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Brown's Chemistry Clinic
Advanced
Level Inorganic Chemistry Revision notes
(e.g. UK Advanced Level Chemistry GCE-AS-A2-IB US K12 grades ~11-12)
The
Periodic Table Part 10
Part
10b "3d block - Transition Metals"
revision notes
1st draft
10b: Fe,
Co, Ni, Cu, Zn (and Ag, Pt)
Part 10a Sc to Mn
GCSE Chemistry
revision notes *
GCSE notes on Transition Metals *
EMAIL
query?comment
Part 10 3d block sub-index: 1.
Introduction * 2.
3d-block data,
general trends and character of Transition Metals * 3.
Scandium * 4. Titanium
* 5. Vanadium * 6. Chromium * 7.
Manganese * 8. Iron * 9. Cobalt 10. Nickel
* 11. Copper * 12. Zinc * 13. Other
Transition Metals e.g. silver Ag or platinum Pt * Appendix 1.
Acidity of hexa-aqua ions * Appendix 2. Complexes
& ligands - the basics * Appendix 3. The
shapes of complexes and isomerism *
Appendix 4. Electron configuration and colour theory *
Appendix 5. Redox equations, feasibility,
calculating Eøreaction * Appendix 6.
Catalysis - examples and theory * Appendix 7. Balancing
redox equations * Appendix 8. Stability
constants of complex ions *
Appendix 9. Colorimetry - quantitative analysis
and determining the formula of a complex ion * Appendix 10.
Preparation of complexes * 3d block -
Transition Metals data * Hydroxide
precipitate 'pictures'
* Extra comparison of 3d-block formulae and
oxidation states
* Extra Electrode Potential Chart for
3d-block
Advanced Periodic Table Index *
Part 1
A brief Periodic Table history *
the modern Periodic Table
* Part 2
Electronic structure of atoms :
Spectroscopy and the H spectrum :
Ionisation energies *
Part 3
Period 1 survey : 1. Hydrogen
:
2. Helium : Summary of
Period 1 : heavier element
formation-stellar nuclear fusion *
Part 7
s-block metals Groups 1/2 Alkali/alkaline Earth Metals *
Part 11
Group and Series data summaries
and links to periodicity plots
 Quick
click to
Introduction *
Sc * Ti
* V * Cr *
Mn
* Fe * Co * Ni
* Cu * Zn * Ag/Pt etc.
10. 8. Chemistry
of Iron Fe, Z=26, 1s22s22p63s23p63d64s2
-
Fe
data table 1 summary
* extended iron data table 2 *
Iron electrode potential chart
3d-block
-
Summary of some
complexes-compounds & oxidation states of iron compare 3d-block
elements
-
Some basic reactions of iron
metal are on the GCSE Reactivity Series of Metals Notes
-
Iron is an
extremely useful metal.
-
Carbon steels like
mild steel (0.1% - 04% carbon) are used for enumerable objects like car
bodies, tin cans, nuts/bolts and piping.
-
Stainless steel, an
alloy with chromium, has extremely good anti-corrosion properties.
-
Tungsten and
manganese steels are very tough and hard wearing and used for cutting
tools and high speed drill bits.
-
Wrought iron is
tough malleable and ductile and good material for a blacksmith to work
with.
-
Cast iron, despite
being brittle, is used for manhole covers, guttering, machinery frames
and drainpipes.
-
An iron/iron(III)
oxide mixture is used as
a catalyst in the Haber synthesis of ammonia from hydrogen and
nitrogen.
-
N2(g) +
3H2(g) ==Fe/Fe2O3==> 2NH3(g)
-
Other catalysis
examples: The iron(II)/iron(III) ion catalysis of the oxidation of
iodide ions by peroxodisulfate ions is described under
homogeneous catalysis in
Appendix 6.
-
All the
details
of iron extraction in a blast furnace are given in the GCSE
notes and there is little point in repeating them here.
-
Starting with impure
iron from blast furnace, the molten iron contains many other
elements and the iron is too brittle initially, so there is a need to
reduce C and remove others like S and P.
-
This is achieved by
the Basic Oxygen Steel making process (BOS) which involves
many redox reactions. It is a 'batch process' and can't be
used as a continuous production line like iron from the blast
furnace.
-
Sulphur is removed
early in the process using magnesium:
-
C, P, Si and others oxidised by molecular
oxygen before scrap iron/steel introduced.
-
After the oxygen blow
the basic oxides CaO/MgO are added to form slag salts with the
weak acidic oxides of Si and P, carbon dioxide gas will
'escape' from the mixture, since any calcium carbonate formed
would thermally decompose at the high temperature of the furnace.
