资源简介

(共148张PPT)
Titanium (Ti)
Electron Configuration
4s23d2 4s23d1 3d2 3d1
Appearance
Hard; lustrous silver colored.
Sources: (Named the titans)
Ilmenite: FeTiO3
Rutile: TiO2 Brilliant white, commonly used as a pigment.
Titanite: CaTiSiO5
Perovskite: CaTiO3
Uses
Resistant to corrosion, lightweight and strong.
Alloys, machinery, aircraft, and missiles.
Bright white pigment TiO2.
Chemistry
Main oxidation state: +4(d0).
Other oxidations states: -1(d5), 0, +2 & +3.
Burns in air and is the only element that burns in nitrogen.
It is ductile only in an oxygen-free atmosphere.
Resistant to dilute hydrochloric and sulfuric acids, most organic acids, chlorine gas and chloride solutions.
Reaction of titanium
with air
Titanium metal is coated with an oxide layer that usually renders it inactive. However once titanium starts to burn in air it burns with a spectacular white flame to form titanium dioxide, TiO2 and titanium nitride, TiN. Titanium metal even burns in pure nitrogen to form titanium nitride.
Ti(s) + O2(g) TiO2(s)
2Ti(s) + N2(g) TiN(s)
with water
Titanium metal is coated with an oxide layer that usually renders it inactive. However, titanium will react with steam form the dioxide, titanium(IV) oxide, TiO2, and hydrogen, H2.
Ti(s) + 2H2O(g) TiO2(s) + 2H2(g)
with acids
Dilute aqueous hydrofluoric acid, HF, reacts with titanium to form the complex anion [TiF6]3- together with hydrogen, H2.
2Ti(s) + 12HF(aq) 2[TiF6]3-(aq) + 3H2(g) + 6H+(aq)
Titanium metal does not react with mineral acids at ambient temperature but does react with hot hydrochloric acid to form titanium(III) complexes.
Reaction of titanium continued
with the halogens
Titanium does react with the halogens upon warming to form titanium(IV) halides. The reaction with fluorine requires heating to 200°C. So, titanium reacts with fluorine, F2, chlorine, Cl2, bromine, I2, and iodine, I2, to form respectively titanium(IV) halides.
Ti(s) + 2F2(g) TiF4(s) [white]
Ti(s) + 2Cl2(g) TiCl4(l) [colourless]
Ti(s) + 2Br2(g) TiBr4(s) [orange]
Ti(s) + 2I2(g) TiI4(s) [dark brown]
Preparation of Ti
Kroll Process
TiO2 + 2C + 2Cl2 TiCl4 +2CO
2FeTiO3 + 7Cl2 + 6C 2TiCl4 + 2FeCl3 + 6CO
TiCl4 + 2Mg Ti + 2MgCl2
Electrolysis:
TiO2 (solid, cathode) => molten salt electrolysis in CaCl2 => Ti (cathode) + O2 (anode)
Chemists have two ways to pry metal from an oxide ore. One, electrolysis, decomposes the ore into its elementary constituents with electricity. Aluminum manufacturing employs this method. The alternative, called chemical reduction, involves reacting the ore with a substance that has a greater affinity for oxygen than the metal to be extracted. This procedure is used to refine iron
Metallurgy
TiCl4
TiCl4 + HCl(c) H2TiCl6
Ziegler-Natta Catalyst
TiCl4 + (CH3CH2)3Al
TiCl4 + 2H2O TiO2 + 4HCl
Partial hydrolysis: TiOCl2
Polyethylene
most popular plastic.
Grocery bags, shampoo bottles, toys, etc.
Simple structure than all polymers.
Branched/low-density = (LDPE)
Easier to make
Linear/high-density = (HDPE)
Ziegler-Natta Catalyst
H
2
C
C
H
2
C
H
2
C
H
2
n




manufacture of TiO2
Reaction of TiO2
rutile(金红石), anatase(锐钛矿), brookite(板钛矿), rutile is the most common form, and the others transform into it on heating.
TiCl4(g) + O2(g) TiO2(s) + 2Cl2(g) @1200 oC, Cl2(g) recycle
TiO2 + H2SO4 (c) TiOSO4 + H2O
no titanyl cation TiO2+, probably [Ti(OH)2(H2O)4]2+
2KOH + TiO2 K2TiO3 + H2O
amphoteric
TESTS FOR TITANIUM
Aqueous solutions containing titanium(IV) give an orange-yellow color on addition of hydrogen peroxide; the color is due to the formation of peroxo-titanium complexes, but the exact nature of these is not known.
