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- Concepts:
- oxidation number (state), redox reaction, reduction, oxidation,
oxidizing agent, reducing agent,
- Skills:
- Be able to define oxidation and reduction in terms of both electron
loss and gain, and change in oxidation number
- Assign oxidation states to atoms in a chemical species
- Deduce whether an element is oxidized or reduced and identify simple
redox reactions using oxidation numbers.
- Identify reduction, oxidation, reducing agent, and oxidizing agent
- Explain the relationship between oxidation states and the names of
compounds
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- 1. Combination/Synthesis
- Combination reactions are ones where the reactants are two different
elements or an element and a compound or two compounds which combine to
produce a single compound.
- Element A + Element B à
Single Compound containing A and B
- A metal and a non-metal can combine to form an ionic compound or two
non-metals can combine to form a covalent compound.
- Two metals cannot combine to form compounds however. (They do mix to
give alloys, which is a mixture and not a compound however.)
- Or
- Element A + Compound BC à
Single Compound containing A, B & C elements
- Or
- Compound AB + Compound CD à Single Compound containing A, B, C & D elements
- Sometimes these reactions are also referred to as a synthesis reaction.
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- 2. Combustion
- Combustion reaction is reaction between a substance (element or
compound) and oxygen.
- This is essentially the reaction that takes place when a substance burns
in air.
- Combustion reaction is the source of most of the energy we use.
- Combustion of natural gas (methane) is the reaction that release heat
energy used in cooking food.
- Combustion of petrol is the reaction that take place in the engine of a
motor vehicle the energy evolved driving it.
- The product(s) of combustion of an element is always the oxide of the
element.
- Mg(s) + O2(g) à MgO(s)
- S(s) + O2(g) à
SO2(g)
- C(s) + O2(g) à
CO2(g)
- N2(g) + O2(g) à NO(g)
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- H2(g) + O2(g) à H2O(l) (the oxide of hydrogen)
- Notice that the reaction between an element and oxygen can be classified
as both combination and combustion.
- When a compound burns the oxides of the elements that make up the
compound are formed.
- When you burned methane in the lab, for instance, the products formed
are carbon dioxide and water.
- The reason those two oxides are formed is because methane is made up of
carbon and hydrogen.
- CH4(g) + O2(g) à CO2 + H2O(l)
- 3. Ion-exchange/double displacement
(replacement)/metathesis/precipitation reaction
- Salt solution 1 + salt solution 2 à precipitate of insoluble salt + solution of other salt
- As you know by now, for the above reaction to occur, one of the
products has to be insoluble.
- 4. A variety of reactions involving acids, namely:
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- i). Acid-metal reactions
- Metal + Acid à Salt +
Hydrogen
- ii). Metal oxide-acid reaction
- Metal oxide + Acid à
salt + water
- iii). Metal hydroxide-acid (neutralization) reaction
- Metal hydroxide + Acid à
salt + water
- iv). Carbonate-, hydrogencarbonate-, sulfite-, hydrogensulfite-
sulfide-, acid reactions
- Metal carbonate (hydrogencarbonate) + Acid à salt + water + carbon dioxide
- Metal sulfite (hydrogensulfite) + Acid à salt + water + sulfur dioxide
- Metal sulfide + Acid à
salt + water + hydrogen sulfide
- 5. Oxide hydration reactions
- When soluble oxides of a metal or a non-metal reacts with water, the
product is either an alkali or an acid.
- Metal oxide + water à
alkali
- Non-metal oxide + water à
Acid
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- Next we will go on to look at Redox reactions.
- Redox reactions are an important type of reaction.
- Photography, respiration, reactions in batteries, wine going off,
bleaching, rusting photosynthesis etc. are all examples of what in
chemistry we call redox reactions.
- Photography involves conversion of silver bromide to silver metal.
- Respiration involves the breakdown of glucose in the presence of oxygen
to get energy.
- Reactions in batteries involves using the different reactivity of metals
to make electrons flow through a circuit which again provides energy.
