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Chapters 11 and12
Chemical Reaction
 One or more substance(s) change into one or more
new substances
 Reactants
Products
 Exothermic- energy is product (on right)
 products are more stable
 Endothermic- energy is reactant (on left)
 Products are less stable
Balancing Chemical Equations
 Includes kinds and parts of substances involved
 **LAW OF CONSERVATION OF MASS
 Mass (matter), charge and energy are always
conserved in a chemical rxn
In a balanced equation,
 Each side of the equation has the same number and
type of atoms
 Remember Dalton’s Theory:

Bonds broken and formed; atoms rearranged
To write a balanced chemical equation:
 First write the skeleton equation using the correct
formulas for elements and compounds (remember
diatomics and criss-cross method)
 H2 + O2 --> H2O
 Then use coefficients to balance the equation so
that it obeys The Law of Conservation of Mass
 2H2 + O2 --> 2H2O
The Rules
1.
Determine the correct formulas for all the reactants
and products

Molecular (covalent)- use prefixes
Ionic- balance charges
Diatomics!!



BrOFINClH
2. Write the skeleton equation by placing the formulas
for the reactants on the left and the formulas for the
products on the right with a “yields” sign in between.
• If two or more reactants or products are involved,
separate their formulas with (+) signs
3. Determine the number of atoms of each element in
the reactants and products.
• TIP: Count a polyatomic ion as a single unit if it appears
unchanged on both sides of the equation
4. Balance the elements one at a time by using
coefficients. Begin by balancing elements that appear
only once on each side of the equation.
 Unwritten coefficients are assumed to be 1
 Once you are certain you have the correct chemical
formula for the substances involved, NEVER change
the subscripts in a chemical formula
5. Check each atom or polyatomic ion to be sure they are
equal on both sides of the equation.
*** Add phase symbols to substances if necessary (s, l, g,
aq)
6. Make sure all of the coefficients are in the lowest
possible ratio
Example #1
 The reaction of zinc with aqueous hydrochloric acid
produces a solution of zinc chloride and hydrogen gas
Example #2
 Sodium reacts with sulfur to produce sodium sulfide
Example #3
 Aqueous nitric acid (HNO3) reacts with aqueous
magnesium hydroxide to produce aqueous magnesium
nitrate and water
Classifying Reactions
 The five general types of reactions are:
 Combination/Synthesis
 Decomposition
 Single-replacement (displacement)
 Double-replacement (displacement)
 Combustion
Synthesis Reactions
 A chemical change in which two or more substances
react to form a single new substance
 Zn(s) + I2(g)  ZnI2(s)
Combination (synthesis) Reactions
 Group A metal + nonmetal  metal cation and
nonmetal anion (ionic compound)
 2 K(s) + Cl2(g)  2 KCl(s)
Combination (synthesis) Reactions
 2 nonmetals  more than one product is often
produced
 S(s) + O2(g) --> SO2(g) (sulfur dioxide)
 2S(s) + 3 O2(g) --> 2SO3(g) (sulfur trioxide)
Combination (synthesis) Reactions
 Transition metal + nonmetal  could produce more
than one product
 Fe(s) + S(s)  FeS(s) (iron (II) sulfide)
 2Fe(s) + 3S(s)  Fe2S3(s) (iron (III) sulfide)
Decomposition Reactions
 A chemical change in which a single compound
breaks down into two or more simpler products
Decomposition Reactions
 One reactant and two or more products
 Difficult to predict products
 Most require energy in the form of heat, light or
electricity (endothermic)
 2HgO(s) --> 2Hg(l) + O2(g)
Combustion Reactions
 A chemical change in which an element or a
compound reacts with oxygen, often producing energy
in the form of heat and light
Combustion
oxygen + hydrocarbon  water + carbon dioxide +
energy
Combustion
 Magnesium and sulfur will also burn in the presence of
oxygen.
2Mg(s) + O2(g) --> 2MgO(s)
 For our purposes, we’ll call these reactions synthesis.
Single Replacement Reactions
 One element replaces a second element in a compound
 NOTE: both the reactants AND the products consist of
an element and a compound
 K + AgCl  Ag + KCl
Single Replacement Reactions
Zn(s) + Cu(NO3)2(aq) --> Cu(s) + Zn(NO3)2(aq)
 Whether one metal will displace another metal from a
compound depends on the relative reactivities of the
two metals
 **TABLE J**
The ACTIVITY SERIES
 Lists elements in order of decreasing reactivity
 A halogen can also replace a halogen- reactivity
decreases as you go down the group
 Metals- want to lose electrons
