Reduction – Oxidation Reactions “REDOX” - Weebly

Balancing Redox Reac
Reactions

Rules for Writing Half-Reactions

1. Write an unbalanced ½ reaction showin
2. Balance all atoms except H and O
3. Balance O by adding H2O(l)
4. Balance H by adding H+(aq)
5. Balance the charge by adding e- and canc

For basic solutions only:

6. Add OH-(aq) to both sides to equal the numbe
7. Combine H+(aq) and OH-(aq) on the same side t

from both sides.

ctions using Half-

ng formulas for reactants and products

cel anything that is the same on both sides

er of H+(aq) present
to form H2O(l). Cancel equal amounts of H2O(l)

Balancing Redox Reac
Half-Re

SUMMARY

1. Use the information provided to start two h
2. Balance each half-reaction equation.

3. Multiply each half-reaction by simple whole

4. Add the two half-reaction equations, cancel
exactly the same on both sides of the equati

ctions by Constructing
eactions

half-reaction equations.

e numbers to balance electrons lost and gained.
lling the electrons and anything else that is
ion.

Balancing Redox Reactions by

• Example: A person suspected of being intoxic
person’s breath reacts with an acidic dichroma
acid) and aqueous chromium(III) ions. Predic

• Create a skeleton equation from the informatio

• Separate the entities into the start of two half-r
• Now use the steps you learned for balancing ha

• Now, balance the electrons lost and gained, an
anything else that is exactly the same on both s

y Constructing Half Reactions

cated blows into this device and the alcohol in the
ate ion solution to produce acetic acid (ethanoic
ct the balanced redox reaction equation.
on provided:

reaction equations
alf reactions

nd add the half reactions. Cancel the electrons and
sides of the equation.

Redox Terms

▫ Review: “LEO the lion says GER”

x Loss of electrons = entity being oxidized
x Gain of electrons = entity being reduced

x BUT…. Chemists don’t say “the reactant b

x Rather, they use the terms OXIDIZING A

x OXIDIZING AGENT: causes oxidation by
another substance in

x REDUCING AGENT: causes reduction by
substance in a redox

What does this mean? Let’s revisit our first exa
Which reactant was reduced?

So…. Which is the Oxidizing Agent (OA)

LEO = Oxidized Zn(s) Æ Zn 2+
GER = Reduced 2 H+(aq) + 2 e

being oxidized” or “the reactant being reduced”
AGENT (OA) and REDUCING AGENT (RA)

removing (gaining) electrons from
n a redox reaction

donating (losing) electrons to another
reaction

ample when zinc and hydrochloric acid reacted.
Which was oxidized?

)? Which is the Reducing Agent (RA)

(aq) + 2 e- Reducing Agent
e- Æ H2 (g) Oxidizing Agent

Redox Terms

▫ Silver ions were reduced to silver metal b
copper metal was oxidized to copper(II) i

▫ If Ag+(aq) is reduced it is the: OXIDIZING
▫ If Cu(s) is oxidized it is the:

REDUCIN

It is important to note that oxidation and redu
and oxidizing agents and reducing agents

by reaction with copper metal. Simultaneously,
ions by reaction with silver ions.

G AGENT (OA)
NG AGENT (RA)

uction are processes,
are substances.

REDOX React

Reduction

• Historically, the formation of a metal from
its “ore” (or oxide)

▫ I.e. nickel(II) oxide is reduced by
hydrogen gas to nickel metal

NiO(s) + H2(g) Æ Ni(s) + H2O(l)

Ni +2 Æ Nio

• A gain of electrons occurs (so the entity
becomes more negative)

• Electrons are shown as the reactant in the
half-reaction

• A species undergoing reduction will be
responsible for the oxidation of another
entity – and is therefore classified as an
oxidizing agent (OA)

tions … so far

Oxidation

• Historically, reactions with oxygen

▫ I.e. iron reacts with oxygen to
produce iron(III) oxide

4 Fe(s) + O2(g) Æ Fe2O3(s)

Fe 0 Æ Fe+3

• A loss of electrons occurs (so the entity
becomes more positive)

• Electrons are shown as the product in the
half-reaction

• A species undergoing oxidation will be
responsible for the reduction of another
entity – and is therefore classified as an
reducing agent (RA)

Redox Terms

• Summary so far:

▫ The substance that is reduced (
is also known as the oxidizing a

▫ The substance that is oxidized (
is also knows as the reducing a

• Question: If a substance is a
does this mean in

The substance has a very

• Question: If a substance is a
does this mean in

The substance has a weak attraction fo

(gains electrons)
agent
(loses electrons)
agent

very strong oxidizing agent, what
terms of electrons?

strong attraction for electrons.

very strong reducing agent, what
terms of electrons?

or its electrons, which are easily removed

Redox Table

• A reaction is considered spontan

• A reduction ½ reaction table is u
of a reaction

▫ Reduction Tables show reduction
therefore all the reactants will be

▫ If we list the OA’s from an experim
we create a reduction ½ reaction

SOA Ag+(aq) +
Cu2+(aq) +
Zn2+(aq) +
Mg2+(aq) +

neous if it occurs on its own
useful in predicting the spontaneity

n ½ reactions in the forward direction,
oxidizing agents

ment in decreasing order of strength,
table:

1 e- Æ Ag(s) SRA
2 e- Æ Cu(s)
2 e- Æ Zn(s)
2 e- Æ Mg(s)

Building Redox Tab

• Consider the following experimental
the redox table you have created

Hg2+(aq) Cu2+(a

Hg(s) ✗ ✗
Cu(s) ✓ ✗
Ag(s) ✓ ✗
Au(s) ✗ ✗

SOA Au3+(aq) +
Hg2+(aq) +
Ag+(aq) +
Cu2+(aq) +
Zn2+(aq) +
Mg2+(aq) +

bles #1

l information and add half-reactions to

aq) Ag+(aq) Au3+(aq)

✗✓
✓✓
✗✓
✗✗

3 e- Æ Au(s) SRA
2 e- Æ Hg(s)
1 e- Æ Ag(s)
2 e- Æ Cu(s)
2 e- Æ Zn(s)
2 e- Æ Mg(s)

Building Redox Tab

SOA Au3+(aq) + 3 e-

Hg2+(aq) + 2 e-
Ag+(aq) + 1 e-
Cu2+(aq) + 2 e-
Zn2+(aq) + 2 e-
Mg2+(aq) + 2 e-

▫ The spontaneity rule!

x A reaction will be spontaneous if

OA = Spontaneous
above Reaction
RA

bles #1

- Æ Au(s) SRA

- Æ Hg(s)
- Æ Ag(s)
Æ Cu(s)

Æ Zn(s)
Æ Mg(s)

f on a redox table:

RA = Non-spontaneous
below Reaction
OA

Building Redox Tab

• Example 2: Use the following information

OA RA

3 Co 2+ (aq) + 2 In(s) Æ 2 In 3+ (aq) + 3 C

OA RA

Cu 2+ (aq) + Co(s) Æ Co 2+ (aq) + Cu(s)

OA RA

Cu 2+ (aq) + Pd(s) Æ no reaction

SOA Pd2+(aq) + 2 e

Cu2+(aq) + 2 e-
Co2+(aq) + 2 e-
In3+(aq) + 3 e-

bles #2

n to create a table of reduction ½ reactions

Co(s) Pd(s)
Co(s)
Cu2+
Co2+

In(s)

e- Æ Pd(s) SRA

- Æ Cu(s)
- Æ Co(s)

Æ In(s)

Building Redox Tab

• Example 3: Use the following information

OA RA

2 A 3+ (aq) + 3 D(s) Æ 3 D2+ (aq) + 2 A(s)

OA RA

G + (aq) + D(s) Æ no reaction

OA RA

3 D 2+ (aq) + 2 E(s) Æ 3 D(s) + 2 E3+(aq)

OA RA

G + + E(s) Æ no reaction
(aq)

SOA A3+(aq) + 3 e-
2 e-
D2+(aq) + 3 e-
E3+(aq) + 1 e- Æ
G+(aq) +

bles #3

n to create a table of reduction ½ reactions

A3+

D2+(aq) D(s)

E(s)

G+

Æ A(s) SRA

Æ D(s)
Æ E(s)
Æ G(s)

Building Redox Tab

• So far we have been using examples wh
and the reducing agents are metal atom

▫ Non-metal atoms I.e. Cl2(g) + 2e- Æ 2 C
▫ Non-metal ions I.e. 2 Br- (aq) Æ Br2(l) +

• Redox Table Trend

▫ OA’s tend to be metal ions and non-m
▫ RA’s tend to be metal atoms and non-

• Also, are there any entities that could a

▫ Multivalent metals

bles

here the oxidizing agents are metal ions
ms. What else could gain or lose electrons?

