|
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
10.11. Chemistry
of Copper Cu, Z=29, 1s22s22p63s23p63d104s1
data comparison of copper
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 COPPER
|
property of copper/unit |
value for Cu |
|
melting point Cu/oC |
1083 |
|
boiling point Cu/oC |
2567 |
|
density Cu/gcm–3 |
8.92 |
|
1st
Ionisation Energy Cu/kJmol–1 |
745 |
|
2nd
IE/kJmol–1 |
1958 |
|
3rd
IE/kJmol–1 |
3554 |
|
4th
IE/kJmol–1 |
5326 |
|
5th
IE/kJmol–1 |
7709 |
|
atomic
radius Cu/pm |
128 |
|
Cu2+
ionic radius/pm |
69 |
|
Relative polarising power Cu2+ ion |
2.9 |
|
oxidation
states of Cu,
less common/stable |
+1, +2, +3 |
|
simple electron
configuration of Cu |
2,8,18,1 |
|
outer electrons of Cu |
[Ar]3d104s1 |
|
Electrode potential Cu(s)/Cu2+(aq) |
+0.34V |
|
Electronegativity of Cu |
1.90 |

The
Chemistry of
COPPER

-
The
electrode potential chart highlights the values for various
oxidation states of copper.
-
COPPER(II) CHEMISTRY
-
When copper(II) salts are
dissolved in water the blue tetraaquacopper(II) ion or the
hexaaquacopper(II) ion is formed.
-
The scope for a variety of
coloured compounds arises from the fundamental electronic configuration
of the Cu2+ ion, namely [Ar]3d9,
giving an incompletely filled 3d sub–shell – criteria for being a
true transition metal.
-
ie there is at least one
electron that can be promoted to a higher level when the 3d sub–shell is
split when the central metal ion interacts with the ligands.
-
The Cu2+
components in the diagrams below illustrates the point.
-
For more details see
Appendix 4.
Electron configuration & complex ion colour theory
-

-

-
Both
the octahedral hexaaquacopper(II) ion [Cu(H2O)6]2+
and the square planar tetraaquacopper(II) ion
-
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.
-
Excess sodium hydroxide
has no significant effect, BUT with excess ammonia, a deep
blue solution is formed of the ammine complex 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 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.
-
At
very high concentration the dark violet hexaamminecopper(II) ion can
be formed.
-
With sodium carbonate
solution, copper(II) ions
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.
-
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
pale 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 complex ion to an anionic complex 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:
-
If you
dissolve copper(II) chloride in water you get a greenish–blue
solution as both copper(II) complexes are present in
equilibrium.
-
By adding
water i.e. dilution, it shifts
the equilibrium to the left, more blue.
-
Increasing the
chloride ion concentration by adding hydrochloric acid or
sodium chloride solution shifts the equilibrium to the
right, more green ==> yellowish brown.
-
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.
-
Summary of some
complexes–compounds & oxidation states of copper compared to other
3d–block elements
-
COPPER(I) CHEMISTRY
-
The colour of copper(I) compounds
-
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.
-
Formation of copper(I)
compounds and examples of copper(I) complexes
-
Copper(I) iodide is
formed on mixing solutions of a soluble copper(II) salt with potassium
iodide solution.
-
2Cu2+(aq)
+ 4I–(aq) ==> 2CuI(s) + I2(aq/s)
-
It is unfortunate, from a
preparation point of view, that iodine is also formed – completely
obscuring the 'white' copper(I) iodide!
-
Copper(I) iodide, like
copper(I) chloride, is white when pure, when left out in air,
they will slowly oxidise to the copper(II) compound eg copper(I)
chloride slowly turns green as copper(II) compounds are formed.
-
I'm not quite sure how you
can isolate the copper(I) iodide from this mixture? I think you can
remove the iodine with sodium thiosulfate and then filter off, wash and
dry the CuI. Try it? You will be lucky if its white and left out in air
will discolour further due to aerial oxidation.
-
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).
-
Biochemistry of Copper