 Nonmetals- want to gain electrons
Double Replacement Reactions
 An exchange of positive ions between two
compounds
 Generally takes place in aqueous solutions and often
produce a precipitate, a gas, or a molecular compound
such as water.
Double Replacement Reactions
 One of the products is insoluble and precipitates from
solution
 **TABLE F**
Na2S(aq) + Cd(NO3)2(aq)  CdS(s) + 2NaNO3(aq)
Double Replacement Reactions
 One of the products is a gas
2NaCN(aq) + H2SO4(aq)  2HCN(g) + Na2SO4(aq)
Double Replacement Reactions
 One product is a molecular compound such as water
(acid-base rxn: neutralization)
Ca(OH)2(aq) + 2HCl(aq)  CaCl2(aq) + 2H2O(l)
Examples
 CaBr2(aq) + AgNO3(aq) 
 FeS(s) + HCl(aq) 
Reactions in Aqueous Solutions
AgNO3(aq) + NaCl(aq)  AgCl(s)+ NaNO3(aq)
 Most ionic compounds dissociate into cations and anions
when dissolved in water
 We can write a complete ionic equation to show the dissolved
ionic compounds as dissociated free ions
Complete Ionic Equation
Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) 
AgCl(s) + Na+(aq) + NO3-(aq)
 Na+(aq) and NO3-(aq) appear unchanged on either side of
the equations; they did not technically participate in
the reaction.
Spectator Ion
 An ion that appears on both sides of an equation and
is not directly involved in the reaction
Net Ionic Equation
 An equation for a reaction in solution that shows only
those particles directly involved in the chemical
change
Ag+(aq) + Cl-(aq) --> AgCl(s)
**You must make sure that the charges on
either side are balanced
Try this one
Pb(s) + AgNO3(aq)  Ag(s) + Pb(NO3)2(aq)
Predicting the Formation of a
Precipitate
Table F- Solubility Guidelines
2Na+(aq) + CO32-(aq) + Ba2+(aq) + 2NO3-(aq) 
 Sodium is an alkali metal- soluble
 Carbonates are generally insoluble
 BaCO3 will precipitate
 Net: Ba2+(aq) + CO32-(aq) --> BaCO3(s)
Types of Reactions
 The number of elements and/or the compounds
reacting is a good indicator of possible reaction types
and thus possible products
Mole-Mole Relationships
 We can learn to use a balanced equation to determine
relationships between moles of reactants and moles of
products
 stoichiometry
 Use equations to determine the number of moles that
can be produced from certain numbers of moles of
reactants
 2H2O  2H2 + O2
 2 moles of water produces 2 moles of H2 + 1 mole of O2
Example
 What number of moles of products will be produced
by the decomposition of 5.8 moles of water?
5.8 moles H2O  5.8 H2 + 2.9 O2
Example
 Calculate the number of moles of O2 required to react
exactly with 4.30 moles of propane, C3H8, in the
reaction described by the following unbalanced
equation:
__ C3H8 + __ O2  __ CO2 + __ H2O
For each of the following unbalanced equations,
indicate how many moles of the second reactant
would be required to react exactly with 0.25 moles
of the first reactant:
1.
2.
CO + O2  CO2
CH4 + Cl2  CCl4 + HCl
Limiting and Excess Reagents
 In a chemical rxn, not enough of any of
the reactants will limit the amount of
product that forms
 Excess means that substance is NOT
the limiting reagent
N2(g) + 3H2(g)  2NH3(g)
 When 1 mol of N2 reacts with 3 moles of H2, 2 mols of
NH3 are produced
 What would happen if two moles of N2 reacted with 3
moles of H2?
In this reaction
 Only the H is completely used up
 Limiting Reagent: reagent that determines
the amount of product that can be formed
by a reaction
 Excess reagent: reactant that is not
completely used up
The first step is always to
 Convert the quantity of each reactant to
number of moles so that the limiting
reagent can be identified
 The amount of product formed in a reaction
can be determined from the given amount of
limiting reagent.
Example
2Cu(s) + S(s)  Cu2S(s)
 What is the limiting reagent when 80.0 g Cu reacts
with 25.0 g S?
 ***The reactant that is present in the smaller amount
by mass or volume is NOT necessarily the l. r.
Percent Yield
 Theoretical yield: the maximum amount of
product that could be formed from the given
amounts of reactants
 Actual yield: the amount of product that
actually forms when the rxn is carried out in
the lab
Percent Yield
 The ratio of the actual yield to the
theoretical yield expressed as a percent
 Percent Yield= actual yield x 100%
theoretical yield
Should normally be near 100%
 But there are many reasons why it could be
less:
 Accuracy of measurement
 Reaction does not go to completion
 Sloppy lab procedures
 Side reactions may occur
Example
What is the percent yield if 13.1 g
CaO is actually produced when
24.8 g CaCO3 is heated?
CaCO3  CaO + CO2