- Wine going off, bleaching, rusting also involve reaction with oxygen
which are classified as redox reaction.
- But before we can begin looking at this type of reaction, first a
description of concecpt essential to the study of redox reactions:
oxidation state.
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- Valence relates to the number of electrons on the outermost shell of the
atom of an element.
- In case of elements in periods 1-3, valence is equal to their group.
- Since the loss of the valence electrons results in a positively charged
ion, in the case of main group metals, the positive charge on the ion
equals the number of valence electrons.
- For example, valence of Na is 1 and charge is 1+, valence of Ca is 2
and charge is 2+, valence of Al is 3 and charge is 3+.
- With non-metals, valence equals the group number.
- Charge on their ion however is negative and is generally equal to the
number of electrons required to complete the valence shell.
- For example, valence of nitrogen is 5 and charge on ion is 3-, oxygen is 6 and charge is 2-, fluorine 7 and charge 1-, etc.
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- Charge on a monatomic ion therefore results from a gain or loss of
electrons.
- Na ® Na+ + e– Cl
+ e– ® Cl–
- Mg ® Mg2+ + 2e–
O + 2e– ®
O–2
- Since ionic compounds are made up of ions, they consist of positively
and negatively charged particles.
- For example, NaCl consists of positive Na ions (Na+) and
negative Cl ions (Cl- ).
- Polyatomic ions such as SO4–2, PO4–3
are negatively charged because they have more electrons than protons,
while NH4+ is positively charged because it has
one less electron than protons.
- Covalent compound however consist of molecules that are made up of
neutral atoms, they are chargeless.
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- Oxidation state however can be assigned to all atoms (elements) in every
kind of chemical environment, whether it is part of an ionic compound,
polyatomic ion or a molecule.
- In the case of binary ionic compounds, oxidation state of the elements
in the compound is the same as the charge on the ions.
- So, oxidation state of a monatomic charged ion is the same as the
charge on it.
- What about oxidation state of an element in a covalent compound or an
element that is part of a polyatomic ion?
- With both covalent substances and polyatomic ions,
- positive oxidation state is assigned to the element that appears first
in the formula (with some exceptions, notably some binary compounds
and polyatomic ions of hydrogen such as NH3, CH4,
NH4+ etc.), and
- negative oxidation state to the one that appears to the right (the
more electronegative one–the one with stronger attraction for the
shared electrons).
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- But where does one start?
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- 1. The oxidation state of an uncombined atom (or atoms in a diatomic
molecule) is zero.
- Element Oxidation state of atom
- 2. In monatomic ion the oxidation state of the element is the same as
the charge on the ion.
- Monatomic ion Oxidation state of atom
- Eg. Na+ 1+ (or +1)
- Cl- 1- (or -1)
- O2-
2-
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- 3. The oxidation states of elements in a compound is zero.
- Compound Oxidation state of atoms
- Eg. NaCl oxidation # of Na + that of Cl = 0
- Na2O 2 (oxidation # of Na) + that of O = 0
- H2SO3 2 (oxidation # of H) + that of S + 3
(that of O) = 0
- 4. The sum of oxidation states in a complex or polyatomic ion equals the
charge on the ion.
- Oxidation state of atoms
- Eg. SO4–2 oxidation # of S + 4 (that
of O) = –2
- MnO4– oxidation # of Mn + 4(that of
O) = –1
- [CoCl4] –2 oxidation # of Co + 4(that
of Cl) = –2
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- 5. Some elements have fixed oxidation states in all or most of their
compounds and using these oxidation state of other atoms in a compound
(whether ionic or covalent) can be determined. And they are:
- Metals Non-metals
- Group IA metals +1 hydrogen +1 (exception: metal hydrides in
which it is -1)
- Group IIA metals +2 fluorine -1
- other halogens (Cl, Br, I) -1 (exception: with oxygen and
other halogens)
- Aluminum +3 oxygen -2 (exception: peroxides in which it is -1,
superoxides and
with fluorine)
- Peroxides are compounds containing two oxygen atoms that have a single
bond between them.
- All binary peroxides contain O2 in their formula.
- Apart from the elements listed above, the rest can be assumed to have
variable oxidation state.
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- However, knowing these oxidation states, those of other atoms in a
compound or ion can be determined very easily.
- Determine oxidation state of
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- So, in most ionic compounds, often oxidation state of the atom is the
same as charge, but not always.
- In covalent compounds the oxidation state is determined by assigning the
shared electrons to the more electronegative element in the molecule.
- Because of this there is a difference between charge and oxidation
state.
- Whereas a charge on a chemical species has a physical basis (difference
in state of electrons and protons) oxidation state does not necessarily
always have one.
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- For example, let’s look at the charge and/or oxidation state of Fe in Fe3O4.
- Assuming oxidation state (or charge) of O is “2-”
that of Fe comes out to 8/3rd.
- Is that the charge on the Fe atom in the oxide?
- Not really (how can an atom have a fraction of a charge?!).
- Is that the oxidation state of Fe in this compound?
- Yes. (It turns out to be the average oxidation state of Fe in the
compound.)
- Oxidation state therefore is not an intrinsic property of an atom; it is
merely a book keeping system more than anything.
- And as such, it is a great tool in balancing redox reactions--reactions
involving reduction and oxidation.
- (Check oxidation state of phosphorus in P4H2,
sulfur in S3O62- and S4O62-.)
- What about oxidation state of C in CH2O and CH3OH?
- While O and H have oxidation states of “2-”
and “1+” respectively, carbon has oxidation state of “0” in CH2O
and “2-” in CH3OH.
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- 1. Provides an electron “bookeeping” device allowing us to recognize a
redox reaction
- 2. Provides a framework within which chemical similarities may be
recognized and chemical properties correlated (for example the acidic
properties of transition metal ions—see AHL Periodicity: d-block
Elements)
- 3. Useful in balancing redox equations
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- The origination of this type of reaction was from reactions of oxygen
with other elements.
- The term oxidation when first used was simply to denote the (chemical)
addition of oxygen to another substance.
- But, when oxygen is added to another element, and if that element is a
metal, then exchange of electrons take place.
- So, the definition of oxidation was expanded to include reactions
involving the loss of electrons, which is a more inclusive term.
- But, when exchange of electrons are presumed to take place, the
oxidation state changes as well.
- What about those reactions that don’t involve electrons exchanges but
are accompanied by changes in oxidation states?
- There are a whole class of reactions between elements (namely
non-metals) no exchange of electrons are involved but is accompanied by
changes in oxidation state.
- It turned out that defining oxidation and reduction (it’s twin process)
in terms of changes in oxidation state was an even more of an inclusive
term.
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- So, redox therefore involves two half processes: reduction and oxidation
- Reduction on the other hand is the other half in which electrons are
gained, which results in a decrease in the oxidation state.
- Oxidation is part of the reaction where loss of electron(s) occurs,
which results in an increase in oxidation state.
- A mnemonic that might help you to remember redox in terms of electron
exchange is to remember the word OILRIG.
- Oxidation Is Loss (of electrons); Reduction Is Gain.
- In terms of oxidation state change, reduction is reduction in oxidation
state, ie. going from a positive to a less positive state or positive to
negative state, or a negative to a more negative state.
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- Some simple reactions can be easily identified as a redox reaction and
they are:
- Reactions that involve formation of a compounds from its constituent
elements, or the reverse, decomposition into elements.
- 2Na (s) + Cl2 (g) ® 2NaCl(s) C (s) + O2 (g) ® CO2 (g)
- 2HgO(s) ® 2Hg (s) + O2 (g)
- Notice that the oxidation numbers change in going from the elemental
state to the combined state.
- 2. Displacement reactions.
- Metal1 + Metal2 salt solution ® Metal1 salt solution + Metal2
- Zn (s) + CuSO4
(aq) ® ZnSO4 (aq)
+ Cu (s)
- Metal + Acid ® Metal salt
solution + Hydrogen gas
- Mg (s) + 2HCl (aq) ® MgCl2 (aq) + H2 (g)
- Metal + Water ® Metal
hydroxide + Hydrogen gas
- 2Na (s) + 2H2O
(l) ® 2NaOH (l)
+ H2 (g)
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- Metal + Water ® Metal oxide
+ Hydrogen gas
- Mg (s) + H2O (g) ® MgO (s) + H2 (g)
- Metal oxide + hydrogen gas ®
Metal + Water
- CuO (s) + H2 (g) ® Cu (s) + H2O (g)
- Metal oxide + carbon ® Metal
+ Carbon dioxide
- 2Fe2O3 (s) + C (s) ® 4Fe (s) + 3CO2
(g)
- Halogen1 + metal halide ®
metal halide + halogen2
- Cl2 (aq) + KBr (aq) ® KCl (aq) + B2 (l)
- I- (aq)
+ Cl2 (aq) ® I2 (aq) + 2Cl- (aq)
- The above set of reactions are also referred to as displacement reaction
for the simple fact that one of the substances takes the place of the
other.
- Let´s look into the oxidation state change for the first reaction.
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- Zn + CuSO4 ® ZnSO4 +
Cu
- Here, Zn goes from a oxidation state of ‘0’ to ‘2+’ and Cu from ‘2+’ to
‘0’
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- There are more complex redox reactions than the simple combination and
displacement reactions we considered thus far.
- Complex redox reaction involve what are referred to as common oxidixing
and reducing agents.
- As you know, oxidizing agents therefore are substances that readily
accept electrons and reducing agents readily donate electrons
- Therefore, oxidizing agents contain elements that are in their highest
oxidation state.
- Two common oxidizing agents in the laboratory are MnO4- (KMnO4-) and Cr2O72- (K2Cr2O7).
- Some other oxidizing agents are: O2, Cl2, F2,
SO3 (SO42- in solution), Fe3+, ClO3- (KClO3),
and NO3- (KNO3).
- Reducing agents on the other hand contain elements in their lower
oxidation state.
- Some common reducing agents are H2, Na, C, CO, SO2
(SO32- in solution), Fe2+, NO2- and I-.
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- Note that oxidizing agents have more oxygen than reducing agents or have
bigger positive charge.
- Those with oxygen are able to liberate oxygen which removes or accepts
electrons from the reducing agents.
- Combining an oxidizing agent and a reducing agent under appropriate
conditions will result in a redox reaction.
- 2MnO4-
(aq) + 6H+ (aq) + 5NO2- ® 2Mn2+ (aq)
+ 5NO3- (aq)+ 3H2O(aq)
- Pink Colorless
- MnO4- + 8H+ + 5Fe2+ ®
Mn2+ + 5Fe3+ + 4H2O
- Cr2O72- (aq) + 3SO32-(aq) + 8H+
(aq) ® 2Cr3+
(aq) + 3 SO42- (aq) + 4H2O (l)
- Orange Green
- These redox reactions are a lot more complex because they involve three
reagents; the third reagent being an acid (H+).
- In other words, these redox reactions require an acidic medium to
occur.
- As a matter of fact, they are so complex that they require a completely
different methos of balancing them! (See Core Reactions - Balancing
Redox rxn in acid.)
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- You will need to memorize the color change accompanying the redox
reactions involving manganate(VII) and dichromate(VI) ions.
- Chlorine is used as a household bleach because of it strong oxidizing
power!
- The active ingredient in bleach however is aqueous chlorate(I) (ClO-(aq)).
- A reaction similar in principle to the one in the previous slide can be
conducted to analyze chlorine bleach.
- 2H+ (aq) + ClO- (aq) + 2I - (aq) ® I2 (aq) + Cl- (aq)
+ H2O (l)
- The analysis is completed by analyzing for the iodine produced by
titrating the reaction mixture against a known concentration of sodium
thiosulfate.
- I2 (aq) + 2S2O32-
(aq) ® 2I- (aq) + S4O62- (aq)
- Similarly, analysis of iron content of a iron supplement tablet (as a
compound of iron(II)) can be conducted also by performing a redox
titration of the acidified tablet soution against potassium
manganate(VII) solution.
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- Oxidizing Agent Reducing Agent
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- Oxidizing Agent Reducing Agent
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- For example, Iodine undergoes both oxidation and reduction in the
following reaction:
- I2 (s) + OH – (aq) ® I – (aq) + IO – (aq)
+ H2O (l)
- In changing to I –, iodine undergoes reduction; oxidation
state goes from ‘0’ to ‘1–’.
- In changing to IO –, it undergoes oxidation; oxidation state
goes from ‘0’ to ‘1+’.
- Another example:
- S2O32– (aq) + 2H+(aq)
à SO2(g) + H2O(l)
+ S(s)
- Yet another one is the decomposition of hydrogen peroxide (into water
and oxygen) in the presence of a catalyst (MnO2). Can you see
why?
- As a matter of fact, decomposition reaction where one of the products
is an element can be classified as disproportionation reaction.
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- Some binary compounds decompose to component elements (redox).
- Nitrate, if it decomposes, either decompose to nitrite and oxygen or to
the oxide, nitrogen dioxide and oxygen (redox).
- Carbonate, if it decomposes, to oxide and carbon dioxide (not redox).
- Sulfate, if it decomposes, either to sulfite and oxygen or to the oxide,
sulfur dioxide and oxygen (redox).
- Chlorates, if they decompose, either to another chlorate and oxygen, or
to chloride and oxygen (redox).
- Hydroxide, if it decomposes, to oxide and water (not redox)
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- The redox—made up of reduction and oxidation—can be represented by
equations like so:
- Oxidation: Zn (s) ® Zn2+(aq) + 2e–
- Reduction: Cu2+ (aq) + 2e– ® Cu (s)
- The equations above showing each half of the redox process involved are
referred to as half-reactions (equations).
- The first one shows oxidation therefore it’s referred to as an oxidation
half-reaction (equation).
- Electrons are lost by Zn (electrons are shown as a product)
- The second one shows reduction and so is referred to as a reduction
half-reaction (equation).
- Electrons are gained by Cu2+ (electrons are shown as a
reactant)
- This electron transfer that take place in a redox reaction can also be
made to flow through an external circuit producing electrical energy,
which is the basis for batteries.
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- For the disproportionation reaction of iodine which you encountered in
one of the previous slides:
- I2 (s) + OH – (aq) ® I – (aq) + OI – (aq)
+ H2O (l)
- Oxidation half-equation: I2 (s) ® 2IO – (aq) + 2e –
- 2 electrons are needed on the right hand side to account for the change
in the oxidation state of iodine from ‘0’ to ‘1+’.
- Reduction half-equation: I2 (s) + 2e – ® 2I – (aq)
- 2 electrons are needed on the left hand side to account for the change
in the oxidation state of iodine from ‘0’ to ‘1–’.
- Similarly for the following reaction:
- Cr2O72- (aq) + Cl- (aq) ® Cr3+ (aq)
+ Cl2 (g)
- Oxidation half-equation: 2Cl- ®
Cl2 + 2e-
- Reduction half-equation: Cr2O72- + 6e- ® 2Cr3+
- NB: In the half equations, the change in the oxidation state of the
oxidized and reduced species must be balanced by adding electrons.
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- Oxidation state does not necessarily have a physical basis.
- Following a set of rules, oxidation state for any atom in a substance
can be determined.
- Oxidation is increase in oxidation state while reduction is decrease
(reduction) in oxidation state.
- That which undergoes oxidation is the reducing agent while that which
undergoes reduction is the oxidizing agent.
- Or that which oxidizes another substance is the oxidizing agent, that
which reduces another substance, the reducing agent.
- Redox reactions can be represented using half-equations.
- Oxidation number allows us to balance redox reactions.