Vanadium (V)
Electron Configuration: 4s23d3 3d4 3d3 3d2
Origin: Named after the Scandinavian goddess 'Vanadis‘ because of its beautiful multi-coloured compounds. .
Ores:
Vanadinite: Pb5Cl(VO4)3
Patronite: VS4
Roscoelite: K2V4Al2Si6O20(OH)4
Carnotite: K2(UO2)2(VO4)2.H2O
Properties
Soft, ductile, bright white.
Main oxidation states: III, IV, & V
Uses
About 80% of the vanadium produced is used as a steel additive.
In this form it produces one of the toughest alloys for armor plate, axles, piston rods and crankshafts.
Less than 1% of vanadium and as little chromium make steel shock- and vibration-resistant.
Vandium(V) oxide is used in ceramics
Vanadium
it is not very useful itself
most of the metal produced is in the form of an alloy ferrovanadium, containing between 40 and 90% vanadium. This is added to steel to produce a very tough 'high-speed' steel.
Ferrovanadium is obtained by reduction of the oxide V2O5 with "ferrosilicon” (Fe + Si).
The pure metal is very difficult to prepare because it combines even more readily with hydrogen, carbon, nitrogen and oxygen than does titanium
Oxides of Group 5B metals
Oxidation state +5 +4 +3 +2
V V2O5 VO2 V2O3 VO
Nb Nb2O5 NbO2 - NbO
Ta Ta2O5 TaO2 - (TaO)
V2O5: an essentially acidic oxide, dissolving in alkalis to give vanadates; however, addition of acid converts the anionic vanadate species to cationic species:
V2O5 (yellow) + H+ VO2+ (pale yellow) + H2O
Redox Reactions
Sulphur dioxide will reduce V(+5) to V(+4) only; zinc in hydrochloric acid will reduce V(+5) step by step to V(+2). The half equations and the overall equation is given for each of these reactions.
The aqueous chemistry of vanadium(V) is complex and depends on pH. Here we take strongly acidic solutions in which the yellow ion VO2+ is the species present. The colour changes for the complete reduction sequence are:
+5 +4 +3 +2
yellow blue green lavender
VO2+ VO2+ [V(H2O)6]3+ [V(H2O)6]2+
The water ligands are not shown in the equations below, but remember that they are there for V(+3) and V(+2).
+5 +4
with aqueous SO2:

2VO2+(aq) + 4H+(aq) + 2e – 2VO2+ (aq) + 2H2O (l)
SO32 – (aq) + H2O (l) SO42 – (aq) + 2H+(aq) + 2e – (aq)
2VO2+(aq) + SO32 – (aq) + 2H+(aq) 2VO2+ (aq) + SO42 – (aq) + H2O (l)

with zinc in HCl:

2VO2+(aq) + 4H+(aq) + 2e – 2VO2+ (aq) + 2H2O (l)
Zn(s) Zn2+ (aq) + 2e –
2VO2+(aq) + 4H+(aq) + Zn(s) 2VO2+ (aq) + Zn2+ (aq) + 2H2O (l)
Redox Reactions Continued
+4 +3 with zinc in HCl

2VO2+ + 2e – + 4H+ (aq) 2V3+ (aq) + 2H2O (l)
Zn(s) Zn2+ (aq) + 2e –
2VO2+ + Zn(s) + 4H+ (aq) 2V3+ (aq) + Zn2+ (aq) + 2H2O (l)


+3 +2 with zinc in HCl

2V3+(aq) + 2e – 2V2+ (aq)
Zn(s) Zn2+ (aq) + 2e –
2V3+(aq) + Zn(s) 2V2+ (aq) + Zn2+ (aq)
TESTS FOR VANADIUM
the reduction of vanadium(V) to vanadium(II) by zinc and acid, gives a very characteristic test for vanadium.
Addition of a few drops of hydrogen peroxide to a vanadate(V) gives a red colour (formation of a peroxo-complex)
Chromium (Cr)
Electron configuration: 4s13d5 3d5 3d4 3d3
Ore: Chromite FeCr2O4
Properties:
Hard, silvery metal with a blue tinge. Capable of taking a high polish.
Metal-metal quadruple bonds are possible.
Gives Ruby its color.
Main oxidation state: III
Uses:
Used in stainless steels (~13% Cr) in increase corrosion resistance.
Used for plating other metals.
Used chemically in pigments (vivid green, yellow, red and orange colors).
Tanning agents, catalysts, and oxidizing agents.
