generated.


 Some ionic compounds have low solubilities in
 water but are still strong electrolytes because what
 does dissolve is 100% dissociated.
• The dissociation of ionic compounds may
  be described with chemical equations
                                       2-
      Na 2SO 4 ( s ) Æ 2 Na (aq ) + SO (aq )
                            +
                                       4


• The hydrated ions, with the symbol (aq),
  have been written separately
• Since physical states are often omitted, you
  might encounter the equation as:
           Na 2SO 4 Æ 2 Na + + SO 2-
                                  4
• Ionic compounds often react when their
  aqueous solutions combine
                                   When a
                                   solution of
                                   Pb(NO3)2 is
                                   mixed with
                                   a solution of
                                   KI the
                                   yellow
                                   precipitate
                                   PbI2 rapidly
                                   forms.
• This reaction may be represented with a
  molecular, ionic, or net ionic equation:
  Molecular: Pb(NO3 ) 2 (aq) + 2KI(aq) Æ PbI 2 ( s) + 2KNO3 (aq)
  Ionic:                   -
         Pb 2+ (aq ) + 2NO 3 (aq ) + 2K + (aq ) + 2I - (aq ) Æ
                                                   -
                       PbI2(s ) + 2K + (aq ) + 2NO 3 (aq )
   Net Ionic: Pb 2+ (aq) + 2I - (aq) Æ PbI 2 ( s)
• The most compact notation is the net ionic
  equation which eliminates all the non-
  reacting spectator ions from the equation
•    Criteria for balanced ionic and net ionic
     equations:
    1) Material balance – the same number of each
       type of atom on each side of the arrow
    2) Electrical balance – the net electrical charge
       on the left side of the arrow must equal the
       net electrical charge on the right side of the
       arrow
    Remember that the charge on an ion must be included
    when it is not in a compound. Adding the charges on all
    the ions on one side of the arrow gives the net electrical
    charge.
• In the reaction of Pb(NO3)2 with KI the
  cations and anions changed partners
• This is an example of a metathesis or
  double replacement reaction
• Solubility rules allows the prediction of
  when a precipitation reaction will occur
• For many ionic compounds the solubility
  rules correctly predict whether the ionic
  compound is soluble or insoluble
 • Solubility rules for ionic compounds in
   water:
     – Soluble Compounds
1) All compounds of the alkali metals (Group IA) are soluble.
                                  -               -
2) All salts containing NH + , NO 3 , ClO -4 , ClO3 , and C 2 H 3O -2
                           4

   are soluble.
3) All salts containing Cl - , Br - , or I - are soluble except when
  combined with Ag + , Pb 2+ , and Hg 2+ .
                                      2

4) All sulfates are soluble except those of Pb 2+ , Ca 2+ , Sr 2+ ,
  Hg 2+ , and Ba 2+ .
     2
      – Insoluble compounds
5) All metal hydroxides and oxides are insoluble except those
  of Group IA and of Ca 2+ , Sr 2+ , and Ba 2+ . When metal
  oxides do dissolve, they react with water to form
   hydroxides. The oxide ion, O 2- , does not exist in water.
                                     2        2
6) All salts that contain PO 3- , CO 3 - , SO 3 - , and S2- are insoluble,
                             4

   except those of Group IA and NH + .
                                   4


  • A knowledge of these rules will allow you
    to predict a large number of precipitation
    reactions
• Acids and bases are another important class
  of compounds
• Acids and bases affect the color of certain
  natural dye substances
• They are called acid-base indicators
  because they indicate the presence of acids
  or bases with their color
• The first comprehensive theory of acids,
  bases, and electrical conductivity appeared
  in 1884 in the Ph.D. thesis of Savante
  Arrhenius
• He proposed that acids form hydrogen ions
  and bases released hydroxide ions in
  solution
• The characteristic reaction between acids
  and bases is neutralization
    HCl(aq) + NaOH(aq) ‡ NaCl(aq) + H2O(l)
• In general, the reaction of an acid and a
  base produces water and a salt
• We can state the Arrhenius definition of
  acids and bases in updated form
Arrhenius Definition of Acids and Bases
 An acid is a substance that reacts with water to produce
 the hydronium ion, H 3O + .
 A base is a substance that produces hydroxide ion in water.