-
e.g. CaO + SiO2
==> CaSiO3 (calcium silicate)
-
The oxides of
Mn/Fe also collect in the slag, so some iron is wasted and the
Mn might be added in a controlled way later for a particular
steel specification.
-
The toxic carbon
monoxide formed must be dealt with and not allowed out into
the atmosphere, it can be burned as a fuel to harmless carbon
dioxide.
-
It is important to
keep track of temperature and composition by
thermocouple probe and atomic
emission spectroscopy.
-
The elements are
oxidised in a sequence in exothermic reactions (no extra heat
needed), so temperature control is essential to avoid
wasting energy and converter lining damage.
-
The added scrap
iron/steel addition acts as coolant because melting is
endothermic.
-
The whole process
must meet the specification for an individual customer requirement.
-
Dissolved oxygen
is removed with aluminium
-
and then C, Mn and Si etc. can be
re-added to a desired specification, plus any other elements,
to make a particular steel.
-
Argon (of
light bulb fame) is bubbled through to stir the mixture
because it so unreactive and most 'stirrers' will melt and
dissolve, and change the composition.
-
In the future
electric arc furnaces maybe used more to recycle steel. Big
carbon electrodes are 'sparked' to melt the scrap iron/steel, lime
added to remove impurities as slag. It is possible to use this
technology on a small scale to produce
-
Steel is an alloy
based on iron.
-
An alloy is a
mixture of a metal with at least one other element (metal or
non-metal) or compound.
-
The
composition of steel, like any other alloy, is crucial in
determining its properties.
-
Small differences
in composition can have significant effect on the properties
of an alloy.
-
Too high a % of C
in iron makes it too brittle, but a low % C makes a
stronger steel.
-
You need to
appreciate the versatile
nature of steel by changing its composition and quote some
examples.
-
There is a need
for excluding impurities eg O, P or S which lead to poor quality
material.
-
The common elements added
to iron to make steel, apart from
carbon, are usually other transition metals.
-
Scrap iron and steel is part of
BOS process and is cost effective, recycling reduces costs of (i)
ore mining extraction, (ii) possibly overseas transport and (iii)
blast furnace reduction of ore. These gains are partly offset by the
cost of collecting scrap metal.
-
In the electric
arc process only scrap steel is used and is handy technology
to produce small batches of particular steel by carefully
controlling what scrap goes in.
-
The composition of
scrap important, needs to be graded and selected to avoid problems
-
When recycling tin
cans, you need to remove the tin and other waste.
-
The cans are
shredded and paper/residual food removed, mechanical shredding
and magnetic separation can be used,
-
and de-tinning is
done by reaction with hot NaOH(aq), after which the valuable
tin can be recovered by electrolysis of the 'waste solution'.
-
A particular scrap
case study, need for steel uncontaminated by radioactive
isotopes from the nuclear and atomic weapon industries, scrap
source from the German ships sunk at Scapa Flow has proved useful
(good geography Q and I don't remember the event!).
-
The
most common oxidation states of iron in its compounds are +2 and +3.
-
IRON(II) and IRON(III)
Chemistry
-
Iron readily dissolves
in dilute hydrochloric or sulphuric acid to form iron(II) chloride
and
iron(II) sulphate respectively. Hydrogen gas is evolved and it is
a redox reaction.
-
Fe (s)
+ 2HCl (aq) ==> FeCl2 (aq) + H2 (g)
-
Fe (s)
+ H2SO4 (aq) ==> FeSO4 (aq) +
H2 (g)
-
The redox-ionic equation
is: Fe (s) + 2H+ (aq) ==>
Fe2+ (aq) + H2 (g)
-
hydrogen ions (H in
+1 ox. state) are reduced by electron gain to hydrogen gas (H
in 0 ox. state) and iron is oxidised from the 0 ox. state to the
+2 ox. state. Note that the lower oxidation state of iron is
formed, since neither acid is a strong oxidising agent.
-
The pale green salts FeCl2.6H2O
and FeSO4.7H2O can
be made by careful evaporation and crystallisation of the
solution. However, they are readily oxidised by dissolved oxygen
to form iron(III) compounds.
-
White anhydrous
iron(II) chloride can be made by passing hydrogen chloride gas
over heated iron.
-
If chlorine is
passed over heated iron, brown anhydrous iron(III) chloride is
formed
-
2Fe (s)
+ 3Cl2 (g) ==> 2FeCl3 (s)
-
An example of 'salt' synthesis by directly combining the constituent elements.
