Chemistry 100
Aqueous Reactions
Solutions



A solution is a homogenous mixture
of two or more substances
One substance (generally the one
present in the greatest amount) is
called the solvent
The other substances - those that
are dissolved - are called the
solutes
The Solution Process



Favourable interactions between the
solute and the solvent drive the
formation of a solution
Example: NaCl (an ionic solid)
dissolving in water
Water is a polar fluid (i.e.,
possesses a permanent dipole)
Electrolytes



Salt is an ionic compound.
NaCl is dissolved in water - the
ions separate.
The resulting solution conducts
electricity . A solute with this
property is called an electrolyte
Strong Electrolytes


Strong electrolytes - completely
dissociated
Some molecular compounds
dissolve in water to form ions.


Dissolve HCl (g) in water.
All the molecules dissociate. So it is
also a strong electrolyte.
Weak and Nonelectrolytes


Weak electrolytes - only some of
the molecules dissociate, i.e.,
acetic acid
Compounds that do not dissociate nonelectrolytes



Sugars
Ureas
Alcohols
Acids


Acid - a substance that ionizes in
water to form hydrogen ions H+.
HCl (aq) H+ (aq) + Cl(aq)
What is H+? A hydrogen atom
without its electron - a bare proton.
Monoprotic, diprotic, triprotic



One molecule of HCl gives one H+
ion:
HCl  H+ + Cl
We say that HCl is monoprotic one proton
One molecule of sulphuric acid,
H2SO4, has two hydrogens to give
away. It is said to be diprotic.
Phosphoric acid, H3PO4 is triprotic.
Some Chemical Structures
O
O
H
S
O
both H's ionize
O
H
H
C
H
O
C
OH
H
only this H ionizes
Acetic Acid

Generally write as CH3COOH, not
HC2H3O2.

Weak acid - doesn’t dissociate
completely
CH3COOH (aq) ⇄ CH3COO- (aq) + H+ (aq)
The double arrow - the system is in chemical
equilibrium!!!!
Bases


Bases are substances that accept
(react with) H+ ions. Hydroxide ions,
OH, are basic. They react with H+
ions to form water:
H+ (aq) + OH (aq)  H2O (l)
Ionic hydroxides like NaOH, KOH,
Ca(OH)2 are basic. When dissolved in
water they form hydroxide ions.
Ammonia solution

When ammonia gas dissolves in
water, some NH3 molecules react
with water:
NH3(aq) + H2O(l) ⇄ NH4+ (aq) +
OH– (aq)
NOTE - only some NH3 molecules
react with water. Ammonia is a
weak electrolyte.
Strong and Weak Acids and Bases



Acids and bases that are strong
electrolytes are called strong acids
and strong bases.
Strong acids are more reactive than
weak acids. Likewise for bases.
Note exception - HF, a weak acid,
is very reactive
Acids you should know
Chloric acid
Hydrobromic acid
Hydrochloric acid
Hydroiodic acid
Nitric acid
Perchloric acid
Sulphuric acid
Acetic acid
HClO3
HBr
HCl
HI
HNO3
HClO4
H2SO4
CH3COOH
(weak)
Bases you should know
Know the following bases:
Strong bases
a) Hydroxides of alkali metals: LiOH,
NaOH, KOH
b) Hydroxides of the heavy alkaline
earth metals: Ca(OH)2, Sr(OH)2,
Ba(OH)2
Weak base: ammonia solution NH3

Metathesis reactions
A metathesis reaction is an aqueous
solution in which cations and anions
appear to exchange partners.
AX + BY  AY + BX
AgNO3 (aq)+ NaCl (aq)  AgCl (s) +
NaNO3 (aq)

Metathesis reactions (cont.)


Three driving forces
Precipitate formation (insoluble
compound)
AgNO3(aq)+ NaCl(aq)  AgCl(s) +
NaNO3(aq)
Metathesis Reactions (Cont’d)
Weak electrolyte or nonelectrolyte
formation
HCl(aq) + NaOH(aq)  NaCl(aq)
+ H2O(l)
 Gas formation
2HCl(aq) + Na2S(aq)  2 NaCl(aq)
+ H2S(g)

Neutralization



Mix solutions of acids and bases - a
neutralization reactions occurs.
acid + base  salt + water
Salt does not necessarily mean
sodium chloride!!!!
Salt - an ionic compound whose
cation (positive ion) comes from a
base and whose anion (negative ion)
comes from an acid
Precipitation Reactions


Some ionic compounds are insoluble
in water.
If an insoluble compound is formed
by mixing two electrolyte solutions,
a precipitate results.
Precipitation (Cont’d)




Solubility - maximum amount of
substance that will dissolve in a
specified amount of solvent.
Saturated solution of PbI2 contains
1 x 10-3 mol/L.
A compound with a solubility of less
than 0.01 mol/L - insoluble.
More accurately - sparingly soluble.
Solubility Fact 1