• In general, acids are molecular compounds
  that react with water to produce ions
• This is called ionization:
                                                   -
   HCl( g ) + H 2 O Æ H 3O (aq ) + Cl (aq )
                                    +
• It is common to encounter the hydrogen ion
  (H+) instead of the hydronium ion
• The previous ionization is also written as
               H 2O
     HCl( g ) ææ Æ H + (aq ) + Cl - (aq )
                æ
• Monoprotic acids are capable of furnishing
  only one hydrogen ion per molecule
• Acids that can furnish more than one
  hydrogen ion per molecule are called
  polyprotic acids
Monoprotic : HCl(aq )+H 2 O Æ H 3O + (aq ) + Cl - (aq )
                                                       -
Diprotic : H 2 CO 3 (aq ) + H 2 O Æ H 3O + (aq ) + HCO 3 (aq )
                -                                   2
            HCO 3 (aq ) + H 2 O Æ H 3O + (aq ) + CO 3 - (aq )
Triprotic : H 3 PO 4 (aq ) + H 2 O Æ H 3O + (aq ) + H 2 PO -4 (aq )
            H 2 PO -4 (aq ) + H 2 O Æ H 3O + (aq ) + HPO 2- (aq )
                                                         4

             HPO 2- (aq ) + H 2 O Æ H 3O + (aq ) + PO 3- (aq )
                 4                                    4
• Some nonmetal oxides react with water to
  produce acids
• They are called acidic anhydrides
  (anhydride means without water)
• Soluble metal oxides are base anhydrides
• Examples include:
 Nonmetal Oxides :
  SO 3 ( g ) + H 2 O Æ H 2SO 4 (aq )     sulfuric acid
  N 2 O 5 ( g ) + H 2 O Æ 2HNO3 (aq ) nitric acid
   CO 2 ( g ) + H 2 O Æ H 2 CO 3 (aq )   carbonic acid
 Metal Oxides :
  CaO( s ) + H 2 O Æ Ca(OH) 2 (aq ) calcium hydroxide
  Na 2 O( s ) + H 2 O Æ 2NaOH(aq )        sodium hydroxide
• Ammonia gas ionizes in water producing
  hydroxide ions
• It is an example of a molecular base
• Many molecules that contain nitrogen can
  act as a base
                                        -
  NH 3 (aq ) + H 2 O Æ NH (aq ) + OH (aq )
                             +
                             4

  For the general base B :
  B(aq ) + H 2 O Æ HB (aq ) + OH (aq )
                        +           -
• Binary compounds of many nonmetals and
  hydrogen are acidic
• In water solution these are referred to as
  binary acids
• They are named by adding the prefix hydro-
  and the suffix –ic to the stem of the
  nonmetal name, followed by the word acid
 Molecular Compound           Binary Acid
 HCl( g ) hydrogen chloride   HCl(aq ) hydrochloric acid
 H 2S( g ) hydrogen sulfide   H 2S(aq ) hydrosulfuric acid
• Acids that contain hydrogen, oxygen, plus
  another element are called oxoacids
• They are named according to the number of
  oxygen atoms in the molecule and do not
  take the prefix hydro-
• When there are two oxoacids, the one with
  the larger number of oxygens takes the
  suffix –ic and the one with the fewer
  oxygen atoms takes the suffix –ous
 H 2SO 4 sufluric acid   HNO3 nitric acid
 H 2SO 3 sulfurous acid HNO 2 nitrous acid
• The halogen can occur with up to four
  different oxoacids
• The oxoacid with the most oxygens has the
  prefix per- the one with the least has the
  prefix hypo-
  HClO hypochlorous acid HClO3 chloric acid
  HClO 2 chlorous acid   HClO 4 perchloric acid
•    Anions are produced when oxoacids are
     neutralized
•    There is a simple relationship between the
     name of the polyatomic ion and the parent
     acid
    1) –ic acids give –ate ions
    2) -ous acids give –ite ions
•    In naming polyatomic anions, the prefixes
     per- and hypo- carry over from the parent
     acid
• Polyprotic acids can be neutralized
• An acidic salt contains an anion that is
  capable of furnishing additional hydrogen
  ions
• The number of hydrogens that can still be
  neutralized is also indicated
  NaHSO 4      sodium hydrogen sulfate
  Na 2 HPO 4 sodium hydrogen phosphate
  NaH 2 PO 4 sodium dihydrogen phosphate
• Naming bases is much less complicated
• Ionic compounds containing metal ions are
  named like any other ionic compound
• Molecular bases are specified by giving the
  name of the molecule
• Acids and bases can be classified as strong
  or weak and so as strong or weak
  electrolytes
• Strong acids are strong electrolytes
• The most common strong acids are:
    HClO 4 (aq ) perchloric acid
    HCl(aq )     hydrochloric acid
    HBr(aq )     hydrobromic acid
    HI(aq )      hydroiodic acid
    HNO3 (aq )   nitric acid
    H 2SO 4 (aq ) sulfuric acid
• Strong bases are the soluble metal
  hydroxides
• These include:
Group IA                Group IIA
LiOH lithium hydroxide
NaOH sodium hydroxide
KOH potassium hydroxide Ca(OH) 2 calcium hydroxide
RbOH rubidium hydroxide Sr(OH) 2 strontium hydroxide
CsOH cesium hydroxide     Ba(OH)2 barium hydroxide

