• Some organic compounds contain only C and H as the constituent elements. They are known as hydrocarbons.

• On the basis of the structure, hydrocarbons are divided into two main groups called aliphatic and aromatic.

• The set of hydrocarbons consisting of open carbon chains only are named as acyclic aliphatic hydrocarbons.

• The aliphatic hydrocarbons are classified as alkanes, alkenes, and alkynes.

• The cyclic organic compounds which are stabilized by forming a cyclic delocalized cloud of π electrons are called aromatic compounds.

• Benzene which is indicated by the molecular formula C6H6 is the simplest of aromatic hydrocarbon compounds.

• Compounds formed by replacing a hydrogen atom of an aliphatic hydrocarbon by a halogen atom are referred to as alkyl halides.

• Compounds formed by replacing a hydrogen atom of a benzene ring by a halogen atom are referred to as aryl halides.

• In many organic compounds, when hetero atoms such as nitrogen and oxygen combine with the carbon chain, due to the difference in electro negativity between the carbon and the combined atoms, this group of atoms will impart to the compound a characteristic reactivity. Such a group of atoms is called a functional group. The compounds are classified according to the functional group present in a molecule.

 

Common functional groups and the names of the corresponding homologous series are given below.

* In the IUPAC nomenclature, halogen is not considered a functional group

 

• The International Union of Pure and Applied Chemistry -IUPAC has developed a method for systematically naming organic compounds.

• The name given to a compound according to the IUPAC nomenclature consists of several parts.

1. The suffix that is used to indicate the main functional group of the structure

2. Name of the chain that is used to identify the main carbon chain of the compound

3. The prefixes that are used to indicate the substituent groups

4. The numbers that are used to indicate the places at which the substituent groups, additional groups and the main functional groups are attached to the chain

• The IUPAC name of an aliphatic compound can easily be developed by following the steps stated below in the given order.

1. Identifying the principal functional group

2. Selecting the main chain

3. Selecting the root name for the principal chain

4. Addition of the suffix for the double/triple bond in the main carbon chain to the name of the chain

5. Addition of the suffix used to indicate the principal functional group to the name of the chain

6. Naming the substituent groups

7. Adding the names of the substituent groups to the name of the chain

8. Numbering the carbon chain

9. Writing the numbers that are used to indicate the positions of the main functional group and the substituent groups in front of these groups.

The series of functional groups arranged in the decreasing order of their priority

 

The root names used for the compounds according to the number of carbon atoms and the name of the corresponding hydrocarbon

No. of Carbons	Root Name	Formula CnH2n+2	Structure
1	methane	CH4	CH4
2	ethane	C2H6	CH3CH3
3	propane	C3H8	CH3CH2CH3
4	butane	C4H10	CH3CH2CH2CH3
5	pentane	C5H12	CH3CH2CH2CH2CH3
6	hexane	C6H14	CH3CH2CH2CH2CH2CH3
7	heptane	C7H16	CH3CH2CH2CH2CH2CH2CH3
8	octane	C8H18	CH3CH2CH2CH2CH2CH2CH2CH3
9	nonane	C9H20	CH3CH2CH2CH2CH2CH2CH2CH2CH3
10	decane	C10H22	CH3CH2CH2CH2CH2CH2CH2CH2CH2CH3

 

Drawing the structural formula of a compound according to the IUPAC nomenclature

• Drawing the structural formula of a compound according to the IUPAC nomenclature can be done by following the steps given below.

1. Identifying the chain and drawing the chain according to that name

2. Numbering the chain

3. Identifying the principal functional group and the remaining groups according to the IUPAC name given and joining these groups to the correct places of the chain according to the number in front of these groups

4. Placing hydrogen atoms in the chain structure so that each carbon atom has a valency of four

Benzene Derivatives

The nomenclature of substituted benzene ring compounds is less systematic than that of the alkanes, alkenes and alkynes. A few mono-substituted compounds are named by using a group name as a prefix to “benzene”, as shown by the combined names listed below. A majority of these compounds, however, are referred to by singular names that are unique. There is no simple alternative to memorization in mastering these names.

When more than one substituent is present on a benzene ring, the relative locations of the substituents must be designated by numbering the ring carbons or by some other notation. In the case of disubstituted benzenes, the prefixes ortho, meta & para are commonly used to indicate a 1,2- or 1,3- or 1,4- relationship respectively. In the following examples, the first row of compounds show this usage in red. Some disubstituted toluenes have singular names and their isomers are normally designated by the ortho, meta or para prefix. Finally, if there are three or more substituent groups, the ring is numbered in such a way as to assign the substituents the lowest possible numbers, as illustrated by the last row of examples. The substituents are listed alphabetically in the final name. If the substitution is symmetrical (third example from the left) the numbering corresponds to the alphabetical order.

• The phenomenon of the existence of compounds with different atomic arrangements while having the same molecular formula is called isomerism.

• The isomers of a compound can show different physical and chemical properties from one another.

• The isomerism of organic compounds can be classified as follows.

Structural isomerism

Structural isomerism is the phenomenon of having compounds with different structural formulae (without considering the spatial orientation of bonds) for the same molecular formula.

Chain isomerism

Chain isomerism occurs when the nature of the carbon chain changes for the same molecular formula in the same homologous series

Position isomerism

Though there is the same molecular formula, the same functional group/substituent and the same carbon chain when there is a change in the carbon atom to which the functional group/substituent is attached or a change in the location of the active position, then there occurs position isomerism.

