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Chemical Reactions
Dr. Schuerch
Describing Chemical Reactions
• A chemical reaction occurs when one or
more reactants changes into one or more
products, example:
H2O2  H2O + O2
•
•
•
•
H2O2 is the reactant
H2O and O2 are the products
+ used to separate two reactants or two products
 “Yields,” separates reactants from products
Symbols Used in Chemical Equations
+
Used to separate two reactants or two products
“Yields,” separates reactants from products
Used in place of  for reversible reactions
(s)
(l)
(g)
(aq)
Δ
Heat
Pt
Designates a reactant or product in the solid state; placed after the
formula
Designates a reactant or product in the liquid state: placed after
the formula
Designates a reactant or product in the gaseous state; placed after
the formula
Designates an aqueous solution; the substance is dissolved in water;
placed after the formula
Indicates that heat is supplied to the reaction
A formula written above or below the yield sign indicates its use as a
catalyst (in this example, platinum).
Word Equations
• Describe what happens in a chemical reaction
– To write a word equation, write the names of the
reactants to the left of the arrow separated by plus
signs; write the names of the products to the right of
the arrow, also separated by a plus sign
– Examples:
Reactants  Products
Iron + oxygen  iron(III) oxide
Hydrogen peroxide  water + oxygen
Methane + oxygen  carbon dioxide + water
Chemical Equations
• Chemical equations are a representation of a
chemical reaction: the formulas of the reactants
on the left are connected by an arrow with the
formulas for the products on the right
– Example: Fe + O2  Fe2O3
• Indicate the physical state of reactants and
products and if a catalyst is used
– Example: Fe(s) + O2(g)  Fe2O3(s)
• Balance the Equation
– Example: 4Fe(s) + 3O2(g)  2Fe2O3(s)
Balancing Chemical Equations
• To write a balanced chemical equation,
first write the skeleton equation. Then use
the coefficients to balance the equation so
that it obeys the law of conservation of
mass
– Example:
H2O2  O2 + H2O (Skeleton Equation)
2H2O2(aq)  O2(g) + 2H2O(l) (Balanced Equation)
Coefficients
Rules for Writing and Balancing Equations
1.
Determine the correct formulas for all the reactants and products
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 yield sign ()
in between. If two or more reactants or products are involved, separate
their formulas with plus signs
3.
Determine the number of atoms of each element in the reactants and
products. 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. When no
coefficient is written, it is assumed to be 1. Begin by balancing elements
that appear only once on each side of the equation. Never balance an
equation by changing 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
Balance the Following Equations
O2(g) 
1.
Al(s) +
2.
AgNO3 +
3.
Zn(OH)2 +
4.
FeCl3 +
H2S 
Al2O3(s)
Ag2S +
H3PO4 
NaOH 
HNO3
Zn3(PO4)2 +
Fe(OH)3 +
H2 O
NaCl
Predicting whether a reaction will occur
• Driving forces of a reaction
– Transfer of electrons
– Formation of a solid
– Formation of water
– Formation of a gas
Types of Chemical Reactions
A. Oxidation - Reduction Reactions
1.
2.
3.
4.
Combination or synthesis Reactions
Decomposition Reactions
Single-Replacement Reactions
Combustion Reactions
B. Double-Replacement Reactions
1. Precipitation Reactions
2. Acid-base Reactions
Oxidation-Reduction
Combination aka Synthesis Reactions
• A chemical change in which two or more
substances react to form a single new substance
– Examples
2Na(s) + Cl2(g)  2NaCl(s)
S(s) + O2(g)  SO2(g)
2S(s) + 3O2(g)  2SO3(g)
Fe(s) + S(s)  FeS(s)
2Fe(s) + 3S(s)  Fe2S3(s)
•Note: When two non
metals or a transition
metal and non metal
are present in the
reaction, more than
one product may result
Oxidation-Reduction
Decomposition Reaction
• Chemical reactions in which a single
compound breaks down into two or more
simpler products
– Examples:
Δ
2HgO(s)  2Hg(l) + O2(g)
Electricity
2H2O(l)  2H2(g) + O2(g)
Δ
2HI  H2(g) + I2(s)
Oxidation-Reduction
Single Replacement Reactions
• A chemical change in which one element
replaces a second element in a compound
– Identified by both reactants and products having
an element and a compound
– Examples:
2K(s) + 2HOH(L)  2KOH(aq) + H2(g)
Br2(aq) + 2NaI(aq)  2NaBr + I2(aq)
Zn(s) + Cu(NO3)2(aq)  Cu(s) + Zn(NO3)2(aq)
Single Replacement Reactions
• Whether one metal will displace another
metal from a compound depends upon the
relative reactivities of the two metals
– Example
Zn(s) + Cu(NO3)2(aq)  Cu(s) + Zn(NO3)2(aq)
Cu(s) + Zn(NO3)2(aq)  No reaction
– Zinc can displace copper because it is more
reactive than copper
– Copper cannot displace zinc because it is less
reactive than zinc
Oxidation Reduction