Chromium compounds are toxic.
CHROMIUM
Pure Metal
Direct reduction of chromate by heating with carbon and calcium oxide gives an alloy of iron and chromium, ferrochrome, which can be added to steel, to make stainless steel (12-15 % chromium).
Chromium Compounds
All compounds of chromium are colored (except Cr(CO)6 ); the most important are the chromates of sodium and potassium and the dichromates and the potassium and ammonium chrome alums. The dichromates are used as oxidizing agents in quantitative analysis, also in tanning leather. Other compounds are of industrial value; lead chromate is chrome yellow, a valued pigment. Chromium compounds are used in the textile industry as mordants, and by the aircraft and other industries for anodizing aluminum.
Chromium(III) halides
Formula Color MP M-X (pm)
CrF3 green 1404 190
CrCl3 red-violet 1152 238
CrBr3 green-black 1130 257
CrI3 black >500d -
Preparations:
CrX3 are prepared from Cr with X2, dehydration of CrCl3.6H2O requires SOCl2 at 650 C
Anhydrous Chromium(III) chloride
peach-coloured solid,
the reaction of chlorine with a heated mixture of chromium(III) oxide and carbon:
the reaction of sulfur dichloride oxide with the hydrated chloride:
hydrated chloride, CrCl3,6H2O
Green colored crystal
[CrIII(H2O)4Cl2]+Cl-.2H2O.
grey-blue
pale green
green
AgNO3
Chromium(II) halides
Formula Color MP m (BM)
CrF2 green 894 4.3
CrCl2 white 820-824 5.13
CrBr2 white 844 -
CrI2 red-brown 868
Preparations:
Reduction of CrX3 with H2/HX gives CrX2
Chromium oxides
Formula Color Oxidation State MP Magnetic Moment
CrO3 deep red Cr6+ 197d -
Cr3O8 - intermediate - -
Cr2O5 - - - -
Cr5O12 etc - - - -
CrO2 brown-black Cr4+ 300d -
Cr2O3 green Cr3+ 2437 -antiferromagnetic < 35 C
Oxides
By heating chromium(III) hydroxide
By heating ammonium dichromate:
green powder, insoluble in water and in acids (cf. aluminum oxide, Al2O3). It is not reduced by hydrogen.
In 1986 the initial drying of the dichromate in a rotary vacuum drier, resulted in a serious explosion in Ohio. The cause was not obvious but the presence of an organic contaminant must be a possibility.
Prep：
Prop：
(greyish green)
(green)
1.Cr2O3
2.Cr(OH)3
△
CHROMIUM(III) HYDROXIDE, Cr(OH)3(HYDRATED)
3
2
2
4
3
4
2
CrCl
O,
12H
KCr(SO
SO
Cr

）
，
）
（
Common salts of Cr(Ⅲ)
Hydrolysis
+
+
+
+
H
]
O
Cr(OH)(H
[
]
O
Cr(H
[
2
5
2
3
6
2
）
）
4
10
-
≈
As reducing agent
As Oxidizing Agent
Blue
Zn
)
(
2Cr
2
2
+
+
+
Zn(s)
Cr
3
+
+
Cr2+(aq) Cr3+(aq)
14H
SO
O
Cr
O
H
7
O
S
3
2Cr
2
4
2
7
2
Ag
2
2
8
2
3
+
+
+
+
+
-
-
-
+
+
O
8H
2CrO
2OH
O
3H
2Cr(OH)
2
2
4
2
2
4
+
+
+
-
-
-
V
33
.
1
)
Cr
/
O
Cr
(
3
2
7
2
=
+
-
Acidic：
E
V
12
.