Functional group isomerism

Functional group isomerism is the existence of structures with different functional groups for the same molecular formula.

Stereoisomerism

Stereoisomerism is the existence of compounds whose structures differ from each other only in the orientation of bonds in three dimensional space (i.e. they have the same molecular formula and the same structural formula).

A pair of stereoisomers whose 3D – structures are mirror images of each other are enantiomers of each other. A pair of stereoisomers whose 3-D structures are not mirror images of each other are diastereomers of each other.

 

Diastereomerism

Geometric isomerism is one occasion where diastereomerism is seen. In a C=C double bond due to the π bond which exists in addition to the σ bond, these carbon atoms cannot freely rotate about the bond. It is possible to have different spatial arrengements of the groups joined to the two carbon atoms. These different arrangements which cannot be interconverted by rotation around carbon – carbon bond axis are known as geomentrical isomers. For geometrical isomers to exist, the two groups attached to each carbon atom of the double bond should not be identical

The words cis and trans are used to indicate the geometrical relationship of two groups attached to different carbon atoms in the same double bond. If the two groups are on the same side with reference to the plane which is perpendicular to the plane of the molecule going through the carbon – carbon axis of the double bond, then the relationship is cis. If the two groups are on opposite sides of the plane then the realtionship is trans.

For example

Enantiomerism

The isomers of which one is the mirror image of the other are known as enantiomers. A compound having a carbon atom which is joined to four different groups shows enatiomerism. Such a carbon atom is known as an asymmetric or chiral carbon atom

When plane – polarized light is passed through a solution containing only one enantiomer, the plane of polarization rotates. One enantiomer rotates the plane of polarization in one direction and the other enantiomer in the opposite direction. As the enantiomers rotate the plane of polarization, they are also known as optically active isomers

 

Reactions with water

• All the elements of the first group react with water liberating hydrogen and become hydroxides. Example : Na reacts rapidly with water liberating hydrogen.

2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂ (g)

• When a small piece of K is added into water it reacts while burning. As K reacts with water more rapidly than Na, it can be concluded that the rate of reaction with water increases down the group.

• A reaction is not seen when a clean piece of Mg is added into water. When the water with Mg is warmed, it is seen to react slowly.

Mg(s) + 2H₂O(l) → Mg(OH)₂ (s) + H₂ (g)

• As the reactivity shown by Mg with water is lower compared to Na, it can be said that the metals of group II compared to metals of group I show a lower reactivity.Be does not react with water. Ca, Sr, and Ba react with water liberating hydrogen and forming the hydroxides.

Ca(s) + 2H₂O(l) → Ca(OH)₂ (aq) + H₂ (g)

• Be and Mg react with steam to form the oxides.

Be(s) + H₂O(g) → BeO(s) + H₂ (g)

Mg(s) + H₂O(g) → MgO(s) + H₂ (g)

 

Reactions with air/O2

There are several reactions of group I metals with air .

4Na(s) + O₂ (g) → 2Na₂O(s)

2Na(s) + Excess O₂(g) → Na₂O₂(s)

2Na(s) + 2H₂O(g) → 2NaOH(s) + H₂ (g)

K, Rb and Cs readily react with O₂ forming superoxides.

Excess K(s) + O₂ (g) → KO₂ (s)

When heated in air only Li of Group I reacts with nitrogen.

6Li(s) + N₂ (g) → 2Li₃N(s)

When a clean piece of Mg ribbon and a small cut piece of Na are exposed to air Na tarnishes faster than Mg. Hence it is clear that the reactivity of Mg is lower than Na.  Accordingly it can be said that relative to metals of group I, the reactivity of group II metals with air is lower.

Metals of the Group II when heated in air burn forming oxides and nitrides.

2Mg(s) + O₂ (g) → 2MgO(s)

3Mg(s) + N₂ (g) → Mg₃N₂ (s)

For Be to react it should be heated to a very high temperature.

 

Reaction with acids

As the metals of the group I react with acids liberating large quantity of heat, an explosion takes place. Therefore it should not be tested.

2Na(s) + dil.H₂SO₄(aq) → Na₂SO₄(aq) + H₂(g)

Group II metals reacts with dilute acids to liberate H₂ rapidly

Mg(s) +  dil.H₂SO₄(aq) → MgSO₄(aq) + H₂(g)

Group II metals can be oxidised by concentrated acids.

Mg(s) + 2 conc.H₂SO₄(aq) → MgSO₄(aq) + SO₂(g) + H₂O(l)

Mg(s) +  4 conc.HNO₃(aq) → Mg(NO₃)₂(aq) + 2NO₂(g) + 2H₂O(l)

Reactions with N₂

Lithium and Group II metals form metal nitrides by reacting with N₂ gas or N₂ presents in air.

6Li(s) + N₂(g) → 2Li₃N(s)

3Mg(s) + N₂(g) → Mg₃N₂(s)

 

Reactions with H₂

s Block elements form metal hydrides by reacting with H₂ gas.

2Li(s) + H₂(g) → 2LiH(s)

2Na(s) + H₂(g) → 2NaH(s)

Mg(s) + H₂(g) → MgH2(s)

 

• Because s block elements easily remove their electrons and form cations they are considered as good reducing agents.

• While atomic radius increases down the group nuclear attraction decreases. Consequently, reducing ability of the elements also increases.

