Chapter 4: Aqueous Reactions and Solution Stoichiometry

• Solution Composition
• Molarity
• Dilution
• Properties of Solutes
• Ionic Compounds
• Molecular Compounds
• Electrolytes--strong/weak
• Acids/Bases/Salts/Neutralization Reactions
• Ionic Equations

• Metathesis Reactions
• Precipitation Reactions
• Solubility
• Products & Prediction
• Oxidation-Reduction Reactions
• Activity Series
• Solution Stoichiometry
• Titration

Solution Composition

• Solutions are composed of a solute and a solvent
• solute--present in smallest quantity
• solvent--present in largest quantity
• Molarity
• concentration gives ratio of solute : solvent/solution
• molarity, M = moles solute
                       1 L solution

Practice Ex. 4.1:

Calculate the molarity of a solution made by dissolving 5.00 g of glucose, C₆H₁₂O₆, in sufficient water to form 100 mL of solution.

5.00 g C6H12O6      x           1 mol         =          0.0278 mol        =         0.278 M
                                             180 g                      0.100 L

Practice Ex. 4.2:

How many grams of Na₂SO₄ are there in 15 mL of 0.50 M Na₂SO₄? How many mL of 0.50 M Na₂SO₄ solution are required to supply 0.038 mol of salt?

15 mL x     1 L       x       0.50 mol Na2SO4     x        142 g        =          1.1 g
               10³ mL                 1 L                                   mol                     Na2SO4


0.038 mol Na2SO4          x          1 L        x        103 mL        =          76 mL
                                               0.50 mol                1 L                        Na2SO4

• Dilution
• Stock solutions are generally concentrated solutions that are diluted before use

General format for diluting a concentrated solution:

Minitial Vinitial = Mfinal Vfinal

or

Mconc Vconc = Mdil Vdil

Practice Ex. 4.3:

How many mL of 5.0 M K₂Cr₂O₇ solution must be diluted in order to prepare 250 mL of 0.10 M solution?

Mconc = 5.0 M K2Cr2O7            Mdil = 0.10 M K2Cr2O7
Vconc = ?                                 Vdil = 250 mL

Mconc Vconc = Mdil Vdil

0.10 mol          x         0.250 L         x        1 L         =         0.0050 L or
    L                                                      5.0 mol                     5.0 mL
Mdil                     x             Vdil                  x            1              =               Vconc
                                                                             Mconc

Properties of Solutes

• Electrolytes
• conduct electricity
• form ions in solution stoichiometrically
• Nonelectrolytes
• do not conduct electricity
• do not form ions in solution
• Weak Electrolytes
• slightly conduct electricity
• form less than stoichiometric amounts of ions
• Ionic Compounds
• dissociate into constituent ions when dissolved
• hence, they are electrolytes if they are soluble
  

ions dissociate stoichiometrically

Na₂SO₄ ® 2Na⁺_(aq) + SO_4(aq)^2-

(NH₄)₂SO₄ ® 2NH_4(aq)⁺ + SO_4(aq)^2-

Ca(NO₃)2(aq) ® Ca²⁺_(aq) + 2NO_3(aq)^-

Practice Ex. 4.4

How many moles of K⁺ ions are present in 0.25 L of 0.015 M K₂CO3 solution?

K₂CO₃ ® 2K_(aq)⁺ + CO_3(aq)^2-


0.25 L      x        0.015 mol       x       2 mol K⁺       =       0.0075
                             1 L 1                mol K2CO3            mol K+

• Molecular Compounds
• structure of the molecules remains intact
• do not separate into ions !
• molecules themselves are separated on the molecular level
• generally not electrolytes
• Strong and Weak Electrolytes
• All soluble ionic compounds are strong electrolytes
• ions are produced stoichiometrically--exist completely or nearly completely as ions in solutions
• Some molecular compounds are weak electrolytes--produce small concentrations of ions when dissolved
• Since molecular compounds do not "contain" ions, they must 'produce' ions through a reaction with water, eg.
• NH3(aq) + H2O Û NH4(aq)⁺ + OH(aq)^-
• HC2H3O2(aq) + H2O Û H3O(aq)⁺ + C2H3O2(aq)^-
• Some molecular compounds are strong electrolytes
• HCl(aq) + H2O ® H3O(aq)⁺ + Cl(aq)^-
• Note the difference in arrows used for chemical equations for weak vs strong electrolytes

