30 most important name reactions which are frequently asked in class 12 chemistry exams are listed below:

1. Dehydrohalogenation reaction (Elimination reaction)

When an alkyl halide is heated with alcoholic solution of KOH, then a molecule of hydrogen halide is eliminated from the haloalkane and alkene is formed. Therefore this reaction is also called dehydrohalogenation reaction. Eg.

Elimination reaction involves the removal of halogen atom of haloalkane and a hydrogen atom from the β- carbon (i.e. adjacent carbon). Therefore, this reaction is also known as β- elimination reaction.

2. Wurtz reaction

When an alkyl halide( haloalkane) is heated with sodium metal in presence of dry ether, a symmetrical alkane containing double number of carbon atoms than in haloalkane is formed. This reaction is called Wurtz reaction.

3. Fittig reaction

Haloarenes when treated with sodium in the presence of dry ether, diaryls are produced. This reaction is known as Fittig reaction.

4. Wurtz-Fittig reaction

Haloarenes when treated with sodium and alkyl halide in presence of dry ether gives toluene. This reaction is known as Wurtz-Fittig reaction.

5. Carbylamine reaction: Test reaction of primary amines

When chloroform is warmed with a primary amine in the presence of alcoholic KOH, an offensive(unpleasant) smell of carbylamines ( i.e. isocyanide) is obtained. This reaction is known as carbylamine reaction.

6. Sandmeyer reaction

Chlorobenzene can be prepared by treating benzene diazonium chloride with cuprous chloride dissolved in HCl. This reaction is called Sandmeyer reaction.

7. Gattermann reaction

Chlorobenzene is prepared by treating benzene diazonium chloride with copper powder dissolved in HCl. This reaction is the modification of Sandmeyer’s reaction and called Gattermann reaction.

8. Diazotization reaction

When aniline is treated with sodium nitrite (NaNO₂) and dil.HCl at temperature below 5⁰C, benzene diazonium chloride is obtained. This reaction is known as diazotization reaction.

9. Dows Process

When chlorobenzene is heated with an aqueous solution of NaOH at 300⁰C and 200 atm gives sodium phenoxide which on acidification gives phenol.

10. Oxo-process (Carbonylation reaction)

Alkenes react with carbon monoxide and hydrogen in the presence of cobalt carbonyl catalyst [Co(CO)₄]₂ at high pressure and temperature to give aldehyde, which on catalytic hydrogenation gives primary alcohol. Eg.

11. Fermentation

Fermentation is a biochemical process of degradation (slow decomposition/ breaking down) of large organic molecules like sugars and starches into simpler compounds by the catalytic action of enzymes.

Eg. Ethanol from sugar: Enzymes invertase and zymase are obtained from yeast. The reactions occurring during the fermentation of sugar are:

12. Esterification reaction

Alcohols react with carboxylic acids in the presence of few drops of conc. H₂SO₄ to form esters. This reaction is known as esterification reaction.

13. Kolbe’s reaction (Carboxylation reaction)

Sodium phenoxide when heated with CO₂ at135⁰C under a pressure of 4-7 atm, sodium salicylate is obtained which when acidified gives salicylic acid.

14. Haloform reaction

Aldehydes and ketones containing CH₃CO- group on reaction with excess halogen in presence of NaOH gives haloform (chloroform’ bromoform, iodoform). Eg.

15. Coupling reaction

Phenol or aniline reacts with benzene diazonium chloride in slightly alkaline medium and at low temperature to form coloured compounds called azo dyes. This reaction is called coupling reaction.

16. Reimer-Tiemann reaction

When phenol is refluxed with chloroform and aq. NaOH at 60⁰C, a mixture of o-hydroxybenzaldehyde and p-hydroxybenzaldehyde is obtained. This reaction is called Reimer-Tiemann reaction.

17. Williamson’s ether synthesis

The reaction in which alkyl halide and sodium or potassium alkoxide are reacted to form ether is known as Williamson’s etherification reaction. Eg.

Both symmetrical and unsymmetrical ether can be prepared from this reaction. Eg.

18. Ozonolysis reaction

Alkene reacts with ozone to give ozonide. On warming ozonide with Zn in water, it breaks down to give two molecules of carbonyl compounds (aldehyde or ketone). This process of formation of ozonide and it’s decomposition to give carbonyl compounds is called ozonolysis.

19. Rosenmund reduction

Aldehydes can be prepared by reducing acid chloride solution with hydrogen in the presence of Palladium(Pd) catalyst deposited on barium sulphate and partially poisoned with sulphur or quinoline. This reaction is called Rosenmund reduction.

20. Clemmensen’s reduction

The reduction of aldehydes and ketones to alkane using zinc amalgam and conc. HCl is Clemmensen’s reduction. In this reaction, carbonyl group (-CO-) is reduced to -CH₂– group. Eg.

21 . Wolff-Kishner reduction

In this method aldehyde and ketone is treated with hydrazine to form hydrazone which is then heated with KOH in presence of glycol to give alkane. Eg.

22. Aldol condensation reaction

Condensation between two molecules of aldehydes or ketones having at least one α – hydrogen atom in presence of dilute alkali to form β-hydroxy aldehyde or β-hydroxy ketone is known as aldol condensation reaction. Examples:

Aldehydes and ketones which do not contain any α – hydrogen atom such as HCHO, (CH₃)₃CCHO, C₆H₅CHO, etc. do not undergo aldol condensation reaction.

23. Cannizzaro’s reaction

Aldehydes which do not contain α-hydrogen like HCHO, C₆H₅CHO,etc. undergo self oxidation and reduction on treatment with conc. alkali. In this reaction one molecule is oxidized to carboxylic acid and other molecule is reduced to alcohol. Thus, a mixture of an alcohol and a salt of carboxylic acid is formed by Cannizzaro’s reaction

24. Perkin’s (condensation) reaction

The condensation of an aromatic aldehyde with an acid anhydride in the presence of sodium or potassium salt of the same acid to produce α,β-unsaturated acid is known as the Perkin’s condensation.

25. Benzoin condensation reaction

Benzaldehyde when heated with alcoholic solution of potassium cyanide, undergoes self condensation between two molecules to form an α-hydroxy ketone known as benzoin. This reaction is called benzoin condensation reaction. Eg.

26. Hell-Volhard Zelinsky [HVZ] reaction

Carboxylic acids (except formic acid) reacts with chlorine or bromine in presence of red phosphorous to give α-chloro or α-bromo acids. The reaction does not stop at monosubstituted product but continues till all α-hydrogen atoms are replaced.

27. Claisen condensation reaction

Condensation between two molecules of esters having α-hydrogen atom in the presence of strong base like sodium ethoxide and an acid to form β-keto ester is known as Claisen condensation reaction. Eg.