 

Alkanes

• In alkanes all the bonds are either C-C or C-H bonds. The polaritiy of those C-C and C-H bonds are low. Therefore, they do not react with common polar reagents (eg. OH^– , CN^– , H⁺ ) under normal conditions.

Chlorination

Alkanes react with the reagents such as Cl₂ and Br₂ that undergo easy homolytic cleavage to generate free radicals.

The mechanism of the chlorination of methane is given below.

Initiation step

Chain propagation steps

The Cl. free radical generated in the initiation step abstracts a H atom from the methane molecule by homolysis of the C-H bond.

 

The .CH₃ free radical then reacts with a Cl₂ molecule forming CH₃Cl and generate
another .Cl free radical which can continue the chain reaction.

Chain termination steps
Combination of free radicals in the reaction mixture to form stable molecules, results in the termination of chains.

Reactions of alkenes

• The loosely bound π electron cloud which lies above and below the plane of the
ethylene molecule is capable of attracting electrophilic reagents.

• A molecule or an ion that can accept a pair of electrons is referred to as an electrophile.

Addition of hydrogen halides

• Here, the hydrogen atom which is the electron deficient pole of HBr molecule acts as an electrophile and attacks the double bond. During these electrophilic addition reactions, intermediate carbocations are formed.

 

Stability of carbocations follows the following order.

• When alkyl groups are attached to the positively charged C atom of the carbocation, the stability of the cation increases. The reason for this is the release of electrons by the alkyl groups through C-C σ bonds towards the positively charged carbon atom to which they are attached. This results in spreading the positive charge thereby stabilizing the ion.

• In the electrophilic addition reactions of hydrogen halides to asymmetric alkenes, two different carbocations can be formed after the bonding of the electrophile. Out of these the more stable carbocation forms more easily.

 

• The more stable carbocation is obtained when the electrophile gets attached to the carbon atom to which the highest number of hydrogen atoms are attached.

• After studying reactions of a large number of alkenes, this observation has been generalized as Markownikoff’s rule.

 

• Hydrogen bromide adds in the opposite way to this rule when there are peroxides in the reaction medium. The reason for this is that in the presence of peroxides the reaction of hydrogen bromide and the alkenes takes place via a free radical mechanism. It is not expected to describe this mechanism.

This change of the direction of addition (Anti-Markownikoff’s addition) is not exhibited by the other hydrogen halides.

 

Addition of bromine to alkenes

 

Mechanism
• When a molecule of Br₂ approaches the π electron cloud of an alkene molecule, it gets polarized . The positive end of the dipole then reacts with the alkene, transferring a Br⁺  ion to it during the reaction (by heterolytic cleavage of the Br-Br bond) forming a cyclic bromonium ion.

• In the second step of the reaction, a Br^– ion acting as a nucleophile, forms a bond to one of the carbon atoms bonded to Br⁺ . The bond formed by that carbon atom to Br⁺  is broken during this step, giving an open chain structure again.

Addition of sulphuric acid to alkenes and the hydrolysis of the product obtained.

• Here, proton acts as an electrophile and HSO4^– ion acts a nucleophile

 

Reaction of alkenes with cold alkaline KMnO₄

Catalytic hydrogenation

eg-Ethene reacts with hydrogen in the presence of a finely divided nickel catalyst at a temperature of about 150°C. Ethane is produced.

 

Reactions of alkynes

• Alkynes have two π bonds which react independantly and undergo electrophilic addition reactions with reagents that add to alkenes.

 

Addition of bromine

 

Addition of hydrogen halides

 

Addition of water

• In the presence of Hg ²⁺ and dilute sulphuric acid, one molecule of water gets added on to alkynes.

 

• The rapid rearrangement of the enol to the aldehyde is due to the high stability of      C = O.

 

Catalytic hydrogenation of alkynes

• In the presence of catalysts such as finely powdered Pt, Pd or Ni alkynes react with hydrogen to give alkanes. The reaction can be stopped at the intermediate alkene stage by using a Pd/BaSO₄ catalyst poisoned by quinoline.

 

• In the alkynes, H-C≡C-Hand R-C≡C-H the H attached to the C that forms the triple bond (terminal hydrogen) shows acidic properties. The acidic H in these alkynes can be displaced by metals.

solubility of the salts of s block salts in water

• Variation of solubility of sulphates and hydroxides down the group can be explained by using lattice enthalpy and hydration enthalpy values.

• Variation of thermal stability of s block carbonates, bicarbonates and nitrates down a group can be explained by using electronegativity values of elements and ionic nature of those compounds.

 

Acidic/basic/amphoteric nature of halides (eg. chlorides) of third period

NaCl – Neutral  

MgCl₂ – very weakly acidic                     

AlCl₃  – acidic                

SiCl₄ – acidic                

PCl₅ – acidic                       

SCl₂ /S₂Cl₂ –  acidic

                                  

 

Acidic/basic/amphoteric nature of oxides of elements of the 3rd period

Na₂O  strongly basic          

MgO  basic          

Al₂O₃  amphoteric                     

SiO₂  very weakly acidic               

P₂O₅  weakly acidic                

SO₂ /SO₃  acidic                    

Cl₂O₇  strongly acidic

 

Nature of oxides down the group 15

Ν₂Ο₃ − acidic

P₂Ο₃ − weakly acidic

As₂Ο₃ − amphoteric

Sb₂Ο₃ − amphoteric

Bi₂Ο₃ − basic

 

