|
 Doc
Brown's Chemistry
Advanced
Level Inorganic Chemistry Periodic Table
Revision Notes – Transition Metals
Part 10. Transition Metals 3d–block:
10.5
Vanadium
Chemistry
Vanadium exhibits
oxidation states of +2, +3, +4 and +5.
principal oxidation states of
vanadium, redox reactions of vanadium, ligand substitution displacement
reactions of vanadium, balanced equations of vanadium chemistry, formula
of vanadium complex ions, shapes colours of vanadium complexes, formula
of compounds
GCSE/IGCSE
Periodic Table Revision Notes *
GCSE/IGCSE Transition Metals Revision Notes
|
|
INORGANIC
Part 10 3d block TRANSITION METALS sub–index: 10.1–10.2
Introduction 3d–block Transition Metals * 10.3
Scandium
* 10.4 Titanium * 10.5
Vanadium * 10.6 Chromium
* 10.7 Manganese * 10.8
Iron * 10.9 Cobalt
* 10.10 Nickel
* 10.11 Copper * 10.12
Zinc
* 10.13 Other Transition Metals e.g. Ag and Pt * Appendix 1.
Hydrated salts, acidity of
hexa–aqua ions * Appendix 2. Complexes
& ligands * Appendix 3. Complexes and isomerism * Appendix 4.
Electron configuration & colour theory *
Appendix 5. Redox
equations, feasibility, Eø * Appendix 6.
Catalysis * Appendix 7.
Redox
equations
* Appendix 8. Stability Constants and entropy
changes *
Appendix 9. Colorimetric analysis
and complex ion formula * Appendix 10 3d block
– extended data
* Appendix 11 Some 3d–block compounds, complexes, oxidation states
& electrode potentials * Appendix 12
Hydroxide complex precipitate 'pictures',
formulae and equations
Advanced
Level Inorganic Chemistry Periodic Table Index *
Part 1
Periodic Table history
* Part 2
Electron configurations, spectroscopy,
hydrogen spectrum,
ionisation energies *
Part 3
Period 1 survey H to He *
Part 4
Period 2 survey Li to Ne * Part
5 Period 3 survey Na to Ar *
Part 6
Period 4 survey K to Kr and important trends down a
group *
Part 7
s–block Groups 1/2 Alkali Metals/Alkaline Earth Metals *
Part 8
p–block Groups 3/13 to 0/18 *
Part 9
Group 7/17 The Halogens *
Part 10
3d block elements & Transition Metal Series
*
Part 11
Group & Series data & periodicity plots * All
11 Parts have
their own sub–indexes near the top of the pages
10.5. Chemistry
of Vanadium V, Z=23, 1s22s22p63s23p63d34s2
data comparison of vanadium
with the other members of the 3d–block and transition metals
|
Z
and symbol |
21
Sc |
22
Ti |
23
V |
24
Cr |
25
Mn |
26
Fe |
27
Co |
28
Ni |
29
Cu |
30
Zn |
|
property\name |
scandium |
titanium |
vanadium |
chromium |
manganese |
iron |
cobalt |
nickel |
copper |
zinc |
|
melting
point/oC |
1541 |
1668 |
1910 |
1857 |
1246 |
1538 |
1495 |
1455 |
1083 |
420 |
|
density/gcm–3 |
2.99 |
4.54 |
6.11 |
7.19 |
7.33 |
7.87 |
8.90 |
8.90 |
8.92 |
7.13 |
|
atomic
radius/pm |
161 |
145 |
132 |
125 |
124 |
124 |
125 |
125 |
128 |
133 |
|
M2+
ionic radius/pm |
na |
90 |
88 |
84 |
80 |
76 |
74 |