-
Iron(III) chloride is
a brown covalently bonded, relatively volatile chloride. Like
aluminium chloride, it exists in the solid form as a dimer Fe2Cl6,
one of the Fe's chlorines acts as a bridge, forming a dative
co-ordinate bond with the other iron atom (see diagram below).
-
Redox
reaction: ox. state changes are Fe (0) to (+3), Cl (0)
to (-1)
-
The iron(III)
chloride reacts very exothermically with water to give pungent
acrid fumes of hydrogen chloride (anhydrous aluminium chloride is
made in the same way and behaves with water in the same way!).
Hence the need for dry conditions in their preparation is
illustrated below. Its also a very good idea to vent the excess
chlorine away safely!
-
FeCl3 (s)
+ 3H2O (l) ==> Fe(OH)3 (s)
+ 3HCl (g)
-
Reducing action of
aqueous iron(II) ions:
These two reactions can be used to quantitatively estimate Fe2+
ions.
-
Oxidising action of
iron(III) ions:
-
With iodide ions,
dark brown solution of iodine (or black solid) formed with
iron(II) ions.
-
2Fe3+(aq)
+ 2I-(aq) ==> 2Fe2+(aq)
+ I2(aq/s)
-
This accounts
for why iron(III) iodide cannot exist.
-
Oxidation state
changes: Fe +3 to +2, I -1 to 0.
-
With zinc,
colourless zinc and pale green iron(II) ions are formed. This
reaction is usually done in the presence of dil. sulphuric acid.
-
Zn(s) + 2Fe3+(aq)
==> 2Fe2+(aq) + Zn2+(aq)
-
Oxidation state
changes: Fe +3 to +2, Zn 0 to +2.
-
The reaction can be used
as part of a process to
titrate and analyse estimate Fe2+ and Fe3+ mixtures.
titration ref
-
-
-
Simple test for
aqueous iron(III) ions: add a few drops of ammonium/potassium
thiocyanate solution (NH4SCN or KSCN). The reaction is NOT
given by hexa-aqua iron(II) ions.
-
Iron/iron(III)
oxide mixture is used as the main
component of the catalyst in the Haber Synthesis of ammonia
from
nitrogen and
hydrogen.
-
 Some
biochemistry of iron
-
The
biological role of iron complexes haemoglobin, myoglobin and ferritin.
-
Oxygen,
O2, molecules co-ordinate to an iron(II) ion in the
haemoglobin (hemoglobin) molecule ('haem'
(porphyrin square
planar complex diagram to do), which acts as a giant complex ion in
transportation systems of the blood. Essential for respiration energy
release, ...
-
 Unfortunately carbon monoxide forms a stronger ligand bond
than oxygen and will displace it to give CO its well deserved toxic
respiration. It only takes a small amount of CO, and a simple ligand
exchange reaction to affect the respiratory system!
-
The enzyme catalase
is extremely efficient at decomposing hydrogen peroxide
molecule in organisms. One proposed mechanism involves a
catalytic cycle of iron(III) and iron(IV) complexes e.g. if somewhat
simplified ....
-
ENZYME-FeIII
+ H2O2 ==> ENZYME-FeIV=O + H2O
-
ENZYME-FeIV=O
+ H2O2 ==> ENZYME-FeIII + H2O
+ O2
-
Other complexes
of Fe2+ and Fe3+ ions
-
Iron(II) ions
complex with the ethanedioate dicarboxylate anion, a bidentate ligand:
-
with cyanide ion
octahedral cyano complexes
-
-
-
RUSTING and anti-corrosion
chemistry
-
The
electrochemical processes of RUSTING which is the
corrosion of iron to form an iron oxide
by an oxidation process which is
energetically favourable, and it is the
opposite of its extraction by reduction of iron oxide.
-
Unfortunately rust
flakes off and so it all eventually corrodes away (later
xref/contrast ZnO, Al2O3, Cr2O3
on metal surface, which do not flake away and offer good
anti-corrosion properties)
-
Factors
affecting rate of rusting
e.g. the following all speed up the
process!
-
decreasing pH,
H+(aq) ions remove OH-(aq)
formed from the reduction of O2(g/aq),
-
increased
concentration of any ions improves the conductivity of the
aqueous media, which is part of 'redox circuit',
-
and if the iron
is in contact with a 'less reactive' metal
(meaning a more +ve half-cell potential), corrosion
rates increase, because the iron is preferentially oxidised
with the more -ve half-cell potential.
