Carbon is the essential constituent of all organic matter, while silicon is an important constituent of inorganic matter.

Electronic Configuration of Group IV(14)  Elements:

Notes:

• The similarities in the properties of these elements &rise due to the similarity in their outer configuration ns² np².
• The differences and gradation in properties arise due to their atomic radii and the number of electrons in the penultimate shell.  Carbon has 2 electrons. Silicon has 8 electrons and other elements have 18 electrons In their penultimate shell.
• This causes a difference in the physical properties of Carbon and Silicon on one hand and Ge, Sn, Pb on the other hand.
• Silicon differs from Carbon having vacant 3d orbitals.  Hence it is able to show covalency greater than four

Position of Silicon in the Periodic Table:

• Atomic Number of silicon is 14. Detailed Electronic Configuration: 1s², 2s² 2p⁶, 3S², 3p_x¹ , 3 p_y¹, 3 p_z⁰
• The last electron enters into ‘p’ sub-shell, therefore, Silicon belongs to ‘p’  block normal non-metallic element.
• The total number of orbits = 3. Therefore, silicon belongs to 3rd period.
• It has four electrons in the valence shell. Therefore, it is placed in the Group IV-A (14) of the periodic table. This is transition position between metals and non-metals.

Electronic Configuration of Silicon:       

Ground State:

Excited state:

Hybridized State:

Oxidation States of Silicon:

• Electronic configuration of ‘Si’ in ground state is 1s², 2s² 2p⁶, 3S²  3p_x¹   3 p_y¹
• Electronic configuration of ‘Si’ in excited state is 1s², 2s² 2p⁶, 3S²  3p_x¹   3 p_y¹  3 p_z¹
• Silicon undergoes sp³ hybridization and forms four half-filled sp3 hybrid orbitals which are used to form four covalent bonds.
• Because of the high Ionization potential existence of (Si⁴⁺) ion is unlikely.  Silicon shows +4 oxidation states by sharing four outer electrons with strongly electronegative elements like halogens and oxygen.

Silicon exhibits -4 oxidation state in silicides.  Silicides are formed at high temperature

Silicon has vacant 3d orbitals. Hence it is possible for silicon to show a covalency greater than four.

Oxidation States of Group IV-A (14) Elements:            

Since they possess high ionization energies, the existence of (4+) ions is unlikely. Their electronegativity values are also low.  Therefore (4-) ions also do not form normally. Only carbon and Silicon exhibit (4-) oxidation state. They show (4+) oxidation state by forming covalent bonds with strongly electronegative elements.  For this purpose sp³ hybrid orbitals are used.

They also show (2+) oxidation state by sharing their two unpaired np² electrons, (ground state).  From carbon to lead, with the increase in nuclear ‘charge, stability of ns² electron pair increases. Hence the stability of lower oxidation state (2+) increases and that of higher oxidation state (4+) decreases from carbon to lead. This is due to inert pair effect.

Metallic and Non-metallic Character of Group IV-A (14) Elements:

The first two elements carbon and silicon are distinctly non-metallic. Germanium is a metalloid. Tin and Lead are well-defined metals. Silicon has most of the characteristics of non-metals. But it is a semiconductor at room temperature, therefore- sometimes it is considered as a metalloid. Metallic nature or electropositive nature increases from carbon to lead.

Electronegativity and Ionisation Potential (Ionization Enthalpy) of Group IV-A (14) Elements:

With the increase in atomic number i.e. from carbon to lead, as the number of shells increases, their atomic radii also increases. Ionization potential and electronegativity decrease with an increase in their atomic radii.

Hardness:

Because of small atomic radii, in carbon and silicon there are strong inter-atomic forces of attraction and hence they are hard solids with high melting and boiling points. Germanium is quite hard, but tin and lead are soft silvery-white metals.

Allotropy:

All the elements except Ge and Pb show allotropy

Catenation:

Carbon can combine with other carbon atoms to form quite large carbon chains. Silicon has a smaller capacity to form chains.  Germanium has still a smaller tendency while tin and lead have hardly any tendency to do so.

Silicon

It derives its name from the Latin word Silex meaning hard stone. Next to oxygen, it is the most abundant element in the earth’s crust. It forms 26% of the mass of the earth’s crust and is widely spread in nature.

Occurrence:

Silicon does not occur in nature in the free state. It occurs in nature as Silica (Si0₂) in the form of sand, quartz, flint, agate, and opal. As Natural complex silicates

• Feldspar : K₂O . A₂O₃ . 6 SiO₂
• Mica:    3K₂0 . 3Al₂O₃ . 6 SiO₂ . 2 H₂O
• Kaolin or China Clay : Al₂O₃ . 2 SiO₂ . 2 H₂0
• Asbestos: CaO. 3MgO. 4SiO₂

Physical Properties of Silicon:

• The physical properties of silicon are as follows.
• It exists in two allotropic forms – amorphous and crystalline.
• The amorphous form is common. It is a brown powder with a specific gravity of 2.35.
• The crystalline form is steel grey coloured with specific gravity 2.50.
• Though inactive, amorphous silicon is more reactive than crystalline form.
• Its melting point is 1693 K (1420°C) and its boiling point is 2873 K (2600°C).
• Both the forms are poor conductors of electricity but conductivity increases with temperature. (it is a semiconductor).
• Its crystals are hard enough to scratch glass.
• It is insoluble in water and in any single acid. But a mixture of hydrofluoric acid and nitric acid reacts with it.  It reacts with fused alkalies or hot boiling solutions of alkalies.

Preparation of Silicon:

Preparation of Amorphous Silicon from Silica:

Finely powdered pure silica or quartz is mixed with a requisite quantity of magnesium powder. The mixture is then heated in a fire clay crucible in the absence of air. The reduction of silica takes place as

SiO₂ +   2Mg   →   2MgO   +    Si

After cooling, the product is treated with dil. HCI which dissolves MgO and unreacted Mg.

MgO    +  2HCI     →   MgCl₂  +   H₂O

Mg       +  2HCI       →    MgCl₂  +   H₂O

The insoluble mass is then treated with hydrofluoric acid. The unchanged SiO₂, is converted into volatile SiF₄.

SiO₂ +    4HF     →    SiF₄   +   2 H₂0

Si remains unaffected by HCl and HF. It is then washed with water and dried in a current of H₂ gas. The process gives silicon as a brown amorphous powder.

Preparation of Amorphous Silicon from Silicon Tetrachloride:

Pure sample of amorphous silicon can be prepared by passing vapours of silicon tetrachloride over molten sodium metal in an inert atmosphere.

SiCl₄        + 4Na      →      Si      +    4NaCl

The product is washed with water to remove NaCl and unreacted sodium. The product is dried in a current of H₂ gas. Silicon is left behind as brown amorphous powder.

Preparation of Crystalline Silicon from Potassium Silicofluoride:

A mixture of potassium fluosilicate or potassium silicofluoride and excess of aluminium powder is heated at 1273 K.

3K₂SiF₆ +  4Al  →   3Si  +  6KF +  4AIF₃ ­

AlF₃ volatilises off. Silicon formed dissolves in molten aluminium. As the molten mass cools, Si crystallises out. The mass is then heated with dil. HCI to remove unreacted aluminium.

2Al   +    6HCl      →  2AICl₃  +   3H₂O

It is then washed with water. (KF is water soluble). The process gives lustrous grey crystalline Si.

Reactions of Silicon:

Reaction with alkalies:

it dissolves in the aqueous hot alkali solution to form alkali silicate and hydrogen.

Si + 2 NaOH  + H₂O →  Na₂SiO₃ (Sodiumsilicate) +   2H₂ ↑

Reaction with Sodium Carbonate or Washing Soda: 

When fused with Sodium carbonate, it forms sodium silicate by displacing carbon and elementary carbon is set free.

Na₂CO₃   +     Si    →  Na₂SiO₃ (Sodiumsilicate)  +  C

Reaction with Metals:

At the temperature of the electric furnace it directly combines with Magnesium, Copper and Chromium forming Silicides.

Reaction with Halogens:

it burns spontaneously in fluorine at room temperature forming Silicon tetrafluoride.

Si       +   2F₂        →            SiF₄   (Silicontetrafluoride)

Si       +   2Cl₂      →           SiCl₄  (Silicontetrachloride)

Action of Steam:

When steam is passed over red-hot silicon, silicon dioxide is obtained with evolution of hydrogen gas

Si       +   2H₂O    →       SiO₂ +   2H₂ ­↑

Uses of Silicon:

• It is used as deoxidiser in metallurgy in making castings of steel, copper and bronze.
• Its major use is in the semiconductors which are the main components of transistors. Transistors are employed in a variety of electronic devices such as Radio, T.V. sets, Computers, Solar batteries.
• It is used in the preparation of important alloys.
• Ferro-silicon:  it is used as a deoxidizer in the manufacture of steel.
• Silicon-steel: Steel containing 5% Si is soft and magnetic. It is used for making magnetic cores.
• Silicon-bronze: It is used in the manufacture of telegraphic and telephone wires.
• It is used in the preparation of refractory materials like Crucibles.

Simple Silicates:

Silicates are compounds containing silicon, oxygen and metal or metals. They contain Si-O-Si linkage or Si-O bond in their structure.

Simple silicates can be considered as the metal salts of orthosilicic acid H₂SiO₄. Simple silicates are orthosilicates consisting of metal cations such as Na⁺ , Mg²⁺, Zn²⁺ etc. and silicate anion (SiO₄)^4-. Example: Mg₂(SiO₄). All silicates contain tetrahedral (SiO₄)^4- units.

The silicates having only discrete (SiO4)^4- units are considered as simple silicates.  The silicates containing two or more (SiO4)^4- units linked through oxygen atom are generally considered as complex silicates. e.g. pyrosilicates having (Si₂O₇)^6- anion.

(SiO₄)^4- Unit Present in Silicates have Tetrahedral Geometry:

Electronic, configuration of silicon atom in its ground state is 1s², 2s² 2p⁶, 3S², 3p_x¹ , 3 p_y¹, 3 p_z⁰. At the time of combination, one of the 3s electron gets promoted to the empty 3p. This excited state of silicon has electronic configuration 1s², 2s² 2p⁶, 3S¹, 3p_x¹ , 3 p_y¹, 3 p_z¹

Si undergoes sp³ hybridisation and forms four sp³ hybrid orbitals which are directed to the corners of regular tetrahedron at the centre of which there is a Si atom.

The electronic configuration of oxygen atom is 1s², 2s² , 2p_x² , 2 p_y¹, 2 p_z¹ . Oxygen has two unpaired electrons in two of its ‘p’ orbitals (2p_y, 2p_z)

Each sp³ hybrid orbital of silicon atom overlaps with one of the 2p orbital (2p_y, 2p_z) of the oxygen atom that has the unpaired electron and forms a covalent Si-O bond. Si-O covalent bond is sp³-p- sigma bond. Si forms such four bonds in (SiO4)^4- units.

Thus silicon completes its octet by forming four covalent bonds with four oxygen atoms. Each oxygen atom is still short of one electron to complete the octet. In order to complete the octet, the oxygen atoms pick up one electron each from some metal. In this process, oxygen atom becomes -1.This gives rise to the anion (SiO4)^4-.

Since the formation of (SiO4)^4- unit involves sp³ hybridisation of Si, it has tetrahedral structure.  Silicon atom lies at the centre and four oxygen atoms lie at the four corners of the regular tetrahedron.

