R EACTIONS IN
AQUEOUS SOLUTIONS
WHAT IS AN AQUEOUS SOLUTION?
WHAT HAPPENS WHEN STUFF
DISSOLVES?


When some ionic compounds are placed in water, they
dissociate, or break apart into their respective
cation/anion pairs.
 In this Aqueous state, they are more readily
available to react chemically with other substances
in the solution than if they were solid.
Ionic compounds are made up of two halves.

Cation




Comes first.
Positive charge.
Can be mono- or poly- atomic.
Anion



Comes second.
Negative charge.
Can be mono- or poly- atomic.
PRACTICE: DISSOCIATE THE FOLLOWING
IONIC COMPOUNDS…

CaCl2 

Pb(NO3)2 

Li3P 

Na2SO4 

(NH4)3PO3 
SOLUBILITY

When a substance readily dissolves in water it is
said to be soluble in water.


Solubility is the mass of substance that will dissolve
in water at a given temperature.
Vocabulary







Solute – What is being dissolved.
Solvent – What is doing the dissolving.
Solution – Solvent with Solute dissolved in it.
Soluble – will dissolve.
Insoluble – will not dissolve.
Slightly Soluble – will only dissolve a little bit.
Moderately Soluble – will dissolve more than a
little but not all the way.
WHY DO SOME THINGS
DISSOLVE?
Different Substances dissolve to different
extents.
 Whether or not something will dissolve in water
depends on two things:

1.
Polarity
The generic rule of solubility is like dissolves like.
 Water is a polar solvent and will only dissolve polar things.
 All ionic compounds, acids, and bases are polar but they do
not all dissolve in water.

2.
Bond Energy
Bond energy holds ions in an ionic compound together.
 In order to dissolve, the polar nature of water must be
strong enough to overcome the attraction between the ions
in a compound.

THE POLARITY OF WATER
Why is water such a good solvent?
•Water is a polar
molecule.
•That means that
the electrons are
unevenly shared.
•The oxygen in
water hogs the
electrons in its
bonds with
hydrogen, giving
it a partial
negative charge.
•The hydrogen
side is left with a
partial positive
charge.
•When ionic
compounds are
placed into water,
they break apart
because the partial
positive and
negative sides of
water attack the
negative and
positive ions
respectively.
•In other words,
the ionic
compounds are
pulled apart by the
polar water
molecules.
WATER BREAKS APART A SOLUBLE SALT
Remember, Salt is another word for Ionic
compounds.
HOW TO TELL IF
SOMETHING DISSOLVES …
ONLY STRONG ELECTROLYTES
DISSOLVE COMPLETELY IN H2O

Strong Electrolytes are completely soluble in
water.
Strong Acids
 Strong Bases
 Soluble Salts



Ionic compounds are sometimes called salts.
Weak electrolytes are only moderately or slightly
soluble in water.
Weak acids
 Weak bases
 Moderately or slightly soluble salts.

STRONG ACIDS & BASES
Strong Acids
HCl
HI
HBr
HNO3
HClO3
H2SO4
If it is not on the list of
strong acids or bases, it
can be considered weak.
Strong Bases
KOH
NaOH
LiOH
CsOH
RbOH
Ba(OH)2
Ca(OH)2
Sr(OH)2
TABLE 5-3 SIMPLE RULES FOR THE SOLUBILITY OF
SALTS IN WATER (MODIFIED FROM TEXT)
1.
2.
3.
4.
5.
6.
Most nitrate (NO3-) salts are soluble.
Most salts containing the alkali metal ions (Li+, Na+,
K+, Cs+ & Rb+) and the ammonium ion (NH4+) are
soluble.
Most chloride, bromide and iodide salts are soluble.
Notable exceptions are salts containing the ions Ag+,
Pb2+, & Hg22+
Most sulfate salts are soluble. Notable exceptions
are BaSO4, PbSO4, Hg2SO4, & CaSO4.
Most hydroxide salts are only slightly soluble. The
important soluble hydroxides are NaOH & KOH.
The compounds Ba(OH)2, Sr(OH)2, & Ca(OH)2 are
marginally soluble.
Most sulfide (S2-), carbonate (CO32-), and phosphate
(PO43-) salts are only slightly soluble.
SIDE BY SIDE COMPARISON
 Soluble
Salts
(aqueous in water)
NO3 Alkali metal ions







