Strong Electrolyte
Sodium chloride dissolved in water. NaCl(s) + H2O(l) ---> NaCl(aq) How it actually exists in solution.
|
Weak Electrolyte
Oxalic acid dissolved in water. H2C2O4(s) + H2O(l) ---> H2C2O4(aq) How it actually exists in solution
|
Nonelectrolyte
Glucose dissolved in water. C6H12O6(s) + H2O(l) ---> C6H12O6(aq) Glucose does not dissociate, but exists in solution exclusively as the hydrated molecule, C6H12O6(aq). |
Solubility Rules for Ionic Compounds | |
Ammonium, Group 1A,
NH41+, Li1+, Na1+, K1+, Rb1+, Cs1+ |
All ammonium and Group 1A (alkali metal) salts are soluble. |
Nitrates, NO31- | All nitrates are soluble. |
Chlorides, bromides, iodides,
Cl1-, Br1-, I1- |
All common chlorides, bromides and iodides are soluble except AgCl, Hg2Cl2, PbCl2, AgBr, Hg2Br2, PbBr2, AgI, Hg2I2, PbI2. |
Sulfates, SO42- | Most sulfates are soluble except CaSO4, SrSO4, BaSO4, Ag2SO4, Hg2SO4, and PbSO4. |
Chlorates, ClO31- | All chlorates are soluble. |
Perchlorates, ClO41- | All perchlorates are soluble. |
Acetates, C2H3O21- | All acetates are soluble. |
Phosphates, PO43- | All phosphates are insoluble except those of NH41+ and Group 1A (alkali metal) salts. |
Carbonates, CO32- | All carbonates are insoluble except those of NH41+ and Group 1A (alkali metal) salts. |
Hydroxides, OH1- | All hydroxides are insoluble except
those of NH41+ and Group 1A (alkali metal) salts.
Ca(OH)2, Sr(OH)2, and Ba(OH)2 are slightly soluble. |
Oxides, O2- | All oxides are insoluble except those of Group 1A (alkali metal) salts. |
Oxalates, C2O42- | All oxalates are insoluble except those of NH41+ and Group 1A metals (alkali metal cations). |
Sulfides, S2- | All sulfides are insoluble except those of NH41+, Group 1A (alkali metal) salts, and Group 2A metals (MgS, Cas, and BaS are sparingly soluble.) |
Exchange or Metathesis Reaction or Double-Replacement
Reaction
An exchange reaction occurs between compounds that, when
written as a molecular equation, appear to involve the exchange of parts
between the two reactants. An exchange reaction will occur when ions in
solution form insoluble products, weak electrolytes, or nonelectrolytes.
These are also called metathesis reactions. The word metathesis comes from
Greek and means, "to change positions".
Exchange reactions have the general form AB + CD -----> AD + CB.
Exchange reactions can be categorized as:
1. Precipitations: one of the products is insoluble (e.g.,
AgCl(s)).
NaCl(aq) + AgNO3(aq) ---> AgCl(s) + NaNO3(aq)
2. Neutralizations: most often one of the products is
water, H2O(l).
HCl(aq) + KOH(aq) --> KCl(aq) + H2O(l)
3. Gas-forming: one of the products is a gas (e.g., SO2(g)).
Na2SO3(aq) + 2 HCl(aq) -----> 2
NaCl(aq) + H2O(l) + SO2(g)
Precipitation Reactions
A precipitate is an insoluble solid compound formed during
a chemical reaction in solution (e.g., CaCO3(s), AgCl(s), BaSO4(s)).
Precipitation reactions are a type of chemical reaction between ions that
produce a precipitate.
Based upon the solubility rules, a reaction can be predicted as to whether it will take place or not. When solutions of electrolytes are mixed, a reaction will take place if one of the products is insoluble.
Reactions can be viewed as rearrangement of atoms. Therefore, swap cation and anions to determine what the possible products might be. If one of the products is insoluble a can be predicted to occur. If both predicted products are soluble then no reaction occurs.
