autumn lecture 6 (redox and oxidation number)
TRANSCRIPT
Redox & Oxidation number Term 1: Week 4
The bronze statue of Great Buddha in Kamakura, Kanagawa, Japan
Outline
• Oxidation and reduction
• Strength of oxidizing and reducing agents
• Ionic equations
• Oxidation number
Oxidation and Reduction Reactions in Daily Life
Corrosion
Oxidation with hydrogen peroxide Batteries
Oxidation of apple
Combustion
Oxidation and reduction in biology
2n CO2 + 2n DH2 + photons → 2(CH2O)n + 2n DO Carbon dioxide + electron donor + light energy → carbohydrate + oxidized electron donor
Oxidation and reduction in medicine
• Oxidation is linked with the effects of aging of humans.
• To forestall the effects of oxidation, doctors recommend antioxidants - natural reducing agents such as vitamin C and vitamin E.
Jacob Tsiperovich
Oxidation and Reduction
• Originally oxidation meant “addition of oxygen” and reduction meant “removal of oxygen”.
• REDOX – a contraction of reduction-oxidation
It implies that two processes always act together
Fe2O3 loses oxygen
and is reduced
Fe2O3 (s) + 3 CO (g) 2Fe (s) + 3 CO2 (g)
CO gains oxygen and is oxidized
Oxidation and Reduction
• Oxidation is the loss of electrons.
• Reduction is the gain of electrons.
2Na0 + Cl20 2Na+Cl-
2Na 2Na+ + 2ē
Cl2 + 2ē 2Cl-
Memory aid
Oxidizing and reducing agents
2 C(s) + O2(g) CO(g)
Oxidizing and reducing agents
Sodium is oxidized – it is the reducing agent
Chlorine is reduced – it is the oxidizing agent
2Na0 + Cl20 2Na+Cl-
2Na 2Na+ + 2ē
Cl2 + 2ē 2Cl-
Active metals: Lose electrons easily Are easily oxidized Are strong reducing agents
Active nonmetals: Gain electrons easily Are easily reduced Are strong oxidizing agents
Oxidizing and reducing agents
1. Free (elemental) nonmetals become negative ions:
F2 + 2ē 2F-
O2 + 4ē 2O2-
2. Positive (usually metal) ions become neutral:
Ag+ + ē Ag
3. Higher oxidation states become lower:
8H+ + MnO4
- + 5ē Mn2+ + H2O
Fe3+ + ē Fe2+
Oxidizing agents
Strong Oxidizing Agents
• Metal Oxyacids
– Chromium Reagents (H2CrO4; K2Cr2O7 + H2SO4;
CrO3 + H2SO4)
– Manganese reagents (KMnO4)
– Osmium Tetroxide (OsO4)
• Nitric Acid and Nitrous Acid
– (HNO3, HNO2)
• Halogens
– (F2 > Cl2 > Br2 > I2)
• Forms of Oxygen and Peroxides
– (O3, H2O2)
HCl H+(aq) + Cl-(aq)
Acids as Oxidizing Agents
Hydrogen ion in hydrochloric acid can be an oxidizing agent because it can be reduced to H2.
2H+(aq) + Zn(s) Zn2+
(aq) + H2(g)
In nitric acid solution, the nitrate ion is a more powerful oxidizing agent than the hydrogen ion.
HNO3 H+(aq) + NO3
-(aq)
Cu(s)+4H+(aq)+2NO3
-(aq) Cu2+
(aq)+2NO2(g)+2H2O(l)
1. Active metals forms ions plus electrons:
Zn Zn2+ + 2ē
Na Na+ + ē
2. Nonmetals combine with other nonmetals, such as F and O, which they take from compounds with metals:
C + [O2-] CO + 2ē
3C + Fe2O3 3 CO + 2 Fe
3. Lower oxidation states become higher:
NO + 2 H2O NO3- + 4H+ + 3ē
Reducing agents
Activity Series
F2(g) + 2ē 2 F- (aq)
Fe3+ (aq) + ē Fe2+(aq)
Cu2+(aq) + 2ē Cu(s)
2H+(aq) + 2ē H2(g)
Ni2+(aq) + 2ē Ni(s)
Fe2+(aq) + 2ē Fe(s)
Zn2+(aq) + 2ē Zn(s)
Al3+(aq) + 3ē Al(s)
Li+(aq) + ē Li(s)
Best oxidizing agent
Worst oxidizing agent Best reducing agent
Worst reducing agent
Comparing oxidizing and reducing strength
𝐬𝐭𝐫𝐨𝐧𝐠𝐞𝐫𝐨𝐱𝐢𝐝𝐢𝐳𝐢𝐧𝐠
𝐚𝐠𝐞𝐧𝐭 +
𝐬𝐭𝐫𝐨𝐧𝐠𝐞𝐫𝐫𝐞𝐝𝐮𝐜𝐢𝐧𝐠
𝐚𝐠𝐞𝐧𝐭
𝐬𝐩𝐨𝐧𝐭𝐚𝐧𝐞𝐨𝐮𝐬
𝐰𝐞𝐚𝐤𝐞𝐫𝐫𝐞𝐝𝐮𝐜𝐢𝐧𝐠
𝐚𝐠𝐞𝐧𝐭 +
𝐰𝐞𝐚𝐤𝐞𝐫𝐨𝐱𝐢𝐝𝐢𝐳𝐢𝐧𝐠
𝐚𝐠𝐞𝐧𝐭
Fe2+(aq) + Cu(s) Fe(s) + Cu2+
(aq)
In which direction will reaction go spontaneously? 18
The formation of hydrogen gas in the reaction of a metal with an acid is a special case of a more general phenomenon – one element displacing (pushing out) another element from a compound by means of a redox reaction.
