TOPIC 9
OXIDATION AND REDUCTION
9.1 INTRODUCTION TO OXIDATION
AND REDUCTION
• Oxidation originally, the addition of
oxygen or loss/ removal of hydrogen
• Reduction originally, the loss/
removal of oxygen or addition of
hydrogen
• Redox reactions when both
oxidation and reduction occur in a
chemical reaction
HISTORY OF CHEMISTRY
• Joseph Preistley discovered oxygen in 1774
and it was called phlogiston until Antoine
Lavoisier
• Antoine Lavoisier determined the role and
name of oxygen in the mid- to late- 1700’s.
He also termed oxidation for reactions
involving oxygen
• Justus von Liebig proposed that reduction
was the removal of oxygen and addition of
hydrogen
OXIDATION AND REDUCTION NOW
• Refer to loss/ gain of electrons in a
chemical reaction
• Many of the reactions considered
as original oxidation and
reduction (in reference to oxygen)
still fall under the category of
redox now
REMEMBER!!
• Oxidation
• Is
• Loss
• Reduction
• Is
• Gain
EQUATIONS
• 2Mg(s) + O2(g) 2MgO(s)
• Oxidation: Mg Mg2+ + 2e• Electrons in product
• Reduction: O2 + 4e- 2O2• Electrons as reactants
HALF- EQUATIONS
• Describe only half of the total reaction
• Broken into oxidation half and
reduction half
• When the two halves are combined to
represent the full reaction, the
equation must be balanced such that
the electrons on the reactant side
equal the electron on the product side
BALANCING HALF-EQUATIONS
• Na Na+ +e• O2+ 4e-  2O2• Have to multiply the whole half-equation for
Na by 4 to gain equal electrons before
combining equations
•
• Sum: 4Na +O2 + 4e-  4Na+ + 2O2- + 4e• Cancel electrons: 4Na +O2 +  4Na+ + 2O2-
OXIDATION NUMBERS
• How the oxidized/ reduced element or compound is
determined
• Rules to determine oxidation state
• Any uncombined element has an oxidation number of 0 (Ex.
O2)
• A simple ion has an oxidation number equal to its charge
• For a compound, the oxidation number is 0
• For an oxoanion, the number is equal to the sum of the
oxidation numbers
• Hydrogen has an oxidation number of +1, unless combined
with a metal, then it will be -1
• Oxygen has an oxidation number of -2, except in H2O2 (-1)
and OF2 (+2)
COVALENT BONDS
• The more electronegative element
gets the negative oxidation number
• Less electronegative gets a positive
oxidation number
• Less electronegative is named first in
the compound
• Ex. ClF
• Cl +1 O.N.
• F -1 O.N.
-/+ O.N.
• Negative oxidation numbers
implies that element has
‘gained control’ of electrons
• Positive O.N. implies that the
element has ‘lost control’ of
electrons
NAMING INORGANIC COMPOUNDS
• O.N. are used to name ionic inorganic
compounds
• Referred to as stock notation
• Oxidation number inserted right after name of
ion and roman numerals after name/ symbol
of element
• Ex. FeCl2
[Fe2+ 2Cl-] iron(II) chloride
• FeCl3
[Fe3+ 3Cl-] iron(III) chloride
• Notation only used for transition metals, tin and
lead
IDENTIFYING REDOX REACTIONS
• Recognized by:
• Deducing all O.N. of atoms in the
molecular, ionic, or half- equation
• Examine all numbers to see if any
O.N. on atoms have changed
• Increase in O.N. is oxidation, decrease is
reduction
REDOX VS. NOT
• 2FeCl2(s) + Cl2(g)  2FeCl3(s)
• Oxidation numbers of iron, +2 and +3
• O.N. of chlorine, 0 and -1
• Redox rxn
• MgO(s) + 2HCl(aq) MgCl2(aq)
+H2O(l)
• O.N. of all remain the same
• Not a redox rxn
DISPROPORTIONATION
• Occurs when a single species is
oxidized and reduced
simultaneously
• Ex. Cl2(g) + H2O(l)  HOCl(aq) +
HCl(aq)
• One Cl decreases from 0 to -1, the
other increases from 0 to +1
9.2 REDOX EQUATIONS
• Constructing half-equations
1. write down formulae of reactant and product
(for one half)
2. Balance with respect to present ion(s)
3. Balance oxygen with water molecule(s)
4. Balance hydrogen with hydrogen ion(s)
5. Determine charges on both sides of the halfequation
6. Balance charges by adding electrons to side
that is more positive
CON’T
• For oxidation half-equation,
electrons will be on left hand side
• For reduction half-equation,
electrons will be on right hand
side
• Use a similar process to
determine each type of halfequation
FORMING REDOX EQUATIONS
• Full redox equations are
constructed by combining two
half-equations
• One for the oxidizing agent and one
for reducing agent
• Ex:
REDOX TITRATIONS
• Similar to acid-base titrations, but
has electron transfer instead of a
hydrogen transfer
• Overall equation used for
stoichiometric calculations is
obtained by combining two halfequations such that, the electrons
lost equal electrons gained
COMMON OXIDIZING AGENTS FOR
REDOX TITRATIONS
• Acidified manganate (VII) ions
• Manganate (VII) ions are purple, but reduced form (Mn(II))
almost colorless. Potassium manganate (VII) difficult to prepare
and reacts slowly
• Acidified dichromate (VI) ions
• Orange in color, reduced form is green (Cr(III)). Can be used as
primary standard
• Iron (III) ions or salts
• Iodine
• Red brown in color colorless when reduced
• Starch improves the color change
• No indicators necessary for these since a color change
is involved
• Acidified hydrogen peroxide
COMMON REDUCING AGENTS FOR
REDOX TITRATIONS
• Iron (II) salts or ions
• Ethanedioic (oxalic) acid and ethanedioate
(oxalate) ions
• Performed at 80°C to catalyze
• Able to determine concentration with high
accuracy
• Hydrogen peroxide
• In presence of a more powerful oxidizing agent
• Iodide ions
• Sodium thiosulfate (VI) or thiosulfate (VI) ions
OXIDIZING AGENTS
• A substance that brings about the oxidation of
substances by accepting electrons from the substance
they oxidize
• Undergo the process of reduction
• Common oxidizing agents
•
•
•
•
•
•
•
•
•
Oxygen
Ozone (O3)
Chlorine
Acidified potassium manganate (VII)
Acidified aqueous potassium dichromate (VI)
Acidified hydrogen peroxide
Metal ions
Hydrogen ions
Manganese (IV) oxide
REDUCING AGENTS
• A substance that brings about reduction of
a substance by donating electrons to the
reduced substance
• Common reducing agents
• Hydrogen atom
• Carbon
• Carbon monoxide
• Metals
• More reactive metals
EXTENTION
• The terms oxidizing and
reducing agents are relative.