Acids/Bases/Salts

• Acids
• have an ionizable hydrogen, H⁺ eg. HCl(aq) or HC2H3O2(aq)
• can be strong or weak electrolytes
• HCl(aq) + H2O ® H3O(aq)⁺ + Cl(aq)^-
• HC2H3O2(aq) + H2O Û H3O(aq)⁺ + C2H3O2(aq)^-
• strong acids are more reactive than weak acids
• can be monoprotic
• HCl(aq) + H2O ® H3O(aq)⁺ + Cl(aq)^-
• or diprotic
• H2SO4(aq) + 2H2O ® 2H3O(aq)⁺ + SO4(aq)^2-
• Strong Acids--know these!
• HCl, HBr, HI, HNO3, H2SO4, HClO4,
  
• Bases
• substances that react with acids
• produce hydroxide ions, OH-, in solution
• can be strong or weak electrolytes
• NaOH(aq) ® Na(aq)⁺ + OH(aq)-
• NH3(aq) + H2O Û NH4(aq)+ + OH(aq)^-
• Strong Bases--know these!
• Group IA metal hydroxides (LiOH, NaOH, KOH, RbOH, CsOH)
• Heavy group IIA metal hydroxides [Ca(OH)2, Sr(OH)2, Ba(OH)2]
  
• Neutralization Reactions
• occur between acids and metal hydroxide bases
• produce water and a salt (any ionic compound)
• HCl(aq) + NaOH(aq) ® H2O(l) + NaCl(aq)

Ionic Equations

• Three ways to express ionic equations
• molecular equation--all species expressed in molecular form
• HCl(aq) + NaOH(aq) ® H2O(l) + NaCl(aq)

HCl(aq) is really H⁺(aq) and Cl^-(aq)

• complete ionic equation--all species expressed in ionic form
• H⁺(aq) + Cl^-(aq) + Na⁺(aq) + OH^-(aq) ® H2O(l)  +  Na⁺(aq)   +   Cl-(aq)
• note that spectator ions undergo no change
• net ionic equation--incl. only species that change
• H⁺(aq) + OH^-(aq) ® H2O(l)
• only soluble, strong electrolytes are written in ionic form
  

• Examples
• H2SO4(aq) + 2NaOH(aq) ® 2H2O(l) + Na2SO4(aq)
  
• 2H⁺(aq) + SO4^2-(aq) + 2Na⁺(aq) + 2OH^-(aq) ®
                                      2H2O(l) + 2Na+(aq) + SO4^2-(aq)
  
• net ionic equation--incl. only species that change
• H⁺(aq) + OH^-(aq) ® H2O(l)
• only soluble, strong electrolytes are written in ionic form
  

• Examples
• H2SO4(aq) + 2NaOH(aq) ® 2H2O(l) + Na2SO4(aq)
  
• 2H⁺(aq) + SO4^2-(aq) + 2Na⁺(aq) + 2OH^-(aq) ®
                                  2H2O(l) + 2Na⁺(aq) + SO4^2-(aq)
  
• net ionic equation--incl. only species that change
• H⁺(aq) + OH^-(aq) ® H2O(l)
• only soluble, strong electrolytes are written in ionic form
  

• Examples
• H2SO4(aq) + 2NaOH(aq) ® 2H2O(l) + Na2SO4(aq)
  
• 2H⁺(aq) + + 2OH^-(aq) ® 2H2O(l)

Metathesis Reactions

• General Form
• ion partner exchange
• AX + BY ® AY + BX
• Pb(NO3)2(aq) + 2KCl(aq) ® PbCl2(s) + 2KNO3(aq)
  
• Driving Forces
• formation of an insoluble solid
• formation of a weak or non-electrolyte
• formation of a gas

• Precipitation Reaction
• formation of an insoluble solid
• know the solubility guidelines!!
• examples:

NaCl(aq) + KNO3(aq) ® NR



AgNO3(aq) + KCl(aq) ® AgCl(s)   + KNO3(aq)


• Formation of a weak or non-electrolyte
• common example is an acid/base reaction--H2O forms
• know/recognize electrolyte vs non-electrolyte
• examples:
  
• HCl(aq) + NaOH(aq) ® H2O(l) + NaCl(aq)
  
  
  
• NiO(s) + 2HNO3(aq) ® Ni(NO3)2(aq) + H2O(l)
• Formation of a gas
• gases exit the reacting solution driving the reaction
• examples:
  
• 2HCl(aq) + Na2CO3(aq) ® 2NaCl(aq) + H2O(l) + CO2(g)
  
                                          (   H₂CO₃      ®    H₂O   +    CO₃)
  
  
  
• 2HCl(aq) + Na2S(aq) ® 2NaCl(aq) + H2S(g)