28. Hoffmann’s Bromamide reaction: Decarbonylation reaction

When an amide is treated with bromine in NaOH or KOH, carbonyl group is removed from made to form primary amine containing one carbon less than that of amide. This reaction is known as Hoffmann bromamide reaction or decarbonylation reaction.

29. Friedel Craft’s reaction

a. Friedel Craft’s alkylation reaction: Introduction of an alkyl group ( – R ) in the benzene ring by treating benzene with an alkyl halide in the presence of anhydrous AlCl₃ is known as Friedel- Craft’s alkylation reaction.

b. Friedel craft’s acylation reaction : Introduction of an acyl group (i.e.keto group) in the benzene ring by treating benzene with acid chloride or acid anhydride in the presence of anhydrous AlCl₃ is known as Friedel- Craft’s acylation reaction.

30. Hydroboration-Oxidation of alkene

In this method, alkene is treated with diborane, (BH₃)₂ or B₂H₆ followed by alkaline oxidation with H₂O₂ to get primary alcohol.

Hydroboration is a reduction process, which is carried out by treating with diborane, (BH₃)₂ to give trialkyl borane, which upon oxidation with alkaline solution of H₂O₂ gives primary alcohol. For example:

Links to get mechanism of some name reactions

• https://chemicalnote.com/aldol-condensation-and-crossed-aldol-condensation-reaction-definition-examples-mechanism-and-uses/
• https://chemicalnote.com/cannizzaros-reaction-and-crossed-cannizzaros-reaction-definition-examples-and-mechanism/
• https://chemicalnote.com/claisen-condensation-reaction-examples-and-mechanism/
• https://chemicalnote.com/diazotization-reaction-mechanism-and-uses/
• https://chemicalnote.com/witting-reaction-examples-and-mechanism/
• https://chemicalnote.com/acetoacetic-ester-synthesis-of-ketones/
• https://chemicalnote.com/malonic-ester-synthesis-of-carboxylic-acids/

The post Name reactions: 30 most important organic named reactions for class 12 appeared first on Online Chemistry notes.

• Father of the periodic table = Dmitri Mendeleev (a Russian chemist)

Mendeleev’s Periodic table

Dmitri Mendeleev arranged the known elements in the increasing order of their atomic weight in the form of the table called Mendeleev’s periodic table.

Mendeleev’s periodic law : Mendeleev’s periodic law states that “the physical and chemical properties of elements are periodic functions of their atomic weights.”

It means that if the elements are arranged in order of increasing atomic weight, the elements with similar properties are repeated after a regular interval.

Features of Mendeleev’s periodic table:

The modified form of Mendeleev’s periodic table consists of:

1. Nine vertical columns- Groups:

There are nine vertical columns numbered as I, II, III, IV, V, VI, VII, VIII and zero (noble gases). Except for groups VIII and zero, each group from I and VII is subdivided into two subgroups A and B.

2. Seven horizontal rows- Periods:

There are seven horizontal rows numbered as 1 to 7. The periods are divided into short periods and long periods. The first three periods are called short periods as they contain fewer elements. Fourth, fifth and sixth periods are long periods.

3. Position of lanthanides and actinides:

The lanthanides (14 elements after Lanthanum) and actinides (14 elements after Actinium) are kept separate in two horizontal rows at the bottom of the periodic table.

Note : Only 63 elements were known when Mendeleev’s periodic table was created. 

Limitations of Mendeleev’s periodic table:

Some main limitations in Mendeleev’s periodic table are as follows:

1. Position of hydrogen: Hydrogen forms both the positive ion like alkali metal (group-IA) and negative ion like halogen (group-VIIA), hence it resembles the elements of group IA and group VIIA. Therefore, the position of hydrogen in the periodic table is controversial.
2. Position of isotopes: Isotopes of the same element have different atomic weight. According to Mendeleev’s periodic law, isotopes of an element should be placed at different places in the periodic table. For example, hydrogen needs three separate positions for protium, deuterium and tritium with atomic weight 1, 2 and 3 respectively. But isotopes were not given a separate place in the periodic table.
3. Position of lanthanides and actinides: Lanthanides and actinides group consists of 14 elements each. All these elements have been grouped in a single place in the III^rd group respectively in the periodic table despite their different atomic masses. They have not been given proper position within the main frame of the table but are placed outside in two separate rows at the bottom of the periodic table.
4. Position of anomalous pairs of elements: Certain elements with higher atomic weight were placed before the elements having lower atomic weight. For example: Ar (at.wt.=39.9) was placed before K (at.wt.=39.1) ; Co (at.wt.=59.9) was placed before Ni (at.wt.=58.6), etc. These pairs of elements do not obey periodic law.
5. Separation of chemically similar elements and grouping of dissimilar elements: In Mendeleev’s periodic table, chemically similar elements like Cu and Hg, Au and Pt, Ag and Tl, Ba and Pb, etc have been placed in separate groups while other dissimilar elements have been placed in same group. For example, coinage metals like Cu, Ag, and Au are lesser reactive metals are placed together with higher reactive metals like Li, Na, K.
6. Cause of periodicity: It does not explain the cause of periodicity. Because of this, atomic weight is not a good basis for the classification of elements.

Modern Periodic Table

• Modern periodic table is based on atomic number.
• Modern periodic law was proposed by Moseley.

Modern periodic law : Modern periodic law states that “the physical and chemical properties of elements are periodic functions of their atomic number.”

It means that if the elements are arranged in order of increasing atomic number, the elements with similar properties are repeated after a regular interval.

Long form/ Extended form/ Present form of periodic table:

It was constructed by Bohr based on atomic number and Bohr’s- Bury electronic configuration concept.

Main features of modern periodic table:

Long form or extended form of periodic table (modern periodic table) was constructed by Bohr and is also called Bohr’s periodic table. This is an improved and extended form of Mendeleev’s periodic table. Main features of the modern periodic table are as follows:

1) There are 18 vertical columns called groups.

• The elements of group 1, 2 and 13 to 17 are called typical or representative elements.
• The elements of group 3 to 12 are known as transition elements.
• Elements of group 18 are called noble gases

2) There are 7 horizontal rows called periods. They are denoted by 1, 2, 3, 4, 5, 6, and 7.

3) 14 elements after Lanthanum(La) i.e. elements with atomic number 58 to 71 (Ce to Lu) are called lanthanides whereas, the 14 elements after Actinium(Ac) i.e. elements with atomic number 90 to 103 (Th to Lr) are called actinides. These elements are collectively called f-block elements or inner transition elements. These elements are given two separate rows below the main periodic table.

4) Elements are classified into s-block, p-block, d-block and f-block.

5) Metals and non-metals are separated from each other.