All the common ionic compounds of
the alkali metals are soluble in
water. The same is true of the
compounds containing the
ammonium ion, NH4+.
NaCl, K2CO3, (NH4)2S are all soluble
Solubility Fact 2
Salts containing the following anions
are soluble
Anion
NO3
CH3COO 
Cl 
Br 
I
SO42
nitrate
acetate
chloride
bromide
iodide
sulphate
exception, salts of
none
none
Ag+, Hg22+,Pb2+
Ag+, Hg22+,Pb2+
Ag+, Hg22+,Pb2+
Ca2+, Sr2+, Ba2+,
Hg22+, Pb2+
Solubility Fact 3

Salts containing the following
anions are insoluble
Anion
S2
sulphide
CO32 carbonate
PO43
phosphate
OH
hydroxide
exception, salts of
alkaline metal cations,
NH4+, Ca2+, Sr2+, Ba2+,
alkaline metal cations,
NH4+
alkaline metal cations,
NH4+
alkaline metal cations,
Ca2+, Sr2+, Ba2+,
Reaction forming gases


A metathesis reaction can occur due
to the formation of a gas which is
not very soluble in water.
Examples involving hydrogen
sulphide and carbon dioxide
Reactions forming H2S


A metathesis reaction occurs when
hydrochloric acid is added to a
sodium sulphide solution.
2HCl(aq) + Na2S(aq)  H2S(g) +
2NaCl(aq)
Net ionic reaction:
2H+(aq) + S2(aq)  H2S (g)
Reactions involving CO2

Carbonates and bicarbonates may
be thought of as the salts of
carbonic acid H2CO3 – unstable!!
H2CO3(aq)  CO2(g) + H2O(l)
Ionic Equations
Consider the reaction
HCl (aq) + NaOH (aq)  NaCl (aq) +
H2O (l)
 The above is known as the
molecular equation
 Note: the compounds are ionic
(except water)!!

Ionic Equations #2
Let’s show ionic compounds as ions
H+(aq) + Cl–(aq) + Na+(aq) + OH–
(aq) 
Na+(aq) + Cl–(aq) + H2O(l)
 Some ions appear on both sides of
the equation.

Out with the spectators!
Remove ions that appear on both
sides
H+ (aq) + Cl– (aq) + Na+ (aq) + OH–
(aq) 
Na+ (aq) + Cl– (aq) + H2O (l)
 The unchanged ions are called
spectators

The Net Ionic Equation

We are left with is the net ionic
equation:
H+(aq) + OH– (aq)  H2O(l)
Note that the equation is balanced
for both mass and charge!!!
Another ionic reaction

Place zinc metal in a hydrochloric
acid solution – hydrogen is
evolved!!
Zn (s) + 2HCl (aq)  ZnCl2 (aq) +
H2 (g)
Why use ionic reactions?

They summarize many reactions.



neutralization of any strong acid by a
strong base is given by H+(aq) + OH–
(aq)  H2O(l)
The chemical behaviour of a strong
electrolyte  behaviour of its
constituent ions.
Ionic equations can be written only
for strong electrolytes which are
soluble.
Concentrations


How do we express the
concentration of a solution?
Percentage is one way.




2% milk
35% cream. (These are not true
solutions)\
Some beer is 5% alcohol
Note: % measurements can be
%w/w, %w/v, %v/v
Molarity


Must work in moles to do chemical
arithmetic.
Chemists - molarity as their unit of
solution concentration
moles of solute
Molarity 
volume of solution (L)
Dilution

Dilute a solution

more solvent is added but the amount
(mass or moles) of solute is
unchanged.
M1V1 = M2V2
The volumes can be either millilitres
(mL) or litres (L).
Ionic Concentration


NaCl in water - totally ionized into
Na+ and Cl ions.
A 2.0 M NaCl solution



Na+ concentration will be 2.0 M
Cl concentration also 2.0 M
A 2.0 M solution of K2CO3,


K+ concentration will be 4.0 M
The concentration of CO32  2.0 M.
Oxidation and reduction
A piece of calcium metal exposed to
the air will react with the oxygen in
the air
2Ca(s) + O2(g)  2 CaO(s)
 Ca has been converted to an ion Ca2+
by losing two 2 electrons.
 Dissolve Ca in acid
Ca(s) + 2H+(aq)  Ca2+(aq) + H2(g)
 Again the Ca has lost 2 electrons —
oxidation

Redox reactions




In the last two reactions, the Ca
atom lost two electrons. Where did
they go?
When one substance is oxidized,
another is reduced. An oxidationreduction reaction occurs. Or a redox
reaction occurs.
Oxidation: loss of electrons (more
positive)
Reduction: gain of electrons (less
positive)
Oxidation of Metals - by air

Many metals react with oxygen in the air.