- Notice that some reaction that we considered before (such as metal-acid
reaction) can also be classified as a redox reaction; in other words,
these classifications aren´t hard a fast.
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- 1. Combination/Synthesis
- 2. Combustion
- 3. Ion-exchange/double displacement (replacement)/metathesis
- 4. Acid reactions
- i) Metal-acid (displacement reaction)
- ii) Metal oxide-acid
- iii) Metal hydroxide-acid—neutralization reaction (acid base reaction)
- iv) Carbonate-, hydrogencarbonate-, sulfite-, hydrogensulfite-,
sulfide-acid
- 5. Acid and base anhydride hydration reaction
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- 6. Redox reactions.
- i) Combination reactions
- ii) Displacement reactions
- iii) Acid-medium redox reactions involving common oxidizing and
reducing agents
- iv) Decomposition reactions where one of the products is an element
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- 1. Relate these concepts to each other by writing a sentence or two.
- (a) oxidizing agent, (b) oxidized substance, (c) reducing agent, (d)
reduced substance.
- 2. For each state whether the change is an oxidation or a reduction.
- (a) MnO4- becomes MnO42-
(b) K becomes K+
- (c) N2 becomes NH3 (d) NH3 becomes
N2O
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- (e) P4O10 becomes P4O6
(f) SO42- becomes SO32-
- (g) HClO4 becomes HCl and H2O (h) O2
becomes O2-
- (I) Cr2O72- becomes Cr3+
- 3. For the following changes, write half reactions (equations):
- (a) MnO4- to MnO42-
- (b) K to K+
- (c) N2 to NH3
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- (d) NH3 to N2O
- (e) P4O10 to P4O6
- (f) SO42- to SO32-
- (g) HClO4 to HCl and H2O
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- (h) O2 to O2-
- (i) Cr2O72- to Cr3+
- 3. Find the oxidation numbers of the underlined elements in these
formulas.
- HClO, (b) KClO3, (c) MnO2, (d) PbO2,
- (e) PbSO4, (f) K2SO4, (g) NH4+,
(h) Na2O2,
- (i) FeO, (j) Fe2O3, (k) NaIO4, (l)
Fe3O4,
- (m) Cr2O72-, (n) MnO42-,
(o) NO3-, (p) ClO3-.
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- 1. N03/30. In which reaction does chromium undergo a change in oxidation
number?
- A. Cr2O3 +3H2SO4 ® Cr2 (SO4)3
+3H2O
- B. Cr2(SO4)3 + 6NaOH ® 2Cr(OH)3 + 3Na2SO4
- C. K2Cr2O7 + 4H2SO4
+ 6HCl ® Cr2(SO4)3
+ K2SO4 + 7H2O +3Cl2
- D. 2K2CrO4 + H2SO4 ® K2Cr2O7
+ K2SO4 + H2O
- 2. N02/31. In the reaction
- 3Br2 + 6CO + 3H2O ® 5Br- + BrO3- + 6HCO
- A. Br2 is only oxidised.
- B. Br2 is only reduced.
- C. Br2 is neither oxidised nor reduced.
- D. Br2 is both oxidised and reduced.
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- 3. N98/26. What is the oxidizing
agent in the following reaction?
- Cl2(aq) + 2 Br- (aq) à 2 Cl- (aq) +
Br2(aq)
- A. Cl2(aq) B. Br-(aq) C. Cl- (aq) D.
Br2(aq)
- 4. 92/34. In the following compounds, sulfur has the highest oxidation
number in:
- A. S8 B. SO2 C. SO3 D. H2S
- 5. 90/47. In the reaction MnO2 + 4 HCl ® Cl2 + 2 H2O + MnCl2
- A. HCl is the oxidizing agent.
B. Cl2 is the oxidization product.
- C. H2O is the reduction product. D. MnO2 is the
reducing agent.
- 6. 90/33. Which one of the
following species could NOT behave as a reducing agent?
- A. H2 B. I- C. Na D. NO3-
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Notes