Merits (Advantages) of modern periodic table:

The modern periodic table has overcome the drawbacks of Mendeleev’s periodic table by choosing atomic number as the basis of classification. The main advantages of modern periodic table are as follows:

1. Atomic number basis: Elements are arranged on the basis of atomic number (number of protons), a fundamental property, rather than atomic mass which can be less consistent for isotopes of elements.
2. Position of isotopes: Atomic number of isotopes is same, so different isotopes can be placed at same place in periodic table. Thus position of isotopes is completely justified.
3. Proper solution of Mendeleev’s misfit points: The position of Ar, Co and Te before K, Ni and I respectively in Mendeleev’s table was not according to his periodic law. Long form of periodic table justified this anomaly by choosing atomic number as the basis of classification.
4. Separate position for subgroups: Separate positions for subgroups A and B is provided in modern periodic table. Separation of subgroups removed the anomalies found in Mendeleev’s periodic table such as grouping of chemically dissimilar elements and separation of chemically similar elements.
5. Separation of metals and non-metals: Metals are placed on the left and non-metals are placed on the right side of the periodic table.
6. Division of elements into four blocks: Division of elements into s, p, d, and f-blocks based on their electronic configuration has made their study easier.
7. Successful to explain periodicity: Modern periodic table has been successful to explain the periodicity in certain atomic properties like atomic radius, ionization potential, etc.

Demerits (Defects) of long form of periodic table:

1. Position of hydrogen is still controversial.
2. Position of helium along with the p-block elements is not completely justified as its electronic configuration is 1s².
3. Lanthanides and actinides are still not placed in main body.
4. Isotopes have different physical properties but have same place in periodic table.

Classification of elements in periodic table

Bohr’s Classification: Depending on the number of incomplete shell in an atom, the elements in the modern periodic table can be classified into four types.

1. Inert gas elements:

• These elements have completely filled ultimate(valence) shell .
• General electronic configuration is ns²np⁶
• Because of most stable configuration, they are very less reactive. Hence, known as noble gas or inert gas elements.
• These elements are present in ‘0’ group or 18^th group.

2. Representative/main group/normal elements:

• These elements have incomplete valence (ultimate) shell.
• These elements lie in group IA to VIIA
• Elements of group IA and IIA are known as alkali metals and alkaline earth metals respectively.
• Elements of group VA, VIA and VIIA are known as pnicogen, chalcogen and halogen family.

3. Transition elements:

• These elements have incompletely filled ultimate (n) and penultimate (n-1) shell.
• These elements lie in group IB to VIIB and group VIII.
• The name transition is due to their properties lies between highly reactive metals on the left side and non-metals on the right side.
• General outer electronic configuration is : (n-1)d^1-10ns^1-2

4. Inner transition elements:

• These elements have incompletely filled ultimate (n), penultimate (n-1) and antepenultimate (n-2) shell.
• These elements lie in group IIB and period 6^th and 7^th.
• These are 28 in number.
• General outer electronic configuration is : (n-2)f^1-14(n-1)d^0-1ns²
• Electronic configuration of Lanthanides is 4f^1-14,5d^0-1,6s² and Actinides is 5f^1-14,6d^0-1,7s²

Classification of elements into blocks:

On the basis of sub-shell (orbital) in which last (differentiating) electron enters, the elements are classified into four blocks: s, p, d and f-blocks.

1. s-block elements:

• Elements in which the last (differentiating) electron enters into the s-orbital of the outermost shell are called s-block elements.
• S-block consists of elements of group IA (alkali metals) and IIA (alkaline earth metals).
• These elements are very reactive metals.
• They are very soft and malleable metals.
• They have low ionization energy and are highly electropositive.
• They have low melting and boiling point.
• General outer electronic configuration is ns^1-2
• Alkali metals have largest atomic size in corresponding periods.
• Their hydroxides are strong bases.
• Most of the s-block elements (except Be and Mg) impart characteristics colour to the flame (i.e. flame test).

2. p-block elements:

• Elements in which the last (differentiating) electron enters into the p-orbital of the outermost shell are called p-block elements.
• General outer electronic configuration is ns²np^1-6
• p-block consists of elements of group IIIA (13), IVA(14), VA(15), VIA(16), VIIA(17) and zero(18).

IIIA (13) = Boron family

IVA (14) = Carbon family

VA (15) = Nitrogen family (pnicogen family)

VIA (16) = Oxygen family ( Chalcogen family >> ore forming family)

VIIA (17) = Halogen family (salt forming family)

Zero (18) = inert gases/ noble gases/ rare gases/ aerogens

3. d-block elements:

• Elements in which the last (differentiating) electron enters into the d-orbital of the penultimate shell are called d-block elements.
• d-block consists of the elements of groups IIIB(3), IVB(4), VB(5), VIB(6), VIIB(7), VIII(8-10), IB(11) and IIB(12).
• These are hard metals with high melting and boiling points.
• They are called transition elements as they exhibit transition behavior intermediate between the properties of s-and p-block elements.
• General outer electronic configuration is : (n-1)d^1-10ns^1-2
• They show variable oxidation state.
• Most of them form coloured salts and their ions or compounds are paramagnetic in nature. These properties are due to presence of unpaired electrons. For example, CuSO₄.5H₂O is blue in colour.
• Most transition elements and their compounds possess catalytic properties. Eg. Fe, Ni, Mo, Pt, etc.

4. f-block elements: Inner transition elements

• Elements in which the last (differentiating) electron enters into the f-orbital of the ante-penultimate shell are called f-block elements.
• These elements lie in group IIB and period 6^th and 7^th.
• These are 28 in number.
• General outer electronic configuration is : (n-2)f^1-14(n-1)d^0-1ns²
• Electronic configuration of Lanthanides is 4f^1-14,5d^0-1,6s² and Actinides is 5f^1-14,6d^0-1,7s²
• They have high melting and boiling point.
• They are heavy metals.
• They show variable oxidation state, commonly +3 state.
• They form coloured salts.
• Lanthanides (4f-series) are called rare earth elements since they occur rarely in the earth crust.
• Actinides (5f-series) are radioactive elements.

Nuclear charge, Shielding effect and Effective nuclear charge

Nuclear charge:

• The total positive charge present in the nucleus is called the nuclear charge.
• Its value is always positive and depends on the number of protons present in the nucleus.
• It is denoted by the letter ‘Z’.
• For example, the value of Z for oxygen is +8.

Shielding effect and Effective nuclear charge:

• In multi-electron atoms, the electrons in the valence shell are pulled by the nucleus and repelled by the electrons of inner shells. Thus, outermost electrons experience less attraction from the nucleus under the combined effect of attractive and repulsive force acting on the valence electrons. This effect is called shielding effect or screening effect.
• Thus larger the number of electrons in inner shell, the larger will be the screening effect.
• The actual charge felt by the valence electron as a result of shielding effect is called an effective nuclear charge (Z_eff).