0
)
Cr(OH)
/
CrO
(
-
4
2
4
-
=
-
Basic：
E
Color of Cr(Ⅲ) Coordination Compds
pH effect
pH<2：Cr2O72- major pH>6：CrO42- major
O
H
O
Cr
2HCrO
2H
2CrO
2
2
7
2
4
2
4
+
+
-
-
+
-
(yellow)
(orange)
2.Interchange between Cr2O72- and CrO42-
Solubility Effect
-
×
=
10
.0
2
)
O
Cr
(Ag
7
7
2
2
-
×
=
10
1.1
)
CrO
(Ag
12
4
2
+
-
+
+
+
+
2H
)
(s,
CrO
2Ag
O
H
O
Cr
4Ag
4
2
2
2
7
2
brick
+
-
+
+
+
+
2H
)
(s,
2BaCrO
O
H
O
Cr
2Ba
4
2
2
7
2
2
lemon
+
-
+
+
+
+
2H
)
(s,
2PbCrO
O
H
O
Cr
2Pb
4
2
2
7
2
2
yellow
K2Cr2O7 PbCrO4
3. Cr2O72- as Oxidant
O
7H
2KCl
2
+
+
7
2
1.33V
)
/Cr
O
(Cr
3
2
=
+
-
O
H
4
Cr
2
3SO
H
8
3SO
O
Cr
2
3
2
4
2
3
2
7
2
+
+
+
+
+
-
+
-
-
O
7H
2Cr
3S
8H
S
3H
O
Cr
2
3
2
2
7
2
+
+
+
+
+
+
-
O
7H
2Cr
3I
14H
6I
O
Cr
2
3
2
2
7
2
+
+
+
+
+
+
-
-
2CrCl
3Cl
)
14HCl(
(s)
O
Cr
K
3
2
7
2
2
+
+
c
O
7H
2Cr
6Fe
14H
6Fe
O
Cr
2
3
3
2
2
7
2
+
+
+
+
+
+
+
+
-
O
7H
2Cr
3Sn
14H
3Sn
O
Cr
2
3
4
2
2
7
2
+
+
+
+
+
+
+
+
-
Dichromate and chromate
Dichromate and chromate equilibria is pH dependent:
HCrO4- -> CrO42- + H+ K=10-5.9 H2CrO4 -> HCrO4- + H+ K=10+0.26 Cr2O72- + H2O -> 2HCrO4- K=10-2.2 HCr2O7- -> Cr2O72- + H+ K=10+0.85
Dichromate preparation
Sodium dichromate is prepared on the large scale by heating powdered chromite with sodium carbonate, with free access of air
Conversion from chromate to dichromate:
2 CrO42-(aq) + 2 H+ (aq) Cr2O72-(aq) + H2O (l)
lower pH: Cr3O102-, Cr4O132-, CrO3.
Strong oxidizing agent.
Polymerization. stronger in oxyacids of Mo and W. polymetallates 多酸
e.g. [Mo36O112]8-, [H2W12O42]10- etc.
Heteropolymetallates
[CoIIW12O40]6-, [NiIVMo9O32]6-, etc.
Heterogeneous catalyst.
Zeolite Linde A: [Na12(Al2Si12O48)].27H2O
Similar behavior for some main group compounds:
aluminosilicate
Chromium (II) Acetate
red
Cr2(OCOCH3)4.2H2O
2.35
cf: 2.58 in metal
MO theory of Metal-Metal Bond Formation



dx2-y2, s, px, py pz
M-L bonds, five
dz2
dxz, dyz

dxy

anhydrous Cr2(OCOCH3)4
d(Cr-Cr) 1.98 :
quadruple bond
4.Identification of Cr(Ⅲ)
pentanol
blue
H+
TESTS FOR CHROMIUM
Fusion of any chromium compound with a mixture of potassium nitrate and carbonate gives a yellow chromate(VI).
CHROMATES AND DICHROMATES
Addition of lead(II) nitrate in ethanoic acid solution gives a yellow precipitate of lead chromate. PbCrO4.
A reducing agent (for example sulphur dioxide) reduces the yellow chromate or orange dichromate to the green chromium(III) state.
Hydrogen peroxide with a chromate or a dichromate gives a blue color.
CHROMIUM(III) SALTS
Addition of alkali gives a green gelatinous precipitate of chromium(III) hydroxide, soluble in a large excess of strong alkali.
Addition of sodium peroxide to a solution gives a yellow color of the chromate.
(NH4)2CrO4(y)
Cr2O3(s,g)
Cr
Cr(OH)3
(grayish green)
Cr(OH)4(g)
-
Cr2O7(o)
2-
CrO4(y)
2-
BaCrO4(s,y)
Ag2CrO4(s,brick)
PbCrO4(s,y)
Cr3+
CrO(O2)2
(blue)
O2，△
△
H+
△
H+
H+
H+
H+
ex.OH -
OH -
OH -
H2O2
Cl2
Br2
ClO-
Sn2+, Fe2+
SO3, H2S
I -(Cl-)
2-
S2O8
2-
H2O2
OH -
O(Et)2
Ag+
Ba2+
Pb2+
Ba2+
Pb2+
Ag+
H+
Cr2+
O2
Zn
Manganese (Mn)
Electron configuration: 4s23d5 4s13d5 3d5 3d4
Properties
Hard, brittle, lustrous silver-blue.