Acidic/basic/amphoteric nature of hydroxides of elements of the 3rd period

NaOH  strongly basic          

Mg(OH)₂  basic            

Al(OH)₃  amphoteric           

Si(OH)₄  →^-H₂O  H₂SiO₃   very weakly acidic        

P(OH)₅   →^-H₂O  H₃PO₄  weakly acidic         

S(OH)₆   →^-2H₂O  H₂SO₄  acidic    

Cl(OH)₇  →^-3H₂O   HClO₄  strongly acidic                                                                                  

 

Acidic/basic/amphoteric nature of hydrides of the 3rd period

NaH  strongly basic                

MgH₂  weakly basic                 

AlH₃  amphoteric                 

SiH₄  very weakly acidic                  

PH₃  very weakly basic                    

H₂S  weakly acidic                 

HCl  very strongly acidic

Aluminium

Aluminum (also called Aluminium) is the third most abundant element in the earth’s crust.  It is commonly used in the household as aluminum foil, in crafts such as dyeing and pottery, and also in construction to make alloys. In its purest form the metal is bluish-white and very ductile. It is an excellent conductor of heat and electricity and finds use in some wiring. When pure it is too soft for construction purposes but addition of small amounts of silicon and iron hardens it significantly.

The electron configuration for Aluminum is 1s²2s²2p⁶3s²3p¹.

Aluminum has three oxidation states.  The most common one is +3.  The other two are +1 and +2.

Reaction of aluminium with air

Aluminium is a silvery white metal. The surface of aluminium metal is covered with a thin layer of oxide that helps protect the metal from attack by air. So, normally, aulumium metal does not react with air. If the oxide layer is damaged, the aluminium metal is exposed to attack. Aluminium will burn in oxygen with a brilliant white flame to form the trioxide alumnium(III) oxide, Al₂O₃.

4Al(s) + 3O₂(l) → 2Al₂O₃(s)

Reaction of aluminium with water

Aluminium is a silvery white metal. The surface of aluminium metal is covered with a thin layer of oxide that helps protect the metal from attack by air. So, normally, aulumium metal does not react with air. If the oxide layer is damaged, the aluminium metal is exposed to attack, even by water.

Reaction of aluminium with the halogens

Aluminium metal reacts vigorously with all the halogens to form aluminium halides. So, it reacts with chlorine, Cl₂, bromine, I₂, and iodine, I₂, to form respectively aluminium(III) chloride, AlCl₃, aluminium(III) bromide, AlBr₃, and aluminium(III) iodide, AlI₃.

2Al(s) + 3Cl₂(l) → 2AlCl₃(s) → Al₂Cl₆(s)

AlCl3 is an electron defficient compound and also highly covalent when anhydrous. AlCl3 exists as a dimer thus attaining an octet of electrons.

2Al(s) + 3Br₂(l) → Al₂Br₆(s)

2Al(s) + 3I₂(l) → Al₂I₆(s)

Reaction of aluminium with acids

Aluminium metal dissolves readily in dilute sulphuric acid to form solutions containing the aquated Al(III) ion together with hydrogen gas, H₂. The corresponding reactions with dilute hydrochloric acid also give the aquated Al(III) ion. Concentrated nitric acid passivates aluminium metal.

2Al(s) + 3H₂SO₄(aq) → 2Al³⁺(aq) + 2SO₄^2-(aq) + 3H₂(g)

2Al(s) + 6HCl(aq) → 2Al³⁺(aq) + 6Cl^–(aq) + 3H₂(g)

Reaction of aluminium with bases

Aluminium dissolves in sodium hydroxide with the evolution of hydrogen gas, H₂, and the formation of aluminates of the type [Al(OH)₄]^–.

2Al(s) + 2NaOH(aq) + 6H₂O → 2Na⁺(aq) + 2[Al(OH)₄]^– + 3H₂(g)

 Carbon

Carbon is a Group 14 element and is distributed very widely in nature. It is found in abundance in the sun, stars, comets, and atmospheres of most planets. Carbon is present as carbon dioxide in the atmosphere and dissolved in all natural waters. It is a component of rocks as carbonates of calcium (limestone), magnesium, and iron.

Coal, petroleum, and natural gas are chiefly hydrocarbons. Carbon is unique among the elements in the vast number of variety of compounds it can form.

Carbon is found free in nature in three allotropic forms: amorphous, graphite, and diamond. Graphite is one of the softest known materials while diamond is one of the hardest

More recently, another form of carbon, buckminsterfullerene, C₆₀, was discovered. This form of carbon is the subject of great interest in research laboratories today.

Pure carbon is available in a number of different forms (allotropes). The most common form of pure carbon is α-graphite. This is also the thermodynamically most stable form. Diamond is a second form of carbon but is much less common. Other forms of carbon include the fullerenes. Whereas diamond and graphite are infinite lattices, fullerenes such as buckminsterfullerene, C₆₀, is a discrete molecular species. Amorphous forms of carbon such as soot and lampblack are materials consisting of very small particles of graphite.

Atom arrangements in the most common allotrops of carbon: α-graphite.

As diamond has a slightly more compact structure its density is greater than that of graphite. The appearance of diamond is well known and it is also one of the hardest materials known. Like graphite, it is relatively unreactive but does burn in air at 600-800°C. Each carbon atom is bound to four neighbours at a distance of 154.45 pm in a tetrahedral fashion and so each diamond crystal is a single giant lattice structure. In principle (and in practice!) graphite may be converted into diamond by the application of heat and pressure.

Crystal strucutres of diamond.

