6 Reactions in Aqueous Solutions
Water is by far the most common medium in which chemical reactions occur naturally. It is not hard to see
this: 70% of our body mass is water and about 70% of the surface of the earth is covered in water.
The types of reactions in aqueous (water) medium that will be studied are as follow:
1. Precipitation Reactions
2. Acid-Base Reactions
3. Oxidation-Reduction Reactions (Special Case of the Reaction of Metals with Nonmetals)
Recall that a solution is a homogeneous (uniform) mixture. A solution consists of two generic components:
• Solute – the component (compound) that is in minor quantities; and
• Solvent – the component that is in major quantity.
Aqueous solutions then have water as the solvent.
6.1 Precipitation Reactions
Precipitation reactions are those reactions in which two aqueous solutions, when mixed, react to form a
solid. The solid thus formed is known as a precipitate. Thus, a precipitate is an insoluble solid (recall
that an insoluble substance is one which does not dissolve). Indeed, if a solid that forms is soluble, then as
soon as it forms it would dissolve back into solution – so if a solid is soluble, it cannot, in general, be a
precipitate.
We have defined precipitates as insoluble solids. More specifically, we know that precipitates are insoluble
ionic compounds. We must now ask what the mechanism for precipitation is.
When an ionic substance is dissolved in water, we indicate this by appending an (aq) after the formula of
the compound. Thus, for example, NaCl (aq) designates an aqueous solution of sodium chloride,
Fe(NO2)3 (aq) designates an aqueous solution of iron (III) nitrite and so on.
It is known that when an ionic substance dissolves in water it dissociates into its component ions. The
evidence for this comes from electrical conductivity studies:
•
Electrical current arises from the movement of charged substances
•
Water does not conduct electricity, therefore it has no ions itself (this is actually a false assertion,
but the concentration of ions in pure water is small)
Thus, if an aqueous solution conducts electricity, it must contain charged substances (ions) that are free to
move – these ions can only come from an ionic solute (since we have made the assertion that water itself
does not provide ions).
For example, when NaCl is dissolved in water, there will be, in principle, no unit of NaCl found anywhere in
the aqueous solution. Rather, we will see a “soup” consisting of Na+ ions and Cl– ions floating around.
Although this is not regarded as a conventional reaction, we may write the process in the form of a
chemical equation:
NaCl (aq)
Na+ (aq) + Cl– (aq)
The (aq) attached to an ionic substance always means the component ions in the correct proportions, and
that the ionic substance is soluble.
NaCl (aq) means the same as Na+ (aq) + Cl– (aq)
It is thus very important to recognize ionic compounds and to know what the component ions are.
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Ionic Compound
copper (II) sulphate
Barium nitrate
Potassium hydroxide
Sodium hydrogen carbonate
Ammonium dihydrogen phosphite
Ions in Solution
CuSO4 (aq)
Means
Cu2+ (aq) and SO42 (aq)
Ba(NO3)2 (aq)
Means
Ba2+ (aq) and 2 NO3 (aq)
KOH (aq)
Means
K+ (aq) and OH– (aq)
NaHCO3 (aq)
Means
Na+ (aq) and HCO3 (aq)
(NH4)H2PO3 (aq)
Means
NH4+ (aq) and H2 PO3 (aq)
What happens when two different aqueous solutions are mixed? Let us look at the specific example of a
solution of an aqueous solution of sodium chloride mixed with an aqueous solution of silver nitrate. We
know that both are ionic. (How?).
Thus:
NaCl (aq)
Na+ (aq) + Cl– (aq)
and
AgNO3 (aq)
Na+ (aq)
and
Cl– (aq)
Ag+ (aq) +
NO3 (aq)
Ag+ (aq)
and
NO3 (aq)
What we have now is a “soup” consisting of Na+, Cl– , Ag+ and NO3 floating around in water. These
atoms are bouncing around the container colliding with each other. If Na+ ions collide with other Na+ ions,
they will repel each other (they have the same charge). Likewise, if Ag+ions collide with Na+ ions or other
Ag+ ions, they will repel each other. Similarly, any combination of collisions between anions will be
repulsive.
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If the Na+ ions collide with Cl– , they will attract each other to form NaCl. However, since NaCl is soluble, it
will immediately break back up into ions. A similar situation arises when Ag+ collides with NO3 ions.
Thus the only new things that can happen is for Na+ ions to collide with NO3 ions to form NaNO3 and for
Ag+ to collide with Cl– to form AgCl. It is an experimental fact that NaNO3 is soluble therefore it cannot form
a precipitate; it is also known experimentally that AgCl is insoluble, therefore it will form a precipitate.