Li +, K +, Na +, Cs +, Rb+
NH4+
Most Cl-, Br-, & IMost SO42NaOH, KOH
Moderately = Ba(OH)2,
Ca(OH)2

Slightly or Insoluble Salts
(solid in water)








AgCl, PbCl2, Hg2Cl2
AgBr, PbBr2, Hg2Br2
AgI, PbI2, Hg2I2
BaSO4, PbSO4, CaSO4
OHS2CO32PO43-
PRACTICE: DETERMINE SOLUBILITY

Write an equation to show what happens to each
substance in water.

CaF2

HF

Mg(OH)2

Ca3(PO4)2

HCl

KOH
IONS IN AQUEOUS SOLUTION ARE
ABLE TO REACT… THE QUESTION IS,
“WILL THEY?”
WHAT HAPPENS WHEN A REACTION TAKES
PLACE?


When a pair of ions comes in contact in an
aqueous environment, there is a chance that they
will react to form a compound that is NOT soluble
in water.
This new compound will form and fall out of
solution.
If a solid forms, a non-uniform cloud will form in the
solution.
 If a gas forms, bubbles will appear and escape the
solution.
 If liquid forms, it can be hard to see easily, unless
there is a color change, or the new liquid is immiscible
in water.

DETERMINING THE FORMULA OF A
POSSIBLE PRODUCT…

Make a list of the ions in solution.

Pair them up, remember make cation-anion pairs.

Some Rules to help decide on products:
1.
2.
The products will NOT be the same as the
reactants.
Solid, liquid, and gas compounds must have a zero
net charge.

3.
Substances with ionic charges are ALWAYS aqueous.
They must contain both cations and anions.

Most compounds contain only ONE OF EACH.
WHAT MAKES CHEMISTRY HAPPEN?
o.O
DRIVING FORCE

A chemical reaction will only take place if there is
a force driving it forward.
Remember entropy?
 Everything wants to attain the lowest possible state
of energetic existence.
 If a reaction leads to a state of lower energy, then
reactants will have a tendency to form products.


Driving Force – a force that leads reactants to
react chemically and form products. The Driving
Force causes a reaction to go in the direction of
the arrow.
4 COMMON DRIVING FORCES
1.
Formation of a solid

2.
Formation of water

3.
In a RedOx Reaction, one substance loses electrons in an
Oxidation Reaction, and another substance gains
electrons in a Reduction Reaction.
Formation of a gas


An Acid Base Reaction occurs when bases act to
neutralize acids, and water forms.
Transfer of electrons

4.
A Precipitation Reaction occurs when two aqueous
solutions are combined and a solid precipitates, or
forms and falls out of solution.
Similar to formation of a solid, except the product is a
gas.
Understanding the driving force behind a chemical reaction
can aid in predicting the products of a chemical reaction.
DRIVING FORCES:
#1 FORMATION OF A SOLID
Solubility Rules and Precipitation Reactions
PRECIPITATION REACTIONS

Precipitation Reactions occur when two
solutions are mixed and an insoluble solid falls
out of solution.


The solid product is called a precipitate.
Precipitation Reactions are also classified as
Double Displacement Reactions.
This is because the reactants contain cation-anion
pairs that swap partners, or displace each other.
 A reaction will occur between two ionic pairs IF one
of the possible products is insoluble.
 Insoluble products precipitate driving the reaction
forward.


