General Chemistry I

Chapter 4

Solution Stoichiometry


n    water d/n conduct electricity very well.

n    electrolytes form ions in solution

n    Can be several types of electrolyte

n    Strong electrolyte

n    Weak electrolyte

n    non-electrolyte

How Compounds Dissolve

n    Ionic compounds

n    ‘Break apart’ to form ions

n   NaCl

n   KMnO4

n    Molecular compounds

n    Do not form ions on dissolving

n   May form ions by subsequent reaction with water!

n   NH3, HC2H3O2


n    Reaction that leaves a significant amount of reactant when complete

n    Rate of forward reaction equals rate of reverse reaction

n    HA º H+ + A-

n    º indicates a ‘reversible reaction’

n   Some (but not all) HA forms H+ + A-

n    Do not confuse with solubility!


n     Strong

n    React or dissociate completely to form ions

n    NaCl, HCl, NaOH…

n     Weak

n    Form ions in solution, but most molecules remain non-ionized

n    NH3, HC2H3O2, …

n     Non-electrolytes

n    No ion formation in solution

n    Sugar, ethanol,…

Problems:  1 – 7 odd

Solubility Rules


n    Nitrate, Acetate - NO EXCEPTIONS

n    Chloride -  EXCEPT Ag+, Pb2+, Hg22+

n   For Br-, I- add Hg2+  to the Cl- solubilities

n    Sulfate - EXCEPT Ca2+, Sr2+, Ba2+, Pb2+, Hg22+, Ag+

n    For now, insoluble means less than 0.01M

n    More detail in Chem II


n    Sulfides - IA, IIA, and NH4+

n    Carbonates, Phosphates - IA and NH4+

n    Hydroxides - IA, Ca2+, Sr2+, Ba2+

Ionic Equations

n    Metathesis reaction

n    Double displacement reaction

n    AB + CD 6 AD + CB

n    Something has to happen on the product side!

n   form precipitate

n   form soluble weak or nonelectrolyte

n   form a gas

Ionic Equations

n    Most familiar form is the molecular equation

n    Pb(NO3)2(aq) + 2KI(aq) 6 PbI2(s) + 2KNO3(aq)

n    Total Ionic equations

n    Show strong electrolytes as ions

n    Pb2+(aq) + 2NO3-(aq) + 2K+(aq) + 2I-(aq) 6 

                  PbI2(s) + 2K+(aq) + 2NO3-(aq)

n    Spectator ions

n    Cancel ions with same formula and charge on both sides of the chemical equation.

n    Pb2+(aq) 2I-(aq) 6  PbI2(s)

n    To write:

n    Write balanced molecular equation

n    Write all strong electrolytes (soluble ionic species, strong acids, and strong bases) as ions

n    Cancel out spectator ions

n   Check to see if stoichiometry is still smallest whole number ratio!

Problems:  9 - 15 odd


n    Donate one or more H+ in water

n    Often called ‘proton donors’

n    May be monoprotic or polyprotic

n    Strong acids

n    Completely dissociate to form H+

n    HCl, HBr, HI, HNO3, HClO3, HClO4, H2SO4

n    Weak acids

n    Form some H+ from molecules

n   HC2H3O2, H3PO4, …

n   Form equilibrium with unionized molecule


n    Accept protons from acids

n    Increase OH- concentration in water

n    Strong bases

n    Soluble hydroxides (most important)

n   (MOH, Ba(OH)2, Sr(OH)2, and Ca(OH)2)

n   M is a Group I cation

n    Weak bases

n    Similar properties to weak acids

n    Nitrogen-containing compounds (NH3)

Strong and Weak Electrolytes

n    Strong electrolytes

n    Strong acids

n    Strong bases

n    Soluble ionic species

n    Weak electrolytes

n    Weak acids

n   H written first in chemical formula

n   Not a strong acid!

n    Weak bases (NH3)

Neutralization Reactions

n    Acidic (H+) and base (normally OH-) recombine to form water

n    Can have other bases

n   NH3, S2-, CO32-, …

n    H+ and OH- will have ions associated with them

n    These ‘other’ ions form a salt

n   ‘Salt’ = ionic compound

n   HCl + NaOH 6 H2O + NaCl

n    Base + Acid often gives salt + water

n    Not always

n    NH3 + HCl 6 NH4Cl

n   More detailed look at acids will show presence of water even here!!!!

n    Formation of a gas

n    2H+ + S2- 6 H2S(g)

n      2H+ + CO32- 6 H2CO3 6 H2O + CO2

n   Vinegar and baking soda!