-
Rust
protection-inhibition
... examples ... are x-ref with
assignment 7 on p174.
-
A plastic or
paint physical barrier to exclude water and oxygen (air),
-
Either
by (i) dipping in molten zinc, or (ii) electrolysis with Zn2+(aq)
solution and the iron/steel object as -ve cathode, galvanising
with Zn layer which results in the formation of ZnO layer,
the redox chemistry is similar to Fe rusting (see Fig 21) but
the layer does not flake away giving a protective layer of
zinc oxide. Even if scratched, the Zn with a more -ve
half-cell potential is preferentially oxidised.
-
Sacrificial
corrosion with blocks of Zn or Mg
and relate their
'sacrifice' to their more negative half-cell potentials, i.e.
preferentially more favourable oxidation.
-
Fe2+(aq) + 2e- 
Fe(s) (EØ =
-0.44V)
-
Zn2+(aq) + 2e-
Zn(s) (EØ =
-0.76V)
-
Mg2+(aq) + 2e-
Mg(s) (EØ =
-2.38V)
-
reminder that
the reduction of oxygen to water is a positive redox potential
-
O2
(aq/g) + 2H2O(l) + 4e-
4OH-(aq)
(EØ =
+0.44V, in alkali)
-
or O2
(aq/g) + 4H+(aq) + 4e-
2H2O(l)
(EØ =
+1.23V, in acid)
-
so all the metal
oxidations are feasible BUT the most negative potential will lead to
the preferential oxidation i.e. Mg > Zn > Fe.
-
Stainless
steel via Cr addition, forms protective layer of
chromium(III) oxide.
-
History lesson in
food preservation: ‘invention’ of the tin can (tin coated
steel) ...
-
Tin plating
steel offers some corrosion protection of the iron because tin is
not a particularly reactive metal (less negative potential).
-
However, early tin cans
suffered from preferential oxidation of Fe due to its more
–ve potential, through any microscopic defect in the tin
layer, or indeed if it got scratched. This was cured by lacquer coating
as an extra protective barrier.
-
Fe2+(aq) + 2e-
Fe(s) (EØ =
-0.44V)
-
Sn2+(aq) + 2e-
Sn(s) (EØ =
-0.14V)
-
Still, fruit juice
was a problem, carboxylic acids complex with Sn2+(aq)
ions, changes Sn(s)/Sn2+(aq)
potential making it more negative than Fe(s)/Fe2+(aq),
so Sn preferentially corrodes, not toxic and contribute to
‘tangy’ taste BUT don’t keep too long as Fe eventually
will dissolve too!
-
Complex
formation affecting corrosion behaviour. Here tin(II) ions form a complex with carboxylic acids
like citric acid (tridentate ligand), by reducing the Sn2+(aq)
concentration, the Sn(s)/Sn2+(aq)
half-cell potential is then made more negative that that of iron! so
the protective thin layer of tin is sacrificially corrode, then its
the iron! Don't worry too much, the rates of reaction are slow, BUT
don't keep tinned fruit on the shelf for too long!
-
Estimation of iron
in iron(II) salts and tablet formulations.
Quick
click to
Introduction *
Sc * Ti
* V * Cr *
Mn
* Fe * Co * Ni
* Cu * Zn * Ag/Pt etc.
10.9.
Chemistry
of Cobalt Co, Z=27, 1s22s22p63s23p63d74s2
-
Co
data table 1 summary
* extended cobalt data table 2 *
Cobalt & electrode potential chart
of 3d-block
-
Summary of some
complexes-compounds & oxidation states of cobalt compared to other
3d-block elements
-
Cobalt is
alloyed with chromium and tungsten to make a metal hard enough,
even at red heat, to be used for high speed cutting tools and valves for
internal combustion engines.
-
COBALT(II) chemistry
-
In aqueous solution,
in the absence of complexing agents,
cobalt forms the stable pink hexaaqua cobalt(II) ion, [Co(H2O)6]2+(aq)
-
The alkalis sodium
hydroxide and ammonia, produce the hydrated
cobalt(II) hydroxide
blue ppt. which turns pink on standing. There is no further reaction with excess
of NaOH or Na2CO3, but see further down for excess NH3.
-
Alkaline aqueous
sodium carbonate solutions produces a precipitate of
pink/blue? cobalt(II) carbonate.
-
When excess ammonia
is added to
a cobalt(II) salt solution, the hexamine complex is formed BUT this is
unstable in the presence of dissolved oxygen and is oxidised to the
cobalt(III) complex. This change in cobalt's oxidation state from +2
to +3 via an oxidising agent is quite common if a complexing agent is
present too.