2H+(aq) + Zn(s) Zn2+
(aq) + H2(g)
Fe is a stronger reducing agent than Cu.
Fe + Cu2+ Fe2+ + Cu
Activity Series
F2(g) + 2ē 2 F- (aq)
Cu2+(aq) + 2ē Cu(s)
2H+(aq) + 2ē H2(g)
Fe2+(aq) + 2ē Fe(s)
Zn2+(aq) + 2ē Zn(s)
Li+(aq) + ē Li(s)
Best oxidizing agent
Worst oxidizing agent Best reducing agent
Worst reducing agent
Fe2+(aq) + Cu(s) Fe(s) + Cu2+
(aq)
Fe(s) + Cu2+(aq) Fe2+
(aq) + Cu(s)
Electron transfer
Zn
Cu
CuSO4(aq) ZnSO4(aq)
Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)
Zn(s) Zn2+(aq) + 2ē
Cu2+(aq) + 2ē Cu(s)
Example of oxidation: corrosion
The rusting of iron is an electrochemical process that begins
with the transfer of electrons from iron to oxygen.
Fe → Fe2+ + 2 e−
Redox reaction in the presence of water:
4 Fe2+ + O2 → 4 Fe3+ + 2 O2−
Rust formation:
Fe2+ + 2 H2O ⇌ Fe(OH)2 + 2 H+
Fe3+ + 3 H2O ⇌ Fe(OH)3 + 3 H+
Dehydration:
Fe(OH)2 ⇌ FeO + H2O
Fe(OH)3 ⇌ FeO(OH) + H2O
2FeO(OH) ⇌ Fe2O3 + H2O
Corrosion
Corrosion is even faster in the
presence of salts and acids,
because these materials make
electrically conductive solutions that
make electron transfer easy.
Zinc is added to any metal
that will be submerged in
water and exposed to stray
currents to provide protection
against galvanic corrosion.
Fe2+ + Zn Fe + Zn2+
Corrosion
Gold and platinum are called
noble metals because they
are resistant to losing their
electrons by corrosion.
Other metals may lose their
electrons easily, but are
protected from corrosion by
the oxide coating on their
surface, such as aluminum
oxide.
Ionic equations
To make the essential processes of redox
reactions clearer ionic equations are employed.
• ions in solution are written separately
• only species that change are shown
(not spectator ions)
SnCl2(aq) + Fe2(SO4)3(aq) FeSO4(aq) + SnCl4(aq) Sn2+(aq) + Fe3+(aq) Fe2+(aq) + Sn4+(aq)
Half Equations
Ag+ (aq) + Cu(s) Ag(s) + Cu2+ (aq)
Ag+ (aq) + ē Ag(s)
Ag+ gains electrons, is reduced, and is the oxidizing agent.
Cu(s) Cu2+ (aq) + 2ē
Cu loses electrons, is oxidized, and is the reducing agent.
Balancing Half Equations
Oxidation: Cu(s) Cu2+ (aq) + 2ē
Reduction: Ag+ (aq) + ē Ag(s) ×2
2Ag+ (aq) + 2ē 2Ag(s)
Net: Cu(s) + 2Ag+ (aq) 2Ag(s) + Cu2+
(aq)
What if there is NO complete
electron transfer from one
substance to another?
C(s) + O2(g) CO2(g)
To overcome this problem, the
concept of OXIDATION NUMBER (ON)
was introduced.