Recently another allotrope of carbon was characterized. Whereas diamond and graphite are infinite lattices, buckminsterfullerene, C₆₀, is a discrete molecular species. The buckminsterfullerene molecule is a net of 12 pentagons and 20 hexagons folded into a sphere. The effect is very similar to the patchwork of 12 pentagonal and 20 hexagonal pieces of leather that sewn together make up an association football (soccer ball). The name buckminsterfullerene (or buckyball was coined because of the relationship between the structure of C₆₀ and R. Buckminster Fuller’s geodesic dome designs. Buckminsterfullerene is now commercially available and has also been identified in interstellar space and soot.

C₆₀, Buckminsterfullerene.

Nanotubes are related to fullerenes. They are tubes giving the appearance of rolled graphite, although they are made from graphite. They are open ended while fullerenes are closed structures.

Oxides of carbon

CO is a colourless, neutral, and poisonous gas. CO is used as an industrial fuel. CO is a Lewis acid.

CO2 is a colourless acidic gas and a non polar molecule. Solid CO2 (dry ice) contains dispersion interactions. Dry ice is used as a freezing agent in food industry, in causing artificial rain, etc.

Carbonic acid

Carbonic acid is weak diprotic acid. There are two salts derived from H2CO3 .

CO₂ + H₂0 → H₂CO₃

H₂CO₃ → H⁺ + HCO₃^–

HCO₃– → H⁺ + CO₃^2-

Reaction of carbon with air

Carbon, as graphite, burns to form gaseous carbon (IV) oxide (carbon dioxide), CO₂. Diamond is a form of carbon and also burns in air when heated to 600-800°C – an expensive way to make carbon dioxide!

C(s) + O₂(g) → CO₂(g)

When the air or oxygen supply is restricted, incomplete combustion to carbon monoxide, CO, occurs.

2C(s) + O₂(g) → 2CO(g)

This reaction is important. In industry, air is blown through hot coke. The resulting gas is called producer gas and is a mixture of carbon monoxide (25%), carbon dioxide (4%), nitrogen (70%), and traces of hydrogen (H₂), methane (CH₄), and oxygen (O₂).

Reaction of carbon with water

Carbon, either as graphite or diamond does not react with water under normal conditions. Under more forsing conditions, the reaction becomes important. In industry, water is blown through hot coke. The resulting gas is called water gas and is a mixture of hydrogen (H₂, 50%), carbon monoxide (CO, 40%), carbon dioxide (CO₂, 5%), nitrogen and methane (N₂ + CH₄, 5%). It is an important feedstock gas for the chemical industry.

C + H₂O → CO + H₂

This reaction is endothermic (ΔH° = +131.3 kJ mol^-1; ΔS° = +133.7 J K^-1 mol^-1) which means that the coke cools down during the reaction. To counteract this, the steam flow is replaced by air to reheat the coke allowing further reaction.

Reaction of carbon with the halogens

Graphite reacts with fluorine, F₂, at high temperatures to make a mixture of carbon tetrafluoride, CF₄, together with some C₂F₆ and C₅F₁₂.

C(s) + excess F₂(g) → CF₄(g) + C₂F₆ + C₅F₁₂

At room temperature, the reaction with fluorine is complex.

The other halogens appear to not react with graphite.

Applications

Carbon has a very high melting and boiling point and rapidly combines with oxygen at elevated temperatures. In small amounts it is an excellent hardener for iron, yielding the various steel alloys upon which so much of modern construction depends. An important (but rare) radioactive isotope of carbon, C-14, is used to date ancient objects of organic origin.

Nitrogen

N₂ is a colourless gas. It contains a triple bond, , with a short bond length of 1.09 °A. It has high dissociation bond energy 946 kJ mol-1 and it is an inert gas.

• Liquid  N₂ (boiling point -196 °C) is used as a coolant.  N₂ gas is used in the manufacture of ammonia.

• Nitrogen shows variable oxidation states.

-3              -2             -1        0               +1            + 2            +3                +4              +5

NH₃          N₂H₄     NH₂OH    N₂           N₂O          NO            N₂O₃           N₂O₄         N₂O₅

Oxo acids of nitrogen

Nitrous acid (HNO₂ ) is unstable except in dilute solutions. HNO₂ is formed when acid reacts with metal nitrites.

NaNO₂(s) + HCl(aq) → NaCl(aq) + HNO₂(aq)

Nitric acid (ΗΝΟ₃ ) is a colourless liquid which boils at 86 °C. ΗΝΟ₃ a stable, strong acid as well as a strong oxidizing agent. Concentrated ΗΝO₃ acid is usually yellow. In the light, it decomposes to form nitrogen dioxide and oxygen.

4HNO₃(l) → 4NO₂(g) + O₂(g) + 2H₂O(l)

Compounds

The two most common compounds of nitrogen are Potassium Nitrate (KNO₃) and Sodium Nitrate (NaNO₃). These two compounds are formed by decomposing organic matter that has potassium or sodium present and are often found in fertilizers and byproducts of industrial waste.

Nitrides

Nitrides are compounds of nitrogen with a less electronegative atom; in other words it’s a compound with atoms that have a less full valence shell. These compounds form with lithium and Group 2 metals. Nitrides usually has an oxidation state of -3

3Mg+ N₂→Mg₃N₂

Ammonium Ions

Nitrogen goes through fixation by reaction with hydrogen gas over a catalyst. This process is used to produce ammonia. As mentioned earlier, this process allows us to use nitrogen as a fertilizer because it breaks down the strong triple bond held by N_2. The famous Haber-Bosch process for synthesis of ammonia looks like this:

 N₂+3H₂→2NH₃

Ammonia is a base and is also used in typical acid-base reactions.