Uses
Added to steels to increase resistance to shock and wear.
Manganese (IV) oxide is used in dry cell batteries.
KMnO4 -> Strong oxidizer.
Essential to all living things.
Has a very wide range of oxidation states.
Ores:
Pyrolucite: MnO2
Manganite: MnO(OH)
Rhodochrosite: MnCO3
Manganese
pyrolusite: MnO2 软锰矿
All steels contain some Mn to prevent brittleness. Higher percentage of Mn will enhance hardness of the steel.
Important compounds: KMnO4(purple), Mn2+(light pink).
MnO2: dry cell industry. From high quality ore or made electrolytically from MnSO4 solution.
Glass decolorizer: presence of Fe(II) greenish color, add MnO2 to molten glass produces red-brown Mn(III), equalizes absorption! Selenium has similar properties.
http://www./
Preparation of the element
The most important ore is pyrolusite, manganese(IV) oxide.
Reduction of this ore by heating with aluminium gives an explosive reaction, and the oxide Mn3O4 must be used to obtain the metal. The latter is purified by distillation in vacuo just above its melting point (1517 K);
the pure metal can also be obtained by electrolysis of aqueous manganese(II) sulphate.
Property of the Metal
The metal looks like iron;
stable over various temperature ranges.
Although not easily attacked by air, it is slowly attacked by water
dissolves readily in dilute acids to give manganese (II) salts.
The stable form of the metal at ordinary temperatures is hard and brittle
hence manganese is only of value in alloys, for example in steels (ferroalloys) and with aluminum, copper and nickel.
Mn2+
2 Mn2+(aq) + 5 NaBiO3(s) + 14 H+(aq) 2MnO4-(aq) + 5 Bi3+(aq) + 5 Na+(aq) + 7H2O
Mn(V)
Disproportionation:
3MnO42-(aq) + 4H+(aq) 2MnO4-(aq) + MnO2(s) + 2H2O(l)
Mn(VI) common oxidizing agent
2 MnO4- + 5 SO32- + 6 H+ 2 Mn2+ + 5 SO42- + 3 H2O acidic
2 MnO4- + 3 SO32- + H2O 2 MnO2 + 3 SO42- + 2 OH- neutral
2 MnO4- + SO32- + 2 OH- 2 MnO42- + SO42- + H2O basic
decomposes (explosively on heating)
MANGANESE(VII) OXIDE,
THE MANGANATES(VII)
disproportionates on heating
manganate(VII) ion slowly oxidizes water
in concentrated alkali, manganese(VI) is more stable than manganese(VII)
in neutral or acid solution manganese(VI) disproportionates
MANGANESE(IV) OXIDE, MnO2
neutral
acidic
Lead (IV) oxide and concentrated nitric acid
Iron (Fe)
Electron configuration: 4s23d6 4s13d6 3d6 3d5
Ores:
Hematite: Fe2O3 (impure)
Limonite: FeO(OH).nH2O
Magnitite: Fe3O4
Siderite: FeCO3
Properties
Very abundant
Main oxidation states: II, III
Can be magnetized.
Uses
Not very useful in “pure” form, used to make steel by mixing with a small amount of carbon.
Iron oxides are used as pigments.
Has immense biological importance.
Preparation of the Element
In dilute nitric acid:
FeCl3
When the anhydrous solid is heated, it vaporizes to form first Fe2Cl6 molecules, then the monomer FeCl3 and finally FeCl2 and chlorine.
It fumes in air (with hydrolysis) and dissolves readily in water to give a yellow (dilute) or brown (concentrated) solution, which is strongly acidic.
FeCl3.6H2O yellow [FeCl2(H2O)4]Cl·2H2O
Coordination Compounds
[Fe(CN)6]3-, [Fe(CN)6]4-,
[Fe(SCN)(H2O)5]2+;
[Fe(H2O)6]2+; green
[Fe(H2O)6]3+; purple
Coordination Compounds II
Fe3+ ions and [Fe(CN)6]4- gets blue precipitate, prussian blue;
Fe2+ ions and [Fe(CN)6]3- gets blue precipitate, Turnbull’s blue.
KFe[Fe(CN)6]
test for a nitrate
When concentrated sulfuric acid (浓硫酸）is added to a nitrate in the presence of aqueous iron(II) sulfate, the nitrogen oxide liberated forms a brown complex [Fe(H2O)5NO]2+ which appears as a “brown ring‘ at the acid-aqueous interface——棕色环反应
Iron is the fourth most abundant element in the earth's crust.
