Transcript Document

Chapters 6-7a
Chemical Reactions
Chapter 6
Table of Contents
6.1
6.2
6.3
7.1
7.2
Evidence for a Chemical Reaction
Chemical Equations
Balancing Chemical Equations
Predicting Whether a Reaction Will Occur
Reactions in Which a Solid Forms
2
Section 6.1
Evidence for a Chemical Reaction
What are the clues that a chemical change has taken place?
•
•
Chemical reactions often give a visual signal.
But reactions are not always visible.
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3
Section 6.1
Evidence for a Chemical Reaction
Some Clues That a Chemical Reaction Has Occurred
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4
Section 6.1
Evidence for a Chemical Reaction
Exercise
What is a clue that a chemical reaction has
occurred?
a)
b)
c)
d)
The color changes.
A solid forms.
Bubbles are present.
A flame is produced.
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5
Section 6.1
Evidence for a Chemical Reaction
Exercise
What is a clue that a chemical reaction has
occurred?
“Colorless hydrochloric acid is added to a red solution
of cobalt(II) nitrate, turning the solution blue.”
a)
b)
c)
d)
The color changes.
A solid forms.
Bubbles are present.
A flame is produced.
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6
Section 6.1
Evidence for a Chemical Reaction
Exercise
What is a clue that a chemical reaction has
occurred?
“A solid forms when a solution of sodium dichromate is
added to a solution of lead nitrate.”
a)
b)
c)
d)
A gas forms.
A solid forms.
Bubbles are present.
A flame is produced.
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7
Section 6.2
Chemical Equations
• Chemical reactions involve a rearrangement of the ways
atoms are grouped together.
• A chemical equation represents a chemical reaction.
 Reactants are shown to the left of the arrow.
 Products are shown to the right of the arrow.
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Section 6.2
Chemical Equations
• In a chemical reaction atoms are not created or destroyed.
• All atoms present in the reactants must be accounted for in
the products.
 Same number of each type of atom on both sides of the
arrow.
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9
Section 6.2
Chemical Equations
Taking 20 kids to the zoo?
At who’s
house
would
you
drop
What
if
you
came
home
with
What
if
you
came
home
with
Parents are funny that
them
off?
22
kids?
only
18
way!kids?
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10
Section 6.2
Chemical Equations
Balancing a Chemical Equation
• Unbalanced Equation:
•
Balancing the Equation:
•
The balanced equation:
CH4 + 2O2  CO2 + 2H2O
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11
Section 6.2
Chemical Equations
Physical States
•
Physical states of compounds are often
given in a chemical equation. These are
sometimes called descriptors.
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12
Section 6.2
Chemical Equations
Example
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13
Section 6.2
Chemical Equations
Exercise
When blue light shines on a mixture of
hydrogen and chlorine gas, the elements react
explosively to form gaseous hydrochloric acid.
What is the unbalanced equation for this
process?
a)
b)
c)
d)
H2(g) + CH4(g)
HCl(g)
HCl(g)
H(g) + Cl(g)
H(g) + Cl(g)
HCl(g)
H2(g) + Cl2(g)
HCl(g)
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14
Section 6.3
Balancing Chemical Equations
•
•
•
The principle that lies at the heart of the
balancing process is that atoms are
conserved in a chemical reaction.
Atoms are neither created nor destroyed.
The same number of each type of atom is
found among the reactants and among the
products.
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15
Section 6.3
Balancing Chemical Equations
•
•
Chemists determine the identity of the
reactants and products of a reaction by
experimental observation.
The identities (formulas) of the compounds
must never be changed in balancing a
chemical equation.
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16
Section 6.3
Balancing Chemical Equations
How to Write and Balance Equations
1. Read the description of the chemical
reaction. What are the reactants, the
products, and their states? Write the
appropriate formulas.
Hydrogen gas (H2) and oxygen gas (O2)
combine to form liquid water (H2O).
2. Write the unbalanced equation that
summarizes the information from step 1.
H2(g) + O2(g)  H2O(l)
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Section 6.3
Balancing Chemical Equations
How to Write and Balance Equations
3. Balance the equation by inspection,
starting with the most complicated
molecule.
Equation is unbalanced by counting the
atoms on both sides of the arrow.
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Section 6.3
Balancing Chemical Equations