To know if either possible product will be a
precipitate, you must know the Solubility Rules
which can be found on Table 7.1 (page 178).
SIDE BY SIDE COMPARISON
 Soluble
Salts
(aqueous in water)
NO3 Alkali metal ions







Li +, K +, Na +, Cs +, Rb+
NH4+
Most Cl-, Br-, & IMost SO42NaOH, KOH
Moderately = Ba(OH)2,
Ca(OH)2

Slightly or Insoluble Salts
(solid in water)








AgCl, PbCl2, Hg2Cl2
AgBr, PbBr2, Hg2Br2
AgI, PbI2, Hg2I2
BaSO4, PbSO4, CaSO4
OHS2CO32PO43-
PREDICTING PRECIPITATION
REACTIONS

When two solutions of aqueous salts are mixed,
the solubility rules can be used to predict
whether or not a reaction will occur:
1.
2.
3.

Write the formulas of the two salts as if they were
reactants.
Determine the formulas of the possible products by
swapping the anions of the two salts.
Check the solubility rules.
If one or both of the possible products is
insoluble (or marginally or slightly soluble) than
a reaction will occur and a precipitate will form.
USE THE SOLUBILITY RULES TO PREDICT
THE PRODUCTS OF A REACTION:

1.
2.
3.



Predict what will happen when the following salts are
mixed in aqueous solution. Write the balanced
equation for any reaction that occurs.
potassium nitrate & barium chloride
sodium sulfate & lead(II) nitrate
potassium hydroxide and iron(III) nitrate
KNO3(aq) + BaCl2(aq)  no rxn
Na2SO4(aq) + Pb(NO3)2(aq)  PbSO4(s) +
2NaNO3(aq)
3KOH(aq) + Fe(NO3)3(aq)  Fe(OH)3(s) + 3KNO3(aq)
USE THE SOLUBILITY RULES TO PREDICT
THE PRODUCTS OF A REACTION:

1.
2.
3.



Predict what will happen when the following salts are
mixed in aqueous solution. Write the balanced
equation for any reaction that occurs.
barium nitrate & sodium chloride
sodium sulfide & copper(II) nitrate
ammonium chloride & lead(II) nitrate
Ba(NO3)2 (aq) + NaCl(aq)  no rxn
Na2S (aq) + Cu(NO3)2(aq)  CuS (s) + 2NaNO3(aq)
2NH4Cl(aq) + Pb(NO3)2(aq)  PbCl2(s) + 2NH4NO3
(aq)
TYPES OF CHEMICAL
EQUATIONS
DESCRIBING REACTIONS IN AQUEOUS
SOLUTION



There are several different ways that a single
chemical reaction can be represented.
Different styles of chemical equations are used to
express relative information as needed.
We will discuss three.
1.
2.
3.
Molecular Equations
Complete Ionic Equations
Net Ionic Equations
CONSIDER THE FOLLOWING
REACTION
An aqueous solution of potassium chromate is
added to an aqueous solution of barium nitrate
and a solid precipitate forms.
 Can you write a chemical equation to represent
the reaction that takes place?


The equation you wrote is a molecular equation.
We will now compare this equation to two new
equations that show the same chemical reaction
in different ways, the complete ionic equation
and the net ionic equation.
1. MOLECULAR EQUATIONS


Molecular equations show the complete formulas
of all reactants and products.
The molecular equation for the potassium
chromate and barium nitrate reaction is:
K2CrO4(aq) + Ba(NO3)2  BaCrO4(s) +
2KNO3(aq)

These are the type of equations studied when
first learning how to balance chemical equations.
(Chapter 6)
2. COMPLETE IONIC EQUATIONS

In a complete ionic equation, all substances that are
strong electrolytes and in the aqueous are represented as
ions.


Reminder: Strong Electrolytes = Strong Acids, Strong Bases, and Soluble Salts
Which compounds dissociate in our example equation?
K2CrO4(aq) + Ba(NO3)2  BaCrO4(s) + 2KNO3(aq)
2K+(aq) + CrO42-(aq) + Ba2+(aq) + 2NO3-(aq)  BaCrO4(s) + 2K+(aq) + 2NO3(aq)

This type of equation better represents the physical state
of the reactants and products as they occur in the reaction
environment.