ΝH₃ acts as an oxidizing agent and as well as an acid

2Na(s) + 2NH₃(l) → 2NaNH2(s) + H2(g)

ΝH₃ acts as a weak reducing agent.

3NH₃(l) + 2Cl₂(g) → 2 N₂(g) + 6HCl(g)

3CuO(s) + 2NH₃(l) →  N₂(g) + 3Cu(s) + 3H₂O(l)

ΝH₃ acts as a weak reducing agent.

2NH₃(aq)+H₂SO₄→(NH₄)₂SO₄(aq)

Ammonium salts

Ammonium salts decompose quite readily on heating.

(NH₄)₂CO₃(s) →Δ  2NH₃(g) + CO₂(g) + H₂O(l)

NH₄Cl(s) →Δ  NH₃(g) + HCl(g)

NH₄NO₂(s) →Δ  N₂(g) + 2H₂O(l)

NH₄NO₃(s) →Δ  N₂O(g) + 2H₂O(l)

(NH₄)₂Cr₂O₇(s) →Δ  N₂(g) + Cr₂O₃(s) + 4H₂O(s)

Oxides of Nitrogen

Nitrides use a variety of different oxidation numbers from +1 to +5 to for oxide compounds.  Almost all the oxides that form are gasses, and exist at 25 degrees Celsius. Oxides of nitrogen are acidic and easily attach protons.

N₂O₅+H₂O→2HNO₃(aq)

The oxides play a large role in living organisms. They can be useful, yet dangerous.

• Dinitrogen monoxide (N₂O) is a anesthetic used at the dentist as a laughing gas.
• Nitrogen dioxide (NO₂) is harmful. It  binds to hemoglobin molecules not allowing the molecule to release oxygen throughout the body. It is released from cars and is very harmful.
• Nitrate (NO₃^–) is a polyatomic ion.
• The more unstable nitrogen oxides allow for space travel.

 

Sulphur

Sulphur dioxide

Sulphuric acid

 

Halogens

• Fluorine is such a powerful oxidising agent that you can’t reasonably do solution reactions with it.
• Chlorine has the ability to take electrons from both bromide ions and iodide ions. Bromine and iodine can’t get those electrons back from the chloride ions formed.That means that chlorine is a more powerful oxidising agent than either bromine or iodine.
• Similarly bromine is a more powerful oxidising agent than iodine. Bromine can remove electrons from iodide ions to give iodine – and the iodine can’t get them back from the bromide ions formed.

This all means that oxidising ability falls as you go down the Group.

 

• Fluoride and chloride ions won’t reduce concentrated sulphuric acid.
• Bromide ions reduce the sulphuric acid to sulphur dioxide. In the process, the bromide ions are oxidised to bromine.
• Iodide ions reduce the sulphuric acid to a mixture of products including hydrogen sulphide. The iodide ions are oxidised to iodine.
• Reducing ability of the halide ions increases as you go down the Group.

 

 

Noble gases

Position of noble gas in periodic table

Noble gases are also known as inert gases and do not take part in chemical reactions. They have their outermost shell complete and thus remain stable. They do not generally combine with other substances, nor are they affected by oxidising agents or by reducing agents. They are placed in the 18 or VIIIA group. Since, the outermost shell is complete, the valency is zero, hence VIIIA group is also referred to as zero group.

occurrence of noble gas

Noble gases always occur in free state because of their inert nature. All the noble gases, except radon are present in air in small amounts. The relative abundance of the noble gases in air is = 1%. Helium is present in natural gas up to the extent of 10 per cent. It is also present in small quantities in the minerals of radioactive elements. Water from certain springs also gives off gases, which are rich in helium and argon.

Uses of noble gas

Argon is used mainly to provide an inert atmosphere in high temperature metallurgical processes (arc welding of metals or alloys) and for filling electric bulbs. It is also used in the laboratory for handling substances that are air-sensitive.

Helium is a non-flammable and light gas. Hence, it is used in filling balloons for meteorological observations. It is also used in gas-cooled nuclear reactors. Liquid helium (boiling point 4.2K) finds use as cryogenic agent for carrying out various experiments at low temperatures. It is used to produce and sustain powerful super conducting magnets, which form essential part of modern NMR spectrometers and Magnetic Resonance Imaging (MRI) systems for clinical diagnosis.

Neon is used in discharge tubes and fluorescent bulbs for advertisement display purposes. There are no significant uses of xenon and krypton. They are used in light bulbs designed for special purposes.

First ionization energy graph of elements from H to Cs

 

Variation of metalic radius, electronnegativity and ionisation energy from Sc to Zn

 

Physical properties of elements

Z and symbol	21 Sc	22 Ti	23 V	24 Cr	25 Mn	26 Fe	27 Co	28 Ni	29 Cu	30 Zn
property\name	scandium	titanium	vanadium	chromium	manganese	iron	cobalt	nickel	copper	zinc
melting point/oC	1541	1668	1910	1857	1246	1538	1495	1455	1083	420
density/gcm–3	2.99	4.54	6.11	7.19	7.33	7.87	8.90	8.90	8.92	7.13
atomic radius/pm	161	145	132	125	124	124	125	125	128	133
M2+ ionic radius/pm	na	90	88	84	80	76	74	72	69	74
M3+ ionic radius/pm	81	76	74	69	66	64	63	62	na	na
common oxidation states	+3 only	+2,3,4	+2,3,4,5	+2,3,6	+2,3,4,6,7	+2,3,6	+2,3	+2,+3	+1,2	+2 only
outer electron configuration [Ar]…	3d14s2	3d24s2	3d34s2	3d54s1	3d54s2	3d64s2	3d74s2	3d84s2	3d104s1	3d104s2

 

 

Oxidation numbers of elements from Sc to Zn (Common oxidation numbers are shown in bold letters)

 

 

Instances where d block elements and their compounds are used in industry as catalysts.