Diagram :

Strength of Si-O Bond in Simple Silicates:

Si and O bond in silicates is covalent. The electronegativity difference between silicon (EN = 1.8) and Oxygen (EN = 3.5) is 1.7 pauling. Due to this Si-O bond becomes polar.  So it gets an ionic character (about 30%) Because of this ionic character the bond has a greater strength and stability and more energy is required to break the bond.

Only HF can break Si-O bond. Fluorine (EN = 4) is more electronegative than oxygen (EN = 3.5) and H-F is more ionic than Si-O bond. As a result, Si-O is broken by H-F only to form a more stable Si-F bond.

The post Silicon appeared first on The Fact Factor.

Double Salts:

These are the addition molecular compounds which are stable in the solid-state but dissociate into constituent ions in the solution. They are simply a combination of two different salts. e.g., Mohr’S salt, [FeSO₄·(NH₄)₂SO₄. 6H₂O gets dissociated into Fe²⁺, NH⁺⁴, and SO^2-₄ ions.

Characteristics of Double Salts:

• They are simply a combination of two different salts. They are prepared by mixing the solution of two different salts in an equimolar ratio.
• They dissociate completely in their constituent ion in their solution and produce simple ions.
• They give a test of all individual species.
• They do not exhibit isomerism.

Complex Salts:

Complex salts are those molecular compound which contains a central metal atom coordinated with ligands within a coordination sphere. e.g., Potassium ferrocyanide, K₄ [Fe(CN)₆].

Characteristics of Complex Salts:

• They contain a central metal atom coordinated with ligands within a coordination sphere. It is prepared by mixing the solution of two different salts in a stoichiometric ratio.
• Complex salt does not dissociate completely into their constituent ions but produces complex-ions.
• They do not give a test of all individual species.
• They exhibit isomerism.

Coordination Compounds:

The transition metals form a large number of complex compounds in which the metal atoms are bound to a number of anions or neutral molecules, these compounds formed are called coordination compounds.

Definition: Coordination compounds are those addition molecular compounds which retain their identity in the solid-state as well as in the dissolved state.

In these compounds. the central metal atom or ion is linked by ions or molecules with coordinate bonds. e.g., Potassium ferrocyanide, K₄ [Fe(CN)₆].

Potassium ferrocyanide undergoes hydrolysis as

K₄ [Fe(CN) ₆]•3H₂O   →   4K⁺  +  [Fe(CN) ₆]^4-

Coordination compounds play an important role in the chemical industry and in life itself. For example, the Ziegler-Natta catalyst which is used for polymerization of ethylene is a complex containing the metals aluminium and titanium. Chlorophyll, which is vital for photosynthesis in plants, is a magnesium complex and haemoglobin, which carries oxygen to animal cells, is an iron complex. Important vitamin B12 is a cobalt complex. Coordination compounds also find many applications in electroplating, textile dyeing and medicinal chemistry.

The complexes tend to retain their identity even in solution, although partial dissociation may occur. Complex-ion may be cationic, anionic or nonionic, depending on the sum of the charges of the central atom and the surrounding ions and molecules.

Terminology of Coordination Compounds:

Complex-ion or Coordination Entity or Coordination Sphere:

The central metal atom and the ligands which are directly attached to it are enclosed in a square bracket and are collectively termed as coordination sphere. The ligands and the metal atom inside the square brackets behave as single constituent unit

For e.g. in Potassium ferrocyanide, K₄ [Fe(CN)₆] the part   [Fe(CN)₆] is a coordination sphere. There are three types of coordination entities

• Cationic complex entity: It is the complex-ion which carries a positive charge. e.g., [Ni(NH₃)₆]²⁺ , [Pt(NH₃)₄]²⁺, etc.
• Anionic complex entity: It is the complex-ion which carries a negative charge. e.g., [PtCl₄]^2- ,  [Ag(CN)₂]^– , [Fe(CN)₆]^4- , etc.
• Neutral complex entity: It is the complex-ion which carries no charge. e.g.,   [Ni(CO)₄] , [Co(NH₃)₄ Cl₂]

Central Atom or Ion:

The atom or ion to which a fixed number of ions or groups are bound is called the central atom or ion. It is a Lewis acid. e.g., in Potassium ferrocyanide, K4 [Fe(CN)6], Fe is the central atom.  The central atom is generally a transition element or inner-transition element.

Ligand:

The molecules or ions that are attached to the metal in a complex-ion are called ligands. A ligand is Lewis base.  Every ligand has at least one unshared pair of valence electrons. The interaction between a metal atom and the ligands can be thought of as Lewis acid-base reaction. Some examples of Ligand are

The atom in the ligand that is bound directly to the metal atom is known as the donor atom. In complex ion [Cu(NH₃)₄]²⁺,  nitrogen is the donor atom while Cu²⁺ is the acceptor ion.

Depending on the number of the donor atoms present, ligands are classified as monodentate, bidentate or polydentate.

• Monodentate or Unidentate: It is a ligand, which has one donor site, e.g. H₂O and NH₃ are monodentate ligands
  with only one donor atom in each.  
• Bidentate or Didentate: It is the ligand. which have two donor sites. ethylenediamine (en) is a bidentate ligand.  The two nitrogen atoms can coordinate with a metal atom.

• Polydentate: It is the ligand, which has several donor sites. e.g., ethylenediaminetetraacetate ion. [EDTA]^4- is a hexadentate ligand.

• Ambidentate Ligands: Some unidentate ligands have more than one donor atom and these may coordinate to the metal ion through either of the two atoms. These types of ligands are called ambidentate ligands.  Complexes of these ligands yield linkage isomers.

M  ← NO₂  (N is the donor atom) or   M    ←   ONO  (O is the donor atom)

M   ←  SCN  (S is the donor atom)  or   M    ←     NCS (N is the donor atom)

M   ← CN   (C is the donor atom) or   M   ←  NC  (N is the donor atom)

• Chelating (Greek: Chele, meaning “claw”) Ligands: Bidentate and polydentate ligands are also called chelating agents because of their ability to hold the metal atom like a claw. Stability of complex depends on the number of the chelating ring.

• Bridging Ligands: Some monodentate ligands can simultaneously coordinate two or more metal atoms. As a result, the ligand acts as a bridge between different metal atoms. It is, therefore, called a bridging ligand. The resulting Compound is known as a bridged complex.

Coordination Polyhedron:

The spatial arrangement of the ligands which are directly attached to the central atom or ion is called coordination polyhedron around the central atom or ion.

Coordination number:

The coordination number in coordination compounds is defined as the number of ligands (donor) atoms/ ions surrounding the central metal atom in a complex ion.

Oxidation number:

The charge of the complex if all the ligands are removed along with the electron pairs that are shared with the central atom, is called oxidation number of the central atom.

The net charge on a complex ion is the sum of the charges on the central atom and its surrounding ligands. In
[Fe(CN)₆]^4- , the net charge is – 4. Each cyanide group has the oxidation state of -1 so the oxidation number of Fe must be +2.

If the ligands do not bear net charges the oxidation number of the metal is equal to the charge of the complex ion. In
[Cu(NH₃)₄]²⁺ each NH₃ is neutral, so the oxidation number of copper is +2.

IUPAC Nomenclature of Coordination Compounds:

Rules Of Nomenclature:

• The name of the compound is written in two parts (i) name of cation, and (ii) name of the anion.
• The cation is named first in both positively and negatively charged coordination complexes. For example, in K3[Fe(CN)6] name K+ first, and in [Co(NH3)4Cl2]Cl compound, name [Co(NH3)4Cl2] first.
• Within complex ligands (complex with dissimilar ligands) are named in alphabetical order first, and the metal
  ion is named last. prefixes are ignored when alphabetizing ligands
• For more than one similar ligands. the prefixes di, tri, tetra, etc are added before its name. For e.g. the ligands in the complex ion [Co(NH₃)₄Cl₂]⁺ are named as “tetraamine dichloro”. If the di, tri, etc already appear in the complex then bis, tris, tetrakis are used. For e.g. the ligand ethylenediamine already contains di, therefore its name is bis(ethylenediamine).
• If the complex part is an anion, the name of the central metal ends with the suffix ‘ate’. Names of the anionic ligands end in ‘0’, names of positive ligands end with ‘ium’, and names of neutral ligands remain as such. But exceptions are we use aqua for H2O, ammine for NH3, carbonyl for CO, and nitrosyl for NO.

Naming of Central Atom

Naming of Anionic Ligands

• The name of the complex part is written as one word.
• The oxidation state for the metal in cation, anion, or neutral coordination compounds is indicated by the Roman numeral in parentheses. For e.g. In complex cation [Cr(NH3)4Cl2]+, the oxidation number of chromium is +3 Hence the name of the ion is written as tetraamminedichlorochromium (III) ion. Similarly, complex anion [Fe(CN)₆]^4- is called hexacyanoferrate(II) ion.
• If the complex ion is a cation, the metal is named the same as the element. The neutral complex molecule is named similar to that of the complex cation. In complex cation [Cr(NH₃)₄Cl₂]⁺ the name of ion is written as tetraamminedichlorochromium (III) ion, similarly cation [Co(NH₃)₆]³⁺ is named as hexaamminecobalt (III)ion. Neutral ion  [Ni(CO) ₄] is named as  tetracarbonyl nickel(0),

Examples of IUPAC Naming:

The post Coordination Compounds appeared first on The Fact Factor.

Moseley’s Criteria:

Henry Moseley, a physicist from England in the year 1913, observed that when certain metals were bombarded with high-speed electrons, X-rays are emitted. He studied the frequencies of the X-rays emitted and found that in all the cases, the square root of the frequency was directly proportional to the atomic number of the atom of the metal.

He plotted a graph between the square root of frequency and atomic numbers of different metals, a straight line was obtained. When a graph was plotted between the square root of frequency and atomic masses of the metals, the graph was not a straight line.

He concluded that there was no correlation between the frequency and the atomic mass. From these observations, Moseley concluded that atomic number and not the atomic mass is, the fundamental property of an element.

On the basis of this hypothesis, he put forward his law of periodicity, commonly known as Modern Periodic Law. It states that “The physical and chemical properties of the elements are the periodic function of their atomic numbers”. It is the basis of “Modern Periodic Table”. Thus Moseley modified Mendeleev’s periodic law.

Explanation of the Modern Periodic Law on the Basis of Electronic Configuration:

In chemical reactions, there is a transfer or sharing of electrons. Thus chemical properties of an element are dependent on the electronic configuration (number of electrons). In an atom, the number of electrons depends on the number of protons (they are equal). The number of protons in an atom depends on the atomic number of the element.

Similarly, physical properties of an element are dependent on the configuration of electrons in an atom of the element. Hence the physical and chemical properties of an element are the function of atomic number.

Periodicity and Magic Numbers:

Periodicity may be defined as the repetition of the similar properties of the elements placed in a group and separated by certain definite gaps of atomic numbers.

When the elements were arranged in increasing order of atomic numbers, it was observed that the properties of elements were repeated after certain regular intervals 01 2, 8, 8, 18, 18 and 32. These numbers are called magic numbers. These numbers are the cause of periodicity in properties due to the repetition of similar electronic configuration.

Features of Modern Periodic Table:

Modern Periodic Table is based on the Modern Periodic Law which states that “The physical and chemical properties of the elements are the periodic function of their atomic numbers”. The elements are arranged in order of increasing atomic numbers in horizontal rows called periods and vertical columns called groups.