The precipitate is not represented as ions because it falls out of
solution as an insoluble solid.
3. NET IONIC EQUATIONS &
SPECTATOR IONS


The Net Ionic Equation shows only those components that are
directly involved in the chemical change that occurs.
The Spectator Ions are those components that are present as
both reactants and products, and remain unchanged by the
reaction.
Consider our Example:
2K+(aq) + CrO42-(aq) + Ba2+(aq) + 2NO3-(aq)  BaCrO4(s) + 2K+(aq) + 2NO3-(aq
1.
What is the net ionic equation?
2.
What are the Spectator Ions?
K+ & NO3-
CrO42-(aq) + Ba2+(aq)  BaCrO4(s)
WRITING EQUATIONS FOR
REACTIONS

Write the molecular, complete ionic, and net ionic
chemical equations for the reaction between aqueous
silver nitrate and aqueous sodium chloride.

Molecular Equation
NaCl(aq) + AgNO3(aq)  AgCl(s) + NaNO3(aq)

Complete Ionic Equation
Na+(aq) + Cl-(aq) + Ag+(aq) + NO3-(aq)  AgCl(s) + Na+(aq) + NO3-(aq)

Net Ionic Equation
Cl-(aq) + Ag+(aq)  AgCl (s)
WRITING EQUATIONS FOR
REACTIONS

Write the molecular, complete ionic, and net ionic chemical
equations for the reaction between aqueous potassium
hydroxide and aqueous iron(III) nitrate.

Molecular Equation
3KOH(aq) + Fe(NO3) 3(aq)  Fe(OH) 3 (s) + 3KNO3(aq)

Complete Ionic Equation
3K+(aq) + 3OH-(aq) + Fe3+(aq) + 3NO3-(aq)  Fe(OH) 3 (s) + 3K+(aq) + 3NO3-(aq)

Net Ionic Equation
Fe3+(aq) + 3OH-(aq)  Fe(OH) 3 (s)
WRITING EQUATIONS FOR
REACTIONS

Write the molecular, complete ionic, and net ionic chemical
equations for the reaction between aqueous sodium sulfide and
aqueous copper(II) nitrate.

Molecular Equation
Na2S(aq) + Cu(NO3)2(aq)  CuS(s) + 2NaNO3(aq)

Complete Ionic Equation
2Na+(aq) + S2-(aq) + Cu2+(aq) + 2NO3-(aq)  CuS(s) + 2Na+(aq) +2 NO3-(aq)

Net Ionic Equation
Cu2+(aq) + S2-(aq)  CuS(s)
WRITING EQUATIONS FOR
REACTIONS

Write the molecular, complete ionic, and net ionic chemical
equations for the reaction between aqueous ammonium
chloride and aqueous lead(II) nitrate.

Molecular Equation
2NH4Cl(aq) + Pb(NO3)2(aq)  PbCl2 (s) + 2NH4NO3(aq)

Complete Ionic Equation
2NH4+(aq) + 2Cl-(aq) + Pb2+(aq) + 2NO3-(aq)  PbCl2 (s) +2 NH4+(aq) + 2NO3-(aq)

Net Ionic Equation
Pb2+(aq) + 2Cl-(aq)  PbCl2 (s)
DRIVING FORCES:
#2 PROTON TRANSFER
Acid-Base Reactions
WHAT ARE ACIDS AND BASES?
Arrhenius (Late 1800s)

Acid – A substance
that produces H+ ions
when it dissolves in
water.

Brǿnsted & Lowry (Early 1900s)

Acid – A proton donor.


HA  H+ + A-
Base – A proton
acceptor.


Base – A substance that
produces OH- ions when
it dissolves in water.
 BOH  B+ + OH-

HA  H+ + A-
B- + H+  BH
This new definition was
necessary to accommodate
substances that behave
chemically as bases, but do
not contain OH- ions.
STRONG ACIDS

Strong Acids are Strong Electrolytes
They dissociate completely in water.
 100% of the HA molecules break up into H+ and A- ions.


There are six common strong acids chemists MUST know:
1.
2.
3.
4.
5.
6.
HCl
HI
HBr
HNO3
HClO3
H2SO4
-
Hydrochloric Acid
Hydroiodic Acid
Hydrobromic Acid
Nitric Acid
Chloric Acid
Sulfuric Acid *1st dissociation only
WEAK ACIDS

Weak Acids are Weak Electrolytes
They DO NOT dissociate completely in water.
 1% of the HA molecules break up into H+ and A- ions.