Sample Problems
1) BaCl2(aq) + ZnSO4(aq) ---> ?
BaCl2(aq) + ZnSO4(aq) ---> BaSO4(s)
+ ZnCl2(aq)
BaSO4 is an insoluble solid (precipitate)
therefore a reaction takes place when aqueous solutions of BaCl2
and ZnSO4 are mixed together.
2) Na2S(aq) + ZnCl2(aq) ---> ?
Na2S(aq) + ZnCl2(aq) ---> 2 NaCl(aq)
+ ZnS(s)
ZnS is an insoluble solid (precipitate) therefore a reaction
takes place when aqueous solutions of Na2S and ZnCl are mixed
together.
3) K3PO4(aq) + NaNO3(aq)
---> ?
K3PO4(aq) + NaNO3(aq)
---> KNO3(aq) + Na3PO4(aq)
All of the proposed products are soluble. Therefore a
reaction does not take place when aqueous solutions of K3PO4
and NaNO3 are mixed.
To help get a better grasp of what is actually occurring in a precipitation reaction, the reaction equation can be written in three different forms.
First, the molecular equation shows compounds as if they
exist in solution in the form of molecules, even though they are soluble
ionic compounds and dissociate in solution.
(NH4)2CO3(aq) + CaCl2(aq)
---> 2 NH4Cl(aq) + CaCO3(s)
Second, the complete ionic equation shows the soluble
ionic compounds as ions in solution and molecules are shown as nondissociated
molecules in solution.
2NH41+(aq) + CO32-(aq)
+ Ca2+(aq) + 2Cl1-(aq) --->
2NH41+(aq) + 2Cl1-(aq)
+ CaCO3(s)
Finally, the net ionic equation shows only the species
that take part in a reaction.
Ca2+(aq) + CO32-(aq)
---> CaCO3(s)
Spectator ions are the ions that do not take part in a reaction and therefore remain in the same form and state on either side of the equation arrow. In the reaction just cited the sodium cation, Na1+(aq), and the nitrate anion, NO31-(aq), are spectator ions.
The molecular equation is useful because it tells us the identity of the reagents. It is necessary in practical work in the lab and industry and useful when you are deciding what chemicals to obtain from the stockroom to carry out a reaction and are usually used for stoichiometric calculations. It identifies the actual reactants and lets us work out the amounts to use.
The complete ionic equation better describes the reaction between electrolytes than the molecular equations. It aids in visualization of all that is happening in the solution when the reaction is taking place.
The net ionic equation is a generalization of the reaction. It focuses the attention most sharply on the chemical change that is actually taking place in the solution. The chemical behavior of a strong electrolyte solution can be attributed to the various kinds of ions it contains. Therefore the net ionic equation focuses on the essential feature of the precipitation reaction. It can be useful in illustrating the similarities between large number of reactions involving electrolytes.
Soluble chloride + soluble silver rxns with same net ionic
eqn:
NaCl(aq) + AgNO3(aq) ---> AgCl(s) + NaNO3(aq)
MgCl2(aq) + AgC2H3O2(aq)
---> AgCl(s) + Mg(C2H3O3)2(aq)
Ag1+(aq) + Cl1-(aq) ---> AgCl(s)
Soluble carbonate + soluble calcium rxns with same net
ionic eqn:
Na2CO3(aq) + Ca(NO3)2(aq)
---> 2 NaNO3(aq) + CaCO3(s)
K2CO3(aq) + Ca(C2H3O2)2(aq)
---> 2 NaC2H3O2(aq) + CaCO3(s)
Ca2+(aq) + CO32-(aq)
---> CaCO3(s)
Acid formulas are most often written with the donatable
proton at the beginning of the formula.
HCl; H2SO3;
HC7H5O2
A base is a compound that:
1. neutralizes the effects of
an acid.
2. increases the concentration
of hydroxide ions (OH1-) when dissolved in water.
3. can accept a proton (H1+).
Base formulas are most often written with the hydroxide
ion at the end of the formula.