Groups:

• There are eighteen (18) groups in the Modern Periodic Table. These are numbered from 1 to 18.
• The elements present in a group are separated by definite gaps of atomic numbers (8, 8, 18, 18, 32).
• The elements present in a group have the same number of electrons in the valence shell of their atoms. i.e. their electronic configuration of valence shell is identical. For example, all the elements of group 1 have valence shell configuration of ns¹. All the elements of group 17 have valence shell configuration of ns²np⁵.
• Due to similar valence shell configuration, the elements present in a group have identical chemical properties.
• The physical properties of the elements in a group such as melting point, boiling point, density vary gradually.
• Elements of groups, 1, 2, 13, 14, 15, 16, and 17 are called normal or representative elements.
• Elements from Gr. 3 to 12 are called transition elements.
• Noble or inert gases are placed in group 18.
• Two series of fourteen elements each are called at bottom of the periodic table. The elements of these series are called Lanthanoids and Actinoids. These elements are also called inner transition elements or rare earth elements.
• Lanthanoids are placed in group 3 and the sixth period.
• Actinoids are placed in group 3 and the seventh period.
• Elements of group 1are called alkali metals.
• Elements of group 2 are called alkaline earth metals.
• Elements of group 16 are called chalcogens [ore forming elements].
• Elements of group 17 are called halogens. [sea salt forming]
• Elements of group 18 are called noble gases.

Periods: 

• Modern Periodic Table has seven (7) horizontal rows called periods. They are numbered 1-7.
• The first (₁H – ₂He) period has two elements and is called the shortest period.
• The second (₃Li – ₁₀Ne) & third (₁₁Na – ₁₈Ar) periods have eight elements and are called short periods.
• The fourth (₁₉K – ₃₆Kr) and the fifth (₃₇Rb – ₅₄Xe) periods have eighteen elements and are called long periods.
• The sixth (₅₅Cs – ₈₆Rn) period contains 32 elements and is called the longest period.
• The seventh (₈₇Fr – )period has 23 elements and is called an incomplete period.
• In all the elements present in a period, the electrons are filled in the same valence shell.
• As the number of electrons in the valence shell change, the chemical properties of the elements present in a period also change.

Merits of the Modern Periodic Table:

• It is quite easy to remember and reproduce.
• The positions of the elements in the periodic table are linked with their electronic configuration.
• Each group is an independent group and the idea of sub-groups as in the case of Mendeleev’s periodic table has been discarded.
• There is one position for all the isotopes of an element because the isotopes have the same atomic number. For example, the three isotopes of hydrogen i.e. protium, deuterium, and tritium have different atomic mass numbers but have the same atomic number one.
• The positions of certain elements which were earlier misfits in Mendeleev’s periodic table are now justified because it is based on an atomic number of the elements.
• This periodic table is linked with the electronic configuration. Hence the further division of the elements into s, p, d, and f blocks has been quite useful in understanding their properties.

Limitations of Long Form of Periodic Table:

• In the long form of the Periodic Table, the position of hydrogen still remains uncertain.
• The inner-transition elements do not find a place in the main body of the table. They are placed separately below the periodic table.

The post Modern Periodic Table appeared first on The Fact Factor.

Need for Classification of Elements:

With the rapid advance and developments in science, the number of discovered elements increased. Due to a large number of elements, it is difficult to study and remember the behaviour and properties of each and every element. Hence attempts have been made to classify these elements into groups of elements having similar characteristics. This method of grouping the elements into different classes is known as periodic calcification classification of elements. This classification of elements makes the study of elements systematic and easy.

Unitary Theory of Periodic Classification:

This theory was proposed by William Prout in 1815. He suggested that the values of atomic masses of all elements were whole numbers or varied only slightly from the whole numbers if hydrogen was considered the basis of all atomic masses. Thus hydrogen was regarded as the ‘central or base’ element around which all other atoms were made. Thus ¹²C (carbon) is made up of 12 units of hydrogen. ²⁰Ca (calcium) is made up of 20 units of hydrogen.

This theory was not able to explain fractional atomic masses of some elements ^63.5Cu (copper) and ^35.5Cl (chlorine), etc.

Dobereiner’s Triads:

This theory was given by a German chemist Dobereiner. According to the theory, If three elements resembling properties are arranged in increasing atomic masses, then in each case, the atomic weight of the middle element is found to be the arithmetic mean (average) of the other two extreme elements. This relation is sometimes referred as the law of triads. The group of such three elements is called a triad.

However, this theory suffered a major limitation as Dobereiner could classify only 9 elements in such a manner, out of all those were known at the time. Therefore, the concept of triads is discarded.

Cooke’s Homologous Series:

 In chemistry, a homologous series of elements is a series of elements arranged in ascending atomic mass such that the atomic mass increases in a regular fashion. Cooke classified elements in several homologous series.

 Newland’s Law of Octaves:

“Octo” means “eight”. In music, every eighth note is similar to that of the first note and of frequency double the frequency of first. These are seven musical notes according to Indian system, these musical notes are sa, re, ga, ma, pa da, ni and again sa. This sa is similar to the first sa. According to the Western system, the corresponding symbols of these notes are do, re, me, fa, so, la, ti again do. This do is similar to first do. In the musical notes, there is a repetition of the same note after a gap of seven.  Thus every eighth note is the repetition of the first note.

In1865, an English chemist, John Alexander Newlands’s observed that when the lighter elements were arranged in order of their increasing atomic masses, the properties of every 8th element were similar to those of the first one like the eighth note of a musical scale. This relation is called Newlands’s law of octaves.

Drawbacks of this concept:

• It was applicable to only lighter elements having atomic weights up to 40 u (calcium). After that, every eighth element did not possess the same properties as by the element lying above it in the same group.
• With the discovery of noble gases, the properties of the 8th element were no longer similar to those of the first one.

Lother Meyer’s Arrangement of Elements:

In 1869, Lothar Meyer, a German chemist, studied the physical properties of the various elements. He Plotted a graph by taking the atomic volume of elements on the y-axis and atomic masses of the elements on the x-axis and observed that the elements with similar properties occupied a similar position on the curve.

Lothar Meyer proposed that the physical properties of the elements are a periodic function of the atomic masses. He arranged the elements in the tabular form in order of their increasing atomic weights.

Characteristics of Graph:

• The elements with similar properties occupied a similar position on the curve.
• The most strongly electropositive alkali metals occupy the peaks on the curve.
• The less strongly electropositive alkaline earth metals occupy the descending position on the curve.
• The most electronegative elements i.e. halogens occupy the ascending position on the curve.

Mendeleev’s Periodic Table:

The Russian Chemist Dmitri Ivanovich Mendeleev tried to co-relate the atomic masses of the elements with their physical and chemical properties. He used the basis of Lother Meyer’s hypothesis. He studied the formulae and chemical properties of several elements compared to Lother Meyer. On basis of his study, he proposed the Periodic Law.

Mendeleev’s Periodic Law: 

It states that the physical and chemical properties of the elements are a periodic function of their atomic masses.

Mendeleev’s Periodic Table:

Mendeleev arranged the elements known at that time in order of increasing atomic masses in horizontal rows and vertical columns. This arrangement was called the periodic table. Elements with similar characteristics were found to present in vertical rows called groups. The horizontal rows were known as periods.

Characteristics of the Periodic Table:

• In the Mendeleev’s periodic table, the elements are arranged in vertical columns called groups and horizontal rows known as periods.
• There are eight groups indicated by Roman Numerals as I, II, III, IV, V, VI, VII, VIII and the elements belonging to the first seven groups have been divided into sub-groups designated as A and B on the basis of similarities in properties. Group VIII consists of nine elements which are arranged in three triads.
• There are six periods (numbered from 1 to 6). The periods 4, 5 and 6 are divided into two halves. The first half of the elements are placed in the upper left corners and the second half occupy lower right corners in each box.
• Mendeleev left some empty spaces in the periodic table for the undiscovered elements which were identified later on and were placed at their respective positions.
• Noble gases are placed in a separate group called a zero group without disturbing the main periodic table.

Merits of Mendeleev’s Periodic Table:

A Systematic study of the elements:

• Mendeleev for the first time arranged a very large number of elements into groups and periods. This made the study of the elements quite systematic.

Prediction of new elements and their properties:

• Mendeleev laid more stress on similarity in properties rather on increasing atomic masses of the elements. So whenever a particular element did not fit in the arrangement, he left a gap in the periodic table. Thus, many gaps for the undiscovered elements were left in the periodic table by Mendeleev.
• He left the gap one place down aluminium and silicon and called these elements as Eka-aluminum and Eka-Silicon. He predicted their physical properties. These elements were discovered afterword and called Gallium and Germanium respectively.

Correction of doubtful atomic masses:

• Mendeleev also corrected the atomic masses of certain elements with the help of their expected positions and properties.

Demerits of the Mendeleev’s Periodic Table:

Position of hydrogen:

• Both hydrogen (H₂O, HCl, H₂S) and alkali metals like Sodium form similar compounds (Na₂O, NaCl, Na₂S) with elements like oxygen, chlorine and sulphur etc., hence Hydrogen (Z = 1) is placed at the top of the alkali metal family because it resembles alkali metals in its properties.
• At the same time, hydrogen (CH₄, SiH₄, GeH₄) also resembles halogens present in group VII A like chlorine (CCl₄ , SiCl₄, GeCl₄) in many of its properties. Both are non-metals and also diatomic in nature. The compounds of both hydrogen and halogens with certain non-metals are of covalent nature.
• Hence there was confusion about the position of hydrogen.

Anomalous positions of some elements:

• The elements in the Mendeleev’s Periodic table have been arranged in order of increasing atomic masses, but in some cases, the element with higher atomic mass precedes the element with lower atomic mass. For example, Ar (Atomic mass =39.9) precedes K (Atomic mass= 39.1) and similarly, Co (Atomic mass 58.9) has been placed ahead of Ni (Atomic mass=58.7).
• No justification has been given by Mendeleev for such anomalous positions of these elements.

Position of isotopes:

• The element whose atoms have different atomic masses but the same atomic number are called isotopes. Since the periodic table has been prepared on the basis of increasing atomic masses of the elements, different positions must have been allotted to all the isotopes of a particular element.
• For example, Hydrogen has three isotopes having atomic masses 1, 2, and 3. Hence three different positions should have been allotted to the isotopes of hydrogen. But they have been assigned only one position.

No co-relation of elements in sub-groups:

• According to Mendeleev, the elements placed in the same group must show similar properties. But there is no similarity among the elements in the two sub-groups of a particular group.
• For example, Li, Na, and K present in group IA are quite different from Cu, Ag, and Au which belong to group IB.

Different groups for similar elements:

• According to Mendeleev, the elements having similar properties are placed in the same group of the periodic table.  But elements with similar properties have been placed in different groups.
• For example,  Both copper and mercury resemble in many properties. But copper has been placed in group IB while mercury has been assigned a position in group IIB

Cause of periodicity:

No proper explanation has been given by Mendeleev as to why the properties of the elements get repeated after gaps of atomic masses in a particular group.

The post Initial Classification of Elements appeared first on The Fact Factor.

Anhydrous HF:

• It is a gaseous compound or dry liquid obtained by condensation of gaseous dry HF.
• Non-conductor of electricity as anhydrous hydrofluoric acid is non-ionised.
• HF stored in steel cylinders

Preparation: 

Dry H₂ gas + Dry F₂ gas and From Fremy’s salt KHF₂

Hydrofluoric Acid:

• It Is an aqueous solution of hydrogen fluoride.
• It is a weak monobasic acid. It is a weak electrolyte
• Stored in ‘gutta-percha’ or Polythene bottles.

Preparation:

By action of H₂SO₄ on CaF₂ (Fluorspar):

When powdered fluorspar i.e. CaF₂ is mixed with 96% H₂SO₄ and distilled in lead retort between 473 K to 573 K, vapours of HF are produced.