There are five common weak acids chemists MUST know:
1.
2.
3.
4.
5.
HF
H3PO4
HC2H3O2
H3BO3
C6H8O7
-
Hydrofluoric Acid
Phosphoric Acid *1st dissociation only
Acetic Acid
Boric Acid *1st dissociation only
Citric Acid
STRONG BASES

Strong Bases are Strong Electrolytes
They dissociate completely in water.
 100% of the BOH molecules break up into B+ and OH- ions.


There are eight common strong bases chemists MUST know:
1.
2.
3.
4.
5.
6.
7.
8.
KOH
NaOH LiOH
CsOH RbOH Ba(OH)2 Ca(OH)2 Sr(OH)2 -
Potassium Hydroxide
Sodium Hydroxide
Lithium Hydroxide
Cesium Hydroxide
Rubidium Hydroxide
Barium Hydroxide
Calcium Hydroxide
Strontium Hydroxide
WEAK BASES

Weak Bases are Weak Electrolytes
They DO NOT dissociate completely in water.
 1% of the BOH molecules break up into B+ and OH- ions.


There are six common weak bases chemists MUST know:
1.
2.
3.
4.
5.
6.
NH3
C5H5N
CH3NH2
C3H5O2NH2
Be(OH)2
Mg(OH)2
- Ammonia
- Pyridine
- Methylamine
- Alanine
- Beryllium Hydroxide
- Magnesium Hydroxide
ACID-BASE REACTIONS

Proton Transfer
This occurs in every acid-base reaction.
 When acids and bases react, a proton is transferred
from the acid to the base.
 This is the driving force of acid-base reactions.
 Once the proton is transferred, it is in a position of
less chemical potential.


Formation of water
If the acid contains H+ and the base contains OH- the
acid base reaction will form water.
 The remaining cation and anion form a salt.


If both the acid and base are weak, no reaction
occurs because no ions are present to initiate the
proton transfer.
NEUTRALIZATION
Acid-Base Reactions are also called Neutralization
Reactions
 Acids and Bases are extremely corrosive materials.

This means that they destroy things they come into
contact with, including skin, plant tissue, and even
metals when the acid or base is strong enough.
 For this reason, chemists should be VERY careful
when handling them.


When acid-base reactions occur, corrosive
materials become harmless materials (water and
salt).

Another way to say this is to say that they are
neutralized.
THE EQUATIONS OF A
STRONG ACID-STRONG BASE REACTION



Consider the reaction between nitric acid and potassium
hydroxide.
Molecular Equation
HNO3(aq) + KOH(aq)  KNO3(aq) + H2O(l)
Complete Ionic Equation – both acid & base dissociate.
H+(aq) + NO3-(aq) + K+(aq) + OH-(aq)  K+(aq) + NO3-(aq) + H2O(l)

Net Ionic Equation – always shows formation of water.
H+(aq) + OH-(aq)  H2O(l)
THE EQUATIONS OF A
WEAK ACID-STRONG BASE REACTION



Consider the reaction between acetic acid and
sodium hydroxide.
Molecular Equation
HC2H3O2(aq) + NaOH(aq)  NaC2H3O2 (aq) +
H2O(l)
Complete Ionic Equation – only the base dissociates.
HC2H3O2(aq) + Na+(aq) + OH-(aq)  Na+(aq) + C2H3O2-(aq) + H2O(l)

Net Ionic Equation
HC2H3O2(aq) + OH-(aq)  H2O(l) + C2H3O2-(aq)
THE EQUATIONS OF A
STRONG ACID-WEAK BASE REACTION



Consider the reaction between hydrochloric acid and
ammonia.
Molecular Equation
HCl(aq) + NH3 (aq)  NH4+(aq) + Cl-(aq)
Complete Ionic Equation – only the acid dissociates.
H+(aq) + Cl-(aq) + NH3(aq)  NH4+(aq) + Cl-(aq)

Net Ionic Equation
H+(aq) + NH3(aq)  NH4+(aq)
DRIVING FORCES:
#3 TRANSFER OF
ELECTRONS
Oxidation – Reduction
Reactions
WHAT ARE REDOX REACTIONS?