-
[Co(H2O)6]2+(aq) + 6NH3(aq)
==> [Co(NH3)6]2+(aq) + 6H2O(l)
-
pink
hexaaquacobalt(II) ion ==> brown hexaamminecobalt(II) ion
-
4[Co(NH3)6]2+(aq) + O2(g/aq) + 4H+(aq)
==> 4[Co(NH3)6]3+(aq) + 2H2O(l)
-
brown ==>
colour? hexaamminecobalt(III) ion
-
+1.82
for [Co(H2O)6]3+(aq) + e-
[Co(H2O)6]2+(aq)
-
+0.10
for [Co(NH3)6]3+(aq) + e-
[
Co(NH3)6]2+(aq)
-
VIEW ppts. with OH-, NH3
and CO32-, & complexes,
if any, with
excess reagent.
-
When hydrogen peroxide is
added to an alkaline cobalt(II) solution, oxidation occurs to give cobalt(III)
complexes.
-
If e.g. sodium chloride
or hydrochloric acid is added to cobalt(II) sulphate solution the
blue tetrachlorocobaltate(II) complex ion is formed.
-
[Co(H2O)6]2+(aq) + 4Cl-(aq)
[CoCl4]2-(aq) + 6H2O(l)
-
This particular
ligand substitution/exchange reaction involves several changes (L
to R):
-
the larger
chloride ion ligand leads to a change in co-ordination number
from 6 to 4,
-
the complex ion
shape changes from octahedral to tetrahedral
-
the colour of the
complex changes from pink to blue,
-
the complex
changes from a cationic to an anionic ion.
-
There is no oxidation
state change at all.
-
This is quite a good
reaction to demonstrate Le Chatelier's equilibrium principles:
-
dilution shifts
the equilibrium to the left, more pink,
-
increasing the
chloride ion concentration shifts the equilibrium to the
right, more blue,
-
increasing the
solution temperature shifts the equilibrium to the right, more
blue
-
or if prepared at
higher temperature, with just enough chloride to turn the
solution blue, on cooling it becomes pink,
-
this shows that
left to right is endothermic and right to left is exothermic.
-
The uncharged ligand
molecules ammonia NH3 and water H2O are similar in
size and ligand exchange occurs without change in co-ordination number.
-
COBALT(III) chemistry
-
Isomerism in
cobalt(III) complexes e.g. with the ligands ammonia + chloride (i)-(iii)
and (iv) ethane-1,2-diamine (ethylenediamine).
-
(i) crystalline
[Co(NH3)6]3+(Cl-)3
is orange-yellow
-
(ii) crystalline
[Co(NH3)5Cl]2+(Cl-)2
is violet
-
(iii) crystalline
[Co(NH3)4Cl2]+Cl-
is violet or green - there are two geometrical isomers (cis and trans)
-
(iv) [Cr(H2NCH2CH2NH2)3]3+,
H2NCH2CH2NH2,
ethane-1,2-diamine (ethylenediamine), is often represented in
shorthand by en,
-
-
Quick
click to
Introduction *
Sc * Ti
* V * Cr *
Mn
* Fe * Co * Ni
* Cu * Zn * Ag/Pt etc.
10.10. Chemistry
of Nickel Ni, Z=28, 1s22s22p63s23p63d84s2
-
Ni
data table 1 summary
*extended nickel data table 2 *
Nickel & electrode potential chart
of 3d-block
-
Summary of some
complexes-compounds & oxidation states of nickel compared to other
3d-block elements
-
Nickel has
many uses from 'silver' coinage metals and monel used for chemical
reactors - both are alloys with copper to give a chemically inert
metal.
-
Nickel is an important
hydrogenation catalyst in converting unsaturated vegetable oils to
saturated fats like margarine.
-
unsaturated oil +
hydrogen ==> low melting solid saturated fat
-
Along the carbon
chain of the oil you get: -CH=CH- + H2 ==> -CH2-CH2-
-
NICKEL(II) CHEMISTRY
-
In aqueous solution
cobalt forms the green stable hexaaqua cobalt(II) ion, [Ni(H2O)6]2+
(aq)
-
The alkalis sodium
hydroxide or ammonia, produce the hydrated
nickel(II) hydroxide
green?
precipitate. There is no further reaction with excess
of NaOH, but see further down for excess NH3.
-
Alkaline aqueous
sodium carbonate solutions produces a precipitate of green ppt. of nickel(II) carbonate.