The vapours of HF are absorbed in water in lead receiver which is kept cooled by immersing it in cold running water. This gives aqueous hydrofluoric acid.

Reaction:

CaF2 +  H₂SO₄   →   CaSO₄ +  2HF­

By Direct Combination of Hydrogen and Fluorine:

Anhydrous HF can be prepared by direct combination of dry hydrogen gas and dry fluorine gas. This reaction occurs even in dark and even at low temperature as low as 63 K

H₂   +    F₂          →      2HF

From Fremy’s salt – Potassium hydrogen fluoride KHF₂:

Pure and perfectly dried Fremy’s salt KHF2 is distilled In a copper retort to about 573 K.

KHF₂    →  KF +    HF

The vapours are collected in a copper vessel cooled by a freezing mixture.  Anhydrous liquid HF is thus obtained.  Copper is not affected by anhydrous-liquid HF. (KF remains in the distilling vessel)

Purification:

• HF is dehydrated using thionyl chloride.

HF + H₂O + SOCl₂ →   HF + 2 HCI ­ + SO₂­

• On cooling gases, only hydrogen fluoride condenses to a liquid.
• Hydrogen chloride is insoluble in liquid hydrogen fluoride.  Sulphur dioxide is removed by fractional distillation.
• Anhydrous HF or aqueous HF are purified by redistillation or fractional distillation.

Storage:

• Both -hydrogen fluoride and hydrofluoric acid attack glass, so they cannot be stored in a glass container.
• Anhydrous hydrogen fluoride can be stored in special steel cylinders.
• Aqueous hydrofluoric acid is stored in ‘gutta-percha’ or Polythene bottles.

Properties of HF:

HF differ from the other halogen hydracids:

• HF is a liquid (B.P. 292.5 K) while the other hydracids are gases.
• The liquid contains associated (HF)n molecules formed due to intermolecular hydrogen bonding. In other hydrogen halides, intermolecular hydrogen bonding does not exist.
• Hydrofluoric acid is the weakest of all the halogen acids.
• It forms acid salts like KHF₂.
• HF attacks the glass and dissolves it. Other hydracids do not react with the glass.
• It is the only acid which reacts with Silica (SiO₂).

Physical Properties of HF:

• Anhydrous hydrogen fluoride is a gas at an ordinary temperature, which condenses to a colourless fuming liquid. Boiling point 292.5 K (19.5°C)
• Both the liquid and the vapours are corrosive and poisonous. They attack the skin causing painful blisters.
• The liquid contains associated (HF)n molecules formed due to Inter-molecular hydrogen bonding.
• It dissolves in water in all proportions.

Hydrogen Bonding in Hydrogen Fluoride:

In hydrogen fluoride, the hydrogen atom is covalently bonded to the fluorine atom which is highly electronegative.  There is a wide difference in the electronegativities between hydrogen (2.1) and fluorine (4.0). The covalent bond is therefore highly polar and has a partial ionic character. The hydrogen atom in HF has a high partial positive charge and fluorine has a high partial negative charge.

H^+δ − F^-δ

Negatively charged F atom from one HF molecule attracts the positively charged H atom from the other molecule through electrostatic attraction. This is called Hydrogen bonding. This causes a large number of H-F molecules to associate together.  The associated molecules form zigzag chains. More amount of energy is thus required to separate these molecules from each other.

Such strong hydrogen bonding is not possible in the case of other hydrogen halides since other halogens are less electronegative than fluorine.

Although HCI, HBr and HI molecules are polar, the magnitude of polarity is not sufficient to form hydrogen bonds between their molecules.  Hence, these molecules exist as gases.

Characteristic Properties of HF due to Intermolecular Hydrogen Bonding:

Hydrogen fluoride is liquid with B.P. 292.5 K, whereas other hydrogen halides are gases. HCI, HBr and HI have B.P. 188 K, 206 K and 238 K respectively. Thus expected B.P. for HF is below 188 K. But it is higher. Thus the unexpectedly higher boiling point of HF is due to intermolecular hydrogen bonding.

Hydrofluoric acid is the weakest of all halogen acids. The degree of dissociation for 0.1 M solutions of HF, HCI, HBr, and HI are 0.08, 0.93, 0.94, and 0.95 respectively at 298 K.

It forms hydrogen-bonded anion like HF₂ ^– and acid salt like KHF₂.

HF is a liquid, while other halogen acids are gases at ordinary temperature:

Due to intermolecular hydrogen bonding, HF molecules are associated as (HF)_n. Formation of hydrogen bonding lowers energy of the system and so HF is liquid. Such hydrogen bonding is absent in other hydrogen halides.  Hence they are gases at ordinary temperature.

Due to hydrogen bonding, the separation of HF molecules requires more energy.  Hence the boiling point of HF will be higher than other hydrogen halides, which do not form hydrogen bonding.

HF is the weakest amongst all hydracids of halogens:

Hydrofluoric acid is a weak monobasic acid having Ka = 6.8 x 10-4. The degree of dissociation for 0.1 M solution of HF, HCI, HBr, and Hi are 0.08, 0.93, 0.94, and 0.95 respectively at 298 K. Fluorine has smallest atomic size, hence in H-F molecule, H-F bond length is short.

The highest electronegativity of fluoride makes the H-F bond highly polar. These two factors make the H-F bond strong and therefore, it needs more energy for ionization. Because of intermolecular hydrogen bonding, HF molecules are associated.  More energy is required to separate HF molecules for ionization.

Other halogens are larger in size and less electronegative than fluorine.  Their halogen acid molecules are not associated. Therefore, HCI, HBr and HI are almost completely ionized in aqueous solution.

Chemical Reactions of Hydrogen Fluoride:

Action of HF on Silica:

A concentrated solution of hydrofluoric acid reacts with Silica (Quartz) to form Silicon tetrafluoride.  But in presence of excess of hydrofluoric acid, hydrofluosilicic acid is formed.  It is soluble in water.

SiO₂  +  4HF  →    SiF₄   +  2 H₂O

SiF4   +   2HF  →   H₂ Si F₆  (hydrofluosilicic acid)

Action of HF on a Glass:

Ordinary glass is essentially a mixture of the silicates of alkali metals, alkaline earth metals and some free Silica. Hydrogen fluoride attacks glass and decomposes it. Hydrofluosilicic acid and metal silicofluoride are formed which are soluble in water. Therefore, glass is slowly eaten up by HF. Thus glass slowly dissolves in HF acid. Hence HF solution is not stored In glass bottles.

Na2SiO₃    +     6 HF  → Na₂ Si F₆   +   3H₂O

CaSiO₃     +     6 HF  → CaSi F₆   +   3H₂O

SiO2         +      6HF →  H₂ Si F₆    +   2H₂O

Uses of HF:

Hydrogen fluoride is used for

• Etching of glass.
• Removal of silica (sand) from graphite.
• Silicate analysis i.e. estimation of silica.
• Preventing dental decay in the form of fluoride ion.
• Liquid HF is used as a non-aqueous solvent.

Etching of a Glass:

The process of making a permanent marking on the glass surface is called etching of the glass. The etching of glass is carried out by treating it with HF (hydrogen fluoride). The glass is a homogeneous mixture of silicates of sodium and calcium along with some free silica. HF reacts with glass-forming metallic fluorides which are water-soluble. Due to this reaction, the glass gets etched.

The reactions involved in the etching of the glass are as follows.

Na2SiO₃    +     6 HF  → Na₂ Si F₆   +   3H₂O

CaSiO₃     +     6 HF  → CaSi F₆   +   3H₂O

SiO2         +      6HF →  H₂ Si F₆    +   2H₂O

Process of etching:

• The glass which is to be etched is cleaned with chromic acid and washed with water.
• Then the glass is coated with thin layer of wax.
• Then the markings or designs are engraved on this coated surface with the help of sharp pointer. Thus the wax in the region is removed and the area to be etched is exposed.
• The exposed surface is then treated with aqueous HF solution or HF vapours.
• The chemical reaction takes place and the glass gets etched in the exposed region.
• The rest of the wax is removed by dissolving it in turpentine.

Use of HF in the Removal of Silica from Graphite:

Graphite is used in preparing electrodes in electrolytic cells.  In place of natural graphite, artificial graphite is used more commonly. Artificial graphite is prepared by heating coke powder with silica in an electric arc furnace. So artificial graphite contains silica as the impurity. The presence of Silica will increase its resistance. To remove silica, graphite is treated with hydrofluoric acid.

SiO₂  +  4HF  →    SiF₄   +  2 H₂O

SiF4   +   2HF  →   H₂ Si F₆  (hydrofluosilicic acid)

The post Hydrogen Fluoride appeared first on The Fact Factor.

Occurrence:

Because of extreme reactivity fluorine does not occur in the free state.  It occurs as the fluorides (F^–) of certain metals such as Calcium and Aluminium. Its major minerals are

• Fluorspar – CaF₂,
• Cryolite – Na₃AlF₆ and
• Fluorapatite -3Ca(PO₄)₂ , CaF₂.

Difficulties in the Isolation of Fluorine:

The following difficulties were involved in the isolation of fluorine.

• Its oxidation is not possible. It exists as fluoride (F^–) in its minerals. The preparation of it from a fluoride is an oxidation reaction.

2 F^– → F₂ +   2e^–

• It is the most electronegative element and itself is a very strong oxidizing agent.  F2 is a more powerful oxidizing agent than O2.  Therefore oxidation of fluoride to fluorine could not be carried out by chemical oxidizing agents. Therefore, no oxidizing agent was able to separate fluorine from hydrogen.
• When hydrofluoric acid is heated with oxidizing agents, like MnO₂, KmnO₄, K₂Cr₂O₇ no fluorine is obtained. This is because of the High affinity of fluorine for hydrogen.
• H-F bond Is considerably strong due to small bond length and high polarity. It cannot be prepared by the electrolysis of HF. Electrolysis of hydrofluoric acid solution gave ozonized oxygen at the anode, instead of fluorine.  This is due to the action of F₂ on water.

2H₂O  +   2F₂  → 4HF +   O2

3H₂O  +   3F₂  → 6HF +   O3

• Then the electrolysis of anhydrous hydrogen fluoride was attempted. It was found to be a poor conductor (non-conductor) of electricity.  HF is highly volatile and poisonous.
• Proper apparatus for preparation and storage was not available. As F2 is extremely reactive, it attacks glass, platinum, carbon, and other materials commonly used for the construction of the apparatus of electrolysis.
• Aqueous HF is an active solvent.  It attacks the glass and various metals. Thus the main need was to find a suitable apparatus end suitable electrolyte.

Preparation of Fluorine:

Dennis Method:

Principle:

Pure, dry and anhydrous potassium hydrogen fluoride (KHF₂) is electrolyzed in the molten state at 523 K when hydrogen is liberated at the cathode while F₂ is liberated at the anode.

Diagram:

Dennis Cell (Apparatus):

This is a heavy V-shaped copper tube. F₂ reacts with copper forming a protective layer of copper fluoride CuF₂. This protects the cell from further attack by fluorine.  The V-shaped tube prevents fluorine liberated at the anode and hydrogen at the cathode from coming in contact with each other.

Electrodes: The electrodes are made of graphite. The copper tube is closed with copper caps through which graphite electrodes are fitted.  Copper caps are cemented to and Insulated from the copper tube by bakelite cement.

Electrolyte: Pure, dry and anhydrous KHF₂, in a molten condition.  Anhydrous HF Is added periodically.