Oxidation-Reduction Reactions are also called Redox
Reactions.
Redox Reactions are reactions in which one or more
electrons is transferred.






Electron transfers involve energy transfers.
Electrons move from positions of higher to lower energy releasing
energy in the process.
Most reactions used for energy production by living things
are Redox Reactions.
Combustion Reactions are Redox Reactions.
Reactions between metals and nonmetals are also Redox
Reactions.
Transfer of electrons can occur even when neither reactants
nor products involve ionic species.
REDOX REACTION
BETWEEN A
METAL & NONMETAL

2Al(s) + 3Cu(OH)2(aq)  2Al(OH)3(aq) + 3Cu(s)

Electrons are lost by Aluminum
Al solid has a charge of zero
 Al in Al(OH)3 has a charge of 3+


Electrons are gained by Copper
Cu in Cu(OH)2 has a charge of 2+
 Cu solid has a charge of zero


In this RedOx Reaction, 6 Electrons are transferred
from 2 Aluminum atoms to 3 Copper ions.
WHY DO ELECTRON TRANSFERS OCCUR IN NONIONIC
COMPOUNDS?

Electronegativity can be thought of as the ability of an atom to
draw electrons towards itself in a covalent bond.


This property increases across the periodic table from left to right, and
going up. (F is the most electronegative element)
The electron effectively belongs to the more electronegative
element in the bond, but there is no charge because the bond is
covalent.

The most electronegative atom takes one or more electrons from the
least electronegative atom.
OXIDATION STATES



Oxidation States provide a way to keep track of
electrons in oxidation-reduction reactions that do
not contain ions.
The oxidation states of the atoms in a covalent
compound is the imaginary charges the atoms
would have if the shared electrons are divided
based on the atoms’ relative abilities to attract
electrons.
Its an imaginary charge and is represented as a
+/-# superscript.

Ionic Charges are represented as #+/-
RULES FOR ASSIGNING OXIDATION STATES








Fluorine’s oxidation state is -1 in all its compounds because it is
the most electronegative element.
Oxygen has an oxidation state of -2 in almost all its compounds.
Hydrogen has a +1 oxidation state in its covalent compounds.
Elemental atoms have an oxidation state of zero.
Elements that occur as diatomic molecules have an oxidation
state of zero in this state, including O2.
The sum of the oxidation states must be zero for an electrically
neutral compound.
Monoatomic ions have the same oxidation state as their ionic
charges in aqueous solution and in ionic compounds.
The sum of the oxidation states of the atoms in polyatomic ions is
equal to the polyatiomic ion’s charge.
ASSIGNING OXIDATION STATES

Assign oxidation states to all atoms in the following:
1.
2.
3.
4.
5.
NaCl
H2O
CO2
SF6
NO3-
1.
+1 + -1 = 0
2.
2(+1) + -2 = 0
3.
+4 + 2(-2) = 0
4.
+6 + 6(-1) = 0
5.
+5 + 3(-2) = -1
OXIDATION-REDUCTION REACTIONS
Review: Chemical Reactions that involve electron
transfer are called oxidation-reduction reactions.
 Oxidation-Reduction reactions actually include
two half reactions, an oxidation and a reduction.
 Oxidation – the half reaction in which one or
more electrons is lost.



The element with an increased oxidation number is
oxidized and the substance that contains it is called
the reducing agent.
Reduction – the half reaction in which one or
more electrons is gained.

The element with decreased oxidation number is
reduced and the substance that contains it is called
the oxidizing agent.
LEO GOES GER
Oxidation
Reduction
X + e -  X M3+ +3e-  M
M  M+ + e  2X-  X2 + 2e-

Loses Electrons

Is or Is Contained in
the Reducing Agent

Electron Donor





Gains Electrons
Is or Is Contained in
the Oxidizing Agent
Electron Acceptor
REDOX HALF REACTIONS

In a reaction between sodium and chlorine, an
electron from sodium is transferred to chlorine.