NaOH; Ca(OH)2; Al(OH)3
A salt is a compound:
1. composed of metal cations or ammonium ion (NH41+),
and nonmetal anions or polyatomic anions other than hydroxide.
2. formed when an acid and a base react. ACID + BASE
---> SALT
NaCl; MgBr2; K2SO4;
NH4NO3
The proton, H1+, cannot exist in water. When
an acid dissolves in water the hydronium ion is formed, H3O1+.
H1+ + H2O ----->
H3O1+
H1+ = H3O1+
For convenience, the hydronium ion, H3O1+, will be indicated by H1+.
Acids can be divided into two broad groups, based upon
the degree which they dissociate into ions, strong and weak.
Strong acids (HZ) are completely dissociated in solution
and are therefore strong electrolytes.
HZ(aq) ---> Hl+(aq) + Z1-(aq)
Weak acids (HA) dissociate very little (i.e., exist primarily
in the form of a molecular species) and are therefore weak electrolytes.
HA(aq) <=====> Hl+(aq)
+ A1-(aq)
There are six (6) strong acids. All other acids are weak
acids.
HCl
Hydrochloric acid
HNO3
Nitric acid
H2SO4
Sulfuric acid
HClO4 Perchloric
acid
HBr
Hydrobromic acid
HI
Hydroiodic acid
The vast majority of acids are weak acids.
Acids can also be classified according to composition
and structure.
Monoprotic acids have only one acidic hydrogen (donatable
proton).
HCl, HNO3, HClO4,
HBr, HI, HCN, HClO, HC2H3O2
Polyprotic acids have more than one acidic hydrogen.
H2SO4, H3PO4,
H2C2O4
Sulfuric acid is a diprotic acid and is a strong acid
in its first ionization and a weak acid in its second ionization.
First ionization: H2SO4(aq)
---> Hl+(aq) + HSO41-(aq)
Second ionization: HSO41-(aq)
<=====> Hl+(aq) + SO42-(aq)
Phosphoric acid is weak in all three of its ionization
steps.
First ionization: H3PO4(aq)
<=====> Hl+(aq) + H2PO41-(aq)
Second ionization: H2PO41-(aq)
<=====> Hl+(aq) + HPO42-(aq)
Third ionization: HPO42-(aq)
<=====> Hl+(aq) + PO43-(aq)
Strong and Weak Bases
Strong bases are completely dissociated in solution and
are therefore strong electrolytes. Most strong bases are inorganic.
NaOH(aq) ---> Nal+(aq)
+ OH1-(aq)
Some bases such as Ca(OH)2 are not very soluble in water, but the part that does dissolve completely dissociates. Ca(OH)2 can be considered to be a strong electrolyte although it is not very soluble.
There are six (6) strong bases. All other bases are weak
bases.
LiOH
Lithium hydroxide
NaOH Sodium
hydroxide
KOH
Potassium hydroxide
Ca(OH)2 Calcium hydroxide
Ba(OH)2 Barium hydroxide
Sr(OH)2 Strontium hydroxide
Ammonia is a weak base. Gaseous ammonia dissolves in water
and then the dissolved ammonia reacts with water to produce hydroxide ions,
but most remains as molecules. Ammonia is less than 5% ionized.
NH3(g) + H2O(l)
---> NH3(aq)
NH3(aq) + H2O(l)
<=====> NH41+(aq) + OH1-(aq)
Amines are organic derivatives of ammonia and like ammonia
are weak bases.
