CHEMICAL REACTIONS
AND EQUATIONS
Chapter 9
Chemical Reaction
• A process in which one or more
substances are converted into new
substances with different chemical
and physical properties.
• Not reversible.
Chemical Reactions
• Examples:
• Breathing
• Burning gasoline
• Baking
• All chemical reactions involve two types of
substances:
• Reactant – a substance that enters into a chemical
reaction
• Product – a substance that is produced by a chemical
reaction
Reasons for Reactions
• The arrangement of electrons in an atom
determines whether it will bond with other
atoms and with which atoms it will bond.
• An atom with a full set of valence electrons will
not form bonds.
• An atom with an incomplete set of valence
electrons will bond.
HOMEWORK
Pg. 281 1-4
CHEMICAL EQUATIONS
9-2
Chemical Equations
• Chemical reactions are represented by
sentences known as chemical equations.
• A chemical equation identifies the reactants
and products in a chemical reaction.
Chemical Equations
• A chemical reaction is the process by which
one or more substances are changed into
one or more different substances
• It can be represented by an equation.
• It shows what changes have taken place
• It also shows the relative amounts of the
various elements and compounds that take part
in these reactions
Chemical Equations
• The starting substances in a chemical
reaction are the reactants
• The substances that are formed by the
chemical reaction are the products
• Reactant(g) + Reactants(g)  Products(g) + Products(g)
Chemical Equations
• The letters in parentheses indicate the
physical state of each substance
involved.
• (g) means gas
• (l) means liquid
• (s) means solid
• (aq) means that it is dissolved in water
Word Equations
• Gives names of the reactants and names of
products.
• Example:
• calcium + oxygen  calcium oxide
• + means “reacts with”
•  means “yields”, indicates direction of reaction.
Formula Equations
• Writing a chemical equation:
• First determine the symbols and formulas that
describe the reactants and products.
• Substitute into the word equation.
• Example: calcium + oxygen  calcium oxide
• Ca + O2  CaO
Practice Problems
• Silver (I) nitrate reacts with copper to form
copper (II) nitrate and silver.
• Hydrogen peroxide (H2O2) decomposes to
form water and oxygen gas.
Balancing Chemical Equations
• Law of conservation of mass:
• Matter is neither created nor destroyed.
• Total mass of the reactants must be equal to the total
mass of the products.
• For mass to remain constant before and after a
chemical reaction, the number of atoms of each
element must be the same before and after a chemical
reaction.
Balancing Chemical Equations
• The same number of atoms of each
element must appear on both sides of the
arrow in a chemical equation.
• Example:
• If reactants have 4 hydrogen and 1 carbon,
products must also have 4 hydrogen and 1
carbon.
Balancing Chemical Equations
• Rules for writing balanced equations
• 1. Determine the reactants and the
products
• 2. Assemble the parts of the chemical
equation
• H2(g) + O2(g)
• Reactants

H2O(l)
Yields
Products
Balancing Chemical Equations
• 3. Write a balanced equation
• Balancing means showing an equal number of
atoms for each element on both sides of the
equation
• Remember nothing is lost or gained!
• Only Transferred
• When balancing equations only change the
coefficients
• Never change the subscripts!!
• Finally make sure that all the coefficients are
in the lowest possible ratio.
Balancing Equations
• Ca + O2  CaO
• Not balanced.
• 2 oxygen on the reactants, 1 oxygen in the
products.
• *Equation cannot be balanced by changing
subscripts*
• Changing a subscript changes the identity of the
substance.
Balancing Equations
• To balance an equation correctly we
need to use coefficients.
• Whole numbers written before the formula
for reactants and products.
•
•
Ca + O2  CaO – not balanced
2 Ca + O2  2 CaO – balanced
Balancing Equations
• Example:
• Methane + oxygen  carbon dioxide + water
• Replace words with symbols:
•
CH4 + O2  CO2 + H2O
Balancing Equations
• To balance:
• Start with those elements that occur in only one
substance on each side of the equation.
• **Number of atoms found by multiplying the subscript by
the coefficient.
• Balance by trial and error.
• Balanced equation should have coefficients in the
lowest whole number ratio possible.
•
CH4
+
O2  CO2 +
H2O
Practice Problems
• Balance the following equation:
• H2(g) + O2(g)  H2O(l)
Practice Problems
• 2H2(g) + O2(g)  2H2O(l)
H=4
O=2
H=4
O=2
• The equation is balanced
Practice Problems
• Balance the following equation:
• C2H6O + O2  CO2 + H2O
Practice Problems
• C2H6O + 3O2  2CO2 + 3H2O
•C = 2
•H = 6
•O = 7
C=2
H=6
O=7
• The equation is balanced
Practice Problems
• Balance the following equations:
• 1. Zn + HCl ----> ZnCl2 + H2
• 2. Al + O2 -----> Al2O3
• 3. Al + CuSO4 ------> Al2(SO4)3 + Cu
• 4. Li + H2O ------> LiOH + H2
Practice Problems
• Aluminum reacts with oxygen to produce aluminum oxide.
