Aqueous Reactions and
Solution Stoichiometry
► Aqueous
Solutions – a
solution which water is
the solvent
► Solution = Solute +
Solvent
► Solute = smaller part
being dissolved
► Solute = larger fraction
doing the dissolving
► Molarity is used to
express conc.
Molarity
►M
= moles of solute/ liters of solution
► A 0.38g sample of sodium nitrate is placed
in a 50.0 ml volumetric flask and is filled
with water to the mark. What is the molarity
of the solution?
Molarity
► How
many grams of
potassium
permanganate are
needed to prepare 2.0
liters of a 0.2 molar
solution of the salt?
Dilution Problems
► M1V1=
M2V2
► How many liters of 12 molar sulfuric acid is
necessary to make 2 liters of a 6 molar
conc. of the acid?
Specific Gravity or Density Problems
► Commercially
available conc. HCl is in a aq
sol cont 38% HCl by mass. D = 1.19g/ml a)
what is the molarity of the solution? b) how
many ml of conc HCl are required to make
1.00 l of a 0.10 molar solution?
Electrolytes
► Solutes
that separate
into tow or more ions
when dissolved in
water are called
electrolytes.
► Solutes that remain
uncharged molecules
are called nonelectrolytes.
► Either may be soluble
in water.
Electrolytes
► NaCl(s)
 Na+ + Cl- ionic solids
dissociate into ions.
► Electrical Conductivity
Strong electrolytes conduct electricity very
well because ions totally dissociate
Weak electrolytes only weakly conduct
because ions only partially dissociate
Nearly all ionic cmps are strong and
molecular cmps are usually non-electrolytes
Acids, Bases, And Salts
► Acids
– are solutes
that can ionize to
produce hydrogen ions
H+(aq)
► Strong acids
dissociate and exist
almost totally as ions
► Ex
HCl(aq)  H+ + Cl-
Weak Acids
► HC2H3O2 H+ +
C2H3O2-

Weak acids only partially
dissociate the double arrows
indicate that the reaction
occurs in significant amounts
in both directions
Acedic Acid or Vinegar is the
classic example.
Acids Cont.
► H+
is simply a proton – acids are often
called proton donors
► Monoprotic acids – contribute one H+ per
molecule of acid ex. HCl or HNO3.
► Diprotic acids contribute two H+ per
moleclue ex. H2SO4
► Polyprotic acids contribute greater than one
H+ ex. H2SO4 or H3PO4.
Bases
► Are
substances that accept H+ ions or a
proton.
► Ex H+(aq) + OH-(aq)  H2O(l)
► Ex H+(aq) + NH3-((aq)  NH4(aq)
► OH- & NH3- are the bases. OH- is the most
common base in aqueous solution.
Sometimes bases are defined as substances
with increased OH- ions in solution.
Note
► Acids
and bases are electrolytes – if a strong acid
or base they are strong electrolytes if a weak acid
or base a weak electrolyte.
► Strong Acids
HCl (aq) Hydrochloric
H2SO4 Sulfuric
HNO3 (aq) Nitric
HBr Hydrobromic
HClO4 (aq) Prechloric
HI
Hyrooiodic
HClO3 (aq) chloric
► Weak
Acids
HF
Hydrofluoric
H2C2H3O2 Acedic
► Strong bases
Grp I Hydroxides
LiOH, NaOH, KOH, RbOH, and CsOH
Grp II Hydroxides
Ca(OH)2, Sr(OH)2, Ba(OH)2
Weak Base NH3 Ammonia
Salts
► Are
ionic compounds
formed from acids by
the replacement of one
or more H+ ions with
some other cation.
► For example replacing
the H+ ion in HCl(aq)
with Na+ ion yields
NaCl.