Cl-(aq) (Cl2(g) could act as a Reducing Agent)
+ 2 e- (2Br-(aq) could act as an Oxidizing Agent)

metal atoms
-metal ions

act as both OA or RA?

• Example 4: Use the following information

RA OA

Ag(s) + Br2(l) Æ AgBr(s)

RA OA

Ag(s) + I2(s) Æ no evidence of rea

OA RA

Cu2+(aq) + I-(aq) Æ no redox reactio

OA RA

Br2(l) + Cl-(aq) Æ no evidence of re

SOA Cl2(g) + 2 e- Æ

Br2(l) + 2 e- Æ
Ag+(aq) + 1 e- Æ
I2(s) + 2 e- Æ
Cu2+(aq) + 2 e- Æ

n to create a table of reduction ½ reactions

action Br2(l) Cl-
on
eaction I2(s) Ag(s)
Cu2+(aq) I-(aq)

Æ 2Cl-(aq) SRA

Æ 2Br-(aq)
Æ Ag(s)
Æ 2I-(aq)
Æ Cu(s)

Predicting Redox Re

• Now that you know what redox reactions ar
if a reaction will occur (is spontaneous) a
reaction equation will be. How do we d

1. The first step is to determine all the entities
▫ Helpful reference: Table 5 pg. 680

▫ Remember: In solutions, molecules and ions
independently of each other.

▫ Example: When a solution of potassium
slowly poured through acidified iron(II)

▫ Does a redox reaction occur and what is th

eactions

re, you will be responsible for determining
and if so, what the
do this?

s that are present.

s behave

permanganate is
) sulfate solution.
he reaction equation?

Predicting Redox Re

2. The second step is to determine all poss

▫ This is a crucial step!! Things to watch out

x Combinations
x (i.e. MnO4-(aq) is an oxidizing agent only i
x To indicate this draw an arc between the
hydrogen ion

x Species that can act as both OA and RA
x Any lower charge multivalent metal i.e. F
x Water (H2O(l))
x Label both possibilities in your list

eactions

sible OA’s and RA’s

t for:

in an acidic solution)
e permanganate and

Fe2+, Cu+, Sn2+, Cr2+

• Before we move on, let’s pract
▫ Pg. 680 #23

tice Step 1 and 2

• Pg. 680 #23



• Pg. 680 #23



Predicting Redox Re

3. The third step is to identify the SOA a

SOA

3. The fourth step is to show the ½ reac

▫ SOA equation straight from table. SRA

▫ Are these equations balanced? Do the n
▫ If not, multiply one or both equations b

eactions

and SRA using Appendix C11 (page 805)

SRA

ctions (from the redox table) and balance

A equation read from right to left

number of electrons lost = electrons gained
by a number then add the balanced equations

Predicting Redox Re

3. The last step is to predict the spontan
ionic equation represent a spontaneo
redox reaction?

If the SOA

above Æ Spontaneo

SRA??

If the SRA

below Æ Nonspontan

SOA

eactions

neity. Does the net
ous or non-spontaneous

ous

neous

Predicting Redox Re

Could copper pipe be used to transport a hyd
1. List all entities
1. Identify all possible OA’s and RA’s

1. Identify the SOA and SRA

2. Show ½ reactions and balance

3. Predict spontaneity

Since the reaction is
nonspontaneous, it should be
possible to use a copper pipe to

carry hydrochloric acid

eactions #2

drochloric acid solution?

REDOX Reacti

Reduction

• Historically, the formation of a metal from
its “ore” (or oxide)

▫ I.e. nickel(II) oxide is reduced by
hydrogen gas to nickel metal

NiO(s) + H2(g) Æ Ni(s) + H2O(l)

Ni +2 Æ Nio

• A gain of electrons occurs (so the entity
becomes more negative)

• Electrons are shown as the reactant in the
half-reaction

• A species undergoing reduction will be
responsible for the oxidation of another
entity – and is therefore classified as an
oxidizing agent (OA)

• Decrease in oxidation number

ions … the end

Oxidation

• Historically, reactions with oxygen

▫ I.e. iron reacts with oxygen to
produce iron(III) oxide

4 Fe(s) + O2(g) Æ Fe2O3(s)

Fe 0 Æ Fe+3

• A loss of electrons occurs (so the entity
becomes more positive)

• Electrons are shown as the product in the
half-reaction

• A species undergoing oxidation will be
responsible for the reduction of another
entity – and is therefore classified as an
reducing agent (RA)

• Increase in oxidation number