Oxidation-Reduction Reactions

• the loss of electrons by a substance
• Ca(s) ® Ca²⁺ + 2e^-
• the gain of electrons by a substance
• Cl2(g) + 2e- ® 2Cl^-
• Oxidation and Reduction always occur together
• Ca(s) + Cl2(g) ® CaCl2(s)
• Ca(s) + O2(g) ® CaO(s)
• 2Na(s) + Cl2(g) ® 2NaCl(s)
• Oxidation of Metals by Acids & Salts
• metal + acid ® salt + hydrogen
• Mg(s) + HCl(aq) ® MgCl2(aq) + H2(g)
  
• 2Al(s) + 6HCl(aq) ® 2AlCl3(aq) + 3H2(g)
  
• Mg(s) + Zn(NO3)2(aq) ® Mg(NO3)2(aq) + Zn(s)
  
• Activity Series
• metals arranged relative to the ease of their oxidation
• most active metals are the easiest to oxidize
• least active metals are the least easy to oxidize
• Examples:
• does a reaction occur between Co(s) and Cu(NO3)2(aq)?
• which is the more active metal?
• the more active metal prefers the oxidized state
  
• Co(s) + Cu(NO3)2(aq) ® Co(NO3)2(aq)    +      Cu(s)
  more        less active
  
• what about Ag(s) and Pb(NO3)2(aq)?
  
• 2Ag(s) + Pb(NO3)2(aq) ® NR
  less            more active

Solution Stoichiometry

• Chemical Analysis of Solutions
• All discussions of stoichiometry apply to solutions as well as solid reactants and products
• use the same format for stoichiometric problems as in chapter 3
• determine the moles of reactant, convert to moles of product
• a solution volume and concentration can give you solute moles
  

• Practice Ex. 4.12:

What volume of 0.500 M HCl(aq) is required to react completely with 0.100 mol of Pb(NO₃)_2(aq), forming a precipitate of PbCl_2(s)?

• write a correct equation:
  2HCl(aq) + Pb(NO3)2(aq) ® PbCl2(s) + 2HNO3(aq)
  

determine amt. HCl to react w/ Pb(NO3)2:
0.100 mol Pb(NO₃)₂             x             2 mol HCl       =         0.200 mol HCl
                                                          1mol Pb(NO₃)₂



convert mol HCl to vol. HCl solution:
0.200 mol HCl        x          1 L sol'n         =            0.400 L or
                                       0.500 mol HCl                   400 mL HCl



• Practice Ex. 4.13:
• What is the molarity of an NaOH solution if 48.0 mL is needed to neutralize 35.0 mL of 0.144 M H2SO4?
  

write the equation for the reaction:
H₂SO_4(aq)       +         2NaOH_(aq)       ®       H₂O       +        Na₂SO_4(aq)
0.144 M                    ? M
0.0350 L                0.0480 L



determine mol of H₂SO_4(aq) :
0.0350 L H₂SO₄          x          0.144 mol H₂SO₄         =             0.00504 mol
                                                       1 L sol'n                                    H₂SO₄




• Practice Ex. 4.13:
• What is the molarity of an NaOH solution if 48.0 mL is needed to neutralize 35.0 mL of 0.144 M H2SO4? H
  
• write the equation for the reaction:
  H2SO4(aq) + 2NaOH(aq) ® H2O + Na2SO4(aq)
  
              
  determine mol of NaOH:
   0.00504 mol H₂SO₄       x             2 mol NaOH          =          0.0101 mol
                                                           1 mol H₂SO₄                         NaOH
  
  

• determine Molarity of NaOH:
  0.0101 mol NaOH            =               0.210 M NaOH
  0.0480 L sol'n

Titrations

• Used commonly, but not exclusively, in neutralization reactions
• the first reactant is "titrated" with the second reactant until stoichiometric equivalence is reached
• the first reactant is added slowly, in small aliquots
  
• this is used to determine:
• the concentration of the first reactant or
• the molarity of the second reactant
  
  

• Practice Ex. 4.14:
• What mass of chloride ion is present in a sample of water if 15.7 mL of 0.108 M AgNO3 is required to titrate the sample?
  
• AgNO_3(aq) + Cl^-_(aq) ® AgCl_(s) + NO₃^-_(aq)
  0.0157 L      ? g
  0.108 M
  
• 0.0157 L     x         0.108 mol Ag⁺        =        0.00170 mol Ag⁺
                                   1 L sol'n
  
• 0.00170 mol Ag⁺     x      1mol Cl^-         =         0.00170 mol Cl^-
                                         1 mol Ag⁺
  
• 0.00170 mol Cl^-       x      35.5 g Cl^-       =          0.0602 g Cl^-
                                             1 mol