General Chemistry I
Chapter
4
Solution
Stoichiometry
Electrolytes
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
Equilibrium
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!
Electrolytes
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
SOLUBLE EXCEPT:
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 INSOLUBLE
EXCEPT:
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
Acids
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
Bases
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
YOU
CAN’T HAVE OXIDATION WITHOUT REDUCTION!
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
Solutions
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
Molarity
n Because
it is a ratio of moles of solute to L of solution, we have another path to get
to moles.
Solutions
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.
Titrations
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.