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HYDROMETALLURGY

Hydrometallurgy is a method for obtaining metals from their ores. It is a technique within the field of extractive metallurgy involving the use of aqueous chemistry for the recovery of metals from ores, concentrates, and recycled or residual materials.
Hydrometallurgy is typically divided into three general areas:

• Leaching
• Solution concentration and purification
• Metal or metal compound recovery

Leaching

Leaching involves the use of aqueous solutions to extract metal from metal bearing materials which is brought into contact with a material containing a valuable metal. The lixiviant solution conditions vary in terms of pH, oxidation-reduction potential, presence of chelating agents and temperature, to optimize the rate, extent and selectivity of dissolution of the desired metal component into the aqueous phase. Through the use of chelating agents, one can selectively extract certain metals. Such chelating agents are typically amines of schiff bases.

Metal Recovery

Metal recovery is the final step in a hydrometallurgical process. Metals suitable for sale as raw materials are often directly produced in the metal recovery step. Sometimes, however, further refining is required if ultra-high purity metals are to be produced. The primary types of metal recovery processes are electrolysis, gaseous reduction, and precipitation. For example, a major target of hydrometallurgy is copper, which is conveniently obtained by electrolysis. Cu²⁺ ions reduce at mild potentials, leaving behind other contaminating metals such as Fe²⁺ and Zn²⁺.

Solution concentration and purification

After leaching, the leach liquor must normally undergo concentration of the metal ions that are to be recovered. Additionally, undesirable metal ions sometimes require removal.^[1]

• Precipitation is the selective removal of a compound of the targeted metal or removal of a major impurity by precipitation of one of its compounds. Copper is precipitated as its sulfide as a means to purify nickel leachates.
• Cementation is the conversion of the metal ion to the metal by a redox reaction. A typical application involves addition of scrap iron to a solution of copper ions. Iron dissolves and copper metal is deposited.
• Solvent Extraction
• Ion Exchange
• Gas reduction. Treating a solution of nickel and ammonia with hydrogen affords nickel metal as its powder.
• Electrowinning is a particularly selective if expensive electrolysis process applied to the isolation of precious metals. Gold can be electroplated from its solutions.

 HYDROMETALLURGY OF COPPER

CHEMISTRY ASSIGNMENT

FOR ROLL NO. 4980-4991

 Diborane

Diborane is the chemical compound consisting of boron and hydrogen with the formula B₂H₆. It is a colorless and highly unstable gas at room temperature with a repulsively sweet odor. Diborane mixes well with air, easily forming explosive mixtures. Diborane will ignite spontaneously in moist air at room temperature. Synonyms include boroethane, boron hydride, and diboron hexahydride.
Diborane is a key boron compound with a variety of applications. The compound is classified as "endothermic", meaning that its heat of formation, ΔH°_f is positive (36 kJ/mol). Despite a high thermodynamic instability, diborane is surprisingly nonreactive for kinetic reasons, and it is known to take part in an extensive range of chemical transformations, many of them entailing loss of dihydrogen.

SYNTHESIS OF DIBORANE

Extensive studies of diborane have led to the development of multiple syntheses. Most preparations entail reactions of hydride donors with boron halides or alkoxides. The industrial synthesis of diborane involves the reduction of BF₃ by sodium hydride, lithium hydride or lithium aluminium hydride:^[8]

8 BF₃ + 6 LiH → B₂H₆ + 6 LiBF₄

Two laboratory methods start from boron trichloride with lithium aluminium hydride or from boron trifluoride ether solution with sodium borohydride. Both methods result in as much as 30% yield:

4 BCl₃ + 3 LiAlH₄ → 2 B₂H₆ + 3 LiAlCl₄

4 BF₃ + 3 NaBH₄ → 2 B₂H₆ + 3 NaBF₄

Older methods entail the direct reaction of borohydride salts with a non-oxidizing acid, such as phosphoric acid or dilute sulfuric acid.

2 BH₄⁻ + 2 H⁺ → 2 H₂ + B₂H₆

Similarly, oxidation of borohydride salts has been demonstrated and remains convenient for small scale preparations. For example, using iodine as an oxidizer:

2 NaBH
4 + I
2 → 2 NaI + B
2H
6 + H
2

Another small-scale synthesis uses potassium hydroborate and phosphoric acid as starting materials.