Iron is obtained from the ores hematite (Fe2O3) and magnetite (Fe3O4) by reduction with coke.
The adult human body contains about 4 grams of iron, mostly in the form of hemoglobin.
Iron
Ferrocene: a sandwich compound
Chemistry of Ferrocene
Metallocenes undergo reactions similar to those of simple aromatic hydrocarbons.
Possible structures of [Fe( 5-C5H5)2]
Cobalt (Co)
Electron configuration: 4s23d7 3d8 3d7 3d6
Properties
Hard, brittle, lustrous blue-white.
Main oxidation state: II
Ores: Cobaltite CoAsS
Uses
Used as an alloy in magnets, jet engine parts.
Used in electroplating because of its attractive appearance, hardness and resistance to oxidation.
Vitamin B12 is a cobalt complex.
Cobalt salts have been used for centuries to produce brilliant blue colors in porcelain, glass, pottery and enamels.
The artificial isotope 60Co is used as a tracer.
The element
Cobalt is a bluish silvery metal, exhibits ferromagnetism.
Chemically it is somewhat similar to iron;
when heated in air it gives the oxides Co3O4 and CoO
but it is less readily attacked by dilute acids.
With halogens, the cobalt(II) halides are formed, except that with fluorine the (III) fluoride, CoF3, is obtained.
[Co(H2O)6]2+, pink
simple cobalt(III) cation cannot exist in aqueous solution (which it would oxidize to oxygen).
Observation:
ammoniacal solution of a cobalt(II) salt changed color on exposure to air. @1798
if cobalt(II) chloride was oxidized in presence of ammonia, the yellow product had the formula CoCl3. 6NH3
Alfred Werner was awarded the
Nobel Prize for chemistry in 1913)
@1890 ~1913
Founding of Coordination Chemistry
charcoal is a catalyst
If hydrogen peroxide is used as the oxidant, a red aquopentamminocobalt (III) chloride, [Co(NH3)5H2O]Cl3, is formed
treatment with concentrated hydrochloric acid gives the red chloropentamminocobalt(III) chloride, [Co(NH3)5Cl]Cl2
Low spin Co(III) compounds are very kinetically inert. We can even separate optical isomers!
Cobalt(III) oxide
Cobalt(III) oxide is obtained as a brown precipitate Co2O3.aq when cobalt(II) hydroxide is oxidized in alkaline conditions
or when a cobalt(III) is decomposed by aqueous alkali.
On heating it gives the black mixed oxide Co3O4. [ CoO· Co2O3 ]
Chemistry of Cobalt (II)
The precipitate
is often blue, but becomes pink on standing; it dissolves in excess
alkali to give the blue [Co(OH)4]2- ion, and in slightly alkaline
solution is easily oxidized by air to a brown solid of composition
CoIIIO(OH).
With sodium hydroxide solution and hydrogen peroxide etc.
excess sodium nitrite is added to a cobalt(II)
salt in presence of ethanoic acid (a strong acid would decompose the nitrite)
Na salt is very soluble, but K+,Rb+,Cs+,NH4+ salt are very insoluble in H2O
CoCl2
Anhydrous cobalt(II) chloride is blue, and the solid contains octahedrally-coordinated cobalt
the hydrated salt CoCl2. 6H2O is pink, with each cobalt surrounded by four water molecules and two chloride ions in a distorted octahedron.
TESTS FOR COBALT
For a cobalt(II) salt the precipitation of the blue pink cobalt(II) hydroxide by alkali, or precipitation of black cobalt(II) sulfide by hydrogen sulphide provide useful tests; the hydroxide is soluble in excess alkali and is oxidized by air to the brown 'CoO(OH)'.
Addition of excess potassium nitrite acidified with ethanoic acid gives a precipitate of the potassium hexanitro-cobaltate(III), K3[Co(NO2)6]
Decomposition of most cobalt(III) complexes by boiling with alkali gives a brown precipitate of the hydrated oxide Co2O3 .aq. This compound will quantitatively oxidize iodide to iodine.
Nickel (Ni)
Electron Configuration: 4s23d8 3d9 3d8 3d7
Nickel occurs more abundantly than cobalt.
Properties
Hard, malleable and ductile, silvery-white, capable of taking on a high polish.
Main oxidation state: II
Soluble in all acids except concentrated nitric acid.
Nickel carbonyl is very toxic.
Uses:
Used extensively in alloys: stainless steel, coinage metals.
The Nickel is 25% nickel and 75% copper
Electroplated Ni gives a protective coat for other metals.