• Effective nuclear charge (Z_eff) = Total nuclear charge (Z) – Screening constant (s)
• Value of effective nuclear charge is always less than that of the nuclear charge.
• For hydrogen, the effective nuclear charge is equal to the nuclear charge as there is no screening effect.

Periodic trends and periodicity (Atomic properties)

Periodicity of elements:

• When the elements are arranged in the modern periodic table in order of increasing atomic number, the occurrence of similar properties of elements after a definite interval is termed as periodicity of an element.
• These properties include atomic radius, ionization potential, electron affinity, electronegativity, etc.

Causes of periodicity:

• The cause of periodicity in properties is due to the same outermost shell electronic configuration coming at regular intervals.
• In the periodic table, elements with similar properties occur at intervals of 2, 8, 8, 18, 18 and 32. These numbers are called magic numbers.

The variation of some properties along with group and period are described below:

Atomic radii (size)

• Atomic radius is the distance from centre of the nucleus to the outermost shell of the electrons.
• Atomic radius cannot be measured directly because the atom cannot be isolated to determine its radius.

It can be measured indirectly from bond length measurement.

• The atomic radii are expressed in terms of covalent radii, metallic radii, Vander Waal’s radii and ionic radii.
• For most of the elements, atomic radii is measured in terms of covalent radii whereas, Vander Waal’s radii is measured for noble gases.

1. Covalent radii: The half of the distance between two nuclei in a homonuclear diatomic molecule attached to single covalent bond is called covalent radii.

For example, the bond length of H2 molecules is 0.74 Å. According to definition, covalent radius is 0.74/2 = 0.37 Å.

2. Metallic radii: The half of the distance between two nuclei of atoms attached by metallic bond in metals is called metallic radii.

3. Vander Waal’s radii: The half of the distance between two non-bonded nuclei of atoms attached by Vander Waal’s force of attraction is called Vander Waal’s radii.

The strength of various bonds is:

Covalent bond > Metallic bond >> Vander Waal’s force of attraction.

Therefore, bond length increases in the order:

Covalent radii < Metallic radii << Vander Waal’s radii

Variation of atomic radii in the periodic table:

In a group:

• In the group, from top to bottom the nuclear charge increases as well as there is increase in principle quantum number or number of shell (orbit).
• Effect of adding the new shell is larger than that of increase in nuclear charge.
• The effective nuclear charge per electron decreases.
• Hence, atomic radius (size) increases.
• Example : Atomic radii of alkali metals is: Li < Na < K < Rb < Cs

In period:

• In the period, from left to right the nuclear charge increases and electrons are added to the same shell.
• The effective nuclear charge per electron increases.
• Hence, atomic radius (size) decreases.
• Same period inert gas has highest atomic radii due to presence of Vander Waal’s radius.

Ionic radii: The ionic radius is the radii of the ions in crystal.

1. Cationic radius:

• The cation is formed by removal of one or more electron from an atom.
• The effective nuclear charge per electron in cation is more than the parent atom.
• Hence, the size of cation is smaller than that of parent atom.

Example: Fe > Fe⁺² > Fe⁺³

2. Anionic radius:

• The anion is formed by addition of one or more electron to an atom.
• The effective nuclear charge per electron in anion is less than the parent atom.
• Hence, the size of anion is larger than that of parent atom.

Example: O^-2 > O^– > O

Size of isoelectronic species:

• Those species having same number of electrons but different nuclear charge are called isoelectronic species.
• In isoelectronic species, size (radii) decreases with the increase in nuclear charge.

Example: N^-3 > O^-2 > F^– > Ne > Na⁺ > Mg⁺² > Al⁺³

Ionization energy (I.E.) / Ionization potential

The minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state to produce a cation is called ionization energy.

M (g) + I.E.  → M⁺ (g) + e^–

Note: It is an endothermic process (∆H = +ve) and measured in electron volt (eV) or KJ/mole

Successive ionization energies:

The term ionization enegy (I.E.) is in place of first ionization energy. The energy required to remove second, third, and fourth electrons are called second, third and fourth ionization energies respectively and are denoted by IE₂, IE₃ and IE₄.

M (g) + IE₁ → M⁺ (g) + e^–

M⁺ (g) + IE₂ → M²⁺ (g) + e^–

M²⁺ (g) + IE₃ → M³⁺ (g) + e^–

Factors affecting I.E.

1. Atomic size: Ionization energy decreases with the increase in atomic size.

2. Nuclear charge: Ionization energy increases with the increase in nuclear charge.

3. Shielding or Screening effect: If the shielding or screening effect of the inner electrons increases then ionization energy decreases.

• In multi-electron atoms, the electrons in the valence shell are pulled by the nucleus and repelled by the electrons of inner shells. Thus, outermost electrons experience less attraction from the nucleus under the combined effect of attractive and repulsive force acting on the valence electrons. This effect is called shielding effect or screening effect.
• The actual charge felt by the valence electron as a result of shielding effect is called an effective nuclear charge (Z_eff).
• With the increase in shielding effect, effective nuclear charge (Z_eff) decreases and hence I.E. decreases.

4. Penetration power of sub shell:

• More penetrating (i.e. more closer) are the sub-shells of a shell to the nucleus more tightly the electrons are held by the nucleus and more is the I.E.
• The penetrating power follows the order : s > p > d > f

5. Electronic configuration:

Half filled and completely filled orbitals are more stable than others and hence more energy is needed to remove an electron from such atoms. Thus, more stable the electronic configuration, the greater will be the I.E.

• Inert gases have highest I.E. due to completely filled orbital. ‘He’ has highest I.E. in the periodic table.
• Elements like Be (1s²,2s²) and Mg (1s²,2s²,2p⁶,3s²) have slightly higher I.E. due to extra stability of fully filled s-orbitals.
• Elements like N (1s², 2s², 2p³) and P (1s²,2s²,2p⁶,3s²,3p³) have higher I.E. due to extra stability of half-filled p-orbitals.

Variation of Ionization energy in the periodic table:

In group: From top to bottom in a group, atomic size increases and shielding effect also increases . Hence, ionization energy gradually decreases.

For example: Ionization energies of alkali metals is : Li > Na > K > Rb > Cs

In period: From left to right in a period, nuclear charge increases and atomic size decreases, so, there is gradual increase in I.E.

However, some elements show irregularities in the general trend. This may due to the extra stability of half-filled and completely filled electronic configurations.

For example: the variation of I.E. among the elements of II^nd period is:

Li < Be > B < C < N > O < F < Ne

Electron Affinity

The amount of energy released when an electron is added to an isolated gaseous atom in its ground state to form a gaseous anion is called electron affinity (EA).

X (g) + e^–  → X^– (g) + energy (EA)

Note: It is measured in electron volt (eV) or KJ/mole.