• A weak reducing agent can
be ‘forced’ to become an
oxidizing agent in the
presence of a more powerful
reducing agent
9.3 REACTIVITY
• Reactions of metals with metal ions in
solution
• Can test reactivity of metals by placing a
piece of Mg metal into a solution of aq
metal ions to observe whether the Mg
changes color to indicate a chemical rxn
• Procedure is repeated with different types
of metals
ACTIVITY SERIES
• With the data collected during experimentation
with the metals, a reactivity or activity series can
be constructed
• Such that, the metals are arranged in order of
increasing reactivity
• The type of reactions that take place are
displacement reactions
• The more reactive metal ‘pushes out’ the less reactive
metal from its salts
• Can also be performed by heating two sold metals
together, or by reacting with water
USING REACTIVITY SERIES
• Able to deduce the probability of a redox
rxn from a reactivity series
• Carbon and hydrogen can be included in
an activity series
• Metals above H will displace H from dilute
acids, those below will not
• Metals above carbon cannot be produced
by reduction of metal ions, those below
can
NON-METALS
• Displacement can occur with
non-metals also
• Especially the halogens
• A reactivity series for halogens
corresponds to their placement
on the periodic table
9.4 VOLTAIC CELLS
• Background
• These attractions of opposite charges produce these
‘simple batteries’
• Electric charge
• Positive and negative charges
• Measured in coulombs (C). Electron = 1.6E-19C
• Carried by electrons from negative terminal to positive
• Electric current
• Rate at which electric charge flows through a circuit
• Measured in amperes (A). 1A= 1C s-1
• Large current produced by a large amount of slowly
moving charge or a small amount of quickly moving
charge
POTENTIAL DIFFERENCE
• When there is a difference in electric
potential between two points in a circuit
• Gives electrical energy to the charge
• The charge transfers energy into other
forms
• Heat, light, chemical energy
• Measured in volts (V)= 1C J
• The charge simply carries energy around
the circuit, is not used
VOLTAIC CELLS
• Daniell cell is a simple voltaic cell
• Constructed by placing Zn electrode in
zinc sulfate solution and a Cu electrode
into a copper (II) sulfate solution
• The two electrode are connected by
wires and a high-resistance voltmeter to
allow electrons to flow
• A salt bridge allows ions to flow and
maintain electrical neutrality
WHAT’S HAPPENING IN A DANIELL
CELL
• Zn undergoes oxidation (higher in reactivity
series)
• Zn ions dissolve into the water
• Electrons flow from surface of Zn electrode
through external circuit to surface of Cu
electrode to reduce it
• Continues until either all Zn electrode or Cu ions
are used up
• Anode where oxidation occurs in a voltaic cell
• Cathode where reduction occurs in a voltaic
cell
ANODE AND CATHODE
• Electric current flows from anode to
cathode
• Difference in electric potential
between anode and cathode,
measured by a voltmeter, is the cell
potential
• The further apart the two metals are in
the activity series, the larger the cell
potential will be
9.5 ELECTROLYTIC CELLS
• A substance that allows electricity to pass
though itself in a conductor
• Metals
• Graphite
• Aqueous solutions of acids, alkalis, ionic compounds
• Semi-conductors only slightly conduct
• A substance not allowing electricity to flow
through itself is an insulator
• Common insulators:
• Non-metallic elements (except graphite)
• Dry samples of covalent compounds
• Solid samples of ionic substances
CONDUCTION OF ELECTRICITY
• Substance must contain electrically
charge particles that move freely
when a potential difference or voltage
are introduced
• Valence electrons are the charged
particles in solid and liquid states of
metals
• Ionic solids do not conduct because
their ions cannot move freely
ELECTROLYSIS OF MOLTEN SALT
• When electric current passes through a metal,
the metal is unaffected
• When electric current passes through an ionic
substance (molten or in a solution), chemical
decomposition occurs
• Substances that undergo this process are electrolytes
• Process of electrolyte decomposition
electrolysis
• Especially used in industry to prepare Al,
Cl,NaOH and H