 

Formation of coloured compounds

Transition ions formed by d block metals have partially filled d orbitals. These ions absorb selected range of wave length in white light and get excited and shows the complementary colours. However, d⁰ and d¹⁰ confingurations are colourless.

Sc³⁺ – Colourless

Co²⁺ – Pink

Ti⁴⁺ – Colourless

Ni²⁺ – Green

Ti³⁺ – Purple

Cu²⁺ – Blue

V³⁺ – Green

Cu⁺ – Colourless

V²⁺ – Violet

Zn²⁺ – Colourless

Cr³⁺ – Purple

Mn³⁺ – Violet

Mn²⁺ – Pale pink

Fe³⁺ – Brown yellow

Fe²⁺ – Pale green

Colours of some oxoanions: Usually, d element oxoanions are coloured.
MnO₄^– – Purple/Violet
MnO₄^2- – Green
CrO₄^2- – Yellow
Cr₂O₇^2- – Orange

 

Complex ions formed by the elements Cr, Mn, Fe, Co, Ni and Cu with the ligands H₂O, NH₃ and Cl^–

 

• OH-acts as a ligand in few d block cations.
eg. [Fe(H₂O)₅OH]²⁺ formed by the hydrolysis of hydrated Fe³⁺ ions. However, most
metal ions with NaOH or NH₃ give insoluble hydroxides and not hydroxo complexes.

eg. Cr(OH)₃ – Green, Fe(OH)₃ – Reddish brown, Fe(OH)₂ – Dirty green, Cu(OH)₂ – Blue, Mn(OH)₂ – White

• The colour of the complex varies depending on the central metal atom.
eg. [Cr(H₂O)₆]³⁺ – Blue – violet

[Fe(H₂O)₆]³⁺ – Yellow

• The colour of the complex varies depending on the oxidation state of the central metal atom.
eg. [Fe(H₂O)₆]²⁺ , Pale green

[Fe(H₂O)₆]³⁺ – Yellow

[Mn(H₂O)₆]²⁺ – Pale pink

[Mn(H₂O)₆]³⁺ – Violet

• The colour changes also when the ligands change.
eg. [Co(H₂O)₆]²⁺ – Pink

[Co(NH₃)₆]²⁺ – Yellow brown

IUPAC nomenclature of complexes

• Here, the information required to develop the IUPAC name of a complex and to write the structural formula when the IUPAC name is given will be discussed. Only complexes formed by elements of the d block are considered.
The complexes are considered simply under two categories.
(i) Cations are simple while the anions are complex.
(ii) Cations are complex while the anions are simple.

• Whatever the complex considered, a common set of rules has to be followed stepwise in their nomenclature.

Writing the name of a complex

(1) As in the case of a simple inorganic compound, first the cation is named and then the anion. A space is left between the name of the cation and the name of the anion.

(2) The complex ion in the compound can be either positively charged or negatively charged. First, identify the metal ion and the ligand/s in the complex ion.

(3) The ligands could be negatively charged, neutral or positively charged (rarely). When naming the ligands, the charge of the ligand is considered.
(i) Neutral ligands have no special ending.
(ii) Some ligands have special names.

Examples:

H₂O aqua
NH₃ ammine
CO carbonyl
NO nitrosyl

 

(iii) For negatively charged ligands an “o” is added to their English name.

Examples:

Cl^– chlorido
CN^– cyanido
NO₂^– nitrito
OH^– hydroxido
SCN^– thiocyanato
H^– hydrido
O^2- oxido

 

(iv) For positively charged ligands the suffix” ium” is added to their English name.

Example: +NH₃-NH₂ hydrazinium

 

(4) When there is more than one ligand of the same type, in order to indicate the number of such ligands, the name of the relevant number is used as a prefix to the name of the ligand. When there are 2, 3, 4, 5 and 6 ligands of the same type, the prefixes di-, tri,- tetrapenta- and hexa- are used respectively.

(5) When several ligand types are present in a complex ion, in naming the ligands they are listed in the alphabetical order (English) of the first letter of the ligand.

Note: The first letter of the prefix used to denote the number of ligands is NOT considered when deciding the alphabetical order.

No space is left between the names of the ligands.
Example: [Fe(CN)₂(NH₃)₄]⁺
Ligands are named as; tetraamminedicyanido

 

(6) When writing the name of the complex ion, the ligands are named first and then the metal ion. The oxidation number of the metal ion is given in capital Roman numerals, within parentheses, after the name.
No space is left between the words when writing the name.
Examples:
[Co(NH₃)₆]³⁺ hexaamminecobalt(III) ion
[Fe(H₂O)₆]²⁺ hexaaquairon(II) ion
[Cu(NH₃)₄]²⁺ tetraamminecopper(II) ion

(7) The complex may be positively charged, negatively charged or neutral. Depending on this, the name also changes.
(i) When the complex is positively charged or neutral, the name of thecomplex ends with the name of the metal.