• Most acids are not completely ionized in
  water
• They are classified as weak electrolytes
                                    The brightness
                                    of light is
                                    experimental
                                    verification of
                                    the
                                    classification
                                    as a strong or
                                    weak
                                    electrolyte.

Weak acids and bases are weak electrolytes
because less than 100% of the molecules ionize.
• Weak acids and bases are in dynamic
  equilibrium in solution
• Consider the case of acetic acid:
                                  Two opposing reactions
                                  occur in solution: the
                                  ionization of the acid,
                                  called the forward
                                  reaction, and the
                                  recombination of ions into
                                  molecules, called the
                                  reverse reaction.

Chemical or dynamic equilibrium results when the rate of the
forward and reverse reaction are equal.
• Neutralization of a strong acid with strong
  base gives a salt and water:
Molecular : HCl(aq) +KOH(aq) Æ KCl(aq)+H 2 O
Ionic : H + (aq) + Cl - (aq) + K + (aq) + OH - (aq) Æ H 2 O + K + (aq) + Cl - (aq)
Net ionic : H + (aq) + OH - (aq) Æ H 2 O
• This net ionic equation applies only to
  strong acids and bases
• The neutralization of a weak acid with a
  strong base involves a strong and weak
  electrolyte
• Consider the neutralization of acetic acid
  with NaOH:
 Molecular : HC 2 H 3O 2 (aq ) + NaOH(aq ) Æ NaC 2 H 3O 2 (aq ) + H 2 O
 Ionic : HC 2 H 3O 2 (aq ) + Na + (aq ) + OH - (aq ) Æ
                            Na + (aq ) + C 2 H 3O -2 (aq ) + H 2 O
 Net ionic : HC 2 H 3O 2 (aq ) + OH - (aq ) Æ C 2 H 3O -2 (aq ) + H 2 O


• Note that in ionic equations the formulas of
  weak electrolytes are written in “molecular”
  form
• The situation is similar when a strong acid
  reacts with a strong base
• For ammonia and HCl the net ionic
  equation is:
  NH 3 (aq ) + H + (aq ) Æ NH + (aq )
                              4

                 or
  NH 3 (aq ) + H 3O + (aq ) Æ NH + (aq ) + H 2 O
                                 4