Complete
Molecular
Equation
+
NaCl (aq)
NaNO3 (aq)
AgNO3 (aq)
Na+ (aq) + Cl– (aq) + Ag+ (aq) +
Complete
Ionic Equation
NO3 (aq)
+
Na+ (aq) + NO3 (aq) +
AgCl (s)
AgCl (s)
Note that in the complete ionic equation, Na+ (aq) and NO3 (aq) appear on both sides of the equation.
This means that these two ions merely serve to ensure that overall, all charges are balanced. They serve
no other function in the chemical reaction. Ions which do not participate in the chemical equation are
known as spectator ions. Spectator ions are ions that appear on both sides of the complete ionic
equation.
If the spectator ions are cancelled as if they were variables in an algebraic equation, then we are left with
the net ionic equation:
AgCl (s)
Ag+ (aq) + Cl– (aq)
The net ionic equation describes the basic chemical process that is occurring by cutting out all nonessential information.
For the following reactions, write the complete molecular equation, the complete ionic equation and the net
ionic equation. Identify the spectator ions.
e.g.
BaCl2 (aq) +Na2SO4 (aq)
Step 1:
Write out the ions in the reactants.
Ba2+ (aq) + 2 Cl– (aq) + 2 Na+ (aq) + SO42
Step 2:
Examine all collisions that can occur and disregard those that cannot lead to a reaction.
[Ba2+ (aq) + 2 Cl– (aq)] + [2 Na+ (aq) + SO42 (aq)]
Step 3:
[Ba2+ (aq) + SO42 (aq)] + [2 Na+ (aq) + 2 Cl– (aq)]
Balance the charges on the product side, being sure to balance them on the reactant
side also. Note that you must be careful in doing this by paying attention to which
chemical produced which ions. Put the complete product molecules together.
[Ba2+ (aq) + 2 Cl– (aq)] + [2 Na+ (aq) + SO42 (aq)]
Step 4:
BaSO4 (s) + 2NaC l(aq)
Write the complete ionic equation
Ba2+ (aq) + 2 Cl– (aq) + 2 Na+ (aq) + SO42 (aq)
Step 7:
BaSO4 (s) + 2NaC l(aq)
Recombine the reactant side ions to give the complete molecular equation.
BaCl2 (aq) +Na2SO4 (aq)
Step 6:
BaSO4 + 2NaCl
From your knowledge of solubilities, determine which (if any) of the products will form a
precipitate. Recall that a precipitate is an insoluble solid.
Ba2+ (aq) + 2 Cl– (aq)] + [2 Na+ (aq) + SO42 (aq)]
Step 5:
?
BaSO4 (s) + 2 Na+ (aq) + 2 Cl– (aq)
Cancel off the spectator ions and write the net ionic equation
Ba2+ (aq) + SO42 (aq)
BaSO4 (s)
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Solubility Rules for Some Common Salts in Water
Note well: When an insoluble salt (ionic compound) forms when two or more solutions are mixed, then the
substance will precipitate out of solution. If the salt is soluble no precipitate will form.
1. Most nitrate ( NO3 )salts are soluble.
+
+
+
+
+
2. Most salts containing the alkali metal ions (Li , Na , K , Cs , Rb ) and the ammonium ion ( NH+4 )
are soluble.
3. Most chloride, bromide and iodide salts are soluble.
Notable exceptions:
Salts containing the ions Ag+, Pb2+ and Hg22+ .
4. Most sulfate salts ( SO24 ) are soluble.
Notable exceptions:
CaSO4 , BaSO4 , PbSO4 and Hg2SO4 .
–
5. Most hydroxide salts (OH )are only slightly soluble. The important soluble hydroxides are NaOH
and KOH. The Group IIA hydroxides are marginally soluble.
2
6. Most sulphide ( S ) , carbonate ( CO32 ) , chromate ( CrO 24 ) and phosphate salts ( PO34 ) are
only slightly soluble.
These solubility rules must be memorized. They will be used to predict whether a precipitate will form.
Exercises:
Fe(NO3)2 (aq) + NaOH (aq)
(NH4)3PO4 (aq) + AlCl3 (aq)
NiCl2 (aq) + Na2CO3 (aq)
Pb(NO3)2 (aq) + NaI (aq)
Precipitation reactions are sometimes known as double displacement reactions because two ions
exchange with each other. Schematically:
AB + XY
AY + XB
Double Exchange
That is, cation A exchanges places with cation X and anion B exchanges with anion Y.
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6.2 Acids and Bases
Arrhenius Definitions:
• Acids: Substances which, when dissolved in water, produce H+.
• Bases: Substances which, when dissolved in water produce OH–.
Strong acids and strong bases are acids that are excellent conductors of electricity.
The list of strong acids (to be memorized) is:
HCl, HBr, HI, HNO3 , H2SO4 , and HClO4
The strong bases are: the alkali metal hydroxides and the alkaline earth hydroxides (even though these
latter are only sparingly soluble.
