Transcript Chapter 4 - Aqueous Reactions

Chemistry 100
Aqueous Reactions
Solutions
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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
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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
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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
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Strong electrolytes - completely
dissociated
Some molecular compounds dissolve
in water to form ions.
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Dissolve HCl (g) in water.
All the molecules dissociate. So it is also
a strong electrolyte.
Weak and Nonelectrolytes
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Weak electrolytes - only some of the
molecules dissociate, i.e., acetic acid
Compounds that do not dissociate nonelectrolytes
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Sugars
Ureas
Alcohols
Acids
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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
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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
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Generally write as CH3COOH, not
HC2H3O2.
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Weak acid - doesn’t dissociate
completely
CH3COOH (aq) ⇄ CH3COO- (aq) + H+ (aq)
The double arrow - the system is in chemical
equilibrium!!!!
Bases
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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
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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
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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
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Metathesis reactions
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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.)
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Three driving forces
Precipitate formation (insoluble
compound)
AgNO3(aq)+ NaCl(aq)  AgCl(s) +
NaNO3(aq)
Metathesis Reactions (Cont’d)
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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
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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
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Some ionic compounds are insoluble
in water.
If an insoluble compound is formed by
mixing two electrolyte solutions, a
precipitate results.
Precipitation (Cont’d)
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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
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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
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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
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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
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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
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Carbonates and bicarbonates may be
thought of as the salts of carbonic acid
H2CO3 – unstable!!
H2CO3(aq)  CO2(g) + H2O(l)
Ionic Equations
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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.
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Out with the spectators!
Remove ions that appear on both
sides
H+ (aq) + Cl– (aq) + Na+ (aq) + OH– (aq)
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Na+ (aq) + Cl– (aq) + H2O (l)
 The unchanged ions are called
spectators
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The Net Ionic Equation
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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
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Place zinc metal in a hydrochloric acid
solution – hydrogen is evolved!!
Zn (s) + 2HCl (aq)  ZnCl2 (aq) + H2
(g)
Why use ionic reactions?
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They summarize many reactions.
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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
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How do we express the concentration
of a solution?
Percentage is one way.
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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
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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
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Dilute a solution
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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
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NaCl in water - totally ionized into Na+
and Cl ions.
A 2.0 M NaCl solution
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Na+ concentration will be 2.0 M
Cl concentration also 2.0 M
A 2.0 M solution of K2CO3,
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K+ concentration will be 4.0 M
The concentration of CO32  2.0 M.
Oxidation and reduction
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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
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In the last two reactions, the Ca atom lost
two electrons. Where did they go?
When one substance is oxidized, another is
reduced. An oxidation-reduction reaction
occurs. Or a redox reaction occurs.
Oxidation: loss of electrons (more positive)
Reduction: gain of electrons (less positive)
Oxidation of Metals - by air
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Many metals react with oxygen in the air.
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Fe rusts - at a cost of $billions each year!
Aluminum oxidizes
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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
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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
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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
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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
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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
136) for more elements
Some observations on the series
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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
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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
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In any elemental form (atom or
molecule), an atom is assigned a 0
oxidation number
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e.g. He, Cu, N in N2, S in S8
For a monatomic ion, the oxidation
number equals the charge
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e.g., -1 for Cl in Cl-, +2 for Ca+2, -2 for S-2
Assigning Ox. Numbers (#2)
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Fluorine’s oxidation number is -1 in
any compound.
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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.
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e.g. -2 for O in H2O and OH-, +2 for O in
OF2, -1 for O in H2O2
Assigning Ox. Numbers (#3)
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For elements in Groups IA, IIA & most
of IIIA, oxidation numbers are positive
and equal to the group number.
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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)
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The sum of the oxidation numbers of
the atoms in a neutral compound is
zero; in a polyatomic ion, the sum
equals the charge.
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e.g. see OH- and H2O above, +6 for S in
SO4-2
Balancing Oxidation-Reduction
(Redox) Equations (#1)
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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)
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Isolate the ATOMS that have
undergone a change of oxidation
number
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A reduction in number indicates a
reduction
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An increase in number, an oxidation
Balancing Redox Equations (#3)
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Isolate the chemical species undergoing
oxidation/reduction (note: separate into
an oxidation and a reduction halfreaction).
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Add the appropriate number of
electrons to the half-reactions

Oxidation – electrons on products side

Reduction – electrons on reactants side
Balancing Redox Equations (#4)
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Remaining steps refer to the individual
half reactions
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Balance for charges
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Add H+ in acidic solution
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Add OH- in basic solution
Balance the H and the O atoms by
adding water
Balancing Redox Equations (#5)
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Balance the number of electrons in the
half-reactions
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Note: electrons lost = electrons
gained
Add the half-reactions, eliminating the
electrons and obtaining the complete
REDOX equation
Titrations
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Volumetric analysis  technique
based on volume measurements
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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
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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)