Reminder
The name of the metal is followed by the oxidation number of the metal ion, in capital Roman numerals, within parentheses. No space is left between the name of the metal ion and the oxidation number given in parantheses.
Example:
[Fe(CN)₃(NH₃)₃]
The complex is neutral. Hence, its name is triamminetricyanidoiron(III).

Example:
[Cu(H₂O)₆ ]²⁺
The complex is positively charged. Its name is hexaaquacopper(II) ion.

(ii) When the complex ion is negatively charged, the suffix ‘ate’ is added to the end of the name of the metal. Here also the oxidation number of the metal ion should be indicated in capital Roman numerals within parentheses. No space is left between the name of the metal ion and the oxidation number given in parentheses.
Examples:
[CoCl₄]^2- tetrachloridocobaltate(II) ion
[Co(CN)₆]^3- hexacyanidocobaltate(III) ion
[CuCl₄]^2- tetrachloridocuprate(II) ion
[Fe(CN)₆]^4- hexacyanidoferrate(II) ion
[Fe(CN)₆]^3- hexacyanidoferrate(III) ion
[Ag(CN)₂]^– dicyanidoargentate(I) ion
[Cr(Br)₆]^3- hexabromidochromate(III) ion
The IUPAC name of any compound can be developed by systematically following the rules studied so far.

 

(8) When writing the name of a coordination compound, a space should be left between the name of the positively charged species and the negatively charged species.

Examples :
• Simple cation and complex anion
K₃[Fe(CN)₅NO] potassium pentacyanidonitrosylferrate(II)
Na₂[ZnCl₄] sodium tetrachloridozincate(II)

• Complex cation and simple anion
[Ag(NH₃)₂]Cl diamminesilver(I) chloride
[Fe(OH)₂(H₂O)₄]Br tetraaquadihydroxidoiron(III) bromide
[CoCl(NH₃)₅](NO₂)₂ pentaamminechloridocobalt(III) nitrate
[CoCl(NH₃)₅](NO₂)₂ pentaamminechloridocobalt(III) nitrite

 

Writing formula of a complex when its name is given

1. The positively charged species is written first, followed by the negatively charged species. No space is left between them.

2. The complex part of the compound is always written within square brackets.

3. When the formula of the complex ion is written, the metal should be indicated first and then the ligands. In writing the ligands, the charge on the ligand is NOT considered. The ligands are written in the alphabetical order of the ligating atoms. (i.e. Atom through which the ligand coordinates to the metal ion).

Note: In multiatomic ligands, where possible, it is recommended that the ligating atom is placed first followed by the other atoms in the ligand.
Example : (:OH₂) rather than H₂O

4. Multi atomic ligands are given in parentheses. The number of each type of ligand is given in Arabic numerals as a subscript on the right hand side immediately after the symbol of the ligand. If parentheses are present, this number is written as a subscript on the right hand side without leaving a space, just outside the parentheses.

5. The formula of a complex ion should be written within square brackets. If the complex ion has a charge, it should be indicated outside the square bracket as a superscript on the right side. The numerical value should be given first followed by the sign of the charge.

Note: No space is left between the formulae of the ligands or between the formulae of the ligands and symbol of the metal ion.

 

Example 1
Write the chemical formula of pentacyanidonitrosylferrate(II) ion.
Step 1: Write the symbol of the metal.
Fe

Step 2: Decide on the order of the ligands. Include the number of ligands of each type when writing their symbols/formulae. Then write the symbols / formulae of the ligands after the symbol of the metal. The ligands are CN^– (ligating atom is C) and NO (ligating atom is nitrogen) hence, the order is CN^– followed by NO. Both ligands are multi-atomic. Hence, their formulae are placed within parentheses. The presence of five CN^– ligands are indicated.
Fe(CN)₅(NO)

Step 3: Place the symbol of the metal ion and the formulae of the ligands within square brackets and show charge of complex.

Oxidation number of metal is +II. Overall charge of the complex ion is + 2 + (-5) = -3
Chemical formula is [Fe(CN)₅NO]^3-

 

Example 2
Write the chemical formula of pentaamminechloridocobalt(III) ion.

Step 1: Write the symbol of the metal.
Co.

Step 2: Decide on the order of the ligands. Include the number of ligands of each type when writing their symbol/formula. Then write the symbols/ formulae of the ligands after the symbol of the metal.

Ligands are Cl^– (ligating atom is Cl) and NH₃ (ligating atom is N). Considering the alphabetical order of the ligating atoms, Cl^– is written first followed by NH₃. The NH₃ ligand is multi-atomic and hence the formula is placed within parentheses. The presence of five NH₃ ligands is indicated.
CoCl(NH₃)₅

Step 3: Place the symbol of the metal and the formulae of the ligands within square brackets and show charge of complex.Oxidation state of the metal ion is +III. Therefore, overall charge of the complex is +3 + (-1) = +2
Chemical formula is; [CoCl(NH₃)₅]²⁺

 

Example 3
Write the chemical formula of pentaamminechloridocobalt(III) bromide.
As shown in Example 2, the chemical formula of the complex ion is;
[CoCl(NH₃)₅]²⁺.

Two Br^– ions are required to neutralize the charge on this complex ion. Therefore, the formula is; [CoCl(NH₃)₅]Br₂