CH3NH2(aq) +
H2O(l) <=====> CH3NH31+(aq)
+ OH1-(aq)
ACID + BASE ----> SALT + WATER
HClO4(aq) + LiOH(aq) -->
LiClO4(aq) + H2O(l)
H2SO4(aq) +
2 NaOH(aq) --> Na2SO4(aq) + 2 H2O(l)
HI(aq) + KOH(aq) --> KI(aq) + H2O(l)
HBr(aq) + Ca(OH)2(aq) -->
CaBr2(aq) + 2 H2O(l)
HCl(aq) + Sr(OH)2(aq) -->
SrCl2(aq) + 2 H2O(l)
HNO3(aq) + Ba(OH)2(aq)
--> Ba(NO3)2(aq) + 2 H2O(l)
Net Ionic Equations of Acid-Base Reactions
HClO4(aq) + LiOH(aq) -->
LiClO4(aq) + H2O(l)
Hl+(aq) + ClO41-(aq)
+ Lil+(aq) + OH1-(aq) ----> Li1+(aq) +
ClO41-(aq) + H2O(l)
Hl+(aq) + OH1-(aq)
----> H2O(l)
The net ionic equation for the reaction between most strong
acids and strong bases are identical
H2SO4(aq) +
2 NaOH(aq) ---> Na2SO4(aq) + 2 H2O(l)
2 HNO3(aq) + Ca(OH)2(aq)
---> Ca(NO3)2(aq) + 2 H2O(l)
H1+(aq) + OH1-(aq)
---> H2O(l)
In the net ionic eqn of a weak acid and a strong base,
the acid is written in its undissociated form.
HC2H3O2(aq)
+ NaOH(aq) ----> Na C2H3O2(aq) + H2O(l)
HC2H3O2(aq)
+ Nal+(aq) + OH1-(aq) ----> Na1+ + C2H3O21-(aq)
+ H2O(l)
HC2H3O2(aq)
+ OH1-(aq) ----> C2H3O21-(aq)
+ H2O(l)
In the net ionic eqn of a strong acid and a weak base,
the base is written in its undissociated form.
Mg(OH)2(s) + 2 H1+(aq)
---> Mg2+(aq) + 2 H2O(l)
The net ionic equation for the reaction of ammonia with
hydrochloric acid is as follows.
H1+(aq) + NH3(aq)
---> NH41+(aq)
Some Common Gas-Forming Reactions
Anion
Reaction with H1+
HCO31-
HCO31- + H1+ ---> CO2(g) +
H2O(l)
CO32-
CO32- + 2 H1+ ---> CO2(g) +
H2O(l)
HSO31-
HSO31- + H1+ ---> SO2(g) +
H2O(l)
SO32-
SO32- + 2 H1+ ---> SO2(g) +
H2O(l)
HS1-
HS1- + H1+ ---> H2S(g)
S2-
S2- + 2 H1+ ---> H2S(g)
Carbonates
Na2CO3(aq) + 2 HCl(aq) ----> 2
NaCl(aq) + H2O(l) + CO2(g)
Sulfites
Na2SO3(aq) + 2 HCl(aq) ----> 2
NaCl(aq) + H2O(l) + SO2(g)
Sulfides
Na2S(aq) + 2 HCl(aq) ----> 2 NaCl(aq) + H2S(g)
Ammonium salts react with strong bases to yield ammonia
gas.
NaOH(aq) + NH4Cl(aq) ---> NaCl(aq) + H2O(l)
+ NH3(g)
Redox reactions involve the transfer of electrons from one reactant to another.
The reducing agent is the species that is oxidized. It
gains electrons.
The oxidizing agent is the species being reduced. It
loses electrons.
The oxidizing agent brings about the reduction of another
element and the reducing agent brings about the oxidation of another element.
SnO2(s) + C(s) ----> Sn(s) + 2 CO(g)
In this reaction carbon brings about the reduction of the tin ore to tin metal so it is referred to as the reducing agent.
The combustion of fuels and the reactions of metals with
oxygen to give oxides were described by the word oxidation. The burning
of magnesium in air is a redox reaction.
2 Mg(s) + O2(g) ----> 2 MgO(s)
In this instance, oxygen is the oxidizing agent because it brings about the oxidation of the metal.
Rules for Assigning Oxidation Numbers
1. The oxidation number of an atom of a pure element
is 0.
2. The oxidation number of a monatomic ion equals
its charge.
3. Some elements have the same oxidation number
in the majority of compounds they form and can be used as a reference.
Hydrogen is +1 unless combined with a metal, in this case it is –1.
Fluorine is -1.