Single Replacement Reactions
Decreasing reactivity
Activity Series of Metals
• Notice that zinc is higher on the
activity series of metals than copper
• A reactive metal will replace any
metal listed below it in the activity
series
Zn + Cu(NO3)2  Cu + Zn(NO3)2
Cu + Zn(NO3)2  No reaction
Oxidation Reduction
• Halogens (Group 7) also participate in
single replacement reactions
– Example:
Br2 + 2NaI  2NaBr + I2
NaBr + I2  No reaction
• Halogens higher in the periodic table will
replace halogens lower in the periodic
table
Decreasing reactivity
Single Replacement Reactions
Oxidation Reduction
Combustion Reactions
• Is a chemical change in which an element or a
compound reacts with oxygen often producing
energy of the form of heat and light
– Examples:
2C8H16(l) + 25O2(g)  16CO2(g) + 18H2O(l)
2Mg(s) + O2(g)  2MgO(s)
S(s) + O2(g)  SO2(g)
– Note that some of the combustion reactions are also
combination reactions
Oxidation-Reduction Reaction
• Electrons are transferred from atom or ion to another
atom or ion.
– The species receiving electrons are oxidizing agents (also the more
electronegative species)
• Also called electron acceptor
– The species donating electrons are reducing agents (also the less
electronegative species)
• Also called electron donor
– Example
Balanced Equation:
2𝐴𝑙 + 𝐹𝑒2 𝑂3 → 𝐴𝑙2 𝑂3 + 2𝐹𝑒
Net Ionic Equation:
2𝐴𝑙 + 2𝐹𝑒 3+ → 2𝐴𝑙 3+ + 2𝐹𝑒
Oxidation Half Reaction:
2𝐴𝑙 → 2𝐴𝑙 3+ + 6𝑒 −
Reduction Half Reaction:
2𝐹𝑒 3+ + 6𝑒 − → 2𝐹𝑒
– In the example above, aluminum is oxidized by the oxidizing agent iron
and iron is reduced by the reducing agent aluminum
Double Replacement Reactions
• A chemical change involving an exchange of positive
ions between two compounds
– Generally occur in aqueous solution and often produce a
precipitate, a gas, or molecular compound such as water
– Examples:
– Precipitation (forms a solid or gas as a product)
• Na2S(aq) +Cd(NO3)2(aq)  CdS(s) + 2NaNO3(aq)
• 2NaCN(aq) + H2SO4(aq)  2HCN(g) + Na2SO4(aq)
– Acid Base (forms water as product)
• Ca(OH)2(aq) + 2HCl(aq)  CaCl2(aq) + 2H2O(l)
Predicting the Formation of a Precipitate
• You can predict the formation of a precipitate by using the general
rules for solubility of ionic compounds (Table 7.1 pg 178 in your text)
Reactions in Aqueous Solution:
Complete Ionic Equations
• Many important reactions take place in water—that is in aqueous
solution
– When in solution, many ionic compounds (called strong
electrolytes) will ionize splitting into their constitute ions
• Weak electrolytes are ionic compounds that are only slightly soluble
in water and when formed in water precipitate out of solution (we
can think of them as being insoluble)
– Equations in aqueous solution can be written as complete ionic
equations
• Example:
A molecular equation
AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)
Can be written as a complete ionic equation:
Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq)  AgCl(s) + Na+(aq) + NO3-(aq)
Reactions in Aqueous Solution:
Net Ionic Equations
• Given the complete ionic equation for the formation of silver chloride
precipitate from the mixture of silver nitrate and sodium chloride
Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq)  AgCl(s) + Na+(aq) + NO3-(aq)
– Notice that the sodium and nitrate ion appear unchanged on
both sides of the equation
• Because they are unchanged in the chemical reaction they are
called spectator ions
– When you rewrite the ionic equation without the spectator ions,
you get a net ionic equation
Ag+(aq) + Cl-(aq)  AgCl(s)
• A net ionic equation shows only those particles involved in the
reaction and is balanced with respect to both mass and charge
•
Reactions in Aqueous Solution:
Net Ionic Equations (Continued)
Consider the following skeleton equation
Pb(s) + AgNO3(aq)  Ag(s) + Pb(NO3)2(aq)
–
Write a complete ionic equation
Pb(s) + Ag+(aq) + NO3-(aq)  Ag(s) + Pb2+(aq) 2NO3-(aq)
–
Write a net ionic equation
Pb(s) + Ag+(aq)  Ag(s) + Pb2+(aq)
–
Is the net ionic equation balanced (Notice the
charge on both sides of the equation)? Balance the
net ionic equation
Pb(s) + 2Ag+(aq)  2Ag(s) + Pb2+(aq)
Reactions in Aqueous Solution
Acid-Base reactions
• Acid - a substance that produces H+ ions when dissolved in
water (Arrhenius definition)
– Strong acids completely disassociate in solution and can be
considered a strong electrolyte

HCl 
 H  Cl
H 2O

• Base - is a substance that produces OH- ions when dissolved in
solution (Arrhenius definition)
– Strong bases completely disassociate in solution and can be considered
a strong electrolyte

NaOH 
 Na  OH
H 2O

Reactions in Aqueous Solution
Acid-Base reactions
• Acids and Bases react to form water and a
ionic compound called a salt
2O
NaOH  HCl H

NaCl  H 2O
2O
Na   OH  H   Cl H

Na   Cl  H 2O


OH  H 
 H 2O
H 2O