How to Write and Balance Equations
3. Balance the equation by inspection,
starting with the most complicated
molecule.
We must balance the equation by adding
more molecules of reactants and/or products.
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Section 6.3
Balancing Chemical Equations
How to Write and Balance Equations
4. Check to see that the coefficients used give
the same number of each type of atom on
both sides of the arrow. Also check to see
that the coefficients used are the smallest
integers that give the balanced equation.
The balanced equation is:
2H2(g) + O2(g)  2H2O(l)
or could be:
4H2(g) + 2O2(g)  4H2O(l)
preferred
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Section 6.3
Balancing Chemical Equations
Another Balancing Example:
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21
Section 6.3
Balancing using the underline method.
Balancing Chemical Equations
2 NaOH(aq)
Na2O(s) + H2O(l)
CH4(g) + 2 O2(g)
CO2(g) + 2 H2O(g)
42 Fe(s) + 3 O2(g)
LiOH(s) + CO2(g)
2 KClO3(s)
D
MnO2
2 Fe2O3(s)
LiHCO3(s)
2 KCl(s) +3 O2(g)
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22
Section 6.3
Balancing Chemical Equations
Exercise
Balance the following equation in standard
form (lowest multiple integers) and determine
the sum of the coefficients?
4 FeO(s) + O2(g) 2 Fe2O3(s)
a)
b)
c)
d)
3
4
7
14
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Section 6.3
Balancing Chemical Equations
Exercise
Which of the following correctly balances the
chemical equation given below? There may be more
than one correct balanced equation. If a balanced
equation is incorrect, explain what is incorrect about
it.
CaO + C  CaC2 + CO2
I.
II.
III.
IV.
CaO2 + 3C  CaC2 + CO2
2CaO + 5C  2CaC2 + CO2
CaO + (2.5)C  CaC2 + (0.5)CO2
4CaO + 10C  4CaC2 + 2CO2
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24
Section 6.3
Balancing Chemical Equations
Exercise
Of the three that are correct, which one is preferred
most (the most accepted convention)? Why?
CaO + C  CaC2 + CO2
I.
II.
III.
IV.
CaO2 + 3C  CaC2 + CO2
2CaO + 5C  2CaC2 + CO2
CaO + (2.5)C  CaC2 + (0.5)CO2
4CaO + 10C  4CaC2 + 2CO2
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Section 6.3
Balancing Chemical Equations
Concept Check
When balancing a chemical equation, which of the
following statements is false?
a) Subscripts in the reactants must be conserved in the
products.
b) Coefficients are used to balance the atoms on both
sides.
c) When one coefficient is doubled, the rest of the
coefficients in the balanced equation must also be
doubled.
d) Phases are often shown for each compound but are
not critical to balancing an equation.
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26
Section 6.3
Balancing Chemical Equations
Notice
•
•
•
•
The number of atoms of each type of element
must be the same on both sides of a balanced
equation.
Subscripts must not be changed to balance an
equation.
A balanced equation tells us the ratio of the
number of molecules which react and are
produced in a chemical reaction.
Coefficients can be fractions, although they are
usually given as lowest integer multiples.
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Section6.3
7.1
Section
Balancing Chemical Equations
Four Driving Forces Favor Chemical Change
1.
2.
3.
4.
Formation of a solid
Formation of water
Transfer of electrons
Formation of a gas
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28
Section6.3
7.1
Section
Ukrainian Wolves
Balancing Chemical Equations
Partially charged water
molecules act like Ukrainian
Wolves in desloving fully
charged NaCl ion.
http://www.northland.cc.mn.us/biology/B
iology1111/animations/dissolve.swf
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29
Section6.3
7.2
Section
Balancing Chemical Equations
Precipitation
•
A reaction in which a solid forms is called a
precipitation reaction.
 Solid = precipitate
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30
Section6.3
7.2
Section
Balancing Chemical Equations
What Happens When an Ionic Compound Dissolves in Water?
•
•
The ions separate and move around independently.
Strong electrolyte – each unit of the substance that
dissolves in water produces separated ions.
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31
Section6.3
7.2
Section
Balancing Chemical Equations
What Happens When an Ionic Compound Dissolves in Water?
•
K2CrO4(aq) + Ba(NO3)2(aq)  Products
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32
Section6.3
7.2
Section
Balancing Chemical Equations
How to Decide What Products Form
•
•
•
K2CrO4(aq) + Ba(NO3)2(aq)  Products
The mixed solution contains four types of ions: K+,
CrO42–, Ba2+, and NO3–.
Determine the possible products from the ions in the
reactants. The possible ion combinations are:
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Section6.3
7.2
Section
Balancing Chemical Equations
How to Decide What Products Form
•
•
•
•
Decide which is most likely to be the yellow solid
formed in the reaction.
K2CrO4(aq)
reactant
Ba(NO3)2(aq) reactant
The possible combinations are KNO3 and BaCrO4.
 KNO3
white solid
 BaCrO4
yellow solid
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34
Section6.3
7.2
Section
Balancing Chemical Equations
Using Solubility Rules
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Section6.3
7.2
Section
Balancing Chemical Equations
Using Solubility Rules
•
Predicting Precipitates
 Soluble solid
 Insoluble solid
 Slightly soluble solid
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36
Section6.3
7.2
Section
Balancing Chemical Equations
Let’s Practice Determining Solubility
Which of the following are soluble in water?
Na2CO3
yes
Cu(OH)2
no
CaCl2
yes
Ba(OH)2
yes
AgCl
no
Ca3(PO4)2
no
BaSO4
no
Pb(NO3)2
yes
(NH4)2S
yes
PbCl2
no
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37
Section6.3
7.2
Section
Balancing Chemical Equations
How to Predict Precipitates When Solutions of Two Ionic
Compounds Are Mixed
1. Write the reactants as they actually exist
before any reaction occurs. Remember that
when a salt dissolves, its ions separate.
2. Consider the various solids that could form. To
do this, simply exchange the anions of the
added salts.
3. Use the solubility rules to decide whether a
solid forms and, if so, to predict the identity of
the solid.
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Section6.3
7.2
Section
Balancing Chemical Equations
Concept Check
Which of the following ions form compounds
with Pb2+ that are generally soluble in water?
a)
b)
c)
d)
e)
S2–
Cl–
NO3–
SO42–
Na+
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Section6.3
7.2
Section
Balancing Chemical Equations
Concept Check
A sodium phosphate solution reacts with a lead(II)
nitrate solution. What precipitate, if any, will form?
a)
b)
c)
d)
Pb3(PO4)2
NaNO3
Pb(NO3)2
No precipitate will form.
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Section6.3
7.2
Section
Balancing Chemical Equations
Concept Check
Consider a solution with the following ions present:
3NO3- , Pb2+, K+, Ag+, Cl-, SO2,
PO
4
4
When all are allowed to react (and there is plenty
available of each), how many different solids will
form? List them.
Five different solids will form.
PbCl2, PbSO4, Pb3(PO4)2, AgCl, Ag3PO4
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Vv
Section 6.3
Balancing Chemical Equations
Chapters 6-7b
Chemical Reactions
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42
Section 6.3
Balancing Chemical Equations
7.3
7.4
7.5
7.6
7.7
Describing Reactions in Aqueous Solutions
Reactions That Form Water: Acids and
Bases
Reactions of Metals with Nonmetals
(Oxidation–Reduction)
Ways to Classify Reactions
Other Ways to Classify Reactions
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43
Section6.3
7.2
Section
Net
Ionic Equations
Balancing
Chemical Equations
AgNO3(aq) + NaCl(aq)
Molecular Equation
Ag++NO3- + Na++ Cl-
1. Divorce
++NO - + Na++ ClAg
3
2. Change Partners
Ag++NO3- + Na++ Cl3. Soluble?
AgCl(s) + NaNO3(aq)
AgCl
+ NaNO3
AgCl(s) + Na+ + NO3-
Total Ionic Equation
4. Cross out
Spectator Ions
Ag++ Cl-
5. Balance
Ag++ Cl-
AgCl(s)
AgCl(s)
Net Ionic Equation
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44
Section6.3
7.2
Section
Net
Ionic Equations
Balancing
Chemical Equations
Pb(NO3)2(aq) + NaI(aq)
Molecular Equation
Pbl2(s) + NaNO3(aq)
+2+NO - + Na++ lPb
3
1. Divorce
+2+NO - + Na++ lPb
3
2. Change Partners
Pb+2+NO3- + Na++ l 3. Soluble?
Pbl2 + NaNO3
Pbl2(s) + Na+ NO3-
Total Ionic Equation
4. Cross out
Spectator Ions
Pb+2+ l-
5. Balance
Pb+2+ 2 l-
Pbl2(s)
Pbl2(s)
Net Ionic Equation
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45
Section6.3
7.2
Section
Net
Ionic Equations
Balancing
Chemical Equations
BaCl2(aq) + Na2SO4(aq)
Molecular Equation
BaSO4(s) + NaCl(aq)
+2 + Cl- + Na++ SO -2
Ba
4
1. Divorce
+2 + Cl- + Na++ SO -2
Ba
BaSO4 + NaCl
4
2. Change Partners
Ba+2 + Cl- + Na++ SO4-2
BaSO4(s)+ Na+ +
3. Soluble?
Cl- Total Ionic Equation
4. Cross out
Spectator Ions
Ba+2+ SO4-2
5. Balance
Ba+2+ SO4-2
BaSO4(s)
BaSO4(s)
Net Ionic Equation
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46
Section6.3
7.3
Section
Balancing Chemical Equations
Types of Equations for Reactions in Aqueous Solutions
1. Molecular Equation
 Shows the complete formulas of all
reactants and products.
 It does not give a very clear picture of
what actually occurs in solution.
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47
Section6.3
7.3
Section
Balancing Chemical Equations
Types of Equations for Reactions in Aqueous Solutions
2. Complete Ionic Equation
 All strong electrolytes are shown as ions.