|
The original
extraction
of copper from copper ores |
- From copper carbonate ores ...
- The ore can be roasted to concentrate the copper as
its oxide.
- Water is driven off and the
carbonate thermally decomposed.
- copper(II) carbonate
==> copper oxide + carbon dioxide
- CuCO3(s) ==> CuO(s) + CO2(g)
- The oxide can be smelted by heating with carbon (coke, charcoal) to
reduce the oxide to impure copper, though this method isn't
really used much these days (the 'bronze age' method
archaeologically!).
- copper(II) oxide +
carbon ==> copper + carbon dioxide
- 2CuO(s) + C(s) ==> 2Cu(s) + CO2(g)
- The carbon acts as the
reducing agent – the 'oxygen remover'.
- From copper sulphide ores ...
- These include
chalcocite/chalcosine = copper(I) sulphide Cu2S
and covellite = copper(II) sulphide CuS
- and chalcopyrite CuFeS2.
which is one of the most important ores for the extraction of
copper.
- This can be roasted in air
to produce copper(I) sulfide which is roasted again in a
controlled amount of air so as not to form a copper oxide (see
below).
- 2CuFeS2 +
4O2 ==> Cu2S + 3SO2 +
2FeO
- Copper sulphide ores can be
rapidly roasted
in heated air enriched with oxygen to form impure copper and
this extraction process is called 'flash
smelting' and is the most widely used and efficient method
of copper extraction.
- Nasty sulphur dioxide gas is
formed, this must be collected to avoid pollution and can be
used to make sulphuric acid to help the economy of the process.
- copper(I) sulphide +
oxygen ==> copper + sulphur dioxide
- Cu2S(s) + O2(g)
==> 2Cu(s) + SO2(g)
- or copper(II) sulphide +
oxygen ==> copper + sulphur dioxide
- CuS(s) + O2(g)
==> Cu(s)
+ SO2(g)
- It is also
possible to dissolve an oxide or carbonate ore in dilute sulphuric acid and extracting copper by ....
- (1) using
electrolysis see purification
by electrolysis below, or
- (2) by adding
a more reactive metal to displace it
e.g. scrap iron or steel is
used by adding it to the resulting copper(II) sulphate solution.
- iron + copper(II)
sulphate ==> iron(II) sulphate + copper
- Fe(s)
+ CuSO4(aq) ==> FeSO4(aq) + Cu(s)
- It is possible to spray acid
onto copper ore waste and leach out the copper compounds prior
to electrolysing the solution or displacing the copper with a
cheap metal like iron AND this can also be achieved with the
help of bacteria for particular ores – see below.
- No industrial process is ever 100%
efficient, and metal extraction processes of create lots of waste
material AND, crucially, that waste may contain some of the desired
metal, or indeed other potentially valuable metals.
- So, any method that can extract the
small percentages of valuable metals from waste will aid the economy
of production of the main product.
- For example, 10% of the copper
produced in the US is derived from bacteria which feed of
chalcopyrite CuFeS2 (this could be waste or very low
grade ores ??).
- The bacteria use the Fe2+
ion and S2– ion to obtain energy needed to live.
- The redox chemistry of the bacteria
via the oxygen from air involves the ...
- oxidation of Fe2+ to Fe3+
and S2– to SO42–
- reduction of O2 to H2O
- overall an energy releasing process to
sustain the bacteria.
- We use sugar, these bacteria use
chalcopyrite! – I don't think we'd like the taste!
- –
- The bacteria effectively break down
the chalcopyrite, releasing copper(II) ions into an acid solution.
- The optimum conditions for this 'bacterial
leaching' are pH 2–3 and 20oC–55oC.
- Extracting copper in this way is
cheaper, quieter, and less polluting than conventional smelting
processes.
- BUT, it is much slower, so it is
primarily being used on waste dumps by spraying dilute acid on them
and the aerated water slowly percolates through rock fragments and
their naturally occurring bacterial colonies.
- The leached solution of copper(II)
ions is very dilute.
- The solution is concentrated and Cu2+
ions separated from other ions e.g. Fe3+
- The copper is displaced using cheap
scrap iron.
- The copper is then further purified
by electrolysis – described below.
- –
|