Rechargeable batteries (Ni-Mh, NiCd)
Preparation of the element
The metal obtained by reduction can be purified by the Mond process, in which it is heated to 320 K with carbon monoxide to give the pure, volatile tetracarbonyl Ni(CO)4; the latter when heated to 500 K gives the pure metal and carbon monoxide is recovered
[Ni(H2O)6]2+, green
Ni(OH)2, green
NiO, black
Fe2+
H+
Fe(OH)2
NH3 · H2O
Fe2+ 淡绿
OH-
Fe(OH)2(s,白)
O2
Fe(OH)3(s,红棕)
HCl
Fe3+
Co(NH3)63+红
O2
Co(NH3)62+黄
NH3·H2O
Co(OH)Cl(s,蓝)
NH3· H2O
Co2+(Cl-) 粉红
OH-
Co(OH)2(s,粉红)
NaClO
Co(OH)3(s,红棕)
浓HCl
Co2+（CoCl4,蓝）
Ni(NH3)62+蓝
NH3 · H2O
Ni2(OH)2SO4浅绿
NH3 · H2O
Ni2+(SO42-) 淡绿
OH-
Ni(OH)2(s,果绿)
NaClO
NiO(OH)(s,黑)
浓HCl
Ni2+
鉴定：Fe(NCS)n3-n 血红 Co(NCS)42- 天蓝 Ni(DMG)2(s,鲜红)
[KFe(CN)6Fe]x(s,蓝) （丙酮）
2-
Copper (Cu)
Electron configuration: 4s13d10 3d10 3d9 3d8
Properties
Malleable and ductile; reddish.
Good conductor of heat and electricity.
Main oxidation state: II
Ores:
Cuprite: Cu2O
Bornite Cu6FeS4
Azurite: Cu3(CO3)2(OH)2
Covellite: CuS
Uses
Electrical wires and switches, plumbing, cooking vessels.
Used as an alloy: brass (Cu/Zn), Bronze (Cu/Sn), Coinage material (Cu/Ni).
Copper sulphate is used widely as an agricultural poison and as an algicide in water purification.
Halides
anhydrous fluoride CuF2 is white, the chloride yellow and the bromide almost black
CuI
Copper(I) oxide, Cu2O
orange-yellow precipitate
Copper(I) chloride, CuCl
white solid, insoluble in water.
Vapor state : dimer of formula Cu2Cl2
concentrated hydrochloric acid
Properties of CuCl
The solid readily dissolves chemically in concentrated hydrochloric acid, forming a complex, and in ammonia as the colorless, linear, complex cation [H3N Cu NH3]+
ammoniacal copper(I) chloride is a good absorber for carbon monoxide, forming CuCl . CO . H2O, and as such is used in gas analysis.
explosive when dry
not in excess)
Copper(l) chloride, bromide and cyanide were used by Sandmeyer
to introduce a chlorine, a bromine atom and a cyanide group
respectively into a benzene ring by addition to the phenyl diazonium
salt.
TESTS FOR COPPER COMPOUNDS
to add hexacyanoferrate(II) (usually as the potassium salt) when a chocolate-brown precipitate of copper(II) hexacyanoferrate(II) is obtained:
chocolate-brown
Inert and Labile Complexes
Complexes in which ligands are rapidly replaced by others are called labile complexes (shorter than 1 min; @ 25 oC, 0.1M)
Slow ---- inert complexes
Labile complexes
All complexes in which the central metal atom contains electrons in eg orbitals
[Ga(C2O4)3]3-, d10; [Cu(H2O)6]2+, d9
[Co(NH3)6]2+,d7 [Ni(H2O)6]2+,d8
[Fe(H2O)6]3+, d5
All complexes that contain less than three d-electrons
[Ti(H2O)6]3+, d1; [V(phen)3]3+, d2
[CaEDTA]2-, d0
Inert Complexes
Octahedral d3 complexes, plus low spin d4, d5 and d6 systems
[Cr(H2O)6]3+,d3
[Fe(CN)6]3-,d5
[Co(NO2)6]3-,d6
[PtCl6]2-,d6
Why
Comparison of the CFSE of a coordination compound and of its activated complex
If the CFSE of the compound is much larger than that of the activated complex, the compound will react slowly
Small fast
Cu Ag Au
Reaction with O2
Au，Ag O2 no reaction，
But if forming precipitate or coordination compound, reaction can happen。
Can not use Cu as NH3 container。
O2
Silverware tarnishes。
Cu,Ag,Au soluble in oxidizing acids。
Ag compounds：
Mostly insoluble
soluble：AgNO3, AgF, AgClO4
insoluble：AgCl, AgBr, AgI, AgCN, AgSCN,
Ag2S, Ag2CO3, Ag2CrO4。
Low thermal stability （light, heat）
color
AgCl AgBr AgI Ag2O Ag2CrO4 Ag2S
white pale yellow dark blood black
yellow brown red
Ag(I) reaction
Silvering mirror reaction：
①
②
③
④
Identification of S2O32-：
)
S(s,
Ag
SO
H
O
H
)
(s,
O
S
Ag
2
4
2
2
3
2
2
b
w
+
+
Relative stability of Ag+ complexes
7
2
3
10
67
.