Fe rusts - at a cost of $billions each year!
Aluminum oxidizes




Na and K do so explosively!
oxide layer forms a skin which prevent further
oxidation. Al hides its reactivity.
Gold and platinum do not react with
oxygen.
Silver tarnishes mainly because of H2S in
the air.
What does copper do?
Oxidation of Metals - by acids


Many metals react with acids:
metal + acid  salt + hydrogen gas
Mg(s) + 2HCl(aq)  MgCl2(aq) +
H2(g)
Metals may also be oxidized by the
salts of other metals. Recall your lab
experiment
Fe(s) + CuSO4(aq)  Cu(s) +
FeSO4(aq)
Activity Series



We has seen that some metals react
with air, some also react with acids
to give hydrogen.
We have seen that some metals can
be oxidized by ions of other metals.
All this is summarized in the activity
series.
Activity Series
Li
K
Ba
Ca
Na
Mg
Zn
Fe
Pb
H
Cu
Ag
Au













Li+
K+
Ba2+
Ca2+
Na+
Mg2+
Zn2+
Fe2+
Pb2+
H+
Cu2+
Ag+
Au3+
+
+
+
+
+
+
+
+
+
+
+
+
+
e
e
2e
2e
e
2e
2e
2e
2e
e
2e
e
3e




A metal can be oxidized
by any ion below it
Metals above H, react
with acids to give H2
The further up the series,
the more readily the
metal is oxidized
See your textbook (p
145) for more elements
Some observations on the series



Lead (Pb) is above H, so is Al. But these
metals are not attacked by 6M HCl. They
form very protective oxides.
Cu reacts with nitric acid (HNO3) because
that acid is a strong oxidizing agent in
addition to being an acid.
Gold (Au) and platinum (Pt) are valuable
because they are (a) rare and (b)
unreactive - they do not tarnish
Oxidation Numbers


Oxidation number - a fictitious
charge assigned to atoms either by
themselves or when combined in
compounds as an electron
bookkeeping device.
There are a number of simple rules
that chemists use to assign
oxidation numbers.
Assigning Oxidation Numbers

In any elemental form (atom or
molecule), an atom is assigned a 0
oxidation number


e.g. He, Cu, N in N2, S in S8
For a monatomic ion, the oxidation
number equals the charge

e.g., -1 for Cl in Cl-, +2 for Ca+2, -2 for
S-2
Assigning Ox. Numbers (#2)

Fluorine’s oxidation number is -1 in
any compound.


e.g. -1 for F in CF4, but 0 for F in F2
Oxygen’s oxidation number is -2
except when combined with fluorine
or in peroxides.

e.g. -2 for O in H2O and OH-, +2 for O
in OF2, -1 for O in H2O2
Assigning Ox. Numbers (#3)

For elements in Groups IA, IIA &
most of IIIA, oxidation numbers are
positive and equal to the group
number.


e.g. +3 for Al in AlCl3, +1 for Na in NaCl,
+2 for Mg in Mg SO4
Hydrogen has a +1 oxidation
number. Exceptions to this rule are
the metallic hydrides, in which it is 1.

e.g., +1 for H in H2O and CH3OH, -1 for
H in NaH
Assigning Ox. Numbers (#4)

The sum of the oxidation numbers
of the atoms in a neutral compound
is zero; in a polyatomic ion, the
sum equals the charge.

e.g. see OH- and H2O above, +6 for S
in SO4-2
Balancing Oxidation-Reduction
(Redox) Equations (#1)

Assign oxidation numbers to all
atoms in the equation.

Note - polyatomic ion that is
unchanged in the reaction may be
treated as a single unit with an
oxidation number equal to its charge.
Balancing Redox Equations (#2)

Isolate the ATOMS that have
undergone a change of oxidation
number


A reduction in number indicates a
reduction
An increase in number, an oxidation
Balancing Redox Equations (#3)


Isolate the chemical species
undergoing oxidation/reduction
(note: separate into an oxidation and
a reduction half-reaction).
Add the appropriate number of
electrons to the half-reactions

Oxidation – electrons on products side

Reduction – electrons on reactants side
Balancing Redox Equations (#4)

Remaining steps refer to the
individual half reactions


Balance for charges

Add H+ in acidic solution

Add OH- in basic solution
Balance the H and the O atoms by
adding water
Balancing Redox Equations (#5)

Balance the number of electrons in
the half-reactions


Note: electrons lost = electrons
gained
Add the half-reactions, eliminating
the electrons and obtaining the
complete REDOX equation
Titrations

Volumetric analysis  technique
based on volume measurements


used to determine the quantity of a
substance in solution.
Titration  a solution of an
accurately known concentration is
added gradually to a solution of an
unknown concentration

Reaction goes to completion.
Other Definitions


Standard solution  solution of
accurately known concentration.
Equivalence point  point at which
unknown substance has completely
reacted with standard solution.

At the equivalence point reagents are
present in stoichiometric amounts.
Gravimetric Analysis

Determine concentration of an
unknown by reacting it with a
second substance to form a ppt.
AgNO3(aq)+ NaCl(aq)  AgCl(s) +
NaNO3(aq)