-
Ni2+(aq)
+ CO32-(aq) ==> NiCO3(s)
-
Its actually a
basic carbonate, a mixture of the hydroxide and carbonate, you can
make the pure carbonate by using sodium hydrogencarbonate solution.
-
Ni2+(aq)
+ 2HCO3-(aq) ==> NiCO3(s) + 4H2O(l) +
CO2(g)
-
With excess aqueous
ammonia the blue hexa-ammine complex is formed:
-
[Ni(H2O)6]2+(aq)
+ 6NH3(aq)
[Ni(NH3)6]2+(aq)
+ 6H2O(l)
-
With lower
concentrations of ammonia the pale blue complex can also have the structure
[Ni(H2O)2(NH3)4]2+
-
VIEW ppts. with OH-, NH3
and CO32-, & complexes,
if any, with
excess reagent.
-
Other complexes of nickel
-
Nickel carbonyl,
Ni(CO)4, is a neutral complex tetrahedrally shaped covalent
molecule. Note (i) nickel is in a zero oxidation state and (ii) the ligand CO also acts as ligand with haemoglobin
(hemoglobin) in carbon monoxide poisoning.
-
Ni2+ forms
the tetrachloronickelate(II) ion, [NiCl4]2-, a
tetrahedral anionic complex
with the chloride ion (Cl-).
-
Ni2+ forms
the tetracyanonickelate(II) ion, [Ni(CN)4]2-, a square planar anionic complex
with the cyanide ion (CN-).
Quick
click to
Introduction *
Sc * Ti
* V * Cr *
Mn
* Fe * Co * Ni
* Cu * Zn * Ag/Pt etc.
 10.11. Chemistry
of Copper Cu, Z=29, 1s22s22p63s23p63d104s1
-
Cu
data table 1 summary
* extra copper data table 2 *
Copper & electrode potential
3d-block
-
Summary of some
complexes-compounds & oxidation states of copper compared to other
3d-block elements
-
Copper is an
important metal in many alloys e.g. brass (with zinc), bronze
(with tin) and coinage metals (with nickel).
-
COPPER(II) CHEMISTRY
-
When alkaline aqueous
ammonia or sodium hydroxide is added to a blue hexa-aqua copper(II) ion solution,
initially a gelatinous blue precipitate of the hydroxide is formed.
-
Note it can be 4 or
6 H2O in the complex ion Cu2+(aq) i.e.
[Cu(H2O)4]2+(aq)
-
[Cu(H2O)6]2+(aq)
+ 2OH-(aq) ==> [Cu(H2O)4(OH)2](s)
+ 2H2O(l)
-
or more simply:
Cu2+(aq)
+ 2OH-(aq) ==> Cu(OH)2(s)
-
(a
precipitation reaction)
-
Excess sodium hydroxide
has no significant effect, BUT with excess ammonia, a deep
blue solution is formed of the ??? ion (ligand substitution is
incomplete), the overall change can be expressed as:
-
[Cu(H2O)6]2+(aq)
+ 4NH3(aq)
[Cu(NH3)4(H2O)2]2+(aq)
+ 4H2O(l)
-
or [Cu(H2O)4]2+(aq)
+ 4NH3(aq)
[Cu(NH3)4]2+(aq)
+ 4H2O(l)
-
or from the
hydroxide precipitate
-
[Cu(H2O)4(OH)2](s)
+ 4NH3(aq)
[Cu(NH3)4(H2O)2]2+(aq)
+ 2OH-(aq) + 4H2O(l)
-
Note: ligand exchange
reaction, not a redox change, co-ordination number remains at 6, both
octahedral complexes, both ligands electrically neutral so the
overall charge of the complex remains at +2, both the ligands are
of similar size but the substitution is incomplete.
-
Kstab
= [ [Cu(NH3)4(H2O)2]2+(aq)
] / [
[Cu(H2O)6]2+(aq) ]
[ NH3 (aq) ]4 = 1.0 x 1012
mol-4 dm12
-
by convention
the term [ H2O(l)
]4 is omitted from the equilibrium expression because water is the
medium and the bulk of the solution, therefore it effectively remains
constant.
-
Sodium carbonate
gives the turquoise? precipitate of copper(II) carbonate,
-
Cu2+(aq)
+ CO32-(aq) ==>
CuCO3(s)
-
Its actually a
basic carbonate, a mixture of the hydrated hydroxide, Cu(OH)2, and carbonate,
CuCO3.
-
You can
make the pure carbonate by using sodium hydrogencarbonate solution.