Factors affecting electron affinity:

1. Atomic size: Electron affinity decreases with the increase in atomic size.
2. Nuclear charge: Electron affinity increases with the increase in nuclear charge.
3. Screening effect: Electron affinity decreases with the increase in screening effect.
4. Electronic configuration: Elements having stable electronic configuration like half filled and completely filled orbitals have EA either very low or almost zero as they do not accept additional electrons so easily.

• EA of inert gases is zero due to completely filled orbitals.
• EA of alkaline earth metals is almost zero due to completely filled s-orbital.
• EA of N, P is very low due to half filled orbital.

Variation of electron affinity in the periodic table:

In period: From left to right in a period, nuclear charge increases and atomic size decreases, so, there is gradual increase in EA.

However, some elements show irregularities in the general trend. This may due to the extra stability of half-filled and completely filled electronic configurations.

For example: the variation of EA among the elements of II^nd and III^rd period is:

In 2^nd period : Ne < Be < N < Li < B < C < O < F

• Halogens possess maximum electron affinity due to small size and maximum effective nuclear charge and after gaining one electron, they attain stable inert gas configuration.

In group: From top to bottom in a group, atomic size increases and shielding effect also increases . Hence, EA gradually decreases.

For example: EA of alkali metals is : Li > Na > K > Rb > Cs

• However, the electron affinity of fluorine is lower than chlorine. Due to small size of fluorine, the incoming electron feels more repulsion and less attraction. In case of chlorine, incoming electron feels less repulsion and more attraction than in fluorine. Hence, the EA of fluorine is lower than that of chlorine.

The EA order of halogens is : Cl > F > Br > I

This type of anomaly is also observed in chalcogens (group 16) i.e. S>O>Se>Te

Anomalous Electron affinity:

• Electron affinity is …………….
• Generally, From left to right in a period, nuclear charge increases and atomic size decreases, so, there is gradual increase in EA. From top to bottom in a group, atomic size increases and shielding effect also increases . Hence, EA gradually decreases.
• However, some exceptions in the general trend in EA are found, which is called anomalous EA. Some of the anomalies are:

1. Zero and very low EA : Elements having stable electronic configuration like half filled and completely filled orbitals have EA either very low or almost zero as they do not accept additional electrons so easily.

• EA of inert gases is zero due to completely filled orbitals.
• EA of alkaline earth metals is almost zero due to completely filled s-orbital.
• EA of N, P is very low due to half filled orbital.

2. Halogens have highest EA: Halogens possess maximum electron affinity due to small size and maximum effective nuclear charge and after gaining one electron, they attain stable inert gas configuration.

3. Electron affinity of fluorine is lower than chlorine: Due to small size of fluorine, the incoming electron feels more repulsion and less attraction. In case of chlorine, incoming electron feels less repulsion and more attraction than in fluorine. Hence, the EA of fluorine is lower than that of chlorine.

The EA order of halogens is : Cl > F > Br > I

This type of anomaly is also observed in chalcogens (group 16) i.e. S>O>Se>Te

Eletronegativity (E.N.)

• Electronegativity of an element is defined as the relative tendency of an atom in a molecule to attract shared pair of electrons towards itself.
• It has no unit.
• Higher the difference in electronegativity, more the polarity in the bond.

Factors affecting Electronegativity:

1. Atomic size: Smaller the size of an atom, the greater is its tendency to attract the shared pair of electrons towards itself. Hence,the electronegativity increases with a decrease in size of the atom.

2. Effective nuclear charge (Z_eff): EN increases with the increase in Z_eff.

3. Ionization energy and Electron affinity: Higher the value of I.E. and EA, higher is the EN.

4. Number and nature of atoms bonded to it: EN of an element depends upon the number and nature of the atoms to which it is bonded. For example, the EN of phosphorous in PCl₅ is higher than in PF₅, since fluorine is more electronegative than chlorine.

5. Type of hybridization: The EN increases as the s-character in hybrid orbital increases.

For example: EN of carbon in methane, ethane and ethyne is:

Ethyne > Ethene > Methane

6. Charge on the ion: Cation is smaller in size while anion has larger size as compared to that of parent atom. Hence, EN increases with the increase in +ve charge and decrease in negative charge.

For example:

Fe < Fe⁺² < Fe⁺³

O^-2 < O^– < O

Variation of electronegativity in the periodic table:

In group: From top to bottom in a group, atomic size increases and shielding effect also increases. Hence, EN gradually decreases.

For example: EN of alkali metals is : Li > Na > K > Rb > Cs

EN of halogens is : F > Cl > Br > I >At

In period: From left to right in a period, nuclear charge increases and atomic size decreases, so, there is gradual increase in EN.

For example: EN order of second period : Li < Be < B < C < N < O < F

Metallic character [Electropositive character]

• The tendency of an element to lose an electron to form a cation is called electropositive character.
• The electropositive character of metal is called a metallic character.
• Lesser the value of ionization energy, more will be the metallic character and vice-versa.

Variation of metallic character in the periodic table:

In group: The metallic character of elements increases in going from top to bottom in a group. This is due to increase in size of the atom.

For example: Group 1: Li < Na < K < Rb < Cs

In period: The metallic character of elements decreases in going from left to right in a period. This is due to increase in effective nuclear charge.

For example: Period 3: Na > Mg > Al

Diagonal relationship

Similarities between certain pairs of elements that are diagonally adjacent to each other in the second and third period of the periodic table is called diagonal relationship.

Diagonal relationship is due to,

1. Almost same electronegativity (main cause)
2. Almost same atomic size.

Note: Bridge elements– Bridge elements are the elements of the second period of the periodic table. These elements show a relationship with the third-period,  which are diagonal.

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EMF (electromotive force) of a cell

The maximum potential difference that exists between two electrodes of a cell is called the electromotive force (EMF) of the cell.

In other words, electromotive force is difference in potential which causes the current to flow from an electrode of higher potential to an electrode of a lower potential.

• It is also known as cell potential
• It is measured in volts.

The emf of a cell measured under standard conditions is called standard emf. It is denoted by ‘E⁰_cell’. The standard conditions are:

1. Concentration of electrolytic solution is 1M.
2. Temperature of the system is 25⁰C.
3. Pressure of the gas is 1 atmosphere.

Q) A galvanic cell is represented as

Zn/Zn⁺⁺//Cd⁺⁺/Cd

Write the ell reaction and calculate the standard emf of the cell.

Given that E⁰_Zn⁺⁺_/Zn = -0.70V and E⁰_Cd⁺⁺_/Cd = -0.40V

Nernst equation

The Nernst equation provides the relation between the cell potential of an electrochemical cell, the standard cell potential, temperature and the equilibrium constant.

Mathematically, the Nernst equation can be expressed as:

For an oxidation half-cell reaction, when the metal electrode M gives Mⁿ⁺ ion,

The Nernst equation takes the form:

The concentration of solid metal [M] is equal to zero. Hence,

At 25⁰C:

On putting the values of R, F and T at 25⁰C, the quantity 2.303RT/F comes to be 0.0591.