Notice: K+ and NO3– ions are present in solution
both before and after the reaction.
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Section6.3
7.3
Section
Balancing Chemical Equations
Types of Equations for Reactions in Aqueous Solutions
2. Complete Ionic Equation
 Spectator ions – ions which do not
participate directly in a reaction in
solution.
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49
Section6.3
7.3
Section
Balancing Chemical Equations
Types of Equations for Reactions in Aqueous Solutions
3. Net Ionic Equation
 Only those components of the solution
that undergo a change.

Notice: Spectator ions are not shown in the net
ionic equation.
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Section6.3
7.3
Section
Balancing Chemical Equations
Concept Check
Write the correct molecular equation, complete ionic
equation, and net ionic equation for the reaction between
cobalt(II) chloride and sodium hydroxide.
Molecular Equation:
CoCl2(aq) + 2NaOH(aq)  Co(OH)2(s) + 2NaCl(aq)
Complete Ionic Equation:
Co2+(aq) + 2Cl(aq) + 2Na+(aq) + 2OH(aq) 
Co(OH)2(s) + 2Na+(aq) + 2Cl(aq)
Net Ionic Equation:
Co2+(aq) + 2OH(aq)  Co(OH)2(s)
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51
Section 6.3
Balancing Chemical Equations
Types of Reactions
1. Combination (one product) A + B
2. Decompostion (one reactant) A
3. Single Replacement A + BC
4. Double Replacement AB + CD
C
B+C
AC + B
AD(s) + CB(l)
5. Acid-Base (Neutralization) HA + BOH
6. Combustion- Organic + O2
H2O + BA
CO2 + H2O
7. No Reaction- both products are (aq).
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52
Section6.3
7.4
Section
Balancing Chemical Equations
Arrhenius Acids and Bases
•
A strong acid is one in which virtually
every molecule dissociates (ionizes) in
water to an H+ ion and an anion.
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Section6.3
7.4
Section
Balancing Chemical Equations
Strong Acids Behave as Strong Electrolytes
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Section6.3
7.4
Section
Balancing Chemical Equations
Arrhenius Acids and Bases
•
A strong base is a metal hydroxide that is
completely soluble in water, giving separate
OH ions and cations.
 Most common examples: NaOH and KOH
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Section6.3
7.4
Section
Balancing Chemical Equations
Arrhenius Acids and Bases
•
The products of the reaction of a strong acid
and a strong base are water and a salt.

•
Net ionic equation

•
Salt  Ionic compound
H+(aq) + OH−(aq)  H2O(l)
Reaction of H+ and OH− is called an acidbase reaction.