Identifying strong and weak
electrolytes
► most
salts are strong
electrolytes
► most acids are weak
electrolytes except for
the strong acids
► most strong bases
hydroxides of metals
are strong electrolytes
except NH3 which is a
weak base and
electrolyte
► Neutralization
reactions are those between an acid
and a base they yield salt and water
► Ionic Equations
a. varying amounts of detail can be included in
chem eq depending on what information is
relevant to the problem at hand
b. molecular equations - have all the species
written as associated molecules even though they
may be disassociated
c. complete ionic eq – strong electrolytes are
written as dissociated ions; molecules, solids, and
weak electrolytes are written as associated ions
d. net ionic equations – an ionic eq in which
the spectator ions have been cancelled.
► Rules for Converting Molecular eq to Ionic
1. Ionic Sub indicated as dissolved in
solution such as NaCl(aq) are normally
written as ions
2. Ionic substances that are insoluble (do
not dissolve readily) either as reactants or
products (precipitates) are rep as formulas
of the compounds
d. Molecular substances that are strong
electrolytes such as strong acids are written
as ions.
e. Molecular substances that are weak
electrolytes are represented as their
molecular formulas
f. Spectator ions – ions in an equation that
do not take part in the equation are
cancelled from both sides of the equation
Examples
► Ba(C2H3O2)2
+ K2CrO4 
► Ba(C2H3O2)2
+ K2CrO4  BaCrO4 + 2KC2H3O2
Ba + 2C2H3O2 + 2K + CrO4  BaCrO4 + 2K
+2 C2H3O2
Ba + CrO4  BaCrO4
Examples
► Ca(C2H3O2)2
+ H2C2O4 
► Ca(C2H3O2)2
+ H2C2O4  CaC2O4 + HC2H3O2
► Ca+
2C2H3O2 + H2C2O4  CaC2O4 + HC2H3O2
► Weak
acids do not split
Examples
► NH4NO3 +
►
NH4NO3 + NaOH  NH4OH + Na + NO3
► NH4 + NO3
►
NaOH 
+ Na + OH  NH3 + H2O + Na
+ NO3
NH4 + OH  NH3 + H2O
Metathesis Reactions
► Reactions
in which two ionic reactants
exchange ion partners
► AX + BY  AY + BX
► Precipitation reactions are those metathesis
reactions in which an insoluble solid product
or precipitate forms
A. Precipitate –insoluble solid product
B. Solubility – the amount of a substance
that can dissolve in a given amount of
solvent
► (expressed
in g/100 ml or in liters/mol) A
substance is considered insoluble if its solubility is
less that 0.01 mol/l.
► Predicting solubility is a matter of know the
solubility rules (Table 4.3 or what I should know
for AP list)
► Some metathesis reactions from water or another
weak electrolyte or non-electrolyte in ionic and net
ionic equations (appear as associated species)
► Some methesis reactions form gases these are
usually binary compounds of non-metals exam
ples H2S, CO2, HCN.
Redox It’s electro-magic
► I.
Oxidization Numbers
A. Oxidization – loss of
electrons. LEO
B. Reduction – gain of
electrons – GER
C. When one reactant
loses e’s another
reactant must gain
them
Intro to electrochemistry
D. reactions in which
electrons are trans are
called oxidization –
reduction or redox
reactions
► E. oxidization number (ON)
rules
1. atoms in the elemental
form – oxidization # = 0
►
2. monatomic ions – the
oxidization # = the charge
of the ion
ex Na+ = 1 S-2 = -2
3. non-metals usually have
neg ON although they can
sometimes be positive
a> Oxygen usually =
-2 except peroxides
O2-2 = each O = -1
b) hygrogen usually = +1 when bonded to
non-metals = -1 when bonded to metals
c) Fluorine = -1 in all cmps – halogens = -1 in
most binary cmps but when combine with
oxygen as in oxyanions they have a +
number
4. The sum of oxidizations #’s
a) in a neutral cmp = 0
b) in a polyatomic ion = charge of ion
Work sample exercise 4.6