Thus, at 25⁰C, the Nernst equation can be written as:

• This equation is for half cell in which oxidation occurs.
• In case of reduction half reaction, the sign of E is reversed.
• If the Nernst equation is applied to the complete cell as a whole, then it will be in the form:

Q) Calculate the emf of the cell.

Zn/Zn⁺⁺(0.001M)//Ag⁺(0.1M)/Ag

The standard electrode potential E0 of Ag/Ag⁺ is 0.80V and Zn/Zn⁺⁺ is -0.76V.

Ans….

Nernst equation applications

Nernst equation can be used to calculate the following:

• Single electrode potential (oxidation or reduction potential) at any conditions.
• Standard electrode potential.
• Comparing the relative ability as an oxidizing or reducing agent.
• Emf of an electrochemical cell
• Unknown ionic concentrations.
• pH of a solution also can be measured using Nernst equation.

References

• Atkins, Peter, Paula, de Julio, Atkin’s Physical Chemistry, Seventh Edition, Oxford University Press, (Printed in India, 2002).
• Gurtu, J.N., Snehi, H., Advanced Physical Chemistry, Seventh Edition, Pragati Prakashan India, 2000.
• Madan, R.L., tuli, G.D., Physical Chemistry, S. Chand and company, New Delhi, 2012.

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• Ionic strength of a solution can be defined as the total concentration of ions present in the solution.
• In another word, ionic strength measures the concentration of ionic atmosphere in a solution.
• It is dimensionless quantity, it has no unit.
• It is denoted by ‘µ’.

Mathematically,

Q) Calculate the ionic strength of 0.1M solution of aluminium sulphate.

Ans.

Q) Calculate the ionic strength of 0.1M KCl and 0.1M CaCl₂ solution.

Activity and Activity coefficient

In electrolytic solution, the experimentally determined value of concentration of ions is less than the actual concentration. The effective concentration of ions or electrolyte in a solution is called as activity. It is denoted by a symbol ‘a’.

Mathematically, activity ‘a’ is taken as the product of actual concentration in molarity or molality and activity coefficient ‘f’.

i.e. a = Cf ——– (i)

Where,

C = Concentration in molarity or molality

f = activity coefficient.

• For very dilute solution, activity coefficient is nearly equal to one. So, the activity becomes equal to the actual concentration.

i.e. a = c

• For concentrated solution, activity coefficient is less than one.

Rearranging equation (i),

F = a/C

Thus, activity coefficient is defined as the ratio of the activity to the actual concentration.

The activity ‘a’ of the electrolyte is taken as the product of activities of cation and anion.

i.e. a = a₊ a_–

Where,

a₊ = activity of cation

a_–= activity of anion

Similarly, activity coefficient of an electrolyle is taken as the product of activity coefficients of cation and anion.

i.e. f = f₊ f_–

where,

f₊ = activity coefficient of cation

f_– = activity coefficient of anion

The activity and activity coefficient can’t be measured experimentally but their mean value can be determined.

Debye-Huckel limiting law- Expression for the activity coefficient of electrolyte in terms of ionic strength

• Debye-Huckel limiting law relates the mean activity coefficient of an electrolyte with valency of ions and ionic strength of the solution.
• This law is applicable only for very dilute solution. So, this law is called limiting law.
• Mathematically, this law can be expressed as:

If -log is plotted against , a straight line passing through the origin having slope equal to AZ₊Z_– is obtained.

Application of Debye-Huckel limiting law:

Debye-Huckel limiting law can be used to calculate the mean activity coefficient of an electrolyte if the ionic strength or concentration of the solution and the value of ‘A’ is known.

References

• Atkins, Peter, Paula, de Julio, Atkin’s Physical Chemistry, Seventh Edition, Oxford University Press, (Printed in India, 2002).
• Gurtu, J.N., Snehi, H., Advanced Physical Chemistry, Seventh Edition, Pragati Prakashan India, 2000.
• Madan, R.L., tuli, G.D., Physical Chemistry, S. Chand and company, New Delhi, 2012.

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An acid-base indicator is a weak organic acid or base that indicates the end point of acid base titration by a visual change in colour. They are weak acid or base having different colour in basic and acidic solution.

Phenolphthalein and methyl orange are two common examples of acid base indicator.

Mechanism of acid-base indicator action – Ostwald’s theory

According to Ostwald’s theory,

1. Acid-base indicators are regarded as organic weak acid or weak base.
2. Acid-base indicators undergo partial ionization in water to furnish H⁺ ions and OH^– ions.
3. Undissociated form of indicator has different colour than the dissociated form of indicator.

The colour produced by indicator depends on the pH of the solution.

• Phenolphthalein is a weak organic acid. It is colourless at its unionized state while it has pink colour in its ionized state.

In acidic medium, equilibrium shifts to the left and solution is colourless. In alkaline medium, equilibrium shifts to the right forming pink coloured solution.

• Similarly, methyl orange is a weak organic base. Its unionized molecule is light yellow in colour but its ions are red in colour.

In alkaline medium, equilibrium shifts to the left and solution becomes light yellow. In acidic medium, equilibrium shifts to the right and solution becomes red.

Selection of indicator in acid-base titration

Indicators are the chemical species that indicates the end point by changing its own colour. Each pH indicator has its own pH range for its colour change. In order to determine the accurate end point of acid-base titration, the indicator should be selected in such a way that the pH range for the colour change of the indicator must coincide with the pH change (jump) at the end point of reaction.

Mainly two types of indicators i.e. methyl orange and phenolphthalein are used during acid-base titration. Methyl orange has pH range 3.1-4.4 and phenolphthalein has 8.2-10.

There are four types of acid base titrations.

1.Strong acid vs strong base titration:

The indicators pH range of both methyl orange (i.e. 3.1-4.4) and phenolphthalein (i.e. 8.2-10) coincide with the pH jump (i.e. 3-11) at end point. Hence, either methyl orange or phenolphthalein can be used as indicator during strong acid vs strong base titration.

2. Strong acid vs weak base titration:

The indicators pH range of only methyl orange (i.e. 3.1-4.4) coincides with the pH jump (i.e. 3-8). Hence, only methyl orange can be used as indicator during strong acid vs weak base titration.

3. Weak acid vs strong base titration:

The indicators pH range of phenolphthalein (i.e. 8.2-10) only coincides with the pH jump (i.e. 6-11). Hence, only phenolphthalein can be used as indicator during weak acid vs strong base titration.

4.Weak acid vs weak base titration:

The indicators pH range of neither methyl orange (3.1-4.4) nor phenolphthalein (i.e. 8.2-10) coincides with the pH jump ( 6-8). Hence, weak acid vs weak base cannot be usually titrated due to lack of suitable indicators.