Oxygen is –2, except in peroxides it is –1.
In binary compounds with metals, Group 7A (halogens) elements have an oxidation
number of –1.
In binary compounds, Group 6A elements have an oxidation number of –2.
In binary compounds, Group 5A elements have an oxidation number of -3.
Group 1A metal cations all have an oxidation number of +1
Group 2A metal cations all have an oxidation number of +2.
4. For the atoms in a neutral species (an isolated
atom, a molecule, or a formula unit) the total of all the oxidation numbers
is 0; the sum of the oxidation numbers in a polyatomic ion equals the charge
on the ion.
Example
What is the oxidation number of nitrogen in sodium nitrate,
NaNO3?
Sodium is a group 1A metal so it has
an oxidation number of +1.
The oxidation number of nitrogen is
calculated as follows:
(1 N)(X) + (3 O)(-2) = 1-
X = +5, therefore the oxidation number
of nitrogen in +5.
Example
What is the oxidation number of phosphorus in magnesium
phosphate, Mg3(PO4)2?
Magnesium is a group 2A metal so it
has an oxidation number of +2.
The oxidation number of phosphorus
is calculated as follows:
(1 P)(X) + (4 O)(-2) = 3-
X = +5, therefore the oxidation number
of phosphorus in +5.
OXIDATION | REDUCTION | "OIL RIG" acronym |
Involves the loss of electrons. | Involves the gain of electrons. | Oxidation
Involves Loss of electrons. Reduction
|
Is characterized by an increase in oxidation number. | Is characterized by a decrease in oxidation number. |
Identification of a redox reaction is carried out by assigning oxidation numbers of the reactants and products. When an oxidation number of an element changes then a redox reaction has taken place.
A redox reaction take place when oxidation number change for the reactants.
Example: ZnSO4 + Mg --> MgSO4 + Zn
Oxidation Number: +2
0 +2
0
Zn2+ + SO42- + Mg --> Mg2+
+ SO42- + Zn
As a reactant the oxidation number of zinc was +2 but as a product the oxidation number of zinc is zero. As a reactant the oxidation number of magnesium was 0 but as a product the oxidation number of magnesium is +2.
Example: 2 HCl + Fe ---> FeCl2 + H2
Oxidation Number: +1
-1 0
+2
-1 0
2 H1+ + 2 Cl1- + Fe ---> Fe2+ + 2 Cl1-
+ H2
As a reactant the oxidation number of hydrogen was +1 but as a product the oxidation number of hydrogen is zero. As a reactant the oxidation number of iron was 0 but as a product the oxidation number of iron is +2.
Since elemental metals have an oxidation number of zero, metals are reducing agents. Strength of reducing agent varies.
The strength of a metal as a reducing agent gives rise to the Activity Series. A metal will displace from solution the ions of any metal that lies below in it the activity series.
Activity Series for Metals (Table 5.5; p. 187)
K Ba Sr Ca Na |
These metals displace H2 from liquid water, steam, or acid. |
Al Mn Zn Cr |
These metals displace H2 from steam or acid |
Ni Sn Pb |
These metals displace H2 from acid. |
|
Including H2 permits the statement that any metal below H2 will not react with acid and any metal above H2 will react with acid. |
Cu Hg Ag Pd Pt Au |
These metals do not react with acids. |
A metal with a higher activity will displace another metal
in a salt.
2 Na(s) + 2 H2O ---> 2
NaOH(aq) + H2(g)
Ca(s) + 2 H2O ---> Ca(OH)2(aq)
+ H2(g)
Mg(s) + H2O ---> NR
Fe(s) + 2 HCl(aq) ---> FeCl2(aq)
+ H2(g)
FeCl2(aq) + H2(g)
----> NR
Cu(s) + 2 AgNO3(aq) --->
Cu(NO3)2(aq) + 2 Ag(s)
Cu(NO3)2(aq)
+ 2 Ag(s) --> NR
Nonmetal Reactivity
2 Cl2(aq) + Na2S(aq)
--> S(s) + 2 NaCl(aq)
Cl2(aq) + 2 NaBr(aq) -->
2 NaCl(aq) + Br2(aq)
Br2(aq) + 2 NaCl(aq) --->
NR.
Example
If 400.0 mL of a solution contains 5.00 x 10-3
moles of AgNO3, what is the molarity of this solution?
5.00 x 10-3 mol AgNO3
/ 0.4000 L = 0.0125 M AgNO3
Example
Calculate the molarity of an HCl solution containing
18.23 g of HCl in 355.0 mL of solution.
A. Calculate the number of moles of
HCl.
(18.23 g HCl) (1 mol HCl/36.46 g HCl) = 0.5000 mol HCl
B. Divide the number of moles of HCl
by the total volume in liters.
0.5000 mol HCl/0.3550 L solution = 1.408 M HCl
Example
How many moles of NaCl are contained in a 27.49 mL sample
of a 0.350 M solution of NaCl?
(0.350 mol/L)(0.02749 L) = 9.62 x
10-3 mol NaCl
Example
What volume of 18.0 M H2SO4 must
is required to make 100 mL of a 5.0 M solution of H2SO4?
(28 mL)
Example
How many milliliters of water are required to dilute
6 M H2SO4 to 175 ml of 0.3 M H2SO4?
(166 mL)
Molarity and Reactions in Aqueous Solutions
Example
How many mL of 6.0 M HCl are needed to digest 3.78 g
of Mg? [52 mL]
Example
How many mL of 0.0487 M Ba(OH)2 are needed
to react with 35.67 mL of 0.0748 M HCl? [27.3 mL Ba(OH)2]
Titrations
Titration is as an analytical procedure of determining
the concentration of one substance in solution by reacting it with a solution
of another substance whose concentration is known, called a titrant (or
standard solution).
To carry out the process, we add the titrant, using a buret, to a known volume of the other solution until the reaction between the two substances is just complete, that is until chemically equivalent amounts of the two reactants are present. We often tell when we have reached this point, called the endpoint or equivalence point, by a change in color of an added chemical substance called an indicator.
The volume of the titrant is used to calculate the number of moles delivered from the buret. The moles of the other reactant are obtained from the coefficients in the balanced reaction equation, and this value is used with the volume of the solution of unknown concentration to calculate its molarity.
Volume of A ßà moles A ßà moles of B ßà Molarity of B
Example 1
In a titration 25.00 mL of sodium hydroxide solution
was neutralized by 32.72 mL of hydrochloric acid. The HCl had a concentration
of 0.129 M. Find the concentration of the NaOH solution.
A. Calculate the number
of moles of HCl delivered.
(0.03272 L HCl)(0.129 mol/L) = 4.22 x 10-3 mol HCl
B. Calculate the
number of moles of NaOH neutralized using the coefficients from the balanced
neutralization reaction equation.
HCl + NaOH ----> NaCl + H2O
(4.22 x 10-3 mol HCl)(1 mol NaOH/1 mol HCl) = 4.22 x 10-3
mol NaOH
C. Calculate the
molarity of the NaOH using the volume of NaOH titrated and the number of
moles of NaOH present in this volume.
4.22 x 10-3 mol NaOH / 0.02500 L NaOH solution = 0.169 M NaOH
Example 2
During a titration, a 20.00 mL portion of a 0.100 M H2SO4
solution was carefully measured into a flask and phenolphthalein was added
to it. The solution was titrated with 18.47 mL of NaOH to reach the endpoint.
What is the concentration of the NaOH?
A. (0.02000
L H2SO4)(0.100 mol H2SO4/L)
= 2.00 x 10-3 mol H2SO4
B. H2SO4
+ 2 NaOH ---> Na2SO4 + 2 H2O
(2.00x10-3 mol H2SO4)(2 mol NaOH/1 mol
H2SO4) = 4.00 x 10-3 mol NaOH
C. 4.00 x
10-3 mol NaOH / 0.01847 L NaOH solution = 0.217 M NaOH