-
Cu2+(aq)
+ 2HCO3-(aq) ==> CuCO3(s) + 4H2O(l) +
CO2(g)
-
VIEW ppts. with OH-, NH3
and CO32-, & complexes,
if any, with
excess reagent.
-
If e.g. sodium chloride
or hydrochloric acid is added to copper(II) sulphate solution the
yellow-brown tetrachlorocuprate(II) complex ion is
formed (seen as green due to the blue from the original Cu2+
ion).
-
[Cu(H2O)6]2+(aq) + 4Cl-(aq)
[CuCl4]2-(aq) + 6H2O(l)
-
This particular
ligand substitution/exchange reaction involves several changes (L
to R):
-
the larger
chloride ion ligand leads to a change in co-ordination number
from 6 to 4,
-
the complex ion
shape changes from octahedral to tetrahedral
-
the colour of the
complex changes from blue to yellow-brown (green due to
residual blue),
-
the complex
changes from a cationic to an anionic ion.
-
There is no oxidation
state change at all, copper is in the +2 state throughout the
reaction.
-
This is quite a good
reaction to demonstrate Le Chatelier's equilibrium principles:
-
dilution shifts
the equilibrium to the left, more blue,
-
increasing the
chloride ion concentration shifts the equilibrium to the
right, more green,
-
The reaction
between copper(II) salts and iodide salts:
-
i.e. the redox
reaction between the copper(II) ion and the iodide ion.
-
On mixing solutions
of a copper(II) salt e.g. blue copper(II) sulphate and an iodide salt
e.g. colourless potassium iodide the dark colour of iodine formation is
seen. Unseen, because it is masked by the iodine, is the formation of a
white copper(I) iodide precipitate. This can be made visible by adding
sodium thiosulphate solution which reduces the iodine back to the
colourless iodide ion.
-
Cu2+(aq)
+ 4I-(aq) ==> 2CuI(s) + I2(aq/s)
-
2S2O32-(aq) + I2(aq) ==>
S4O62-(aq) + 2I-(aq) (black/brown
==> colourless)
-
This reaction
between the released iodine and sodium thiosulfate can be used to
estimate oxidising agents like copper(II) ions. The iodine is titrated
with standardised sodium thiosulphate (e.g. 0.10 mol dm-3)
using a few drops of starch solution as an indicator. Iodine gives a
blue colour with starch, so, the end-point is very sharp change from the
last hint of blue to colourless.
-
Copper analysis eg. in brass
-
Brass can be dissolved in
acid and potassium iodide solution added.
-
The resulting
iodine formed can be titrated with sodium thiosulfate using starch
indicator.
-
Need more details and an example calculation.
-
COPPER(I) CHEMISTRY
-
Disproportionation reactions:
-
If solid copper(I)
oxide is dissolved in dil. sulphuric acid a pinky-brown precipitate of
copper and a blue solution of copper(II) sulphate solution is obtained.
-
If solid copper(I)
sulphate is dissolved in water the observations and oxidation number
changes are identical to the reaction above.
-
Cu2SO4(s)
+ aq ==> Cu(s) + CuSO4(aq)
-
Cu2SO4(s)
+ aq ==> Cu(s) + Cu2+(aq) + SO42-(aq)
-
Oxidation state
changes: 2Cu(+1) ==> Cu (0) + Cu (+2)
-
These two reactions
suggest that Cu+(aq) has no stability in aqueous
media and spontaneously undergoes a redox change and an electrode
potential argument predicts this potential for instability and therefore
the observations.
-
Note: A chemical
change in which 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, is called a disproportionation reaction. The argument is as follows ....
-
(i) Cu+ + e-
Cu (EØCu+/Cu = +0.52V)
-
(ii) Cu2+ + e-
Cu+ (EØCu2+/Cu+ =
+0.15V)
-
(i) with the more
positive redox potential represents the reduction half-cell reaction and
(ii), reversed, with the less positive potential, will represent the
oxidation half-cell reaction.
-
EØreaction
= EØreduction - EØoxidation =
(+0.52) - (+0.15) = +0.37V
-
showing the
disproportionation is thermodynamically feasible, i.e. EØreaction
must be greater than zero.
-
See
manganese(VI) chemistry for
another example of disproportionation.
-
Copper(I)/Cu+(aq)
can be stabilised by making complexes from suitable ligands e.g.
copper(I) chloride dissolves in conc. hydrochloric acid to form the
stable dichlorocuprate(I) complex ion (NOT a redox reaction).
