Chapter 8: Chemical Equations and Reactions
I.
Describing Chemical Reactions
A. A chemical reaction is the process by which one or more substances
are changed into one or more different substances. A chemical
equation represents, with symbols and formulas, the identities and
relative amounts of the reactants and products in a chemical reaction.
B. Indications of a Chemical Reaction
There are 4 indicators that show a reaction has taken place.
1. Evolution of heat and light.
2. Production of a gas.
3. Formation of a precipitate.
4. Color change.
C. Characteristics of Chemical Equations
1. There are 3 requirements for any chemical equations.
a. The equation must represent known facts. All products and
reactants must be identified.
b. The equation must contain the correct formulas for the
reactants and the products. Remember diatomics (H2, etc),
sulfur (S8), and phosphorous (P4).
c. The law of conservation of mass must be satisfied. To satisfy
this law, coefficients are added where necessary.
2. Word and Formula Equations.
a. Word Equations: reactants and products are represented by
words. Only a descriptive meaning. For example:
methane + oxygen  carbon dioxide + water
b. Formula Equations: reactants and products are represented
by their symbols or formulas. For example:
CH4(g) + O2(g)  CO2(g) + H2O(g)
c. In order for the law of conservation of mass to be enforced, we
must balance the equation. To do this we add coefficients in
front of the formulas to make the number of each type of atom
on both sides of the arrow equal. Back to the example:
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) Notice the number of
each type of atom:
L
R
C
1
1
H
4
4
O
4
4
d.
e.
There are many other symbols used in reactions. Know yields,
reversible, solid, liquid, gas, aqueous, heat, and catalyst.
See Table 2 on page 266.
D. Significance of a Chemical Equation
1. Chemical equations are useful in doing quantitative work. The
following is a list of qualitative information in a chemical equation.
a. The coefficients of a chemical reaction indicate relative, not
absolute, amounts of reactants and products.
b. The relative masses of the reactants and products of a chemical
reaction can be determined from the reaction’s coefficients.
c. The reverse of the chemical reaction has the same relative
amounts of the substances as the forward reaction.
II. Types of Chemical Reactions
A. Synthesis Reactions
1. In a synthesis (combination, composition) reaction two or more
substances combine to form a single, new compound.
2. The general form is A + X  AX
3. Reactions of Elements with Oxygen and Sulfur
a. Produces and oxide or sulfide depending on which is present.
b. An example is 2Mg(s) + O2(g) 2MgO(s).
c. Another example is 16Rb(s) + S8(s)  8Rb2S(s)
4. Reactions of Metals with Halogens
a. A would be a metal, X would be a halogen.
b. An example is 2Na(s) + Cl2(g)  2NaCl(s)
5. Synthesis Reactions with Oxides
a. Active metals are highly reactive metals.
b. Oxides of active metals react with water to produce metal
hydroxides.
c. An example would be CaO(s) + H2O(l)  Ca(OH)2(s)
d. Many oxides of nonmetals in the upper right portion of the table
react with water to produce oxyacids.
e. An example would be SO2(g) + H2O(l)  H2SO3(aq)
f. Certain metal oxides and nonmetal oxides react with each other in
synthesis reactions to form salts.
g. An example would be CaO(s) + SO2(g)  CaSO3(s)
B. Decomposition Reactions
1. In a decomposition reaction a single compound undergoes a
reaction that produces two or more simpler substances.
2. The general form is AB  A + B.
3. Decomposition of Binary Compounds
4.
5.
6.
7.
a. Simplest kind is the decomposition of a binary compound into its
elements.
b. One way to do this is electrolysis (adding an electric current). Ex.
H2O(l)  H2(g) + O2(g).
c. Oxides of the less active metals, located in the lower center
portion of the table, decompose into their elements when heated.
d. For example: 2HgO(s)  2Hg(l) + O2(g).
e. Another way is to add a catalyst. A catalyst speeds up the
reaction without being consumed in the reaction.
f. For example: 2H2O2(l)  2H2O(l) + O2(g). Cat=KI.
Decomposition of Metal Carbonates
a. When a metal carbonate is heated, it breaks down to produce a
metal oxide and CO2 gas.
b. Ex: CaCO3(s)  CaO(s) + CO2(g)
Decomposition of Metal Hydroxides
a. All metal hydroxides, except those containing Group 1A metals,
decompose when heated to yield metal oxides and water.
b. Ex: Ca(OH)2(s)  CaO(s) + H2O(g)
Decomposition of Metal Chlorates
a. When a metal chlorate is heated, it decomposes to produce a
metal chloride and oxygen.
b. Ex: 2KClO3(s)  2KCl(s) + 3O2(g)
Decomposition of Acids
a. Certain acids decompose into nonmetal oxides and water.
b. Ex: H2CO3(aq)  CO2(g) + H2O(l)
C. Single Replacement (Displacement) Reactions
1. In a single replacement reaction one element replaces another
element in a compound.
2. The general form is A + BX  AX + B or Y + BX  BY + X.
3. Replacement of a Metal in a Compound by Another Metal
a. Al is more active than Pb.
b. 2Al(s) + 3Pb(NO3)2(aq)  3Pb(s) + 2Al(NO3)3(aq)
4. Replacement of Hydrogen in Water by a Metal
a. The most active metals react vigorously with water to produce
metal hydroxides and hydrogen.
b. Ex: 2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g).
5. Replacement of Hydrogen in an Acid by a Metal
a. The more active metals react with certain acidic solutions
replacing the hydrogen in the acids.
b. The reaction products are a metal compound (salt) and hydrogen
gas.
c. Ex: Mg(s) + 2HCl(aq)  H2(g) + MgCl2(aq).
6. Replacement of Halogens
a. One halogen replaces another halogen in a compound.
b. Fluorine is the most active halogen, it will replace any other
halogen in compounds.
c. Ex: Cl2(g) + 2KBr(aq)  2KCl(aq) + Br2(l)
d. Ex: Br2(l) + KCl(aq0  No Reacton
D. Double Replacement (Displacement) Reactions
1. In double replacement reactions the ions of two compounds
exchange places in an aqueous solution to form new compounds.
2. The general form is AX + BY  AY + BX.
3. Formation of a Precipitate
a. The formation of a precipitate occurs when the cations of one
reactant combine with the anions of another reactant to form an
insoluble or slightly soluble compound.
b. Ex: 2KI(aq) + Pb(NO3)2(aq)  PbI2(s) + 2KNO3(aq)
c. Guidelines for determining which precipitates form are in Ch 14.
4. Formation of a gas
a. In some double replacement reactions, one of the products is an
insoluble gas that bubbles out of the mixture.
b. Ex: FeS(s) + 2HCl(aq)  H2S(g) + FeCl2(aq)
5. Formation of water
a. In some double replacement reactions, a very stable molecular
compound, such as water, is one of the products.
b. Ex: HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
E.
Combustion Reactions
1. In combustion reactions a substance combines with oxygen,
releasing a large amount of energy in the form of heat or light.
a. Ex: 2H2(g) + O2(g)  2H2O(g)
2. Many times the substance that reacts with oxygen is a
hydrocarbon(CH).
3. For example, C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g).
F.
Compounds in Aqueous Solutions
1. Dissociation
a. When a compound that is made of ions dissolves in water, the
ions separate from one another. This separation is called
dissociation.
b. For example, NaCl in water and CaCl2 in water.
c. We will assume 100 % dissociation if it is said to occur.
d. Precipitation reactions
1. No compound is completely insoluble, but ones with very low
solubility are considered insoluble for most practical
applications.
2.
2. Though it is impossible to make rules to cover all situations
of solubility, we make some general guidelines for solubility.
We follow them to determine if a substance is soluble.
3. General Solubility Guidelines
a. Most Na, K, and NH4+ compounds are soluble in water.
b. Most nitrates, acetates, and chlorates are soluble.
c. Most chlorides are soluble, except those of Ag, Hg(I),
and Pb. Lead (II) chloride is soluble in hot water.
d. Most sulfates are soluble, except those of Ba, Sr, and
Pb.
e. Most carbonates, phosphates, and silicates are insoluble,
except those of Na, K, and ammonium.
f. Most sulfides are insoluble, except those of Ca, Sr, Na,
K, and ammonium.
4. We use these rules to predict whether substances formed in
reactions are soluble or not. If it is not soluble in water,
since we are usually using aqueous solutions, it is the
precipitate.
Ionic Equations: 2 types, complete and net ionic.
a. Easiest way to describe them is to practice them. Usually they
are double replacement reactions.
b. For example, (NH4)2S(aq) + Cd(NO3)2(aq)  ? + ?
c. Steps for writing CIE
1. Write balanced formula equation with states.
2. Rewrite, breaking all aqueous substances into their
component ions. Not aqueous, just rewrite it.
d. To write a NIE we add a couple of steps.
3. Determine and cross out spectator ions. Spectator ions
are ions that do not take part in a chemical reaction and are
found in solution both before and after the reaction.
4. Rewrite what is left after spectators are removed.
III. Activity Series of the Elements
A. An activity series is a list of elements organized according to the ease
with which the elements undergo certain chemical reactions. (Table 3 on
page 286)
B. The order is determined by single replacement reactions.
C. The most active element in the series is placed on top. It can replace
any of the elements below it. An element further down the series can
replace any element below it, but not any above it.