72 |
69 |
74 |
|
M3+
ionic radius/pm |
81 |
76 |
74 |
69 |
66 |
64 |
63 |
62 |
na |
na |
|
common oxidation
states |
+3
only |
+2,3,4 |
+2,3,4,5 |
+2,3,6 |
+2,3,4,6,7 |
+2,3,6 |
+2,3 |
+2,+3 |
+1,2 |
+2
only |
|
outer electron config. |
3d14s2 |
3d24s2 |
3d34s2 |
3d54s1 |
3d54s2 |
3d64s2 |
3d74s2 |
3d84s2 |
3d104s1 |
3d104s2 |
|
Electrode
potential M(s)/M2+(aq) |
na |
–1.63V |
–1.18V |
–0.90V |
–1.18V |
–0.44V |
–0.28V |
–0.26V |
+0.34V |
–0.76V |
|
Electrode
potential M(s)/M3+(aq) |
–2.03V |
–1.21V |
–0.85V |
–0.74V |
–0.28V |
–0.04V |
+0.40 |
na |
na |
na |
|
Electrode
potential M2+(aq)/M3+(aq) |
na |
–0.37V |
–0.26V |
–0.42V |
+1.52V |
+0.77V |
+1.87V |
na |
na |
na |
Extended data table for VANADIUM
|
property of vanadium/unit |
value for V |
|
melting
point V/oC |
1910 |
|
boiling
point V/oC |
3380 |
|
density V/gcm–3 |
6.11 |
|
1st
Ionisation Energy V/kJmol–1 |
650 |
|
2nd
IE/kJmol–1 |
1414 |
|
3rd
IE/kJmol–1 |
2828 |
|
4th
IE/kJmol–1 |
4507 |
|
5th
IE/kJmol–1 |
6294 |
|
atomic
radius V/pm |
132 |
|
V2+
ionic radius/pm |
88 |
|
Relative polarising power M2+ ion |
2.3 |
|
V3+
ionic radius/pm |
74 |
|
Relative polarising power V3+ ion |
4.1 |
|
V4+
ionic radius/pm |
60 |
|
Polarising power V4+ ion |
6.7 |
|
oxidation
states of V,
less common/stable |
+2, +3, +4, +5 |
|
simple electron
configuration of V |
2,8,11,2 |
|
outer electrons of V |
[Ar]3d34s2 |
|
Electrode potential V(s)/V2+(aq) |
–1.18V |
|
Electrode potential V(s)/V3+(aq) |
–0.85V |
|
Electrode potential V2+(aq)/V3+(aq) |
–0.26V |
|
Electronegativity of V |
1.63 |

The Chemistry
of VANADIUM

-
The
electrode potential chart highlights the values for various
oxidation states of vanadium.
-
Vanadium shows a
'classic' display of variable oxidation states
of varying colours
when a solution of e.g. ammonium vanadate(V), is reduced by a
zinc/dilute sulphuric acid mixture. do picture
-
You go from the
vanadium(V) vanadate(V) ion ==> vanadium(IV) oxovanadate(IV) ion ==>
vanadium(III) ion ==> vanadium(II) ion
-
Acidification changes
the vanadate(V) ion into the pale yellow oxo–cation VO2+
(oxovanadium(V) ion)
-
VO43–(aq) + 4H+(aq)
VO2+(aq) + 2H2O(l) [an
acid–base reaction, NOT a redox change]
-
Note: Highly charged cations >3+ rarely exist as the
simple 'hydrated' tetra or hexa–aqua ion.
-
The theoretical polarising power
of the 'central metal ion' is so strong that
they form oxocations (see above) or oxyanions e.g.
-
orange
dichromate(VI) Cr2O72–, yellow
chromate(VI) CrO42–, purple manganate(VII)
MnO4– etc.
-
For transition metals they may be coloured
even if electronically the theoretical 'central metal ion'
has a noble gas structure e.g. [Ar] in its maximum
oxidation state like V(V), Cr(VI) and Mn(VII).
-
These oxyanions are
called charge transfer complexes and the theory is beyond
pre–university chemistry.
-
Three
successive reduction
steps then follow to eventually give V2+ ions,
shown as half–cell equations:
-
(i) V(V, +5) ==> V(IV,
+4): VO2+(aq) + 2H+(aq) +
e–
VO2+(aq) + H2O(l)
-
(ii) V(IV, +4) ==> V(III,
+3): VO2+(aq) + 2H+(aq) +
e–
V3+(aq) + H2O(l)
-
EØhalf–cell
potential = +0.34V,
blue to
the
green vanadium(III) ion
-
Here the vanadium(III) ion, V3+, is actually the
green hexaaquavanadium(III) ion,
-
Both V(IV) and V(III) species
are slowly oxidised by dissolved oxygen back to the V(V) compound in
acid solution.
-
(iii) V(III,
+3) ==> V(II, +2): V3+(aq) + e–
V2+(aq)
-
EØhalf–cell
potential = –0.26V,
green to the purple–violet vanadium(II) ion.
-
V2+(aq)
is powerful reducing agent and is unstable in the presence of
air.
-
Any dissolved oxygen will oxidise V2+(aq)
back to the vanadium(III) cation.
-
V2+ is actually the
purple–violet hexaaquavanadium(II) ion, [V(H2O)6]2+
-
Note
-
The standard
electrode potential EØZn(s)/Zn2+(aq)
is –0.76V, so the reducing power of zinc is sufficient
to effect any of the three vanadium oxidation state reduction
changes described above.
-
The reduction occurs on the surface of the zinc
metal i.e. the
site of electron transfer and you can write the above reductions
as a fully balanced complete equations ...
-
(i) 2VO2+(aq) +
4H+(aq) + Zn(s) ==> 2VO2+(aq) + 2H2O(l)
+ Zn2+(aq)
-
(ii) 2VO2+(aq) + 4H+(aq)
+ Zn(s) ==> 2V3+(aq) +
2H2O(l) + Zn2+(aq)
-
(iii) 2V3+(aq) + Zn(s
==>
2V2+(aq) + Zn2+(aq)
-
EØreaction
= –0.26 – (–0.76) = +0.50V
-
BUT
the vanadium(II) cation is unstable in the presence of
dissolve oxygen in air.
-
1/2O2(g) + 2H+(aq)
+ 2e–
H2O(l) has a standard electrode
potential of +1.23V,
-
so,
for the vanadium(II) oxidation reaction ...
-
1/2O2(g) + 2H+(aq)
+ 2V2+(aq) ==> 2V3+(aq)
+
H2O(l)
-
EØreaction
= EØreduction – EØoxidation
= +1.23 – (–0.26) = +1.49V
-
hence the if left standing open to air, the violet V2+(aq)
solution will gradually change to a green V3+(aq)
solution and in turn V3+(aq) will
revert back to VO2+(aq) in the
presence of air because of oxidation by dissolve oxygen
unless protected by an inert atmosphere. (see
Redox Electrode
Potential Chart, V2+/V3+
and V3+/VO2+ potentials are less
positive (below)
that for O2/H2O/H+
potentials).
-
You will see
hydrogen formed simultaneously from the unavoidable metal–acid
reaction.
-
Does vanadium
chemistry show an example of disproportionation?
-
This is just a little academic
exercise using standard electrode potential data.
-
A disproportionation
reaction is where a species in one oxidation state spontaneously and
simultaneously changes into two species of different oxidation states –
one higher and one lower in oxidation number.
-
Examples:
disproportionation in manganese(VI)
chemistry and
disproportionation in copper(I) chemistry
-
Question: In
terms of aqueous ions, is the disproportionation of vanadium(III) into
vanadium(II) and vanadium (IV) feasible?
-
(i) VO2+(aq) + 2H+(aq) + 2e–
V3+(aq) + H2O(l)
(EØVO2+/V3+ = +0.34V)
-
(ii) V3+(aq) + e–
V2+(aq) (EØV3+/V2+
= –0.26V)
-
The
disproportionation equation would be (iii) 2V3+(aq)
+ H2O(l)
V2+(aq) + VO2+(aq) + 2H+(aq)
-
For equation (iii),
(ii) will be the reduction half–cell equation and (i) reversed will be
the oxidation half–cell reaction.
-
EØreaction
= EØreduction – EØoxidation =
= EØV3+/V2+ – EØVO2+/V3+ =
(–0.26) – (+0.34) = –0.60V
-
showing the
disproportionation is thermodynamically NOT feasible i.e. EØreaction
is less than zero.
-
In fact what can
actually happen is if you mix salt solutions of vanadium(IV) and
vanadium(II) on an equimolar basis, you end up with a solution of
vanadium(III) salts, a sort of 'anti–disproportionation' reaction!
-
Summary of some
complexes–compounds & oxidation states of vanadium compared to other
3d–block elements
-
–
Scandium
* Titanium * Vanadium
* Chromium
* Manganese * Iron * Cobalt
* Nickel
* Copper *
Zinc
* Silver & Platinum
keywords redox reactions ligand
substitution displacement balanced equations
formula complex ions complexes ligand exchange reactions redox reactions ligands
colours oxidation states: vanadium ions V2+ V(+2) V(II) V3+ V(+3) V(III) V4+ V(+4) V(IV)
V5+ V(+5)
(V) SO2 + V2O5 ==> SO3 + V2O4 V2O4 + 1/2 O2 ==> V2O5 VO43– + 4H+ VO2+ + 2H2O
V(V, +5) ==> V(IV, +4): VO2+ + 2H+ + e– VO2+ + H2O V(IV, +4) ==> V(III, +3):
VO2+ + 2H+ + e– V3+ + H2O [V(H2O)6]3+ V(III, +3) ==> V(II, +2): V3+ + e– V2+
VO3+/VO2+ (+1.00V), VO2+/V3+ (+0.34V) and V3+/V2+ (–0.26V) 2 VO2+ + 4H+ + Zn ==>
2 VO2+ + 2H2O + Zn2+ 2VO2+ + 4H+ + Zn ==> 2V3+ + 2H2O + Zn2+ 2V3+ + Zn(s ==>
2V2+ + Zn2+ 1/2O2 + 2H+ + 2V2+ ==> 2V3+ + H2O V2+/V3+ and V3+/VO2+ potentials
VO2+ + 2H+ + 2e– V3+ + H2OEØVO2+/V3+ = +0.34V) (ii) V3+ + e– V2+ (EØ V3+/V2+ =
–0.26V) EØ V3+/V2+ – EØ VO2+/V3+ oxidation states of vanadium, redox reactions
of vanadium, ligand substitution displacement reactions of vanadium, balanced
equations of vanadium chemistry, formula of vanadium complex ions, shapes
colours of vanadium complexes
A level Revision notes for GCE Advanced
Subsidiary Level AS Advanced Level A2 IB
Revise AQA GCE Chemistry OCR GCE Chemistry Edexcel GCE Chemistry Salters
Chemistry CIE Chemistry, WJEC GCE AS A2 Chemistry, CCEA/CEA GCE AS A2 Chemistry revising courses for pre–university students
(equal to US grade 11 and grade 12 and AP Honours/honors level courses)
 Website
content copyright © Dr W P Brown 2000–2012 All rights reserved
on
revision notes, puzzles, quizzes, worksheets, x–words etc. * Copying of website
material is not permitted
chemhelp@tiscali.co.uk

Alphabetical Index for Science
Pages Content
A
B C D
E F
G H I J K L M
N O P
Q R
S T
U V W
X Y Z
Scandium
* Titanium * Vanadium
* Chromium
* Manganese * Iron * Cobalt
* Nickel
* Copper *
Zinc
* Silver & Platinum
Introduction 3d–block Transition Metals * Appendix
1.
Hydrated salts, acidity of
hexa–aqua ions * Appendix 2. Complexes
& ligands * Appendix 3. Complexes and isomerism * Appendix 4.
Electron configuration & colour theory *
Appendix 5. Redox
equations, feasibility, Eø * Appendix 6.
Catalysis * Appendix 7.
Redox
equations
* Appendix 8. Stability Constants and entropy
changes *
Appendix 9. Colorimetric analysis
and complex ion formula * Appendix 10 3d block
– extended data
* Appendix 11 Some 3d–block compounds, complexes, oxidation states
& electrode potentials * Appendix 12
Hydroxide complex precipitate 'pictures',
formulae and equations |