• Note that water only appears as a product if
  the hydronium ion is used
• Both strong and weak acids react with
  insoluble hydroxides and oxides
• The driving force is the formation of water
• Magnesium hydroxide has a low solubility
  in water, but reacts with strong acid
• The net ionic equation is:
  Mg(OH) 2 ( s ) + 2 H + (aq ) Æ Mg 2+ (aq ) + 2H 2 O
• Magnesium hydroxide is written as a solid
  because it is insoluble
• A number of metal oxides also dissolve in
  acids
• For example, iron(III) oxide reacts with
  hydrochloric acid:
 Molecular : Fe 2 O 3 ( s ) + 6HCl(aq ) Æ 2FeCl3 (aq ) + 3H 2 O
 Net ionic : Fe 2 O 3 ( s ) + 6H + (aq ) Æ 2Fe3+ (aq ) + 3H 2 O
• Some reactions with acids or bases produce
  a gas
• The reactions are driven to completion
  because the gas escapes and is unavailable
  for back reaction
Gas Compounds               Net Ionic Equation
H 2S Sulfides               2H + + S2- Æ H 2S
HCN Cyanides                H + + CN - Æ HCN
CO 2   Carbonates                     2
                            2H + + CO 3 - Æ CO 2 + H 2 O
                                     -
       Hydrogen Carbonates H + + HCO 3 Æ CO 2 + H 2 O
SO 2   Sulfites                        2
                             2H + + SO 3 - Æ SO 2 + H 2 O
       Hydrogen Sulfites               2
                             H + + HSO 3 - Æ SO 2 + H 2 O
NH 3   Ammonium Salts        NH + + OH - Æ NH 3 + H 2 O
                                4
   (CO2 and SO2 are produced by the decomposition
     of H2CO3 and H2SO3, respectfully)
• Solutions are characterized by their
  concentration
• The molar concentration or molarity (M)
  is defined as

    molarity (M) =      moles of solute
                       liters of solution

• The molarity of a solution gives an
  equivalence relation between the moles of
  solute and volume of solution
• Solutions provide a convenient way to
  combine reactants in many chemical
  reactions
  – Example: How many grams of AgNO3 are
    needed to prepare 250 mL of 0.0125 M AgNO3
    solution?
  ANALYSIS: Find moles, then mass of solute.
  SOLUTION:
                         0.0125 mol AgNO3       169.9 g AgNO3
   0.250 L AgNO3 sol ¥   1.00 L AgNO3 sol   ¥     mol AgNO3

            = 0.531 g AgNO3
• Solutions of high concentration can be
  diluted to make solutions of lower
  concentration
• Conservation of solute mass requires:
  Vdil ¥ M dil = Vconcd ¥ M concd
• Where dil labels the diluted and concd the
  concentrated solution
• Stoichiometry problems often require
  working with volumes and molarity
– Example: How many mL of 0.124 M NaOH are
  required to react completely with 15.4 mL of
  0.108 M H2SO4?
   2 NaOH + H2SO4 ‡ Na2SO4 + 2H2O
ANALYSIS: Use the mole-to-mole ratio to
  convert.
SOLUTION:
                               0.108 mol H 2SO 4
0.0154 L H 2SO 4 sol ¥         1.00 L H 2SO 4 sol
                                                        mol NaOH
                                                    ¥ 12mol H 2SO 4
        1.00 L NaOH sol
    ¥   0.124 mol NaOH    ¥ 1000LmL = 26.8 mL NaOH sol
                              1
• Limiting reagent problems are also common
   – Example: How many moles of BaSO4 will form
     if 20.0 mL of 0.600 M BaCl2 is mixed with
     30.0 mL of 0.500 M MgSO4?
       BaCl2 + MgSO4 ‡ BaSO4 + MgCl2
   ANALYSIS: This is a limiting reagent problem.
   SOLUTION:
0.0200 L BaCl 2 sol ¥ 0.600Lmol BaCl2 ¥ 11mol BaSO24 = 0.0120 mol BaSO 4
                      1.00 BaCl 2 sol     mol BaCl
                                         1 mol BaSO
0.0300 L MgSO 4 sol ¥ 0.500Lmol MgSO 4 ¥ 1 mol MgSO44 = 0.0150 mol BaSO 4
                      1.00 MgSO 4 sol

    \ 0.0120 mol BaSO 4 formed
• Titration is a technique used to make
  quantitative measurements of the amounts
  of solutions
• The end-point is often determined visually
                          The long tube is
                          called the buret. The
                          valve at the bottom of
                          the buret is called the
                          stopcock. The
                          titration is complete
                          when the indicator
                          changes color.
• Paths for working stoichiometry problems
  may be summarized with a flowchart: