Chapter 11 Chemical Reactions
11.1 Describing Chemical Reactions
11.2 Types of Chemical Reactions
11.3 Reactions in Aqueous Solutions
The objective of this chapter is for you to identify a type of equation,
predict the product from the reactants, and balance the final equation.
Writing Chemical Equations
There are two parts to a chemical equation
a. Reactants – those elements or compounds that will
combine together to form new compounds or molecules.
Always on the left side of the equation.
b. Products – those new elements or compounds that form in
a chemical reaction. Always on the right side of the equation.
Reactants
Product
The arrow means “yields”
Word Equations:
Consist of writing the names of the reactants on the left side
of the yield sign and the products on the right side of the yield
sign
Iron + Oxygen
Iron (III) oxide
Chemical Equations:
Is a representation of a chemical reaction using the
formulas of the reactants and products.
Fe + O2
Fe2 O3
+ Used to separate reactants and products
Used for reversible reactions
(s) Designates a solid
(l) Designates a liquid
(g) Designates a gas
(aq) Designates an aqueous solution; the substance is dissolved in water
D
Indicates heat is supplied to the reaction
MnO2
A formula written above or below the yield sign indicates its use as
a catalyst – ( a substance that speeds a reaction)
Skeleton Equation:
Is a chemical equation that does not indicate the relative amounts of
the reactants and products
Balanced Equation:
Is a chemical equation in which each side of the equation has the
same number of atoms of each element and mass is conserved. To
balance the equation, whole numbers called coefficients are used on
both sides of the equations to help balance the number of atoms in the
reactants and product.
The coefficient actually represents how many moles are present in the
reactant and product that are necessary to run the reaction.
Link
Steps for writing Chemical Equations:
1. Write a skeleton equation
2. Use subscripts to balance charges in the product
3. Use coefficients to balance atoms on each side of the equation
Try
Potassium + Oxygen
K + O2
K + O2
4K + O2
Now Try:
P + O2
AgNO3 + Cu
P4O10
Ag + Cu(NO3)2
Potassium Oxide
K+1O+2
K2O
2K2O
Law of the Conservation of Mass:
The mass of the atoms present in the reactants must equal the
mass of the atoms present in the product
11.2 Types of Chemical Reactions
The following is a list of the four major types of reactions
By knowing the type of reaction, the products can be predicted.
1. Composition Reaction
2. Decomposition Reaction
3. Replacement Reaction
4. Combustion
Combination Reaction:
Also called a Synthesis Reaction, occurs when two or more
substances combine to form a more complex substance.
Composition reactions have the general form;
A+X
AX
Examples:
Iron and Sulfur combine to form Iron(II) Sulfide
Fe + S
FeS
Magnesium and Oxygen gas form Magnesium Oxide
2Mg + O2
2MgO
Water and Sulfur Trioxide
H2O + SO3
H2SO4
Two Special Combination Reactions:
1. Metal Oxide + Water
Na2O + H2O
Hydroxides (which are bases)
2 NaOH (Sodium Hydroxide)
Try, CaO + H2O
?
2. Nonmetal oxide + Water
SO3 + H2O
Try, Cl2O5 + H2O
Acids
H2SO3 (Sulfurous Acid)
?
Don’t remember your acids and bases then review chapter 9!
Decomposition Reaction:
Reactions that are in reverse to decomposition reactions. Here
one substance breaks down to form two or more simpler substances.
Decomposition reactions have the general form;
AX
A+X
Examples:
Water decomposes, yielding hydrogen and oxygen
2H2O
2H2(g) + O2(g)
Potassium Chlorate decomposes, yielding potassium chloride and oxygen
2KClO3
2KCl + 3O2(g)
Mercury(II) Oxide decomposes to form metallic mercury and oxygen
2HgO
2Hg(l) + O2(g)
There are six types of decomposition reactions:
1. Metallic carbonates, when heated, form metallic oxides and
carbon dioxide CaCO3
CaO + CO2(g)
2. Many metallic hydroxides, when heated, decompose into metallic
oxides and water. Ca(OH)2
CaO + H2O(g)
3. Metallic chlorates, when heated, decompose into metallic
chlorides and oxygen. 2KClO3
2KCl + 3O2(g)
4. Some acids, when heated, decompose into nonmetallic oxides and
water.
H2CO3
H2O + CO2(g)
5. Some oxides, when heated, decompose though most are stable.
2HgO
2Hg + O2(g)
6. Some decomposition reactions are produced by an electric current
2H2O (electricity)
2H2(g) + O2(g)
Replacement Reaction:
Occur when one substance is replaced in its compound by
another substance.