1
)
)
Ag(NH
(
×
=
+
13
3
=
-
2
3
2
10
9
.
2
)
)
O
Ag(S
(
×
20
2
10
48
.
2
)
Ag(CN)
(
×
=
-
Identification of Ag+
I-
HNO3
Mercury differs from zinc and cadmium in a number of significant ways.
It is used in thermometers, barometers, electrical contacts, and in some types of electrochemical cells. Mercury vapor is used in fluorescent lamps.
Dental amalgam is an alloy of mercury, silver, and tin, and is used for dental restorations.
Long-term exposure to mercury can present a serious health hazard.
Mercury can react with and inactivate sulfur-containing enzymes.
Mercury
Mercury - Hg
Mercury is a heavy, silvery, liquid metal.
Mercury occurs very rarely free in nature, but can be found in ores.
Mercury easily forms alloys with other metals such as gold, silver and tin, which are called amalgams.
Its ease in amalgamating with gold is made use of in recovering gold from its ores.
The metal is widely used in making advertising signs, mercury switches, and other electrical apparatus.
It is used in laboratory work for making thermometers, barometers, diffusion pumps and many other instruments.
Other uses are in pesticides, dental work, batteries and catalysts.
It is a virulent poison, readily absorbed through the respiratory tract, the gastrointestinal tract or through the skin. It is a cumulative poison and dangerous levels are readily attained in air. It is always handled with the utmost care.
Mercury is stable with air and water.
It is a rather poor conductor of heat compared to other metals, and a fair conductor of electricity
Reactions of Hg
2. S2-
1. OH-
If Hg2+ comes from Hg(NO3)2：
HgS -0.758V Hg2S -0.598V Hg
HgS： =1.6×10-52
Therefore：Hg(I) is stable in solution，but disproportionate when precipitate or complex forms。
Aminomercuric chloride
3. Reaction with NH3
In general, all these products are obtained in proportions which depend on the concentrations of NH3 and NH4+ and on the temperature, but more or less pure products can be prepared by suitably
adjusting the conditions
[Hg(NH3)2Cl2], descriptively known as “fusible white precipitate”,
the amide [Hg(NH2)Cl], “infusible white precipitate”
adding will shift eq to right
-
2
4
Nessler
HgI
4
NH
+
reagent，
is
-
I
reaction with I-，SCN-
2
-
2
)
(s,
HgI
2I
Hg
+
+
gold
+
-
2
4
-
2
)
(aq,
HgI
2I
HgI
colorless
2
2
2
-
2
2
Hg
HgI
)
(s,
I
Hg
2I
Hg
+
+
+
green
2
-
O
3H
7I
+
+
2
-
-
2
4
4
)
I(s,
NH
Hg
Hg
4OH
]
2[HgI
NH
+
+
+
brown
O
Can also use HgCl2 to identify Sn2+
Identification of Hg2+
)
(aq,
Hg(SCN)
2SCN
Hg(SCN)
-
2
4
-
2
colorless
+
(s)
Hg(SCN)
2SCN
Hg
2
-
2
+
+
Ruthenium and Osmium
All the platinum group metals are isolated from “platinum concentrates” which are commonly obtained either from “anode slimes” in the electrolytic refining of nickel and copper,
or as “converter matte” from the smelting of sulfide ores.
Preparation of the elements
Ru and Os are usually removed by distillation of their tetroxides immediately after the initial dissolution with hydrochloric acid and chlorine.
Collection of the tetroxides in alcoholic NaOH and aqueous HCl respectively yields OsO2(NH3)4Cl2 and (NH4)3RuCl6 from which the metals are obtained by ignition in H2.
The oxides of ruthenium and
osmium
Oxidation state +8 +4
Ru RuO4 RuO2
Os OsO4 OsO2
Rhodium and Iridium
Pt square planar complexes

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