• Sodium nitrate reacting with calcium chloride to produce
sodium chloride and calcium nitrate
• Dinitrogen pentoxide reacts with water to produce nitric
acid (HNO3)
Writing Complete Chemical Equations
• A complete reaction must be balanced and
include the physical state of each reactant
and product.
• Gas (g), solid (s), liquid (l)
• Example:
• CH4 (g) + 2O2 (g)  CO2 (g) + 2H2O (g)
HOMEWORK
Pg. 290 1-4
CLASSIFYING
CHEMICAL REACTIONS
9-3
Type of Reactions
• Four types of chemical reactions.
• Direct combination (Synthesis)
• Decomposition
• Single-replacement
• Double-replacement
Synthesis
• Two or more reactants come together
to form a single product.
• Synthesis reactions.
• A + B  AB
• Product is always more complex than
either of the reactants.
Synthesis
• Synthesis – two or more substances combine
to form one new substance
• It follows this general form:
• Element or compound + Element or compound  Compound
Synthesis
• Examples of synthesis:
• NH3 + HCl  NH4Cl
• CaO + SiO2  CaSiO3
• 2H2 + O2  2H2O
• 2Na + Cl2  2 NaCl
• CO2 + H2O  H2CO3
Decomposition Reactions
• Something breaking down into smaller parts.
• A reaction in which a single compound breaks
into two or more smaller compounds or elements.
• Identified by only one reactant:
• AB  A + B
• AB represents a compound
• A and B represent elements or simpler
compounds.
Decomposition Reactions
• Decomposition – when a compound is breaking
apart or decompose into simpler substances
when energy is supplied
• Energy may be supplied in the form of heat,
light, mechanical shock, or electricity
• It follows this general form:
• Compound  Two or more elements or compounds
Decomposition Reactions
• Examples of Decomposition:
• 2H2O2 → 2H2O + O2
• CaCO3 → CaO + CO2
• H2CO3 → H2O + CO2
• 2KClO3 → 2KCl + 3O2
• 2 H2 O  2 H2 + O 2
• CaCO3  CaO + CO2
Single-Replacement Reaction
• Single Replacement – one element
displaces another in a compound.
• It follows this general form:
 AC + B
• Element + Compound  Element +
Compound
• Example of single displacement:
• A + BC
–2AgNO3(aq) + Zn(s) → 2Ag(s) + Zn(NO3)2(aq)
Single-Replacement Reaction
Single-Replacement Reaction
• An uncombined element displaces an
element that is part of a compound.
• Reactants: one element and one
compound
• A + BX  AX + B
• Generally ionic compounds.
Single-Replacement Reaction
• Example:
• Mg + CuSO4  MgSO4 +Cu
• A more active element will replace a less active element.
• Figure 9-19 – activity series:
• Used to predict whether or not a single-replacement reaction will
occur.
• An element can replace any element that is below it.
• Metals replace metals or hydrogen
• Nonmetals replace nonmetals
• Example:
• Cl2 + 2 KI  ?
• Cl2 + 2 KI  KCl + I2
Double-Replacement Reaction
• Double Replacement – is similar to single
displacement but it uses two compounds
instead of one element
• It follows this general form:
• Compound + Compound  Compound +
Compound
• AB + CD → AD + CB
Double-Replacement Reaction
• Atoms or ions from two different
compounds replace each other.
• Identifying feature: two compounds as
reactants and two compounds as products.
• AX + BY  AY + BX
• Example:
• CaCO3 + 2HCl  CaCl2 + H2CO3
Double-Replacement Reaction
• Examples of Double Displacement:
• AgNO3 + HCl  AgCl + HNO3
• Fe2O3 + HCl  FeCl3 + H2O
Double-Replacement Reaction
• Conditions:
• Most reactions will not occur unless the
reactants are dissolved in water so that the ions
can separate into ions.
• Reactions are more likely to take place if one of
the products is a molecular compound, a
precipitate, or a gas.
Combustion
• Combustion – occurs when a compound burns in
air, it is actually reacting with the oxygen in the air.
• The products of the oxidation of the hydrocarbon under
normal conditions are carbon dioxide and water vapor.
• Combustion reactions are refered to as oxidation
reactions.
• They follow this general form:
• Hydrocarbon + Oxygen  Carbon Dioxide + Water (MOST USED)
• Hydrocarbon + Flourine  Carbon Flouride + Hydrogen Flouride
Combustion
• Examples of Combustion/Oxidation:
• CH4 + 2O2 → CO2 + 2H2O
• CH2S + 6F2 → CF4 + 2HF + SF6
• CH4 + 2O2 + 7N2 → CO2 + 2H2O + 7N2 + heat
Combustion
• Not all reactions take one of these five
general forms.
• There are other classes of reactions,
but those we will talk about in a later
chapter.
HOMEWORK
Pg. 297 1-4