Structure and bonding

Diborane adopts a D_2h structure containing four terminal and two bridging hydrogen atoms. The model determined by molecular orbital theory indicates that the bonds between boron and the terminal hydrogen atoms are conventional 2-center, 2-electron covalent bonds. The bonding between the boron atoms and the bridging hydrogen atoms is, however, different from that in molecules such as hydrocarbons. Having used two electrons in bonding to the terminal hydrogen atoms, each boron has one valence electron remaining for additional bonding. The bridging hydrogen atoms provide one electron each. Thus the B₂H₂ ring is held together by four electrons, an example of 3-center 2-electron bonding. This type of bond is sometimes called a 'banana bond'. The lengths of the B-H_bridge bonds and the B-H_terminal bonds are 1.33 and 1.19 Å respectively, and this difference in the lengths of these bonds reflects the difference in their strengths, the B-H_bridge bonds being relatively weaker. The weakness of the B-H_bridge vs B-H_terminal bonds is indicated by their vibrational signatures in the infrared spectrum, being ~2100 and 2500 cm⁻¹, respectively. The structure is isoelectronic with C₂H₆²⁺, which would arise from the diprotonation of the planar molecule ethene. Diborane is one of many compounds with such unusual bonding.
Of the other elements in Group IIIA, gallium is known to form a similar compound, digallane, Ga₂H₆. Aluminium forms a polymeric hydride, (AlH₃)_n, although unstable Al₂H₆ has been isolated in solid hydrogen and is isostructural with diborane.


VIKAS BHATI


NOTES ON GENERAL PRINCIPLES OF METALLURGY

CHEMISTRY ASSIGNMENT 
ROLL NO. 4940-4955

 ALLOTROPHY

Allotropy  is the property of some chemical elements to exist in two or more different forms, in the same physical state, known as allotropes of these elements. Allotropes are different structural modifications of an element;^[1] the atoms of the element are bonded together in a different manner. For example, the allotropes of carbon include diamond (where the carbon atoms are bonded together in a tetrahedral lattice arrangement), graphite (where the carbon atoms are bonded together in sheets of a hexagonal lattice), graphene (single sheets of graphite), and fullerenes (where the carbon atoms are bonded together in spherical, tubular, or ellipsoidal formations). The term allotropy is used for elements only, not for compounds

ALLOTROPHS OF CARBON

1. Diamond – an extremely hard, transparent crystal, with the carbon atoms arranged in a tetrahedral lattice. A poor electrical conductor. An excellent thermal conductor.
2. Lonsdaleite – also called hexagonal diamond.
3. Q-carbon – a ferromagnetic, tough, and brilliant crystal structure that is harder and brighter than diamonds.
4. Graphite – a soft, black, flaky solid, a moderate electrical conductor. The C atoms are bonded in flat hexagonal lattices (graphene), which are then layered in sheets.
5. Linear acetylenic carbon (Carbyne)
6. Amorphous carbon
7. Fullerenes, including Buckminsterfullerene, a.k.a. "buckyballs", such as C₆₀.
8. Carbon nanotubes – allotropes of carbon with a cylindrical nanostructure

FULLERENE

These are small molecules of carbon in which the giant structure is closed over into spheres of atoms (bucky balls) or tubes (sometimes caled nano-tubes). The smallest fullerene has 60 carbon atoms arranged in pentagons and hexagons like a football. This is called Buckminsterfullerene.
The name 'buckminster fullerene' comes from the inventor of the geodhesic dome (Richard Buckminster Fuller) which has a similar structure to a fullerene. Fullerenes were first isolated from the soot of chimineys and extracted from solvents as red crystals.
The bonding has delocalised pi molecular orbitals extending throughout the structure and the carbon atoms are a mixture of sp2 and sp3 hybridised systems.
Fullerenes are insoluble in water but soluble in methyl benzene. They are non- conductors as the individual molecules are only held to each other by weak van der Waal's forces.