Temperature: 523 K

Heating unit: The copper tube is covered from outside, with asbestos cement insulator of current but conductor of heat resistance wire for electrical heating legging – to prevent loss of heat by radiation.

Outlets: Two outlets are provided for hydrogen and fluorine in the upper portions of the two arms.

Current and Voltage: A current of 5 amperes and at 12 volts is employed.

Reactions:

2KHF₂ → 2KF       +    2HF

2KF  →    2K⁺  +    2 F^–

(HF acts as an ionizing liquid in which KF ionizes.  KF is a carrier of current)

At Anode

2F^–  →    F₂ +       2e^–

At Cathode

2K    +    2e^– →   2K

2K  +   2HF  → 2KF          + H₂ ↑

As only HF is used up in the process, it is added periodically.

Drawbacks of Dennis Method:

• F₂ liberated at the anode is usually contaminated with HF and CF4
• The graphite electrodes are continuously eaten up by fluorine and hence electrodes have to be periodically replaced.
• Molten electrolyte froths or foams. This foam or froth rises or creeps up the slanted V-shaped tube and chokes the outlets. This leads to an explosion due to the mixing of H₂ and F₂.
• Current efficiency in this method is low i.e. about 30%.
• At the high temperature (523 K) at which electrolysis is carried out, the attack of fluorine on the cell is vigorous.

Note:

Graphite is not suitable electrode at anode because it combines with fluorine formed and produces CF4.  Hence it is necessary to replace graphite electrode time to time. The reaction is as follows

C   +   2F₂   →    CF4

(Carbon)        (Carbon tetrafluoride)

Whitlaw Gray’s Method:

Principle:

Pure, dry and anhydrous potassium hydrogen fluoride is electrolyzed in the molten state at 523 K. When hydrogen is liberated at the cathode while F₂ is liberated at the anode.

Diagram:

Whytlaw-Gray Cell (Apparatus):

The cell construction is modified to overcome some disadvantages of the ‘Dennis method”. Three main modifications are The cylindrical vessel, the vertical walls of this vessel minimize the creeping of the electrolyte. The diaphragm between electrodes prevents the mixing of Hydrogen and Fluorine. Blocking of exit tubes is avoided since they are with a large bore.

Electrodes: The cell consists of a rectangular copper vessel wound for electrical heating.  This copper vessel act as a cathode. A pure thick graphite rod suspended at the centre serves as an anode. Anode and cathode are separated by a diaphragm.  The anode is cemented to and insulated from the diaphragm cell by a paste of fluorspar CaF₂. and water glass Na₂SiO₃. The diaphragm is closed at the bottom and it is laterally perforated below the level of the electrolyte.  The upper solid non-porous part leads to a delivery tube (with large bore) for F2, gas.

Electrolyte: Pure, dry and anhydrous KHF₂, in a molten condition.  Anhydrous HF Is added periodically.

Temperature: 523 K

Heating unit: The copper vessel is covered from outside, with asbestos cement insulator of current but conductor of heat. A resistance wire for electrical heating and lagging – to prevent loss of heat by radiation.

Outlets: The outer copper vessel has an outlet for hydrogen and Inner copper diaphragm vessel has an outlet for fluorine gas. The outlets have large bore.

Current and Voltage: A current of 12 to 15 amperes and at 15 volts is employed.

Reactions:

2KHF₂ → 2KF       +    2HF

2KF  →    2K⁺  +    2 F^–

(HF acts as ionising liquid in which KF ionises.  KF is a carrier of current)

At Anode

2F^–  →    F₂ +       2e^–

At Cathode

2K    +    2e^– →   2K

2K  +   2HF  → 2KF          + H₂ ↑

As only HF s used up in the process, it is added periodically.

Purification:

F₂ gas obtained by Dennis method contains two impurities  (a)  HF and (b) CF₄. To purify fluorine, the gas is passed through copper U tubes containing dry NaF which absorbs HF vapours to form sodium hydrogen fluoride.

NaF_(s)     + HF_(g)   →    NaHF₂

Carbon tetrafluoride is condensed by cooling the gas in a trap cooled in liquid oxygen (90 K). Only CF₄ gets condensed and is removed. Fluorine (BP 86 K) remains uncondensed and is collected in pure form.

Collection and Storage:

Fluorine gas is heavier than air. It is collected by upward displacement of air in copper cylinders. Fluorine is then compressed and stored In Cu-Ni alloy cylinders.

The Advantages of Whytlaw Gray’s Method over Denis Method:

• Fluorine liberated at the anode and hydrogen at the cathode are prevented from coming In contact with each other by the use of diaphragm cells.
• The walls of the electrolytic cell are vertical and broad outlets provided at the top of the containers for fluorine and hydrogen gas.  The molten electrolyte cannot choke the outlets.
• The current efficiency is as high as 80%.
• This process is continuous.

Properties of Fluorine:

Physical properties of Fluorine:

• Colour: Fluorine is a pale yellow gas.
• Odour: It has an Irritating pungent odour.
• Nature: It Is highly poisonous and has a corrosive action on skin.
• Density: It Is heavier than air.
• Boiling point: 86 K

Chemical Reactions of Fluorine:

Fluorine is one of the most chemically active elements known since it reacts with almost all elements with the exception of nitrogen, oxygen, helium, and argon.

Reaction with Water:

Fluorine decomposes water at ordinary temperature forming hydrofluoric acid and liberating oxygen and ozone.

2H₂O  +   2F₂  → 4HF +   O2

3H₂O  +   3F₂  → 6HF +   O3

Reaction with cold dilute sodium hydroxide:

Cold dilute (2%) aqueous solution of NaOH gives sodium fluoride and oxygen difluoride.

2 NaOH +   2F₂  →   2NaF  + H20 +  OF₂

Sodium fluoride             Oxygen difluoride

Reaction with hot concentrated sodium hydroxide:

Hot concentrated, aqueous solution of NaOH gives sodium fluoride and oxygen.

4 NaOH + 2F₂  →  4NaF  + 2H₂0 +  2O₂

Reaction with Metals:

• With Sodium: Sodium burns spontaneously in fluorine to form sodium fluoride.

2Na   +  F2  →      2NaF

• With Magnesium: Magnesium reacts with fluorine on worming to form magnesium fluoride.

Mg   +   F₂    →     MgF₂

• With Copper: Copper forms a thin layer of CuF₂, on its surface which resists a further attack of fluorine.

Cu   +   F₂   →     CuF₂

• With Mercury: Mercury reacts with fluorine to give mercury fluoride

Hg   +   F₂   →     HgF₂

Reaction with Nonmetals:

• With Hydrogen: Fluorine reacts with Hydrogen with explosive violence even in dark and even at a low temperature as low as 63 K to form hydrogen fluoride.

H₂ +  F₂ → 2HF

• With Boron : 

2B  +  3F₂   →  2 BF₃  (Boron trifluoride)

• With Carbon :

C    +   2F₂ →   CF₄ (Carbon tetrafluoride)

• With Silicon (Quartz):  

Si   +   2F₂ →   SiF₄ (Silicon tetrafluoride)

• With Phosphorous : 

2P  +  3F₂ → 2PF₃  (Phosphorus trifluoride)  and 

2P +    5F₂ →   2PF₅ (Phosphorus pentafluoride)

• With Sulphur: 

S   +  3F2  →    SF₆  (Sulphur hexafluoride)

With chlorine Sulphur forms SCl4  and with fluorine, it forms SF₆.

• Fluorine has the highest electronegativity and the smallest atomic size, due to which it is the strongest oxidizing agent. Hence it brings out the maximum oxidation state of any element.
• Hence with fluorine, every element shows a higher oxidation state. Hence with fluorine Sulphur shows its highest oxidation state of +6 while chlorine shows the +4 oxidation state.

Action of hydrocarbon (Methane): 

Because of strong affinity of fluorine for hydrogen, fluorine decomposes hydrocarbons forming hydrogen fluoride and leaving behind carbon. Excess of fluorine react with carbon to form gaseous carbon tetrafluoride.

CH₄ + 2F₂ → 4 HF +  C

Oxidising properties:

• Potassium Chlorate (KClO₃):

KClO₃   +   F₂ +   H₂O  →  2HF + KClO₄  (Potassium perchlorate)

• Potassium Sulphate (K₂SO₄)

K₂SO₄   +   F₂ →  2KF +  K₂S₂O₈ (Potassium persulphate)

• Potassium Carbonate (K2CO3) Potash ash:

K2CO₃   +   F₂ →   2KF +  K₂C₂O₆  (Potassium percarbonate)

lnterhalogen Compounds of Fluorine:

Binary compounds of fluorine with other halogens are called interhalogen compounds of fluorine. In these compounds, fluorine shows -1 oxidation state.  The other halogens show positive oxidation states such as +1, +3, +5, +7.

Fluorine can form four different types of binary compounds (fluorides) with other halogen and they are as follows.

Action of Fluorine on Chlorine:

When chlorine and fluorine are made to react with each other in equal volumes, chlorine monofluoride is formed

Cl₂ +   F₂ →   2CIF  (Chlorine monofluoride)

When chlorine and excess fluorine are made to react, chlorine trifluoride is formed

Cl₂ +  3 F₂  →  2CIF₃  (Chlorine trifluoride)

Action of Fluorine on Bromine:

When fluorine is made to react with bromine diluted with nitrogen, bromine monofluoride is formed

Br₂ + 3F₂     →         2BrF₃    (Bromine trifluoride)

When excess fluorine is made to react with bromine, bromine pentafluoride is formed.

Br₂ +  5 F₂ →  2BrF₅  (Bromine pentafluoride)

Action of Fluorine on Iodine:

When excess fluorine is made to react with iodine, iodine pentafluoride is formed

I₂ +   5F₂   →    2IF₅  (Iodine pentafluoride)

When much excess fluorine is made to react with iodine, iodine heptafluoride is formed

I₂ +  7 F₂ →      IF₇    (Iodine heptafluoride)

Uses of Fluorine:

• Freon: It is used in the preparation of dichloro-difluoro-methane, CCl₂F₂, known as ‘Freon’.  Freon is used in refrigeration and air-conditioning.
• Teflon: It is used to prepare Teflon. Teflon is polymerized tetrafluoroethylene. Teflon is a chemically inert and electrical insulator. It is used for coating frying pans, cooking pots, and reaction vessels.
• UF₆: It is used for the preparation of UF₆.  UF₆ is helpful in the separation of isotopes of Uranium.
• SF₆: It is used for the preparation of SF₆.  SF₆ is used in the vulcanization of rubber and for high voltage insulation.
• Rocket Fuel: In combination with hydrazine (N₂H₄), fluorine is used as rocket fuel.

Prevention of Tooth Decay or Dental Decay:

The tooth enamel is composed of Ca₅(OH)(PO₄)₃ which slowly converts to Ca₅F(PO₄)₃ due to the regular brushing of teeth with the fluoride toothpaste. (The fluoride ion is added in the form of soluble SnF₂. or sodium mono-fluoro-phosphate to toothpaste).

The fluoro compound Ca₅F(PO₄)₃  resists the attack of germs and mild organic acids associated with foodstuff and prevents tooth decay.

The post Fluorine appeared first on The Fact Factor.

Among metals, aluminium is the most abundant. It is the third most abundant element in earth’s crust (8.3% approx. by weight). It is a major component of many igneous minerals including mica and clays. Many gemstones are impure forms of Al₂O₃ and the impurities range from Cr (in ‘ruby’) to Co (in ‘sapphire’). Iron is the second most abundant metal in the earth’s crust. It forms a variety of compounds and their various uses make it a very important element. The principal ores of aluminium, iron, copper, magnesium, and zinc have been given in the following table

Terminology Used in Metallurgy:

Metallurgy:

The scientific and technological process used for the isolation of the pure metal from its ores is known as metallurgy.

