-
In
the reactions given below, identify the species undergoing oxidation and
reduction :
(i) \[{{H}_{2}}S(g)+C{{l}_{2}}(g)\to 2HCl(g)+S(s)\]
(ii)\[3F{{e}_{3}}{{O}_{4}}(s)+8Al(s)\to
9Fe(s)+4A{{l}_{2}}{{O}_{3}}(s)\]
(iii) \[2Na(s)+{{H}_{2}}(g)\to
2NaH(s)\]
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Justify
that the reaction :
\[2Na(s)+{{H}_{2}}(g)\to 2NaH(s)\]is
a redox change.
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Using
stock notation, represent the following compounds
\[HAuC{{l}_{4}},T{{l}_{2}}O,F{{e}_{2}}{{O}_{3}},CuI,CuO,MnO\]
and
\[Mn{{O}_{2}}\].
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Justify
that the reaction :
\[2C{{u}_{2}}O(s)+C{{u}_{2}}S(s)\to 6Cu(s)+S{{O}_{2}}(g)\]
is a redox reaction. Identify the
species oxidised/reduced, which acts as an oxidant and which acts as a
reductant.
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Which
of the following species, do not show the disproportionation reaction and why?
\[Cl{{O}^{-}},ClO_{2}^{-},ClO_{3}^{-}\]and\[ClO_{4}^{-}\]
Also write reaction for each
species that disproportionates.
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Suggest
a scheme of classification of the following redox reactions:
(a) \[{{N}_{2}}(g)+{{O}_{2}}(g)\to 2NO(g)\]
(b) \[2Pb{{(N{{O}_{3}})}_{2}}(s)\to
2PbO(s)+2N{{O}_{2}}(g)+\frac{1}{2}{{O}_{2}}(g)\]
(c) \[NaH(s)+{{H}_{2}}O(l)\to NaOH(aq)+{{H}_{2}}(g)\]
(d)\[2N{{O}_{2}}+2O{{H}^{-}}(aq)\to
NO_{2}^{-}(aq)+NO_{3}^{-}(aq)+{{H}_{2}}O(l)\]
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Why
do the following reactions proceed differently?
\[P{{b}_{3}}{{O}_{4}}+8HCl\to 3PbCl+C{{l}_{2}}+4{{H}_{2}}O\]and
\[P{{b}_{3}}{{O}_{4}}+4HN{{O}_{3}}\to
2Pb{{(N{{O}_{3}})}_{2}}+Pb{{O}_{2}}+2{{H}_{2}}O.\]
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Write the net ionic equation for
the reaction of potassium dichromate (VI) \[{{K}_{2}}C{{r}_{2}}{{O}_{7}}\]with
sodium sulphite\[N{{a}_{2}}S{{O}_{3}}\]in acid solution to give chromium (III)
ion and sulphate ion.
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Permanganate ion reacts with
bromide ion in basic medium to give manganese dioxide and bromate ion. Write
the balanced ionic equation for the reaction.
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Permanganate (VII) ion, \[MnO_{4}^{-}\]
in basic solution oxidises iodide ion, \[{{I}^{-}}\] to produce molecular
iodine \[({{I}_{2}})\] and manganese (IV) oxide\[(Mn{{O}_{2}})\]. Write the
balanced ionic equation to represent this redox reaction.
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1, 2 Assign oxidation number to the underlined elements in
each of the following species:
(a) \[\text{K}{{\underset{\text{-}}{\mathop{\text{I}}}\,}_{\text{3}}}\] (b)
\[Ca{{\underset{-}{\mathop{O}}\,}_{2}}\]
(c)\[Na{{H}_{2}}\underset{-}{\mathop{P}}\,{{O}_{4}}\] (d)\[{{H}_{4}}{{\underset{-}{\mathop{P}}\,}_{2}}{{O}_{7}}\]
(e) \[Na\underset{-}{\mathop{B}}\,{{H}_{4}}\] (f) \[{{H}_{2}}{{\underset{-}{\mathop{S}}\,}_{2}}{{O}_{7}}\]
(g)\[KAl{{(\underset{-}{\mathop{S}}\,{{O}_{4}})}_{2}}l2{{H}_{2}}O\]
(h)\[{{H}_{2}}{{\underset{-}{\mathop{S}}\,}_{4}}{{O}_{6}}\]
(i) \[\underset{-}{\mathop{C}}\,{{H}_{3}}\underset{-}{\mathop{C}}\,{{H}_{2}}OH\] (j)
\[\underset{-}{\mathop{C}}\,{{H}_{3}}\underset{-}{\mathop{C}}\,OOH\]
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1, 2
Assign oxidation number to the underlined elements in each of the following
species:
(a) \[\text{K}{{\underset{\text{-}}{\mathop{\text{I}}}\,}_{\text{3}}}\] (b)
\[Ca{{\underset{-}{\mathop{O}}\,}_{2}}\]
(c)\[Na{{H}_{2}}\underset{-}{\mathop{P}}\,{{O}_{4}}\] (d)\[{{H}_{4}}{{\underset{-}{\mathop{P}}\,}_{2}}{{O}_{7}}\]
(e) \[Na\underset{-}{\mathop{B}}\,{{H}_{4}}\] (f) \[{{H}_{2}}{{\underset{-}{\mathop{S}}\,}_{2}}{{O}_{7}}\]
(g)\[KAl{{(\underset{-}{\mathop{S}}\,{{O}_{4}})}_{2}}l2{{H}_{2}}O\]
(h)\[{{H}_{2}}{{\underset{-}{\mathop{S}}\,}_{4}}{{O}_{6}}\]
(i) \[\underset{-}{\mathop{C}}\,{{H}_{3}}\underset{-}{\mathop{C}}\,{{H}_{2}}OH\] (j)
\[\underset{-}{\mathop{C}}\,{{H}_{3}}\underset{-}{\mathop{C}}\,OOH\]
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Justify
that following reactions are redox reactions:
(a) \[CuO(s)+{{H}_{2}}(g)\to Cu(s)+{{H}_{2}}O(g)\]
(b) \[F{{e}_{2}}{{O}_{3}}(s)+3CO(g)\to
2Fe(s)+3C{{O}_{2}}(g)\]
(c) \[4BC{{l}_{3}}(g)+3LiAl{{H}_{4}}(s)\to
2{{B}_{2}}{{H}_{6}}(g)+LiCl(s)+3AlC{{l}_{3}}(s)\](d) \[2K(s)+{{F}_{2}}(g)\to
2{{K}^{+}}{{F}^{-}}(s)\]
(e) \[4N{{H}_{3}}(g)+5{{O}_{2}}(g)\to
4NO(g)+6{{H}_{2}}O(g)\]
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Fluorine reacts with ice and results in the change:
\[{{H}_{2}}O(s)+{{F}_{2}}(g)\to HF(g)+HOF(g)\]
Justify that this reaction is a redox reaction.
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Calculate the oxidation number of sulphur, chromium and
nitrogen in \[{{H}_{2}}S{{O}_{5}},C{{r}_{2}}O_{7}^{2-}\] and\[NO_{3}^{-}\].
Suggest structure of these compounds. Count for the fallacy.
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Write down the formulae of following compounds:
(a) Mercury (II) chloride
(b) Nickel (II) sulphate
(c) Tin (IV)oxide
(d) Thallium (I) sulphate
(e) Iron (III) sulphate
(f) Chromium (III) oxide
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Suggest a list of substances where
carbon can exist in oxidation states from - 4 to + 4 and nitrogen from -3 to+5.
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While sulphur dioxide and hydrogen
peroxide can act as oxidising as well as reducing agents in their reactions,
ozone and nitric acid act only as oxidants why?
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Justify that following reactions are redox reactions:
(a) \[CuO(s)+{{H}_{2}}(g)\to Cu(s)+{{H}_{2}}O(g)\]
(b) \[F{{e}_{2}}{{O}_{3}}(s)+3CO(g)\to
2Fe(s)+3C{{O}_{2}}(g)\]
(c) \[4BC{{l}_{3}}(g)+3LiAl{{H}_{4}}(s)\to
2{{B}_{2}}{{H}_{6}}(g)+LiCl(s)+3AlC{{l}_{3}}(s)\](d) \[2K(s)+{{F}_{2}}(g)\to
2{{K}^{+}}{{F}^{-}}(s)\]
(e) \[4N{{H}_{3}}(g)+5{{O}_{2}}(g)\to
4NO(g)+6{{H}_{2}}O(g)\]
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The compound \[Ag{{F}_{2}}\] is
unstable compound. However, if formed, the compound acts as a very strong oxidizing
agent?
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Whenever a reaction between an
oxidising agent and a reducing agent is carried out, a compound of lower oxidation
state is formed if the reducing agent is in excess and a compound of higher
oxidation state is formed if oxidising agent is in excess. Justify this statement
giving three illustrations.
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Whenever a reaction between an
oxidising agent and a reducing agent is carried out, a compound of lower oxidation
state is formed if the reducing agent is in excess and a compound of higher
oxidation state is formed if oxidising agent is in excess. Justify this statement
giving three illustrations.
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Identify the substances oxidised, reduced, oxidizing agents
and reducing agents for each of the following reactions.
(a) \[2AgBr(s)+{{C}_{6}}{{H}_{6}}{{O}_{2}}(aq)\to
2Ag(s)+2HBr(aq)\]
\[+{{C}_{6}}{{H}_{4}}{{O}_{2}}(aq)\]
(b) \[HCHO(l)+2{{[Ag{{(N{{H}_{3}})}_{2}}]}^{+}}(aq)+3O{{H}^{-}}(aq)\]
\[\to
2Ag(s)+HCO{{O}^{-}}(aq)+4N{{H}_{3}}(aq)+2{{H}_{2}}O(l)\](c) \[HCHO(l)2C{{u}^{2+}}(aq)+5O{{H}^{-}}\to
C{{u}_{2}}O(s)\] \[+HCO{{O}^{-}}(aq)+3{{H}_{2}}O(l)\]
(d) \[{{N}_{2}}{{H}_{4}}(l)2{{H}_{2}}O(l)\to
{{N}_{2}}(g)+4{{H}_{2}}O(l)\]
\[(e)Pb(s)Pb{{O}_{2}}(s)+2{{H}_{2}}S{{O}_{4}}(aq)\to
2PbS{{O}_{4}}(s)+2{{H}_{2}}O(l)\]
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Consider the reactions:
\[2{{S}_{2}}O_{3}^{2-}(aq)+{{I}_{2}}(s)\to
{{S}_{4}}O_{6}^{2-}(aq)+2{{I}^{-}}(aq)\]
\[{{S}_{2}}O_{3}^{2-}(aq)+2B{{r}_{2}}(l)\to
2SO_{4}^{2-}(aq)\]
\[+4B{{r}^{-}}(aq)+10{{H}^{+}}(aq)\]
Why does same reductant, thiosulphate react differently with
iodine and bromine?
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Justify giving reactions that
among halogens, fluorine is the best oxidant and among hydrohalic compounds, hydroiodic
acid is the best reductant?
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Why does the following reaction occur?
\[XeO_{6}^{4-}+2{{F}^{-}}(aq)+6{{H}^{+}}(aq)\to
Xe{{O}_{3}}(g)+{{F}_{2}}(g)+3{{H}_{2}}O(l)\]
What conclusion about the compound \[N{{a}_{4}}Xe{{O}_{6}}\]
(of which \[XeO_{6}^{4-}\] is a part) can be drawn from the reaction?
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Consider the reactions:
(a) \[{{H}_{3}}P{{O}_{2}}(aq)+4AgN{{O}_{3}}(aq)+2{{H}_{2}}O(l)\]
\[\to
\,{{H}_{3}}P{{O}_{4}}(aq)+4Ag(s)+4HN{{O}_{3}}(aq)\]
(b) \[{{H}_{3}}P{{O}_{2}}(aq)+2CuS{{O}_{4}}(aq)+2{{H}_{2}}O(l)\]
\[\to
{{H}_{3}}P{{O}_{4}}(aq)+2Cu(s)+{{H}_{2}}S{{O}_{4}}(aq)\]
(c) \[{{C}_{6}}{{H}_{5}}CHO(l)+2{{[Ag{{(N{{H}_{3}})}_{2}}]}^{+}}(aq)+3O{{H}^{-}}\to
\]\[{{C}_{6}}{{H}_{5}}CO{{O}^{-}}(aq)+2Ag(s)+4N{{H}_{3}}(aq)+2{{H}_{2}}O(l)\](d)
\[{{C}_{6}}{{H}_{5}}CHO(l)+2C{{u}^{2+}}(aq)+\underset{No\,change\,observed.}{\mathop{5O{{H}^{-}}(aq)\to
}}\,\]
What inference do you draw about the behaviour of and \[C{{u}^{2+}}\]from
these reactions?
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Balance the following redox reactions by ion electron method:
(a)\[MnO_{4}^{-}(aq)+{{I}^{-}}(aq)\to \underset{\left(
in\text{ }basic\text{ }medium
\right)}{\mathop{Mn{{O}_{2}}(s)+{{I}_{2}}(s)}}\,\]
(b) \[MnO_{4}^{-}(aq)+S{{O}_{2}}(g)\to \underset{\left(
in\text{ acidic }medium \right)}{\mathop{M{{n}^{2+}}(aq)+HSO_{4}^{-}(aq)}}\,\]
(c) \[{{H}_{2}}{{O}_{2}}(aq)+F{{e}^{2+}}(aq)\to
F{{e}^{3+}}\underset{\left( in\text{ }acidic\text{ }medium
\right)}{\mathop{(aq)+{{H}_{2}}O(l)}}\,\]
(d)\[C{{r}_{2}}O_{7}^{2-}+S{{O}_{2}}(g)\to
C{{r}^{3+}}\underset{\left( in\text{ }acidic\text{ }medium
\right)}{\mathop{(aq)+SO_{4}^{2-}(aq)}}\,\]
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Balance the following equations in basic medium by ion-electron
method and oxidation number method :
(a) \[{{P}_{4}}(s)+O{{H}^{-}}(aq)\to
P{{H}_{3}}(g)HPO_{2}^{-}(aq)\]
(b) \[{{N}_{2}}{{H}_{4}}(l)+ClO_{3}^{-}(aq)\to
P{{H}_{3}}(g)+HPO_{2}^{-}(aq)\]
(c) \[{{N}_{2}}{{H}_{4}}(l)+ClO_{3}^{-}(aq)\to
P{{H}_{3}}(g)+HPO_{2}^{-}(aq)\]
\[+{{O}_{2}}(g)+{{H}^{+}}\]
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What sorts of information can you
draw from the following reaction?
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The \[M{{n}^{3+}}\] ion is unstable in a solution and
undergoes disproportionation to give \[M{{n}^{2+}}\]\[Mn{{O}_{2}}\] and \[{{\text{H}}^{+}}\]ion.
Write a balanced ionic equation
for the reaction.
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Consider the elements:
Cs, Ne, I and F
(a) Identify the element that exhibits only negative oxidation
state.
(b) Identify the element that exhibits only positive oxidation
state.
(c) Identify the element that exhibits both positive and negative
oxidation states.
(d) Identify the element that exhibits neither the negative
nor does the positive oxidation state.
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Chlorine is used to purify
drinking water. Excess of chlorine is harmful. The excess of chlorine is
removed by treating with sulphur dioxide. Present a balanced equation for this
redox change taking place in water.
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Refer to the periodic table given in your book and now
answer the following questions:
(a) Select the possible non-metals that can show disproportionation.
(b) Select three metals that can
show disproportionation reaction.
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In
Ostwald's process for the manufacture of nitric acid, the first step involves
the oxidation of ammonia by oxygen to give nitric oxide gas and steam. What is
the maximum weight of nitric oxide that can be
obtained starting only with 10 g
ammonia and 20 g oxygen?
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Using standard electrode; predict if the reaction between
the following is feasible:
(i) \[F{{e}^{3+}}(aq)and\,{{I}^{-}}(aq)\]
(ii) \[A{{g}^{+}}(aq)and\,Cu\]
(iii) \[F{{e}^{3+}}and\,B{{r}^{-}}\]
(iv) \[Ag\,and\,F{{e}^{3+}}(aq)\]
(v) \[B{{r}_{2}}(aq)and\,F{{e}^{2+}}(aq)\]
(Given, \[E_{{{I}_{2}}/{{I}^{-}}}^{{}^\circ
}=0.541V,\,E_{C{{u}^{2+}}/Cu}^{{}^\circ }=+0.34V,\]
\[E_{B{{r}_{2}}/B{{r}^{-}}}^{{}^\circ
}=1.09V,E_{A{{g}^{+}}/Ag}^{{}^\circ }=+0.80V,\]
\[E_{F{{e}^{3+}}/F{{e}^{2+}}}^{{}^\circ }=+0.77V)\]
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Predict the products of electrolysis in each of the following:
(i) An aqueous solution of\[AgN{{O}_{3}}\] with silver
electrodes.
(ii) An aqueous solution of \[\text{AgN}{{\text{O}}_{\text{3}}}\]with
platinum electrodes.
(iii) A dilute solution of sulphuric acid using platinum electrodes.
(iv) An aqueous solution of \[CuC{{l}_{2}}\]
with platinum electrodes.
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Arrange
the following metals in the order in which they displace each other from
solution of their salts
Al, Cu, Fe, Mg, Zn
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Given
the standard electrode potentials:
\[{{K}^{+}}/K=-2.93V\,\,\,\,\,\,;\,\,\,\,\,\,\,\,\,\,\,\,\,\,A{{g}^{+}}/Ag=0.80\]
Arrange these in increasing
reducing power.
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Depict the galvanic cell in which
\[Zn(s)+2A{{g}^{+}}\to Z{{n}^{2+}}(aq)+2Ag(s)\]takes place.
Further show:
(i) Which electrode is negatively charged?
(ii) Carriers of current in the cell.
(iii) Individual reaction at each electrode.
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question_answer41)
In the following questions only one option
is correct.
Which of the following is not an
example of redox reaction?
(a)\[CuO+{{H}_{2}}\to
Cu+{{H}_{2}}O\] (b) \[F{{e}_{2}}{{O}_{3}}+3CO\to 2Fe+3C{{O}_{2}}\]
(c)\[2K+{{F}_{2}}\to 2KF\] (d)
\[BaC{{l}_{2}}+{{H}_{2}}S{{O}_{4}}\to BaS{{O}_{ & 4}}+2HCl\]
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question_answer42)
The more positive the value of\[{{E}^{O-}}\],
the greater is the tendency of the species to get reduced. Using the standard electrode
potential of redox couples given below find out which of the following is the
strongest oxidising agent?
\[{{E}^{O-}}Values\,:\,\,\,\,F{{e}^{3+}}/{{F}^{2+}}=+0.77;{{I}_{2}}(s)/{{I}^{-}}=+0.54;\]\[C{{u}^{2+}}/Cu=+0.34;A{{g}^{+}}/Ag=+0.80V\]
(a)\[F{{e}^{3+}}\] (b)
\[{{I}_{2}}(s)\] (c) \[C{{u}^{2+}}\] (d)\[A{{g}^{+}}\]
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question_answer43)
\[{{E}^{O-}}\]values of some
redox couples are given below. On the basis of these values choose the correct
option.
\[{{E}^{O-}}\]values:
\[B{{r}_{2}}/B{{r}^{-}}=+1.90;A{{g}^{+}}/Ag(s)=+0.80;\]
\[C{{u}^{2+}}/Cu(s)=+\,0.34;{{I}_{2}}(s)/{{I}^{-}}=+0.54\]
(a) Cu will reduce \[B{{r}^{-}}\] (b) Cu will reduce\[\text{Ag}\]
(c) Cu will reduce \[{{I}^{-}}\] (d)
Cu will reduce \[B{{r}_{2}}\]
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question_answer44)
Using the standard electrode potential,
find out the pair between which redox reaction is not feasible.
\[{{E}^{O-}}values:\,\,\,F{{e}^{3+}}/F{{e}^{2+}}=+0.77;{{I}_{2}}/{{I}^{-}}=+0.54;\] \[C{{u}^{2+}}/Cu=+0.34;Ag/Ag=+0.80V\]
(a) \[F{{e}^{3+}}\]and\[{{I}^{-}}\] (b)
\[A{{g}^{+}}\] and Cu (c) \[F{{e}^{3+}}\]and\[\text{Cu}\] (d)
\[\text{Ag}\]and \[F{{e}^{3+}}\]
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question_answer45)
Thiosulphate reacts differently
with iodine and bromine in the reactions given below :
\[2{{S}_{2}}O_{3}^{2-}+{{I}_{2}}\to
{{S}_{4}}O_{6}^{2-}+2{{I}^{-}}\]
\[{{S}_{2}}O_{3}^{2-}+2B{{r}_{2}}+5{{H}_{2}}O\to
2SO_{4}^{2-}+2B{{r}^{-}}+10{{H}^{+}}\]
Which of the following statements justifies the
above dual behaviour of thiosulphate?
(a) Bromine is a stronger
oxidant than iodine.
(b) Bromine is a weaker oxidant
than iodine.
(c) Thiosulphate undergoes
oxidation by bromine and reduction by iodine in these reactions.
(d) Bromine undergoes oxidation
and iodine undergoes reduction in these reactions.
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question_answer46)
The oxidation number of an
element in a compound is evaluated on the basis of certain rules. Which of the following
rules is not correct in this respect?
(a) The oxidation number of
hydrogen is always +1.
(b) The algebraic sum of all the
oxidation numbers in a compound is zero.
(c) An element in the free or
the uncombined state bears oxidation number zero.
(d) In all its compounds, the
oxidation number of fluorine is ?1.
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question_answer47)
In which of the following
compounds, an element exhibits two different oxidation states?
(a) \[N{{H}_{2}}OH\] (b)
\[N{{H}_{4}}N{{O}_{3}}\] (c) \[{{N}_{2}}{{H}_{4}}\] (d)
\[{{N}_{3}}H\]
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question_answer48)
Which of the following
arrangements represent increasing oxidation number of the central atom?
(a) \[CrO_{2}^{-},ClO_{3}^{-},CrO_{4}^{2-},MnO_{4}^{-}\] (b) \[ClO_{3}^{-},CrO_{4}^{2-},MnO_{4}^{-},CrO_{2}^{-}\]
(c) \[CrO_{2}^{-},ClO_{3}^{-},MnO_{4}^{-},CrO_{4}^{2-}\] (d)
\[CrO_{4}^{2-},MnO_{4}^{-},CrO_{2}^{-},ClO_{3}^{-}\]
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question_answer49)
The largest oxidation number
exhibited by an element depends on its outer electronic configuration. With
which of the following outer electronic configurations the element will exhibit
the largest oxidation number?
(a) \[3{{d}^{1}}4{{s}^{2}}\] (b)\[3{{d}^{3}}4{{s}^{2}}\] (c)\[3{{d}^{5}}4{{s}^{1}}\] (d)\[3{{d}^{5}}4{{s}^{2}}\]
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question_answer50)
Identify disproportionation
reaction :
(a) \[C{{H}_{4}}+2{{O}_{2}}\to
C{{O}_{2}}+2{{H}_{2}}O\] (b) \[C{{H}_{4}}+4C{{l}_{2}}\to
CC{{l}_{4}}+4HCl\]
(c) \[2{{F}_{2}}+2O{{H}^{-}}\to 2{{F}^{-}}+O{{F}_{2}}+{{H}_{2}}O\] (d)
\[2N{{O}_{2}}+2O{{H}^{-}}\to NO_{2}^{-}+NO_{3}^{-}+{{H}_{2}}O\]
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question_answer51)
Which of the following elements
does not show disproportionation tendency?
(a)\[\text{Cl}\] (b)
\[\text{Br}\] (c)\[\text{F}~\] (d)
\[\text{I}\]
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question_answer52)
In the following questions
two or more options may be correct.
Which of the following
statement(s) is/are not true about the following decomposition reaction?
\[2KCI{{O}_{3}}\to
2KCl+3{{O}_{2}}\]
(a) Potassium is undergoing
oxidation.
(b) Chlorine is undergoing
oxidation.
(c) Oxygen is reduced.
(d) None of the species are
undergoing oxidation or reduction.
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question_answer53)
Identify the correct statement(s) in relation
to the following reaction :
\[Zn+2HCl\to
ZnC{{l}_{2}}+{{H}_{2}}\]
(a) Zinc is
acting as an oxidant. (b) Chlorine is acting
as a reductant.
(c) Hydrogen ion is acting as an oxidant. (d)
Zinc is acting as a reductant.
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question_answer54)
The exhibition of various
oxidation states by an element is also related to the outer orbital electronic
configuration of its atom. Atom(s) having which of the following outer most electronic
configurations will exhibit more than one oxidation state in its compounds.
(a)\[\text{3}{{\text{s}}^{\text{1}}}\] (b)\[\text{3}{{\text{d}}^{\text{l}}}\text{4}{{\text{s}}^{\text{2}}}\] (c)\[\text{3}{{\text{d}}^{\text{2}}}\text{4}{{\text{s}}^{\text{2}}}\] (d)\[\text{3}{{\text{s}}^{\text{2}}}\text{3}{{\text{p}}^{\text{3}}}\]
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question_answer55)
Identify the correct statements
with reference to the given reaction:
\[{{P}_{4}}+3O{{H}^{-}}+3{{H}_{2}}\to
P{{H}_{3}}+3{{H}_{2}}PO_{2}^{-}\]
(a) Phosphorus is undergoing
reduction only.
(b) Phosphorus is undergoing
oxidation only.
(c) Phosphorus is undergoing oxidation
as well as reduction.
(d) Hydrogen is undergoing
neither oxidation nor reduction.
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question_answer56)
Which of the following
electrodes will act as anodes, when connected to standard hydrogen electrode?
(a) \[Al/A{{l}^{3+}}\] \[{{E}^{O-}}=1.66\]
(b) \[Fe/F{{e}^{2+}}\] \[{{E}^{O-}}=-0.44\]
(c) \[Cu/C{{u}^{2+}}\] \[{{E}^{O-}}=+0.34\]
(d) \[{{F}_{2}}(g)2{{F}^{-}}(aq)\] \[{{E}^{O-}}=+2.87\]
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question_answer57)
The reaction:
\[C{{l}_{2}}(g)+2O{{H}^{-}}(aq)\to
Cl{{O}^{-}}(aq)+C{{l}^{-}}(aq)+{{H}_{2}}O(l)\]
represents the process of bleaching.
Identify and name the species that bleaches the substances due to its oxidizing
action.
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question_answer58)
\[MnO_{4}^{2-}\]undergoes
disproportionation reaction in acidic medium but\[MnO_{4}^{2-}\] does not. Give
reason.
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question_answer59)
\[\text{PbO}\]and \[\text{Pb}{{\text{O}}_{\text{2}}}\]react
with \[\text{HCl}\]according to following chemical equations :
\[2PbO+4HCl\to
2PbC{{l}_{2}}+2{{H}_{2}}O\]
\[Pb{{O}_{ & 2}}+5HCl\to
PbC{{l}_{2}}+C{{l}_{2}}+2{{H}_{2}}O\]
Why do these compounds differ in
their reactivity?
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question_answer60)
Nitric acid is an oxidising
agent and reacts with \[\text{PbO}\]but it does not react with\[~\text{Pb}{{\text{O}}_{\text{2}}}\].
Explain why?
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question_answer61)
Write balanced chemical equation
for the following reactions :
(i) Permanganate ion \[(MnO_{4}^{-})\]reacts
with sulphur dioxide gas in acidic medium to produce \[\text{M}{{\text{n}}^{\text{2}+}}\]and
hydrogen sulphate ion. (Balance by ion electron method)
(ii) Reaction of liquid
hydrazine \[\left( {{\text{N}}_{\text{2}}}{{\text{H}}_{\text{4}}} \right)\]with
chlorate ion\[(ClO_{3}^{-})\] in basic medium produces nitric oxide gas and chloride
ion in gaseous state.
(Balance by oxidation number
method)
(iii) Dichlorine heptaoxide\[(C{{l}_{2}}{{O}_{7}})\]in
gaseous state combines with an aqueous solution of hydrogen peroxide in acidic
medium to give chlorite ion \[(ClO_{2}^{-})\]and oxygen gas.
(Balance by ion electron method)
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question_answer62)
Calculate the oxidation number
of phosphorus in the following species :
(a) \[HPO_{3}^{2-}\]and
(b) \[PO_{4}^{3-}\]
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question_answer63)
Calculate the oxidation number
of each sulphur atom in the following compounds :
(a) \[N{{a}_{2}}{{S}_{2}}{{O}_{3}}\] (b)
\[N{{a}_{2}}{{S}_{4}}{{O}_{6}}\] (c)\[\text{N}{{\text{a}}_{\text{2}}}\text{S}{{\text{O}}_{\text{3}}}~~\] (d)\[\text{N}{{\text{a}}_{\text{2}}}\text{S}{{\text{O}}_{\text{4}}}\]
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question_answer64)
Balance the following equations by the
oxidation number method :
(i) \[F{{e}^{2+}}+{{H}^{+}}+C{{r}_{2}}O_{7}^{2-}\to
C{{r}^{3+}}+F{{e}^{3+}}+{{H}_{2}}O\]
(ii) \[{{I}_{2}}+NO_{3}^{-}\to
N{{O}_{2}}+IO_{3}^{-}\]
(iii) \[{{I}_{2}}+{{S}_{2}}O_{3}^{2-}\to
{{I}^{-}}+{{S}_{4}}O_{6}^{2-}\]
(iv)\[Mn{{O}_{2}}+{{C}_{2}}O_{4}^{2-}\to
M{{n}^{2+}}+C{{O}_{2}}\]
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question_answer65)
Identify the redox reactions out of the
following reactions and identify the oxidising and reducing agents in them :
(i) \[3HCl(aq)+HN{{O}_{3}}(aq)\to
C{{l}_{2}}(g)+NOCl(g)+2{{H}_{2}}O(l)\]
(ii) \[HgC{{l}_{2}}(aq)+2KI(aq)\to
Hg{{I}_{2}}(s)+2KCI(aq)\]
(iii) \[F{{e}_{2}}{{O}_{3}}(s)+3CO(g)\xrightarrow{\Delta
}2Fe(s)+3C{{O}_{2}}(g)\]
(iv)\[PC{{l}_{3}}(l)+3{{H}_{2}}O(l)\to
3HCl(aq)+{{H}_{3}}P{{O}_{3}}(aq)\]
(v) \[4N{{H}_{3}}+3{{O}_{2}}(g)\to
2{{N}_{2}}(g)+6{{H}_{2}}O(g)\]
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question_answer66)
Balance the following ionic
equations :
(i) \[C{{r}_{2}}O_{7}^{2-}+{{H}^{+}}+{{I}^{-}}\to
C{{r}^{3+}}+{{I}_{2}}+{{H}_{2}}O\]
(ii) \[C{{r}_{2}}O_{7}^{2-}+F{{e}^{2+}}+{{H}^{+}}\to
C{{r}^{3+}}+F{{e}^{3+}}+{{H}_{2}}O\]
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question_answer67)
Match column I with column II
for the oxidation states of the central atoms.
|
Column l
|
Column ßß
|
|
(i) \[C{{r}_{2}}O_{7}^{2-}\]
|
(a) \[+3\]
|
|
(ii) \[MnO_{4}^{-}\]
|
(b) \[+4\]
|
|
(iii) \[VO_{3}^{-}\]
|
(c) \[+5\]
|
|
(iv) \[FeF_{6}^{3-}\]
|
(d) \[+6\]
|
|
|
(e) \[+7\]
|
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question_answer68)
Match the items in Column I with
relevant items in Column II.
|
|
Column I
|
Column II
|
|
(i)
|
Ions having positive charge
|
(a)
|
+7
|
|
(ii)
|
The sum of oxidation number of all atoms in a
neutral molecule
|
(b)
|
- 1
|
|
(iii)
|
Oxidation number of hydrogen ion \[({{H}^{+}})\]
|
(c)
|
+ 1
|
|
(iv)
|
Oxidation number of fluorine in \[NaF\]
|
(d)
|
0
|
|
(v)
|
Ions having negative charge
|
(e)
|
Cation
|
|
|
|
(f)
|
Anion
|
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question_answer69)
In the following questions a
statement of Assertion (A) followed by a statement of Reason (R) is given.
Choose the correct option out of the choices given below each question.
Assertion (A): Among halogens
fluorine is the best oxidant.
Reason (R): Fluorine is the
most electronegative atom.
(a) Both A and R
are true and R is the correct explanation of the A.
(b) Both A and R
are true but R is not the correct explanation of the A.
(c) A is true but R
is false.
(d) Both A and R
are false.
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question_answer70)
Assertion (A): In the reaction
between potassium permanganate and potassium iodide, permanganate ions act as
oxidising agent.
Reason (R): Oxidation state
of manganese changes from +2 to +7 during the reaction.
(a) Both A and R
are true and R is the correct explanation of the A.
(b)Both A and R are true but R is
not the correct explanation of the A.
(c) A is true but R
is false.
(d) Both A and R are false.
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question_answer71)
Assertion (A): The decomposition
of hydrogen peroxide to form water and oxygen is an example of disproportionation
reaction.
Reason (R): The oxygen of
peroxide is in -1 oxidation state and it is converted to zero oxidation sate in
\[{{O}_{2}}\] and -2 oxidation state in\[{{H}_{2}}O\].
(a) Both A and R
are true and R is the correct explanation of the A.
(b)Both A and R
are true but R is not the correct explanation of the A.
(c) A is true but R
is false.
(d) Both A and R
are false.
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question_answer72)
Assertion (A): Redox couple is
the combination of oxidised and reduced form of a substance involved in an oxidation
or reduction half cell.
Reason (R): In the
representation\[E_{F{{e}^{3+}}/F{{e}^{2+}}}^{O-}\] and
\[E_{C{{u}^{2+}}/Cu}^{O-},F{{e}^{3+}}/F{{e}^{2+}}\]and\[C{{u}^{2+}}/Cu\]
are redox couples.
(a) Both A and R
are true and R is the correct explanation of the A.
(b)Both A and R
are true but R is not the correct explanation of the A.
(c) A is true but R
is false.
(d) Both A and R
are false.
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question_answer73)
Explain redox reactions on the
basis of electron transfer. Give suitable examples.
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question_answer74)
On the basis of standard
electrode potential values, suggest which of the following reactions would take
place?
(Consult the book for \[{{E}^{O-}}\]
value).
(i) \[Cu+Z{{n}^{2+}}\to
C{{u}^{2+}}+Zn\]
(ii) \[Mg+F{{e}^{2+}}\to
M{{g}^{2+}}+Fe\]
(iii) \[B{{r}_{2}}+2C{{l}^{-}}\to
C{{l}_{2}}+2B{{r}^{-}}\]
(iv)\[Fe+C{{d}^{2+}}\to
Cd+F{{e}^{2+}}\]
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question_answer75)
Why does fluorine not show
disproportionation reaction?
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question_answer76)
In question 34, write redox
couples involved in the reactions (i) to (iv).
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question_answer77)
Find out the oxidation number of
chlorine in the following compounds and arrange them in increasing order of oxidation
number of chlorine.
\[NaCl{{O}_{4}},NaCl{{O}_{3}},NaClO,KCl{{O}_{2}},C{{l}_{2}}{{O}_{7}},Cl{{O}_{3}},C{{l}_{2}}O,\]\[NaCl,C{{l}_{2}},Cl{{O}_{2}}.\]
Which oxidation state is not present in any of
the above compounds?
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