|
The
Purification of Copper by Electrolysis |
|
 
|
-
The impure copper from a smelter is cast into
a block to form the positive anode. The cathode is made of previously
purified copper. These are dipped into an electrolyte
of copper(II) sulphate solution.
-
When the d.c electrical current is passed through the
solution electrolysis takes place. The copper anode dissolves forming blue copper(II) ions Cu2+.
-
These positive ions are attracted to the negative
cathode and become
copper atoms. The mass of copper dissolving at the anode exactly equals
the mass of copper deposited on the cathode. The concentration of the
copper(II) sulphate remains constant.
-
Any impurities present in the impure
copper anode fall to the bottom of the electrolysis cell tank. This 'anode
sludge' is not completely mineral waste, it can contain valuable metals such
as silver!
-
See section below
for extraction of impure copper from an ore.
|
|
Raw materials for the
electrolysis process:
Electrolysis
is using
d.c. electrical energy to bring about chemical changes at the electrolyte
connections called the anode and cathode electrodes.
An
electrolyte is a conducting melt or solution of ions which carry the
electric charge as part of the circuit.
Scrap copper
can be
recycled and purified this way too ,and is cheaper than starting
from copper ore AND saves valuable mineral resources. |
The redox details of the electrode processes:
- At the positive (+) anode, the process is an oxidation, electron
loss, as the copper atoms dissolve to form copper(II) ions.
Cu(s) ==> Cu2+(aq)
+ 2e–
- at the negative (–) cathode, the process is a reduction, electron
gain by the attracted copper(II) ions to form neutral copper atoms.
Cu2+(aq) + 2e–
==> Cu(s)
- Note: Reduction and Oxidation
always go together, hence the use of the term redox change or
reaction.
- Electroplating is mentioned on
the Industrial
Chemistry and Electrochemistry
pages.
|
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: copper ions Cu(0) Cu+ Cu(+1) Cu(I) Cu2+ Cu(+2) Cu(II) Cu(+3) Cu(III)
CuSO4 Cu2O CuSO4.5H2O [Cu(H2O)4]2+ [Cu(H2O)6]2+ + 2OH– ==> [Cu(H2O)4(OH)2] + 2H2O
[Cu(H2O)6]2+ + 4 NH3 [Cu(NH3)4(H2O)2]2+ + 4 H2O [Cu(H2O)4]2+ + 4NH3
[Cu(NH3)4]2+ + 4H2O [Cu(H2O)4(OH)2] + 4NH3 [Cu(NH3)4(H2O)2]2+ + 2 OH– + 4H2O
Cu(OH)2] + 4NH3 [Cu(NH3)4]2+ + 2OH– Kstab = [ [Cu(NH3)4(H2O)2]2+ ] / [
[Cu(H2O)6]2+ ] [ NH3 ]4 = 1.0 x 1012 mol–4 dm12 Cu2+ + CO32– ==> CuCO3
Cu2+ + 2HCO3– ==> CuCO3 + H2O + CO2 [Cu(H2O)6]2+ + 4Cl– [CuCl4]2– + 6H2O units
of Kstab = [ [CuCl4]2+ ] / [ [Cu(H2O)6]2+ ] [ Cl– ]4 = ? mol–4 dm12 Cu2+
+ 4I– ==> 2 CuI + I2 Cu2O + H2SO4 ==> Cu + CuSO4 + H2O Cu2O + 2H+ ==> Cu + Cu2+
+ H2O Oxidation number changes: 2 Cu(I) ==> Cu(0) + Cu(II) Cu2SO4 + aq ==> Cu +
CuSO4 Cu2SO4 + aq ==> Cu + Cu2+ + SO42– Oxidation state changes: 2Cu(+1) ==> Cu
(0) + Cu (+2) 2Cu+ ==> Cu2+ +Cu 2Cu2+ + 4 I– ==> 2 CuI + I2 CuCl + Cl– ==>
[CuCl2]– 2Cu + 2H+ + 4Cl– ==> 2[CuCl2]– + H2 CuCl + 4CN– ==> [Cu(CN)4]3– + Cl–
oxidation states of copper, redox reactions of copper, ligand substitution
displacement reactions of copper, balanced equations of copper chemistry,
formula of copper complex ions, shapes colours of copper complexes Na2CO3
NaOH NH3
Advanced Level Inorganic Chemistry
of Copper – A level Revision notes to help
revise for GCE Advanced
Subsidiary Level AS Advanced Level A2 IB
Revise AQA GCE Advanced Level Chemistry OCR GCE Advanced Level Chemistry Edexcel GCE
Advanced Level Chemistry Salters
AS A2 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
|