Structure As the molecule is totally symmetrical with all bond lengths and angles being equal, it is likely/inevitable that the hybridisation of the carbon atoms is somewhere between that of sp2 and sp3. Another example of a theory (hybridisation in this case) having to be modified to accomodate the observed experimental data.

Property	Explanation	Fullerene structure
Soft and slippery	Few covalent bonds holding the molecules together but only weak Vander Waals forces between molecules.	Click on the image with the left mouse button and drag to get a different view. (if you can't see the image you have to download chime)
Brittle	Soft weak crystals typical of covalent substances
Electrical insulator	No movement of electrons available from one molecule to the next. The exception could be the formation of nano-tubes that are capable of conducting electricity along their length. These are the subject of some experiments in micro electronics
Insoluble in water.	There are only very weak Van der Waal's attractions between the carbon atoms and the water molecules whereas the carbon atoms are bonded very tightly to one another in the molecules.
Low m.p. solids	Typical of covalent crystals where only Van der Waal's interactions have to be broken for melting.

^{ALLOTROPHY IN SULPHUR}

^{The allotropes of sulfur refers to the many allotropes of the element sulfur. In terms of large number of allotropes, sulfur is second only to carbon.[1] In addition to the allotropes, each allotrope often exists in polymorphs, delineated by Greek prefixes (α, β, etc.).}

1. Rhombic Sulphur:

It is an allotropic form of sulphur which is stable below 96 one molecule of rhombic sulphur contains 8-atoms i-e 8_g. The crystal of rhombic sulphur has octahedral structure.

 Properties:

• It is consist of pale yellow crystals.
• It melts at 110℃.
• It is insoluble in water and soluble in carbon disulphide.
• It is stable at room temperature.
• Its specific gravity is 208g/cm3.

2. Monoclinic Sulphur:

It is the allotropic form of sulphur which is stable between 96 to 119 a molecule of monoclinic sulphur consists of eight sulphur atoms i-e 8g, but is different from rhombic sulphur in the arrangement of atoms.

Properties:

• It is stable from 96℃-119℃.
• Its melting point is 119℃.
• It is soluble in carbon disulphide.
• Its one molecule consist of 8 atoms.
• It is found as pale yellow needle shaped crystals.

3. Plastic Sulphur:

It is a non crystalline allotropic form of sulphur, it can be stretched like a rubber, it is unstable and changes into rhombic sulphur on slight heating even at room temperature it also changes.

4.Colloidal Sulphur

 

This type of sulphur is prepared by passing hydrogen sulphide through a cooled saturated solution of sulphur dioxide in water, or by adding a solution of sulphur and alcohol in water. Colloidal sulphur is soluble in carbon disulphide. It is used in medicine.

5.Milk of Sulphur

Milk of sulphur is prepared by the action of dilute hydrochloric acid on ammonium sulphide. Milk of sulphur is also prepared by boiling roll sulphur with an aqueous solution of calcium hydroxide. The mixture is then filtered and dilute hydrochloric acid is added to the filtrate to get milk of sulphur.

Milk of sulphur is non-crystalline and white in color.

 ALLOTROPHS OF PHOSPHORUS

Elemental phosphorus can exist in several allotropes, the most common of which are white and red solids. Solid violet and black allotropes are also known. Gaseous phosphorus exists as diphosphorus and atomic phosphorus

White phosphorus

White phosphorus, yellow phosphorus or simply tetraphosphorus (P₄) exists as molecules made up of four atoms in a tetrahedral structure. The tetrahedral arrangement results in ring strain and instability. The molecule is described as consisting of six single P–P bonds. Two different crystalline forms are known. The α form, which is stable under standard conditions, has a body-centered cubic crystal structure. It transforms reversibly into the β form at 195.2 K. The β form is believed to have a hexagonal crystal structure.