Note : Some acid and base examples

Strong acid – HCl, HNO₃, H₂SO₄, etc.

Strong base – NaOH, KOH, Ca(OH)₂, etc.

Weak acid – formic acid, acetic acid, oxalic acid, etc.

Weak base – NH₄OH, Cu(OH)₂, Fe(OH)₃, etc.

References

• Negi, A.S., Anand, S.C., A Text Book of Physical Chemistry, Seventh Edition, New Age International Pvt. Ltd. Publishers, 1999.
• Atkins, Peter, Paula, de Julio, Atkin’s Physical Chemistry, Seventh Edition, Oxford University Press, (Printed in India, 2002).
• Gurtu, J.N., Snehi, H., Advanced Physical Chemistry, Seventh Edition, Pragati Prakashan India, 2000.
• Madan, R.L., tuli, G.D., Physical Chemistry, S. Chand and company, New Delhi, 2012.

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A solution which can resist the change in pH even on addition of small amount of acid or base is called buffer solution.

Example: In a biological system, blood is an example of buffer and its pH remains almost constant to 7.4 (i.e. 7.35-7.45). Blood contains a buffer of carbonic acid (H₂CO₃) and bicarbonate anion (HCO₃^–).

Types of buffer solution

1. Acidic buffer solution: The buffer solution prepared by mixing equimolar quantities of weak acid and its salt with strong base is known as acidic buffer solution.

The pH value of acidic buffer solution is less than 7.

Example: A solution containing equimolar quantities of acetic acid (CH₃COOH) and sodium acetate (CH₃COONa) is acidic buffer.

2. Basic buffer solution: The buffer solution prepared by mixing equimolar quantities of weak base and its salt with strong acid is known as basic buffer solution.

The pH value of acidic buffer solution is more than 7.

Example: A solution containing equimolar quantities of ammonium hydroxide (NH₄OH) and ammonium chloride (NH₄Cl) is basic buffer.

Mechanism of buffer action

Mechanism of buffer solution can be explained by considering acidic buffer solution as:

• Suppose, an acidic buffer prepared by adding equimolar quantities of CH₃COOH and CH₃COONa, where CH₃COOH is weakly ionized and CH₃COONa is strongly ionized.

• If we add few drops of acid (HCl) to the buffer solution then it provides H⁺ ions to the buffer. These H⁺ ions would combine with CH₃COO^– ions present in the buffer solution to form weakly ionized acetic acid as,

• Similarly, if we add small amount of NaOH to the buffer solution then it provides OH^– ions to the buffer. These OH^– ions would combine with H+ ions present in the buffer solution to form undissociated H₂O molecules.

• Thus, addition of H⁺ ions is neutralized by CH₃COO^– ions and OH^– is neutralized by H⁺ ions. Therefore, pH of the solution remains unchanged.

Buffer capacity and Buffer range

Buffer capacity : Buffer capacity can be defined as the ability of a solution to resist rapid changes in pH.

In other words, buffer capacity is defined as the number of moles of acid or base added per litre of the buffer required to cause a unit change in pH.

Buffer range: A buffer solution can be designed for any pH range, but a given buffer solution will work effectively only over a particular range.

The pH range over which a buffer solution is effective is termed as buffer range. For acidic buffer solution approximate pH range is pH=pK_a±1 and pOH=pK_b±1 for basic buffer solution.

Applications of buffer solution

Buffer solutions are important in various fields of science and industry due to their ability to maintain a stable pH level. Some of their key applications are:

1. Laboratory work: Buffers are essential in chemical and biological laboratories to maintain constant pH during experiments and reactions. This helps to ensure the accuracy and reproducibility of results.
2. Analytical chemistry: Instruments like ph meter depend on buffer solutions to calibrate and accurately measure the pH of other solutions.
3. Medicinal and clinical diagnostics: Buffers are used in various medical diagnoses, including blood tests.
4. Pharmaceuticals: Buffers are used in the formulation of drugs to ensure the stability and effectiveness of drugs, especially in intravenous medicines.
5. Water treatment: Buffer solutions are sometimes used to adjust and maintain the pH of water in swimming pools, aquariums and industrial water treatment processes.

Handerson-Hasselbalch equation – Determination of pH of buffer

Handerson-Hasselbalch equation is a formula used to calculate the pH of a buffer solution. It is expressed as:

Derivation of Handerson-Hasselbalch equation :

Preparation of buffer solution

Handerson-Hasselbalch equation can be used for the calculation while preparing buffer solution. Example:

Q) The dissociation constant of acetic acid is 1.75×10^-5. How will you prepare a buffer solution of 4.4 by mixing sodium acetate and acetic acid.

Ans..

References

• Atkins, Peter, Paula, de Julio, Atkin’s Physical Chemistry, Seventh Edition, Oxford University Press, (Printed in India, 2002).
• Gurtu, J.N., Snehi, H., Advanced Physical Chemistry, Seventh Edition, Pragati Prakashan India, 2000.

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When aniline is treated with sodium nitrite (NaNO₂) and dilute hydrochloric acid (HCl) at ice cold temperature (0-5⁰C), benzene diazonium chloride salt is obtained. This reaction is called diazotization reaction.

At first nitrous acid(HNO₂) is obtained by the action of NaNO₂ and dil.HCl and then the nitrous acid reacts with with aniline to give benzene diazonium chloride.

Mechanism of diazotization reaction

Step-I : Nitrous acid (HNO₂) is obtained by the action of NaNO₂ and dil. HCl in-situ.

Step-II : Protonation of nitrous acid followed by loss of water gives nitrosonium ion.

Step-III : Nucleophilic attack of aniline on nitrosonium ion gives nitrosoaniline.

Step-IV : Protonation of the oxygen in the nitrosoaniline, followed by elimination of water gives diazonium ion ,which gets chloride ion and forms benzenediazonium chloride.

Uses of Diazotization reaction

Diazotization reactions have several important uses in organic chemistry. Here are some of the main applications:

1. Synthesis of azo dyes: Diazotization reactions are widely employed in the production of azo dyes. Azo dyes are an important class of organic compounds used extensively in the textile, printing, and coloring industries. By diazotizing an aromatic amine and coupling it with a suitable coupling agent, various colored azo compounds can be synthesized.

2. Preparation of aryl halides: Diazotization reactions can be used to convert aromatic amines into aryl halides, such as aryl chlorides or aryl bromides. The diazonium salt can be treated with a copper halide or cuprous salt, resulting in the replacement of the diazonium group with a halogen atom.

Some other uses are :

3. Pharmaceutical synthesis: Diazotization reactions are utilized in the synthesis of pharmaceutical compounds. The diazonium salts obtained can be reacted with a range of nucleophiles, such as phenols, amines, and thiols, to produce pharmaceutical intermediates or active pharmaceutical ingredients (APIs).