Depending upon the nature of metal and the nature of ore, different methods are used in the extraction process. The different metals used are a) Pyrometallurgy, b) Hydrometallurgy and c) Electrometallurgy

• Pyrometallurgy: A process in which the ore is reduced to metal at high temperature using a suitable reducing agent like carbon, hydrogen, aluminium, etc. is called pyrometallurgy.
• Hydrometallurgy: A process of extraction of metals from aqueous solutions of their salts using suitable reducing agents is called hydrometallurgy.
• Electrometallurgy: A process of extraction of metals by electrolytic reduction of molten (fused) metallic compounds is called electrometallurgy.

Minerals:

A naturally occurring chemical substance obtained in mining which contains the metal in a free state or combined state is called mineral.

Ores:

A mineral containing a high percentage of metal from which the metal can be profitably extracted is called an ore. Ex: Copper glance (Cu₂S), Haematite ( Fe₂O₃). Ores may be divided into four groups

• Native Ores:  These ores contain the metal in free state eg. Silver gold etc.
• Oxidized Ores: These ores consist of oxides or oxy-salts (eg. carbonates, phosphate) and silicate of metal. e.g. Haematite, the ore of iron Fe₂O₃, Bauxite the ore of aluminium Al₂O₃.2H₂O, Cassiterite SnO₂, Cuprite Cu₂O, Zincite  ZnO, Ilmenite FeTiO₃, Rutile, TiO₂ etc.

• Important carbonate ores are limestone (CaCO₃), Calamine (ZnCO₃), Siferite, (FeCO₃), Cerrusite, (PbCO₃)  etc.
• Important sulphate ore is Anglesite (PbSO₄).

• Sulphurised Ores: These ores consist of sulphides of metals like iron, lead, mercury etc. e.g. iron pyrites (FeS₂). galena (PbS), Cinnabar (HgS). They are also called pyrites
• Halide ores: Metallic halides are very few in nature. Chlorides are most common examples include horn silver (AgCl) carnallite KCl. MgCl₂.6H₂O and fluorspar  (CaF₂), Cryolite, (Na₃AlF₆).
• Silicate Ores: Hemimorphite, (2ZnO.SiO₂.H₂O)

Gangue:

ores are usually contaminated with unwanted earthly(sandy) or undersides materials known as gangue. For example in the extraction of iron, silica is gangue present in haematite (Fe₂O₃).

Mineral Wealth in India:

Steps Involved in the Extraction of Metals:

The extraction and isolation of metals from ores involve the following major steps:

• Concentration of the ore (or) Purification of the Ore
• Conversion of ores into oxides or other desired compounds.
• Reduction of ores to form crude metals
• Refining of metals

Step – 1: The concentration of the ore:

The removal of earthy and siliceous impurities (i.e. gangue or matrix) from the ores is called concentration, dressing or benefaction. of ores. This process increases the percentage of the desired metal or metal compound in the ore.

It involves several steps and selection of these steps depends upon the differences in physical properties of the compound of the metal present and that of the gangue. The type of the metal, the available facilities and the environmental factors are also taken into consideration.

Hand-picking:

In this process the earthy impurities (heavy impurities) which are present in the ore like rocky materials, pellets are picked by hand.

Pulverizing:

Ore obtained from mines is in a lump form. It is finely ground to form powdered ore, the process is called pulverizing.

For pulverizing the ore,  jaw crushers or grinders are used.  The big lumps of the ore are brought in between the plates of a crusher forming a jaw. One of the plates of the crusher is stationary while the other moves to and fro and the crushed pieces are collected below.

The crushed pieces of the ore are then pulverized (powdered) in a stamp mill. The heavy stamp rises and falls on a hard die to powder the ore.

Hydraulic washing or Gravity Separation:

It is done by two ways a) Wilfley table method and b) Hydraulic classifier method

• Wilfley Table Method: The pulverized ore is made to fall through the hopper on the slanting floor of Wilfley’s table. The floor of Wilfley’s table is made up of fixed wooden triangular strips called cleats or riffles. This floor is continuously vibrating. A running stream of water is passed across the table. The lighter gangue particles are washed away and the heavier ore particles settle between wooden cleats of the table.
• Hydraulic Classifier Method: Hydraulic classifier consists of a tank having a hopper at the top and a pipe for carrying pressurized water. The ore is fed into the tank from the top, and the water is allowed with high pressure from the bottom of the tank. During this process the lighter impurities which are adhered to the ore float over the water which is removed by flowing water. The ore particles will settle down at the bottom of the tank. In this process, the lighter impurities are removed.

Magnetic separation:

In this method, electromagnetic separators are used. This method of concentration is used when either the ore or impurities associated with it are magnetic in nature.

This process is used for the Iron ore only. → The ore is passed through a belt which is connected by two rotating wheels, one among is made up of magnetic material. Once the ore is passed through the belt, the ore or impurity particles having magnetic properties are attracted to the magnetic wheel and fall near to it. While the impurities or ore particles which are nonmagnetic are not attracted by the magnetic wheel and fall away from the magnetic wheel as shown in the figure. Thus two separate heaps are formed.

Example: Cassiterite is ore of tin (Sn) contains non-magnetic stannic oxide SnO₂, the magnetic impurity tungstate or wolframite FeWO₄. They can be separated by this method.

Froth floatation:

Principle:  The surface of sulphide ores is preferentially wetted by oils while that of gangue is preferentially wetted by water. Hence this method is used for concentration of sulphide ores.

Process: In this process, a suspension of the powdered ore is made with water. Certain oils like pine oil, eucalyptus oil, xanthates or fatty acids are added as collectors along with froth stabilisers. Collectors are the chemical substances which enhance non-wettability of the mineral particles. Froth stabilizers are the chemical substances stabilise the froth. e.g.: cresols, aniline.

A rotating paddle agitates the mixture and draws air in it. As a result, a froth is formed which carries the mineral particles. The froth is light and is skimmed off. It is then dried for recovery of the ore particles

Depressants are the chemical substances which help to separate two sulphide ores by adjusting the proportion of oil to water. For e.g., in case of an ore containing ZnS and PbS, the depressant used is NaCN. It selectively prevents ZnS from coming to the froth but allows PbS to come with the froth.

Leaching:

Leaching is a process in which powdered ore is treated with a suitable reagent which is selectively dissolved the ore but not the impurities. This process is often used if the ore is soluble in some suitable solvent. It is a chemical method used for purification of ore.

Leaching of Alumina from Bauxite:

The principal ore of aluminium bauxite consists of three chemical impurities ferric oxide (Fe₂O₃), silica (SiO₂) and titanium oxide (TiO₂).

Bayer’s Process:

This method is used when major impurity is ferric oxide (Fe₂O₃)

The ore is treated with Sodium hydroxide solution (NaOH), at about 423 K. The impurity ferric oxide (Fe₂O₃) does not react with NaOH and can be removed by filtration.

Al₂O₃.2 H₂O_(s) +  2NaOH_(aq)  →  2NaAlO₂ (aq) + 3 H₂O_(l)

The solution is filtered to remove insoluble impurities and is agitated with freshly prepared Al(OH)₃, so that the aluminium in NaAlO₂ get precipitated to Al(OH) ₃.

2NaAlO₂ + 2 H₂O →  NaOH +  Al(OH) ₃

The precipitate obtained in the process is washed, dried and heated to get Al₂O₃.

Al(OH₃  →Heat  Al₂O₃  +   3 H₂O

Hall’s Process:

In this process, the ore and sodium carbonate are fused together to convert aluminium oxide into soluble sodium meta aluminate.

Al₂O₃.2 H₂O_(s)  +  Na₂CO_3(aq)   →  2NaAlO_{2 (aq)}  + CO_2(g)  + 2H₂O_(l)

The insoluble residue contains the impurities of silica and iron oxide. The filtrate is warmed and neutralized by passing carbon dioxide through the solution to produce aluminium hydroxide.

2NaAlO₂ + 3 H₂O   +  CO_2(g) →  2Al(OH) ₃   +  Al(OH)₃  +  Na₂CO_3(aq)

The precipitate obtained in the process is washed, dried and heated to get Al₂O₃.

Al(OH₃  →Heat  Al₂O₃  +   3 H₂O

Leaching of Silver and Gold:

In the metallurgy of silver and that of gold, the respective metal is leached with a dilute solution of NaCN or KCN in the presence of air (for O₂) from which the metal is obtained later  by replacement:

4M_(s) + 8CN^– _(aq) + 2H₂O  + O_2(g)   → 4[M(CN) ₂] ^– _(aq) + 4OH^–_(aq)  (M= Ag or Au)

2[M (CN) ₂] ^–  _(aq)   + Zn_(s) → [Zn(CN)4] ^{– –} _(aq)   + 2 M_(s)

Step – 2: Extraction of Crude Metal from Concentrated Ore:

The concentrated ore must be converted into a form which is suitable for reduction. Usually, the sulphide ore is converted to oxide before reduction. Oxides are easier to reduce. Thus isolation of metals from concentrated ore involves two major steps viz., (a) conversion to oxide, and (b) reduction of the oxide to metal.

Roasting:

Roasting is a process in which ores are heated to a high temperature below their melting point in the presence of an excess of air.

Calcination is generally used for ores containing carbonates and hydrated oxides.

Roasting is carried out in a reverberatory furnace  This process is usually used for sulphide ores. During this process the moisture is removed, organic matter is destroyed and non-metallic impurities like that of S, P and As are oxidized and removed as volatile gases.

S₈ + 😯₂    →  8SO₂­­↑ (Sulphur dioxide)

P₄ + 5O₂  → 2P₂O₅­ ↑ (Phosphorus pentaoxide)

4As + 3O₂   → 2As₂O_3­↑ (Arsenious oxide)

Ores are generally converted into metallic oxides

2ZnS (Zinc sulphide) + 3O₂   →  2ZnO (Zinc oxide) + 2SO₂­↑

2PbS (Lead oxide) + 3O₂    →   2PbO (Lead oxide) + 2SO₂­­↑

2Cu₂S (Cuprous sulphide) + 3O₂    →  2Cu₂O (Cuprous oxide) + 2SO₂­­↑

Calcination:

Calcination is a process in which the ore is heated to a high temperature below its melting point in the absence of air or in a limited supply of air.

It is carried out in a reverberatory furnace  In this process moisture, volatile impurities like carbon dioxide, sulphur dioxide are expelled from the ore. The ore is made porous. Carbonate ores decompose to form a metal oxide and carbon dioxide.

ZnCO₃ + Heat →ZnO_(s) + CO_2(g) ↑

CaCO₃ + Heat →CaO_(s) + CO_2(g) ↑

CaCO₃.MgCO₃ + Heat →CaO_(s) + MgO_(s) + 2CO_2(g) ↑

2Fe₂O₃ . 3H₂O + Heat →2Fe₂O₃ + 3H₂O_(g) ↑

The post Basic Principles of Metallurgy appeared first on The Fact Factor.

In this article, we shall study hydroxides of third row elements.

Different types of hydroxides of third-row elements are classified according to their mode of dissociation. According to Arrhenius’s theory, the acid is the substance which gives H⁺ ions in an aqueous medium, while the base is the substance which gives OH^– ions in an aqueous medium.

Let us use M-O-H be the general formula to represent a hydroxy compound of third row elements. The mode of ionization decides the nature of hydroxide whether it is acidic or basic.