6.2.1 Dissociation of Acids
As with other ionic substances, acids dissociate into their component ions when dissolve in water.
However, the dissociation involves the removal of only one of the protons to produce, H+ and the
conjugate acid anion.
Examples:
H2SO4 (aq)
H+ (aq) + HSO4 (aq)
H3PO4 (aq)
H+ (aq) + H 2 PO42 (aq)
H2S (aq)
NOT
[H2SO4 (aq)
2 H+ (aq) + SO42 (aq) ]
H+ (aq) + HS (aq)
Obviously this is not an issue for monoprotic acids (acids having only one ionizable proton, e.g., HCl, HNO3
etc.)
6.2.2 Acid-Base Reactions
Acid-Base reactions, reactions between acids and bases, are sometimes referred to as neutralization
reactions. The product of neutralization reactions are always a salt (an ionic compound) and water.
e.g.
HCl (aq) + NaOH (aq)
[H+ (aq) + Cl– (aq)] + [Na+ (aq) + OH– (aq)]
NaCl (aq) + H2O (l)
Double Exchange
Complete Molecular Equation: HCl (aq) + NaOH (aq)
Complete Ionic Equation:
Net Ionic Equation:
e.g.
+
–
NaCl (aq) + H2O (l)
+
[H (aq) + Cl (aq)] + [Na (aq) + OH– (aq)]
Na+ (aq) Cl– (aq) + H2O (l)
H+ (aq) + OH– (aq)]
H2O (l)
This is the net ionic equation for all acid-base reactions
H2SO4 (aq) + NaOH (aq)
Ans:
H2SO4 (aq) + 2 NaOH (aq)
2 H2O (l) + Na2SO4 (aq)
Note that Acid-Base reactions are double displacement reactions.
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6.3 Oxidation Reduction Reactions
In the double displacement reactions discussed before (precipitation and acid-base reactions), we see a
transfer of ions. Oxidation-Reduction reactions (Redox reactions) there is a transfer of electrons.
Let us examine the reaction of sodium with chlorine.
2 Na (s) + Cl2 (g)
2 NaCl (s)
Here, initially the oxidation number of Na = 0 and of Cl = 0 on the reactant side of the equation. On the
product side, NaCl is formed from the attraction of Na+ and Cl– ions. Clearly, in the process of forming NaCl
from its elements, sodium has lost an electron and chlorine has gained an electron:
Na
Na+ + e–
Cl + e–
Cl–
Thus it can be observed that NaCl forms because of a transfer of electrons from sodium to chlorine. In the
process, sodium changes its oxidation state from 0 to +1, whereas chlorine changes its oxidation state
from 0 to –1. The change in oxidation state requires a transfer of electrons.
Oxidation is defined to be the loss of electrons. The substance becomes more positive.
Reduction is defined to be the gain of electrons. The substance becomes more negative.
Thus we see in the above reactions that:
•
Sodium is oxidized
•
Chlorine is reduced.
************************
For all reactions, you should always check the oxidation numbers of the reactants and of product. If there
is a change of oxidation state then the reaction is a redox reaction. If you can observe that the reaction is a
double displacement or an acid base reaction, it is not necessary to do this – double displacement
(precipitationand acid-base reactions) are never redox reactions also.
6.4 Other Ways of Classifying Chemical Reactions
There are several other ways of classifying chemical reactions. Very often, a particular reaction can be
characterized in more than one ways. Three of those ways will be shown below.
6.4.1 Combustion Reactions
Combustion reactions usually refer to chemical reaction with molecular oxygen to form oxides (other
substances, such as Cl, will also support combustion but such cases are not often considered).
Examples:
The combustion of magnesium: 2 Mg(s) + O2 (g)
2 MgO (s)
3 CO2 (g) + 4 H2O (l)
The combustion of propane:
C3H8 (g) + 5 O2 (g)
Note:
1. Organic molecules are molecules that are composed of at least C and H atoms. Their combustion
products are always CO2 and H2O.
2. Combustion reactions are also redox reactions.
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6.4.2 Synthesis Reactions
Synthesis reactions are reactions that produce more products that are more complex than the reactants.
Examples are:
2 Mg(s) + O2 (g)
2 MgO (s)
2 H2 (g) + O2 (g)
2 CH4 (g)
2 H2O (l)
C2H6 (g) + H2 (g)
Note that each of these can also be classified in other ways. The three examples for example are also
redox reactions; the first two are combustion reactions.
6.4.3 Decomposition Reactions
Reactions that result in products that are simpler than the reactants are termed decomposition reactions.
Examples include:
C (s) + O2 (g)
CO2 (g)
CO2 (g) + H2O (l)
C2H2 (g) + O2 (g)
Note once more that reactions can be classified in more than one way.
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