Oxidation number (ON)
• Oxidation number – is a number assigned
to atom or an ion to describe its relative
state of oxidation or reduction.
H2O ((H+)2O-2) ON of H in H2O = +1
ON of O in H2O = -2
HCl (H+Cl-) ON of H in HCl = +1
ON of Cl in HCl = -1
Don’t misunderstand!
• Oxidation number has no structural or physical significance. It is not a charge of atom! Oxidation number is relative value of oxidation, which can be equal to the charge.
Recommended name Common (trivial) name HClO chloric (I) acid hypochloric acid FeSO4 iron (II) sulphate Fe2(SO4)3 iron (III) sulphate
used in systematic nomenclature
Advantages of Oxidation Numbers
• useful in balancing equations Oxidation: Cu(s) Cu2+ (aq) + 2ē
Reduction: Ag+ (aq) + ē Ag(s) ×2
2Ag+ (aq) + 2ē 2Ag(s)
Cu(s) + 2Ag2+ (aq) 2Ag(s) + Cu2+ (aq)
Guidelines for Determining Oxidation Numbers
1. The algebraic sum of oxidation numbers
in neutral compound must be zero; in a
polyatomic ion, the sum must be equal to
the ion charge (Al2O3 or MnO4-)
Guidelines for Determining Oxidation Numbers
2. Each atom in a pure element has an
oxidation number of zero (Cu, I2 or S8)
3. Elements of Group 1A – 3A form monoatomic
ions with positive charge and the oxidation number is equal to the group number.
Guidelines for Determining Oxidation Numbers
Example
K: 1s2 2s2 2p6 3s2 3p6 4s1 K+: 1s2 2s2 2p6 3s2 3p6 4s0
K+: [Ar] – octet stability
Guidelines for Determining Oxidation Numbers
4. The oxidation number of H is +1 and
fluorine is always -1 in compounds with
other elements.
Exceptions:
When H forms a binary
compound with a metal, the
metal forms positive ion and H
becomes a hydride ion H-
Guidelines for Determining Oxidation Numbers
5. The oxidation number of O is -2 in
most compounds
Exceptions:
Oxygen can have an oxidation
number -1 in a class of compounds called peroxides
Guidelines for Determining Oxidation Numbers
6. Cl, Br and I are always -1 in
compounds except when combined with
oxygen and fluorine
( Cl has an oxidation number -1 in NaCl, but in the ion
ClO- has an oxidation number +1)
Guidelines for Determining Oxidation Numbers
7. When there is a conflict between two of these
rules or an ambiguity in assigning an oxidation
number, apply the rule with the lower number
and ignore the conflicting rule. 1. The algebraic sum of oxidation numbers in neutral compound must
be zero; in a polyatomic ion, the sum must be equal to the ion
charge
2. Each atom in a pure element has an oxidation number of zero
3. Elements of Group 1A – 3A form monoatomic ions with positive
charge and the oxidation number is equal to the group number
4. The oxidation number of H is +1 and fluorine is always -1 in
compounds with other elements
5. The oxidation number of O is -2 in most compounds
6. Cl, Br and I are always -1 in compounds except when combined with oxygen and fluorine
1. The algebraic sum of oxidation numbers in
neutral compound must be zero; in a
polyatomic ion, the sum must be equal to
the ion charge
2. Each atom in a pure element has an
oxidation number of zero
3. Elements of Group 1A – 3A form
monoatomic ions with positive charge and
the oxidation number is equal to the group
number
4. The oxidation number of H is +1 and
fluorine is always -1 in compounds with
other elements
5. The oxidation number of O is -2 in most
compounds
6. Cl, Br and I are always -1 in compounds
except when combined with oxygen and fluorine
Example
Mg(NO3)2
(+2)+2×(ON(N)+3×(-2))=0
ON(N) = +5
Exercise
• H3PO4 ON (P) =
• Cr2O72- ON (Cr) =
• H2C2O4 ON (C) =
• NaClO3 ON (Cl) =
+5
+6
+3
+5
(+1)×3 + ON(P) + (-2)×4 = 0
2×ON(Cr) + (-2)×7 = -2
(+1)×2 + 2×ON(C) + (-2)×4 = 0
(+1) + ON(Cl) + (-2)×3 = 0
WELL DONE!!!
October's MONTHLY QUIZ is at 8 a.m. Wednesday. Groups F and G in 2/302 D, E, H, I in 5/103 A, B, C in 3/143 Turn up in good time with calculator, pencil and ruler. Periodic tables will be provided.
When the student is ready, the master appears.
Buddhist Proverb