-
CuCl(s)
+ Cl-(aq) ==> [CuCl2]-(aq)
-
The same
complex ion is formed if copper metal is boiled with conc.
hydrochloric acid when the redox reaction,' surprisingly'
produces hydrogen.
-
2Cu(s)
+ 2H+(aq) + 4Cl-(aq)
==> 2[CuCl2]-(aq) + H2(g)
-
The Cu2+/Cu
potential is +0.34V and the Cu+/Cu potential is +0.15V, so
hydrogen shouldn't be formed (EØH+/H2 =
0.00V), BUT the actual redox potential involved is for the [CuCl2]-/Cu
half-cell system which is <0.00V.
-
Copper(I)
compounds dissolve in an excess of potassium cyanide solution to
give the tetracyanocuprate(I) complex ion.
-
CuCl(s)
+ 4CN-(aq) ==> [Cu(CN)4]3-(aq)
+ Cl-(aq)
-
E argument
for Cu(CN)2]-/Cu is -0.44V
-
[Cu(CN)2]-(aq)
==> Cu(s) + Cl-(aq) ????
-
Copper(I)
oxide Cu2O is formed as a dark red-brown precipitate
when an aldehyde or reducing sugar reacts with Fehlings solution (a
copper(II) complex with a carboxylic acid).
-
Biochemistry of Copper
Quick
click to
Introduction *
Sc * Ti
* V * Cr *
Mn
* Fe * Co * Ni
* Cu * Zn * Ag/Pt etc.
10.12. Chemistry
of Zinc
Zn, Z=30, 1s22s22p63s23p63d104s2
Quick
click to
Introduction *
Sc * Ti
* V * Cr *
Mn
* Fe * Co * Ni
* Cu * Zn * Ag/Pt etc.
10.13.
Other Transition Metals
10.13a.
4d
block 2nd row elements Y to Cd
-
Yttrium, Y,
Z=39, [Kr]4d15s2, is not a true transition
metal and is like scandium Z=21 forming the colourless Y3+(aq)
ion, e.c. = [Kr].
-
Rhodium, Rh, Z=45, [Kr]4d105s1 :
Metal used as catalyst in car exhaust (see
platinum below
for more details).
-
Silver, Ag, Z=47, [Kr]4d105s1
-
?
-
 The silver(I) ion
forms linear complexes with several ligands. The bond angle
is 180o and
co-ordination number 2 e.g.
-
The water
molecule ligands in the aqueous silver ion [Ag(H2O)2]+(aq)
can be replaced e.g. with (i) with ammonia,
NH3,
(neutral ligand) giving a cationic complex, (ii) with the negative cyanide
ion CN-
giving an anionic complex, and (iii) the negative thiosulfate
ion S2O32- forming an
anionic complex:
-
[Ag(NH3)2]+(aq)
solution is used as Tollen's reagent (ammoniacal silver nitrate) in
organic chemistry. It is readily reduced by aldehydes (NOT
ketones) to form a 'silver mirror' on the side of the test
tube on warming the mixture to 60oC.
-
[Ag(CN)2]-(aq)
solution is used in the electrolyte in silver electroplating. The
object to be coated in silver is made the negative cathode
electrode.
-
[Ag(H2O)2]+(aq)
+ 2CN-(aq)
[Ag(CN)2]-(aq)
+ 2H2O(l)
-
The
equilibrium is well over to the right but a very low
concentration of silver ions gives a good even and strongly
adhering surface deposit of silver metal on the conducting
negative cathode electrode.
-
At the
cathode (-): [Ag(H2O)2]+(aq) + e- ==> Ag(s)
+ 2H2O(l)
-
The
silver is replenished using a silver anode
-
At the
anode (+): Ag(s) + 2CN-(aq)
- e- ==> [Ag(CN)2]-(aq)
-
[Ag(S2O3)2]3-(aq) is formed when sodium thiosulphate is used to
remove unreacted silver bromide (AgBr) or silver iodide (AgI) crystals in developing
photographic films.
-
AgBr(s) +
2S2O32-(aq) ==> [Ag(S2O3)2]3-(aq)
+ Br-(aq)
-
NOT a redox
reaction, Ag is +1 and Br is -1 throughout the reaction. The
thiosulfate ion is here acting as a ligand and not a
reducing agent e.g. like with iodine.
-
-
-
The use of
silver nitrate and ammonia for the halide test is described in the
Chemical Tests Notes for anions
-
-
13b.
5d
block 3rd row elements La, Hf to Hg
|
where L-L represents H2NCH2CH2NH2