Problems:  17 - 31 odd

Reaction of Metals

n    Oxidation-reduction

n    First redox observed were metals with oxygen, so metal was ‘oxidized’

n    Exchange of electrons

n    Redox reactions

n    After discovery of ‘electrons’ a better understanding of the process was obtained

n   Don’t need oxygen for oxidation!

Oxidation Numbers

n    Electron ‘bookkeeping’

n   Oxidation number (ON) assigned by a set of rules.

n    Elemental form – ON = 0

n    Monoatomic ion – ON = ionic charge

n   Fe2+ means ON = +2

n    Nomenclature – number in parenthesis is the oxidation number!

n   Cl- means ON = -1

n    Nonmetals

n    Fluorine – ON = -1 in all compounds

n   Other halogens usually -1

n    Oxygen – ON usually -2

n   -1 in ‘peroxide’

n    Hydrogen

n   +1 when bonded to nonmetals

n   -1 when bonded to metals

n    ‘hydride’ ion

n    Calculation

n    For a neutral compound, the sum of all oxidation numbers is zero.

n    For a polyatomic ion, the sum of all oxidation numbers is equal to the charge on the polyatomic ion.


Redox Reactions of Metals

n     Much more detail in Chapter 20!

n     Displacement reactions

n    Single replacement reactions

n    Reaction with metallic ion or H+

n    A + BX 6 AX + B

n    Zn + 2AgCl 6 ZnCl2 + 2Ag

n    Zn + 2HCl 6 ZnCl2+ H2


Predicting Displacement Reactions

n    Activity series

n    Used to predict if reaction will occur.

n   Table 4.5, pg 124

n   Shows ‘1/2 reactions’ for the metals

n    If a reactant metal is listed above a product ion , the reaction will occur.


Displacement Reactions

n      Predict if the following reactions will occur!

Fe + 2HCl 6 FeCl2 + H2

n     YES

Cu + 2HCl 6 CuCl2 + H2

n     NO

Mg + FeO 6 MgO + Fe

n      Yes – Sacrificial anode

Problems:  33 - 45 odd


n    Concentration

n    Ratio of solute to amount of solution or solvent.

n    Homogeneous mixture.

n    Molarity

n    Moles of solute to liters of solution.

n    symbol M

n    conversion factor


n    Because it is a ratio of moles of solute to L of solution, we have another path to get to moles.


n    Electrolyte concentrations

n    Express as conc. of salt or individual ions

n   0.020 M Na2SO4 or 0.040 M Na+ and 0.020 M SO42-

n    Dilution

n    addition of more solvent.

n   C1V1 = C2V2

n   volumes can be any unit if they are the same.

n   volumes are volumes of solution.


n    Standard solution

n    solution for which the concentration of the solute is known very accurately.

n    Equivalence point

n    stoichiometrically equivalent amounts of reactants (neither reactant is the limiting reactant, or both reactants are the limiting reactant)

n    use M  as a conversion factor.

Problems:  47 - 73 odd

Learning Goals

n     Calculate molarity, solution volume, or moles of solute given any two of these quantities.

n     Calculate the volume of a more concentrated solution that must be diluted to obtain a given quantity of a more dilute solution.

n     Identify substances as acids, bases, or salts.

n     Predict whether a substance is a nonelectrolyte, a weak electrolyte, or a strong electrolyte from its formula.

n     Predict the ions formed by electrolytes when they dissociate or ionize.

n     Identify the spectator ions and write the net ionic equations for solution reactions starting with their molecular equations.

n    Predict the products of metathesis reactions and write balanced chemical equations for them.

n    Identify the driving force for any metathesis reaction.

n    Use solubility rules to predict whether a precipitate will form when electrolytes are mixed.

n     Use the activity series to predict whether a reaction will occur when a metal is added to an aqueous solution of either a metal salt or an acid; write the balanced chemical molecular and net ionic equations for the reaction.

n     Calculate the concentration or mass of solute in a sample from titration data.