Types of Hydroxy Compounds:

Basic hydroxy compound:

The hydroxy compounds which give OH^– ions in an aqueous medium are called basic hydroxides.

M — O — H _(aq.)    →    M⁺   +   OH^–

This is possible when the element has very low ionization potential. Valence electrons are loosely held by the atom. Due to very low ionization enthalpy and electronegativity metal atom cannot hold valence electrons. Electron pair between M and O is pulled more towards more electronegative oxygen. M-O bond becomes weak while the O-H bond becomes strong. Thus the bond between M and OH breaks

e.g. hydroxide NaOH of sodium and hydroxide Mg(OH)₂ of magnesium are basic hydroxy compounds.

NaOH _(aq.)    →  Na ⁺    +   OH^–

Mg(OH) _{2  (aq.)} →    Mg ^{2 +}  +  2 OH ^–

Both hydroxides give OH^– ions in aqueous medium.

Acidic Hydroxy Compounds (Oxyacids): 

The hydroxy compounds which give OH – ions in an aqueous medium are called basic hydroxides.

M — O — H _(aq.)    →   MO ^–   +   H ⁺

This is possible when the element has greater ionization potential. Valence electrons are strongly held by atom. Due to higher ionization enthalpy and electronegativity, metal atom hold valency electrons strongly. Electron pair between M and O is pulled more towards more electronegative element M. As a result of which bond M-O-H behaves as a base.

M-O bond becomes strong while the O-H bond becomes weak. Thus the bond between MO and H breaks, which results in the production of H + ions in aqueous medium.

The hydroxy compounds which give H ⁺ ions in an aqueous medium are called acidic hydroxy compounds or oxyacids. If M-O-H is assumed oxyacid then O-H bond break in an aqueous medium and it will give H ⁺ ions.

e.g. Si(OH)₄ of Silicon, P (OH)₃ and PO (OH)₃ of Phosphorous, SO (OH)₂ and SO₂(OH)₂ of sulphur , ClOH, ClO(OH) , ClO₂(OH) and ClO₃(OH) of chlorine are acidic hydroxy compounds.

Amphoteric Hydroxy Compound:

Hydroxy compound which acts as acid as well as base and can neutralize the acid, as well as base producing salt and water, is called amphoteric hydroxide.

Al(OH)₃ of aluminium is an amphoteric oxide. It neutralizes acid as well as base producing a salt and water.

Al (OH)₃ (as a base) +   3HCl    →  AlCl ₃  + 3 H₂O

Al(OH)₃ (as an acid) + NaOH  → NaAlO₂   (Sodium meta-aluminate) +2 H₂O

Thus the nature of hydroxides or oxyacids is mainly governed by the ionization potential of elements.

Trends in acid -base behaviour of hydroxy compounds:

It is seen that as we move from Na to Cl along the third row, the basic character of hydroxy compounds gradually decreases while acidic character gradually increases.

The trend in acid-base behaviour of hydroxy compounds of the third row can be summarized as follows.

Explanation:

The nature of the hydroxy compound mainly depends upon the ionization potential of elements. The trend is so because along third-row ionization potential increase, electronegative character increases, atomic size decreases. Along the third row difference in electronegativity of the element M and that of Oxygen decreases.

If ionization potential of elements is low, then such hydroxy compound give OH ^– ions in an aqueous medium and hence is basic in nature.

M — O — H _(aq)        →   M ⁺  +  OH ^–

NaOH and Mg(OH)₂ are basic. Na and Mg have low ionization potential. Na – O and Mg – O bonds are weaker than O-H bond. And thus the bond between M and O breaks to give OH^– ions.

NaOH _(aq.)    →  Na ⁺    +   OH^–

Mg(OH) _{2  (aq.)} →    Mg ^{2 +}  +  2 OH ^–

If ionization potential of an element is greater, then such hydroxy compound gives H ⁺ ions in an aqueous medium and hence is acidic in nature.

M — O — H _(aq)        →   M O^–  +  H ⁺

Si(OH)₄ of Silicon, P(OH)₃ and PO(OH₃ of Phosphorous, SO(OH)₂ and SO₂(OH)₂ of sulphur, ClOH, ClO(OH), ClO₂(OH) and ClO₃(OH) of chlorine are acidic hydroxy compounds.

Al(OH)₃ is amphoteric. It neutralizes acid as well as base producing salt and water. Hence it exhibits both the properties hence it is an amphoteric oxide.

Scientific Reasons:

Hydroxy compound of Sodium NaOH is a strong base.

The hydroxy compounds which give OH – ions in an aqueous medium are called basic hydroxides.

M — O — H _(aq)        →   M ⁺  +  OH ^–

This is possible when the element has very low ionization potential. Valency electrons are loosely held by atom. Due to very low ionization potential and electronegativity, metal atom cannot hold valency electrons. Electron pair between M and O is pulled more towards more electronegative oxygen. M-O bond becomes weak while the O-H bond becomes strong.

Sodium is the strongest electropositive element. Na has lower ionization potential and lower electronegativity. Na -O bond breaks more readily in an aqueous medium and Sodium hydroxide ionizes as

NaOH _(aq.)    →  Na ⁺    +   OH^–

Mg (OH)₂ is weakly basic than NaOH.

The hydroxy compounds which give OH – ions in an aqueous medium are called basic hydroxides.

M — O — H _(aq)        →   M ⁺  +  OH ^–

This is possible when the element has very low ionisation potential. Valency electrons are loosely held by atom. Due to very low ionisation potential and electronegatively metal atom cannot hold valency electrons. Electron pair between M and O is Pulled more towards more electronegative oxygen. M-O bond becomes weak while the O-H bond becomes strong.

Na has lower ionisation potential and lower electronegativity than that of Mg. So Na – O bond is relatively weak than Mg – O bond.  Na -O bond breaks more readily than the Mg-O bond in an aqueous medium. Thus Mg(OH)₂ is weakly basic than NaOH.

Aluminium hydroxide is amphoteric compound.

Hydroxy compound which acts as acid as well as base and can neutralise acid, as well as base producing salt and water, is called amphoteric hydroxide.

Al(OH) ₃ of aluminium is an amphoteric oxide. It neutralises acid as well as base producing a salt and water.

Al (OH)₃ (as a base) +   3HCl    →  AlCl ₃  + 3 H₂O

Al(OH)₃ (as an acid) + NaOH → NaAlO₂   (Sodium meta-aluminate) + 2 H₂O

It acts as a base when treated with strong acid. It acts as an acid when treated with a strong base. Due to its dual character, Al(OH)₃ is amphoteric in nature.

Atomic size of Aluminium is smaller than Sodium and Magnesium and larger than Silicon, Phosphorous, Sulphur and Chlorine. The ionisation potential of Aluminium is larger than Sodium and Magnesium and lesser than Silicon, Phosphorous, Sulphur and Chlorine. Electronegativity of Aluminium is larger than Sodium and Magnesium and lesser than Silicon, Phosphorous, Sulphur and Chlorine. Thus in aluminium hydroxide, both Al – O and O – H bonds have equal strength. Hence the fission of bond depends on the attacking reagent. Hence Aluminium hydroxide is amphoteric compound.

Orthosilicic acid Si(OH)₄ is very weak acid. 

Electronegativity of Silicon is 1.8 units. The ionisation potential of Silicon is higher than that of Aluminium.

Si – O is covalent bond with the ionic character. Hence Si – O bond is a strong bond. Oxygen atom pulls the shared electron in O – H bond towards itself and on fission produces H⁺ ions. However, O-H bond is not broken in water.

The orthosilicic acid reacts with strong alkali on heating.

H₄SiO₄  +  2 NaOH  →  Na₂Si ₃   +  3 H₂O

Hydroxy compounds of Phosphorous, Sulphur And Chlorine are acidic. 

Phosphorous, Sulphur And Chlorine are non-metals with smaller atomic size, high nuclear charge, High ionisation potential and high electronegativity. These elements have very little or practically no tendency to give an electron to Oxygen.

In the structure of M – O – H Oxygen, therefore, tries to pull electron pair between O – H towards itself, This releases H⁺ ions in aqueous solution.

Structures of Hydroxy Compounds of Phosphorous or Oxyacids of Phosphorous:

Phosphorous acid (P(OH)₃ OR H₃PO₃) and  Phosphoric  acid (PO(OH)₃ OR (H₃PO₄)

H₃PO₃  +  2 NaOH  →  Na₂HPO₃  +  2 H₂O

H₃PO₄  +  3 NaOH  →  Na₃PO₄  +  3 H₂O

Phosphoric acid has an un-hydrogenated oxygen atom so O-H bond breaks readily. Hence phosphorous acid is stronger than phosphorous acid.

Structures of Hydroxy compounds of Sulphur or Oxyacids of Sulphur:

Sulphurous acid SO(OH)₂ OR (H₂SO₃) and Sulphuric  acid SO₂(OH)₂ OR (H₂SO₄)

H₂SO₃  +  2 NaOH  → Na₂SO₃  +  2 H₂O

H₂SO₄  +  2 NaOH  → Na₂SO₄  +  2 H₂O

These oxyacids of sulphur are strongly acidic due to greater IP of sulphur.  Sulphuric acid has two un-hydrogenated oxygen atoms while in sulphurous acid there has one un-hydrogenated oxygen atom hence, sulphuric acid is stronger than sulphurous acid.

” Greater the Oxidation no. More the Acidic Nature”. In H₂SO₄ the oxidation no. of the central atom Sulphur is +6 and that in H₂SO₃ is +4. The oxidation number of sulphur is greater in sulphuric acid than the sulphurous acid. Hence sulphuric acid is stronger than sulphurous acid.

Hydroxy compounds of Chlorine or Oxyacids of Chlorine:

Hypochlorous acid Cl(OH) OR (HOCl), Chloric acid ClO₂(OH) OR (HClO₃), Perchloric  acid ClO₃(OH) OR (HClO₄)

HOCl   +  NaOH  →   NaOCl    +   H₂O

HClO₄  +  NaOH  →  NaClO₄  +  —– H₂O

These are oxyacids of chlorine which are very strongly acidic. Chlorine has greater I.P and electronegativity. Due to the presence of 3 un-hydrogenated oxygen atoms attached to Chlorine atom,  HClO₄ is strongest amongst these oxyacids of chlorine.

The strengths are in the order are ClO₃(OH) > ClO₂ (OH)   > ClO(OH)  >  ClOH.

Science > Chemistry > Third Row Elements > Hydroxides of Third Row Elements

The post Hydroxides of the Third Row Elements appeared first on The Fact Factor.

A binary compound of an element with oxygen, in which the oxygen atom is electronegative is called an oxide. e.g. MgO, Al₂O_3, etc. The oxide in which Oxygen exhibits the normal oxidation state of -2 is called normal oxide. In this article, we shall study the oxides of third-row elements.

Classification of Oxides:

Depending upon the chemical behaviour, oxides of third-row elements are Classified into three types, namely i) acidic oxides  Na₂O, MgO ii) basic oxides SiO₂, SO₃, Cl₂O_7, P₂O₅  and iii) amphoteric oxides Al₂O₃.

Basic Oxides:

The oxides which react with water and produce alkali and can neutralise acids forming salt and water are called basic oxides.

Na₂O of sodium and MgO of magnesium of the third row are basic oxides. Because both oxides produce alkali when treated with water and can neutralise acid producing salt and water.