Replacement reactions have the general form;
Single Replacement- A + BX
AX + B
OR
Y + BX
Double Replacement – AY + BX
BY + X
AX
+ BY
There are four specific types of replacement reactions:
1. Replacement of Hydrogen in water by metals
2. Replacement of a metal in a compound by a more active metal
3. Replacement of Hydrogen in acids by metals
4. Replacement of Halogens
Reactivity of the elements determine if the reactions will occur.
One atom must be more reactive then the element that is being
replaced in the equation.
Replacement of Hydrogen in water by metals:
The very active metals such as potassium, calcium, and
sodium, react vigorously with water. They replace half the
hydrogen to form metallic hydroxides. At elevated
temperatures less active metals such as magnesium, zinc,
and iron react with steam to replace hydrogen. Because of
the high temperature involved, oxides rather than
hydroxides are formed. Metals less active than iron do not
react measurably with water.
Example: Ca + 2H2O
Ca(OH)2 + H2(g)
Replacement of a metal in a compound by a more active metal:
A more reactive metal replaces the less reactive metal in a
compound
-It is important to understand the periodic trends for the
reactivity of metals
Example: Zn + CuSO4
ZnSO4
+ Cu(s)
Replacement of Hydrogen in acids by metals:
Many metals react with certain acids to replace the hydrogen in
the acid to form a metallic compound. Metals from Li to Na will
Replace hydrogen from water and acids. Metals from Mg to Pb
Will replace hydrogen from acids only.
Example: Zn + H2SO4
ZnSO4 + H2(g)
Replacement of Halogens:
Replacement of a halogen with another halogen depends on the
reactivity of the two halogens involved. A more reactive halogen
always replaces a less active halogen.
Cl2 + 2KBr
2KCl + Br2
How does the reactivity of the halogens progress?
Double Replacement Reaction:
An exchange of positive ions between two compounds. There
are generally three rules that govern this type of reaction.
These reactions are generally ionic in nature and take place
in an aqueous solution. To occur, one of the products must be
a. An insoluble precipitate b. A gas c. A molecular compound
Double Replacement – AY + BX
AX
+ BY
1. One of the products is only slightly soluble and precipitates
from solution. Na2S + Cd(NO3)2(aq)
CdS + 2NaNO3(aq)
2. One of the products is a gas
2NaCN(aq) + H2SO4(aq)
2HCN(g) + Na2SO4(aq)
3. One product is a molecular compound such as water
Ca(OH)2(aq) + 2HCl(aq)
CaCl2(aq) + 2H2O(l)
Combustion Reaction:
Occurs when an element or compound reacts with oxygen, often
producing energy in the form of light and heat. The reaction involves
oxygen as a reactant while the other reactant is often a hydrocarbon.
In this case the complete combustion of a hydrocarbon produces
carbon dioxide and water.
2C8H18(l) + 25O2(g)
16CO2(g) + 18H2O(l)
Other elements can be combusted with oxygen and look much
like a combination reaction
2Mg(s) + O2(g)
2MgO(s)
A hydrocarbon is a compound composed only of hydrogen and carbon
Many are used as fossil fuels – methane, propane, butane, and octane
11.3 Reactions in Aqueous Solution:
Net Ionic Reactions –
-The earth is 70% water
-Your body is 66% water
-Many important chemical reaction take place in water (an
aqueous solution) causing the compounds to separate into ions
Example: When sodium chloride and silver nitrate are placed
in solution, the ions dissociate. You can use these ions to
write a complete ionic equation.
-Complete ionic equation – an equation that shows dissolved
ionic compounds as dissociated ions.
Ag+(aq)+NO3-(aq)+Na+(aq)+ Cl-(aq)
AgCl(s)+ Na+(aq)+NO-3(aq)
Ag+(aq)+NO3-(aq)+Na+(aq)+ Cl-(aq)
AgCl(s)+ Na+(aq)+NO-3(aq)
Note that the sodium and nitrate ion are unchanged,
the equation can be simplified by eliminating these
ions because they do not participate in the reaction
Ag+(aq) + Cl-(aq)
AgCl(s)
This is called the net ionic reaction – an equation that
shows only the particles involved in the chemical change
Spectator Ions – an ion that appears on both sides of an
equation that is not directly involved in the reaction.
Predicting the Formation of Precipitate:
You can predict the formation of a precipitate by using the general
rules for solubility of ionic compounds.
Salts of alkali metals and
Ammonia
Nitrate salts and chlorate salts
Soluble
Soluble
Sulfate salts, except
compounds with Pb(II), Ag(I),
Hg(II), Ba+2, Sr+2, Ca +2
Soluble
Chloride salts, except
compounds with Ag+, Pb+2,
Hg+2
Soluble
Carbonates, phosphates,
chromates, sulfides,
hydroxides
Most are Insoluble