Red phosphorus

Red phosphorus may be formed by heating white phosphorus to 300 °C (482 °F) in the absence of air or by exposing white phosphorus to sunlight. Red phosphorus exists as an amorphous network. Upon further heating, the amorphous red phosphorus crystallizes. Red phosphorus does not ignite in air at temperatures below 240 °C, whereas pieces of white phosphorus ignite at about 30 °C. Ignition is spontaneous at room temperature with finely divided material. Heating red phosphorus in the presence of moisture creates phosphine gas, which is both highly flammable and toxic

Black phosphorus 

Black phosphorus is the thermodynamically stable form of phosphorus at room temperature and pressure. It is obtained by heating white phosphorus under high pressures (12,000 atmospheres). In appearance, properties, and structure, black phosphorus is very much like graphite with both being black and flaky, a conductor of electricity, and having puckered sheets of linked atoms. Phonons, photons, and electrons in layered black phosphorus structures behave in a highly anisotropic manner within the plane of layers, exhibiting strong potential for applications to thin film electronics and infrared optoelectronics. Light absorption in black phosphorus is sensitive to the polarization of incident light, film thickness, and doping. Black phosphorus photo-transistors exhibit hyper-spectral detection attributes in the infrared and visible spectrum.
Black phosphorus has an orthorhombic structure and is the least reactive allotrope, a result of its lattice of interlinked six-membered rings where each atom is bonded to three other atoms.Black and red phosphorus can also take a cubic crystal lattice structure. A recent synthesis of black phosphorus using metal salts as catalysts has been reported.

Diphosphorus

The diphosphorus allotrope (P₂) can normally be obtained only under extreme conditions (for example, from P₄ at 1100 kelvin). In 2006, the diatomic molecule was generated in homogenous solution under normal conditions with the use of transition metal complexes (for example, tungsten and niobium).^[23]
Diphosphorus is the gaseous form of phosphorus, and the thermodynamically stable form between 1200 °C and 2000 °C.

THANKING YOU 

VIKAS BHATI

 

CHEMISTRY ASSIGNMENT

ROLL NO. 4929-4939

 Electronegativity: 

 The concept of electronegativity was first introduced by Linus Pauling in 1930 as a means of describing bond energies. 
1 When we consider the formation of a covalent bond the attraction of the two nuclei for the electrons is not the same and the electron pair is closer to one of the nuclei.

  2 This tendency of a atom in a molecule to attract the electron density towards itself (or the reluctance to release the electron density) is called electronegativity. 

 3 The electronegativity of a atom depends on the size, effective nuclear charge, oxidation state and the hybridization of the atom in the molecule. It therefore depends on the structure of the molecule and the atom. 

 4 If the size of the atom is small and it has almost closed shell electronic configuration the tendency to attract electrons increase and the atom is highly electronegative.  

 Variation of electronegativity

  1.Electronegativity increases from left to right in a period. 

 2 Electronegativity decreases down the group.   

Many scales of electronegativity have been proposed. One of the most commonly used   scale is the Pauling’s scale. 

 Pauling’s Electronegativity:

 For any covalent bond, the bond energy of the heteronuclear bond E(A-B) is greater than the bond energy of the sum of homonuclear bonds E(A-A) and E(B-B). This excess bond energy can be attributed to ionic contribution in the bonds. He treated this ionic contribution by the equation.
       E(A-B) = [E(A-A)×E(B-B)]1/2 + 96.48(ΧA - ΧB)2
E(A-B) is expressed in kJ mol-1  and ΧA - ΧB is  the difference in "electronegativity" between the two elements. The largest electronegativity difference exists between Cs and F. The value of F was set arbitrarily at 4.0and electronegativity values of all other elements found relative to it

 Allred and Rochow electronegativity:

 Allred and Rochow gave the electronegativity values by considering the electrostatic force exerted by effective nuclear charge, Zeff, on the valence electron. They gave the equation: XAR = (3590 x Zeff/r2cov) + 0.744 

 Mullikan electronegativity: Mullikan proposed that two energies associated with the atom i.e. the electron affinity EAv and the ionization potential IEv  should be a measure of electronegativity. The Mulliken electronegativity, ΧM is related to the electron affinity EAv and the ionization potential IEv by the equation:
ΧM = (IEv + EAv)/2
          The subscript v denotes a specific valence state.
          The Mulliken electronegativity ΧM can be expressed on the Pauling scale by the relationship given below if the values of IE and EA are in MJ mol-1:
ΧM = 3.48[((IEv + EAv)/2) - 0.602] 

BY

VIKAS BHATI

  C.B.C.S. SYLLABUS

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