Na₂O    +     H₂O        →    2 NaOH (strong base)

Na₂O    +   2 HCl        →   2 NaCl    +   H2O

MgO    +   2 H₂O        →   Mg (OH)₂ (base)

MgO    +   2 HCl        →   MgCl₂    +   H₂O

Na and Mg have bigger atomic size. 1 and 2 valence electrons respectively and very low ionisation potential value.

Amphoteric Oxide:

The oxide which acts as an acid as well as base and neutralizes acid as well as base to give salt and water is called an amphoteric oxide.

Al₂O₃ of aluminium is an amphoteric oxide. It neutralizes acid like HCl as well as a base like NaOH.

Al₂O₃ is acidic in nature because it reacts with a base to give salt and water.

Al₂O₃    +    2 NaOH  →   2 NaAlO₂  (sodium aluminate)   +      H₂O

Al₂O₃ is acidic in nature because it reacts with a base to give salt and water.

Al₂O₃    +    6 HCl  → 2 AlCl₃    +   3 H₂O

Aluminium has higher I.P. than sodium and its electronegativity is greater than Mg and Na metals. So Al-O bond in  Al₂O₃  shows amphoteric in nature.

Acidic Oxides:

The oxides of electronegative elements which give acid when treated with water and can neutralise bases to produce salt and water called acidic oxides.

Oxides of Phosphorous, Sulphur, Chlorine are acidic.

SiO₂   +  2 NaOH →   Na₂SiO₃   + H₂O

P₄O₁₀ +   6 H₂O  → 4 H₃PO₄

SO₃   +  H₂O    →     H₂SO₄

SO₃    +  2 Na OH  → Na₂SO₄  +  H₂O

Cl₂O₇   + H₂O →  2 HClO₄ (perchloric acid)

Cl₂O₇  + 2 Na OH  →  2 NaClO₄ +   H₂O

Si, P ,S, and Cl have greater ionisation potential , electronegativity , tendency to attract electrons.  They have smaller atomic size and more valency electrons.

Trend in Acidic and Basic Character:

Trend:

Generally, the oxides of metals are basic in nature while that of non-metals are acidic. As we move from Na to Cl along the third period, the acidic character of elements goes on gradually increasing while basic character goes on decreasing.

Reasons:

• In third row elements, those elements who have low ionisation potential, bigger atomic size, less electronegativity, less number of valence electrons form basic oxides as their oxides give alkalies when treated with water.
• On the other hand, those elements who have greater ionisation potential, smaller atomic size, greater electronegativity, more valency electrons form acidic oxides, Such oxides give acids when treated with water.
• The trend is so because from Na to Cl while going along the third period, ionisation potential gradually increases, atomic size decreases, the number of valence electrons increases.

Examples:

Na₂O and MgO of Na and Mg respectively are highly basic in nature. They produce alkalies when treated with water.  They neutralise acids. Na and Mg have very low ionisation potential amongst third-row elements.

Na₂O    +     H₂O        →    2 NaOH (strong base)

Na₂O    +   2 HCl        →   2 NaCl    +   H2O

MgO    +   2 H₂O        →   Mg (OH)₂ (base)

MgO    +   2 HCl        →   MgCl₂    +   H₂O

The oxides of electronegative elements such as Si, P, S. and Cl are acidic in nature.  Their oxides give acids when treated with water can neutralise base producing salt and water. Si, P, S and Cl elements have greater ionisation potential, more valency electrons, greater electronegativity.

SiO₂   +  2 NaOH →   Na₂SiO₃   + H₂O

P₄O₁₀ +   6 H₂O  → 4 H₃PO₄

SO₃   +  H₂O    →     H₂SO₄

SO₃    +  2 Na OH  → Na₂SO₄  +  H₂O

Cl₂O₇   + H₂O →  2 HClO₄ (perchloric acid)

Cl₂O₇  + 2 Na OH  →  2 NaClO₄ +   H₂O

Al₂O₂ of aluminium is amphoteric because it neutralises acid as well as base producing salt and water. Due to its dual nature, it is called amphoteric.

Al₂O₃ is acidic in nature because it reacts with a base to give salt and water.

Al₂O₃    +    2 NaOH  →   2 NaAlO₂  (sodium aluminate)   +      H₂O

Al₂O₃ is acidic in nature because it reacts with a base to give salt and water.

Al₂O₃    +    6 HCl  → 2 AlCl₃    +   3 H₂O

The oxides of third-row elements are summarised in the following table.

Scientific Reasons:

Sodium oxide is more basic than Magnesium oxide. 

A binary compound of an element with oxygen, in which Oxygen atom is electronegative is called an oxide. The oxides which react with water and produce alkali and can neutralise acids forming salt and water are called basic oxides.

Na₂O, MgO are basic oxides. These oxides produce alkali when treated with water and can neutralize acid-producing salt and water. Sodium is a more electropositive element than Magnesium. Sodium has lower ionization potential and lower electronegativity than Magnesium.

Due to the above reasons Sodium loses its electron very easily to oxygen than Magnesium and acts as more basic than Magnesium.

Science > Chemistry > Third Row Elements > Oxides of the Third Row Elements

The post Oxides of the Third Row Elements appeared first on The Fact Factor.

in this article, we shall study the crystal structure of molecular solids of third perid of periodic table.

Molecular solid:

The substance in which lattice points are molecules which are held together by means of weak physical forces (van der Waal’s forces) are called molecular solid. Phosphorous, sulphur, chlorine, and argon are molecular solids because lattice points are molecules.

They have greater ionization potential and vacant valency orbitals are not available.

Characteristics of Molecular Solids:

• In the crystal structure of these elements, the units occupying lattice points are molecules
• They have a greater ionization potential
• The molecules are attached to each other by weak van der Wall’s forces of attraction.
• In these solids, the atoms are joined together within the molecule by strong covalent bonds.
• In these solids, vacant valency orbitals are not available. All the valence orbitals are used for intra-molecular strong covalent bonding.

Scientific Reasons:

Phosphorous, Sulphur, Chlorine, and Argon as solid are soft easily compressible and distorted. They are volatile their boiling points and melting points are very low.

In the crystal structure of these elements, the units occupying lattice points are molecules. They have a greater ionization potential. The molecules are attached to each other by weak van der Wall’s forces of attraction. Within the molecules, the atoms are joined together by strong covalent bonds. Thus they are soft easily compressible and distorted.

Similarly less energy is required to separate the molecules from each other in a molecular crystal. Hence they are volatile and possess low boiling points and melting points.

Phosphorous and Sulphur are solids at room temperature while Chlorine and Argon are gases at room temperature.

In the crystal structure of these elements, the units occupying lattice points are molecules. They have a greater ionization potential. The molecules are attached to each other by weak van der Wall’s forces of attraction. Within the molecules, the atoms are joined together by strong covalent bonds.

Sulphur is Octaatomic, Phosphorous is tetra atomic, Chlorine is diatomic, Argon is monoatomic. The size of the molecule decreases in the order S₈; P₄ > Cl₂ > Ar. van der Wall’s forces of attraction decrease in the same order. They are stronger in phosphorous and sulphur while negligible in chlorine and argon. Thus Chlorine and Argon have very low boiling and melting points. Hence Phosphorous and Sulphur are solids at room temperature while Chlorine and Argon are gases at room temperature.

Structure of Molecular Solids:

Structure of Phosphorous:

Phosphorous is molecular solid because lattice points are P₄ molecules which are held together by means of van der Wall’s forces of attraction. Within the molecules, the atoms are joined together by strong covalent bonds.

The atomic number of Phosphorous is 15. Electronic configuration of phosphorous is, 1s2, 2s22p6, 3s2 , 3px1 3py1 3pz1. Phosphorous has greater ionization potential. Due to the non-availability of vacant valency orbitals, it is a molecular solid. Phosphorous atoms are sp³ hybridized.

One phosphorous atom forms three covalent bonds with three other phosphorous atoms while one sp³ hybrid orbital contained paired electrons (lone pair) remain non-bonded. Thus tetrahedral P₄ molecule is formed.  The P-P-P bond angle is 60°. The phosphorous molecule consists of four phosphorous atoms in it.

In white phosphorous, these tetrahedral molecules are joined together by means of weak van der Waal’s forces of attraction due to smaller molecular size.  So white phosphorous has a low melting point.

Red phosphorous is a polymeric form. In red phosphorous one P-P bond of P₄ unit gets ruptured and freed bonds link up to form a chain. Thus in red phosphorous, the tetrahedral molecules are joined to one another by means of strong covalent bonds, forming a chain like structure. Hence red phosphorous exists in polymeric form.

In white phosphorous, the molecules are joined together by weak van der Wall’s forces while in red phosphorous molecules are bound by strong covalent bonds. Therefore more energy is required for breaking bonds between molecules of red phosphorous than that in white phosphorous. Hence red phosphorous shows a high melting point and less reactivity than white phosphorous.

Structure of Sulphur:

Sulphur is molecular solid because lattice points are S₈ molecules which are held together by means of van der Waal’s forces of attraction. Within the molecules, the atoms are joined together by strong covalent bonds.

The atomic number of Sulphur is 16. Electronic configuration of phosphorous is 1s2, 2s22p6, 3s2 , 3px2 3py1 3pz1. Sulphur has greater ionization potential. Its two half-filled orbitals are used for intermolecular bonding. Due to the non-availability of vacant valency orbitals, it is a molecular solid. Each Sulphur molecule consists of eight sulphur atoms in it.

The sulphur molecule has a puckered ring structure or crown structure. Sulphur has two half-filled 3p orbitals in the valency shell. Within each S₈ molecule, each s atom is linked to two adjacent S atoms by the covalent single bond. Thus if seen from the top it forms a ring-like structure with four sulphur atoms in one plane and alternate other four in a parallel plane. Each Sulphur atom has a lone pair of electrons. The S-S-S bond angle is 107.8 ⁰ and S-S bond length is 2.04 ^o

Structure of Chlorine:

Chlorine is molecular solid because lattice points are Cl₂ molecules which are held together by means of van der Waal’s forces of attraction. Within the molecules, the atoms are joined together by strong covalent bonds.

The atomic number of Chlorine is 17. Electronic configuration of phosphorous is 1s², 2s²2p⁶, 3s² , 3p_x² 3p_y² 3p_z¹. Chlorine has greater ionization potential. Hence it forms covalent bonding. Its one half-filled orbital is used for intermolecular bonding. Due to the non-availability of vacant valency orbitals, it is a molecular solid. Each chlorine molecule consists of two Chlorine atoms in it bonded by a covalent bond.

Chlorine in solid state consists of layers of chlorine molecules which are held by weak Wander Wall’s forces. Hence Chlorine is gas at room temperature.

Structure of Argon:

The substance in which lattice points are molecules which are held together by means of weak physical forces (vander waal’s forces) are called molecular solids.

Argon is a molecular solid or atomic solid because lattice points are Ar atoms (molecules) which are held together by means of van der Waal’s forces of attraction. The atomic number of Argon is 18. Electronic configuration of Argon is 1s2, 2s22p6, 3s2 , 3px2 3py2 3pz1. Thus its octet is complete. It has no unpaired electrons. Argon has greater ionization potential. Due to the non-availability of vacant valency orbitals, it is a molecular solid. Each Argon molecule consists of one Argon atom.

Argon crystal in its solid-state consists of a continuous pattern of atoms giving rise to a face centred closed pack cubic lattice like aluminium. But the difference is that in the case of aluminium lattice points are aluminium ions while in case of argon lattice points are Argon atoms. Melting and boiling points of argon are very low.

Science > Chemistry > Third Row Elements > Molecular Solids of the Third Row

The post Molecular Solids of the Third Row appeared first on The Fact Factor.