Think about the ions that Na and Cl form…Which
one is going to take the electron?


2Na + Cl2  2NaCl
Chlorine will; it forms an anion.
The half reactions show which species loses and
which gains the electron(s) being transferred.

Remember one loses and become a cation (LEO), and the
other gains and becomes an anion (GER):



Sodium is oxidized:
Chlorine is reduced:
Na  Na+ + eCl + e-  Cl-
The electron is transferred from the metal (sodium) to
the more electronegative nonmetal (chlorine).
MORE EXAMPLES OF REDOX RXNS

Magnesium reacts with Oxygen to form Magnesium
Oxide: 2Mg(s) + O2(g)  2MgO(s)

The Half Reactions:
Mg  Mg2+ + 2e2 O + 2e  O

Magnesium is oxidized and is the reducing agent.
 Oxygen is reduced and is the oxidizing agent.


Aluminum reacts with Iron(III) Oxide to form Iron and
Aluminum Oxide: Al(s) + Fe2O3(s)  Al2O3(s) + Fe(s)

The Half Reactions:
Al  Al3+ + 3e Fe3+ + 3e-  Fe

Aluminum is oxidized and is the reducing agent.
 Iron is reduced and is the oxidizing agent.

BALANCING REDOX EQUATIONS
•
•
•
•
•
Oxidation-Reduction Reactions that take place in
aqueous solution can be complex and very difficult to
balance by the inspection or the try & try again
method.
To balance Redox reactions it is useful to separate the
reaction into two half reactions.
The halves include a reaction that involves the
substance being reduced, and a reaction that involves
the substance being oxidized.
The two halves are balanced individually, following a
set of balancing rules the are unique to either acidic
or basic aqueous solutions because different ions are
available to react in these two types of solutions.
Once balanced, the halves are put back together to
form one balanced net ionic equation for the Redox
Reaction in question.
OXIDATION-REDUCTION REACTIONS

Identify the atoms that are oxidized and reduced and
specify the oxidizing and reducing agents in the
following reaction:
2Al(s) + 3I2(s)  2AlI3(s)

Assign oxidation states
Aluminum is oxidized and is the reducing agent.
 Iodine is reduced and is the oxidizing agent.

OXIDATION-REDUCTION REACTIONS

Identify the atoms that are oxidized and reduced, and
specify the oxidizing and reducing agents in each of the
following equations:
2PbS(s) + 3O2(g)  2PbO(s) + 2SO2(g)
PbO(s) + CO(g)  Pb(s) + CO2(g)

1.
2.
Determine the change in oxidation state for each atom.
Sulfur is oxidized & Oxygen is reduced. O2 is the
oxidizing agent & PbS is the reducing agent. Lead’s
oxidation state does not change.
Lead is reduced & Carbon is oxidized. PbO is the
oxidizing agent & CO is the reducing agent. Oxygen’s
oxidation state does not change.
IDENTIFYING ELECTRON
TRANSFER IN REDOX REACTIONS


Write the half reactions for the following RedOx
reactions. Identify which species is being
oxidized and which is being reduced. Indicate
which is the oxidizing agent and which is the
reducing agent.
2Al(s) + 3I2(s)  2AlI3(s)
Half Reactions:
Al  Al3+ + 3eI + e-  I-
Aluminum is oxidized and is the reducing agent.
 Iodine is reduced and is the oxidizing agent.

IDENTIFYING ELECTRON
TRANSFER IN REDOX REACTIONS


Write the half reactions for the following RedOx
reactions. Identify which species is being
oxidized and which is being reduced. Indicate
which is the oxidizing agent and which is the
reducing agent.
2Cs(s) + F2(g)  2CsF(s)
Half Reactions:
Cs  Cs2+ + 2eF + e -  F-
Cesium is oxidized and is the reducing agent.
 Fluorine is reduced and is the oxidizing agent.