H+  acidic ion
OH−  basic ion
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Section6.3
7.4
Section
Balancing Chemical Equations
Summary of Strong Acids and Strong Bases
1. The common strong acids are aqueous
solutions of HCl, HNO3, and H2SO4.
2. A strong acid is a substance that completely
dissociates (ionizes) in water (into H+ ions
and anions).
3. A strong base is a metal hydroxide
compound that is very soluble in water (and
dissociates into OH– ions and cations).
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Section6.3
7.4
Section
Balancing Chemical Equations
Summary of Strong Acids and Strong Bases
4. The net ionic equation for the reaction of a
strong acid and a strong base is always the
same: it shows the production of water.
5. In the reaction of a strong acid and a strong
base, one product is always water and the
other is always an ionic compound called a
salt, which remains dissolved in the water.
This salt can be obtained as a solid by
evaporating the water.
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Section6.3
7.4
Section
Balancing Chemical Equations
Summary of Strong Acids and Strong Bases
6. The reaction of H+ and OH– is often called an
acid-base reaction, where H+ is the acidic ion
and OH– is the basic ion.
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59
Section6.3
7.4
Section
Balancing Chemical Equations
Concept Check
The net ionic equation for the reaction of HNO3 and
LiOH is
a)
b)
c)
d)
H+ + NO3– + LiOH → H2O + LiNO3
HNO3 + LiOH → H2O + LiNO3
H+ + OH– → H2O
Li+ + NO3– → LiNO3
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Section6.3
7.5
Section
Balancing Chemical Equations
Oxidation–Reduction Reaction
•
•
Reactions between metals and nonmetals
involve a transfer of electrons from the metal
to the nonmetal.
A reaction that involves a transfer of electrons.

2Mg(s) + O2(g)  2MgO(s)
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Section
Section6.3
7.4
7.5
Section
Balancing Chemical Equations
Concept Check
Which of the following best describes what is
happening in the following representation of an
oxidation–reduction reaction:
a)
b)
c)
d)
Metal Al gains 3 e– and O2 – in Fe2O3 loses these 3e–.
Metal Al gains 3 e– and Fe3+ in Fe2O3 loses these 3e–.
Metal Al loses 3 e– and O2 – in Fe2O3 gains these 3e–.
Metal Al loses 3 e– and Fe3+ in Fe2O3 gains these 3e–.
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Section
Section6.3
7.4
7.5
Section
Balancing Chemical Equations
Characteristics of Oxidation–Reduction Reactions
1. A metal–nonmetal reaction can always be
assumed to be an oxidation–reduction
reaction, which involves electron transfer.
2. Two nonmetals can also undergo an
oxidation–reduction reaction. At this point we
can recognize these cases only by looking
for O2 as a reactant or product. When two
nonmetals react, the compound formed is
not ionic.
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Section
Section6.3
7.4
7.6
Section
Balancing Chemical Equations
Precipitation Reaction
•
Formation of a solid when two solutions
are mixed.
•
Notice this is also a double–
displacement reaction.

AB + CD  AD + CB
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Section
Section6.3
7.4
7.6
Section
Balancing Chemical Equations
Acid–Base Reaction
•
Involves an H+ ion that ends up in the
product water.


H+(aq) + OH−(aq)  H2O(l)
HCl(aq) + KOH(aq)  H2O(l) + KCl(aq)
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65
Section
Section6.3
7.4
7.6
Section
Balancing Chemical Equations
Oxidation–Reduction Reaction
•
Transfer of electrons

2Li(s) + F2(g)  2LiF(s)
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Section
Section6.3
7.4
7.6
Section
Balancing Chemical Equations
Formation of a Gas
•
•
Oxidation–reduction reaction
Single–replacement reaction

A + BC  B + AC
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Section
Section6.3
7.4
7.7
Section
Balancing Chemical Equations
Combustion Reactions
•
Involve oxygen and produce energy (heat) so
rapidly that a flame results.


CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
Special class of oxidation–reduction reactions.
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Section
Section6.3
7.4
7.7
Section
Balancing Chemical Equations
Synthesis (Combination) Reactions
•
A compound forms from simpler materials.


C(s) + O2(g)  CO2(g)
Only one product!
Special class of oxidation–reduction reactions.
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69
Section
Section6.3
7.4
7.7
Section
Balancing Chemical Equations
Decomposition Reactions
•
Occurs when a compound is broken down
into simpler substances.


Only one reactant!
2H2O(l)  2H2(g) + O2(g)
Special class of oxidation–reduction reactions.
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Section
Section6.3
7.4
7.7
Section
Balancing Chemical Equations
Summary
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