4. Formation of heterocyclic compounds: Diazotization reactions are employed in the synthesis of various heterocyclic compounds. The diazonium salt can react with compounds containing active methylene groups, such as β-ketoesters or malonates, leading to the formation of fused or isolated heterocyclic rings.

5. Preparation of aromatic fluorides: Diazotization reactions can also be used to produce aromatic fluorides. By reacting the diazonium salt with hydrogen fluoride (HF), the diazonium group is replaced by a fluorine atom, resulting in the formation of aryl fluorides.

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Organometallic compounds are chemical compounds in which at least one metal atom like Li, Mg, Fe , etc. is bonded with carbon atom of organic molecule. Examples: C₂H₅MgBr, CH₃Li, (C₂H₅)₂Zn, (C₂H₅)₃Al, etc.

[Note: Compounds like Ni(CO)₄, NaCN, RCOONa, (C₃H₇O)₄Ti, etc. are not organometallic compounds as is not metal-carbon bond.]

The study of compounds containing metal-carbon bonds and their reaction s is called organometallic chemistry.

Henry Gilman was an American organic chemist known as the father of organometallic chemistry.He, discovered Gilman reagent , diorganolithium copper (R₂CuLi).

The first synthetic organometallic compound is Zeise’s salt , i.e. K[PtCl₃(C₂H₄)].

Classification of Organometallic compounds

On the basis of nature of the metal to carbon bond, organometallic compounds are classified as:

1. Sigma(σ) bonded organometallic compounds: Organometallic compounds having metal-carbon σ –bond are called sigma bonded organometallic compounds. Examples: Ethylmagnesium bromide (C₂H₅MgBr), Ethyl lithium (C₂H₅ Li), etc.

2. Pi(π) bonded organometallic compounds: Organometallic compounds having metal-carbon π –bond are called pi-bonded organometallic compounds. Examples: Zeise’s salt, Ferrocene i.e. Fe(C₅H₅)₂, etc. Examples:

Nature of metal-carbon bond

The relative electronegativities of carbon and metal suggests that the C-M bond will be highly polar. Electropositive metal atom gives up electrons to the carbon atom, resulting in a polar C-M bond with a partial positive charge on the metal and negative charge on the carbon.

Partial negative charge of an organic group bonded to a highly reactive metal results in a special reactivity which is as nucleophilic character.

Preparation of organometallic compounds

1. Organolithium compound (alkyl lithium) can be prepared as:

2. Organocopper compound (lithium dialkylcuprate) can be prepared as:

3. Organo cadmium compound (dialkyl cadmium) can be prepared as:

Grignard reagent

Grignard reagent is the alkyl magnesium halide. It is organometallic compound which is represented by the general formula RMgX. Eg.

CH₃MgBr ( Methyl magnesium bromide)

CH₃CH₂MgI ( Ethyl magnesium iodide)

Preparation : Grignard reagent can be prepared by heating haloalkanes or haloarenes with magnesium in the presence of dry ether.

Precautions :

⇒Grignard reagent is very sensitive to water. When it comes in contact with water, it converts to alkane.

Therefore, during preparation of Grignard reagent there should not be the presence of water molecules i.e. all the reagents should be anhydrous and apparatus oven dried.

⇒There should not be naked flames nearer.

Reactions of Grignard reagents

1. With water: When Grignard reagent is hydrolyzed with water, alkanes are obtained.

2. With alcohol: Alcohols react with Grignard reagent to form alkane.

3. With aldehydes and ketones: Aldehydes and ketones (i.e carbonyl compounds) when treated with Grignard reagent gives addition product, which upon acidic hydrolysis give alcohols.

a. Formaldehyde gives primary alcohol. Eg.

b. Aldehydes other than formaldehyde give secondary alcohol. Eg.

c. Ketones give tertiary alcohol. Eg.

4. With carbon dioxide: Carboxylation(carbonation) reaction:

Grignard reagent reacts with carbon dioxide in presence of dry ether to give addition product and hydrolysis of addition product in presence of dilute acid gives carboxylic acid. Eg.

5. With acid chlorides:

Acid chlorides react with Grignard’s reagent to give ketones, which further react with Grignard’s reagent to give tertiary alcohols. Eg.

6. With esters:

Esters (except those of formic acid) react with Grignard’s reagent to give ketones, which further react with Grignard’s reagent to give tertiary alcohols. Eg.

7. With hydrogen cyanide (HCN):

Reaction of Grignard reagent with HCN followed by acid hydrolysis gives aldehydes. Eg.

8. With alkane nitrile (RCN):

Reaction of Grignard reagent with RCN followed by acid hydrolysis gives ketones. Eg.

EXERCISE

1. Organometallic compounds are chemical compounds in which at least one metal atom like Li, Mg, Fe, etc. is bonded with carbon atom of organic molecule.

a. Write two examples of organometallic compounds.

b. What is the nature of carbon-metal bond of organometallic compounds?

c. Write one example each of sigma and pi bonded organometallic compound.

d. Potassium acetate and sodium ethoxide are not organometallic compounds, why?

2. The carbon-magnesium bond in a Grignard reagent is polar covalent with carbon being the negative end of the dipole, which explains its nucleophilicity and the magnesium-halogen bond is largely ionic.

a. Define Grignard reagent. How aliphatic and aromatic Grignard reagent is prepared?

b. Mention the precautions for its preparation.

c. What is the role of ether in Grignard reaction?

d. Ether should be pure and dry in Grignard reaction, why?

e. Write the products obtained by reacting methyl magnesium bromide with (a)HCHO (b)CO₂ (c)CH₃CHO (d)CH₃CN (e)CH₃COOCH₃ (f)CH₃COCl. Write their complete reactions.

3. Convert:

a. Hydrogen cyanide to acetaldehyde

b. Ethanenitrile to acetone

c. Phenyl magnesium bromide to benzoic acid

d. Benzonitrile to acetophenone

e. Ethyl methanoate to propan-2-ol

f. Ethanoyl chloride to 2-methyl propan-2-ol

4. Grignard reagent is an organometallic compound which is used to synthesize various organic compounds. Starting from CH₃MgBr, how would you prepare:

a. Methane

b. Ethanoic acid

c. Ethanol

d. Propan-2-ol

e. 2-methylpropan-2-ol

5. An aromatic hydrocarbon ‘X’ is treated with halogen in dark to give compound ‘Y’. the compound ‘Y’ on heating with magnesium metal gives compound ‘Z’. How can you convert benzoic acid and acetophenone from compound ‘Z’. Write a suitable reaction.

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Spectroscopy is the study of the interaction (absorption, emission, and scattering) between electromagnetic radiation and matter.

The term “spectroscopy” defines a large number of techniques that use electromagnetic radiation to obtain information about the structure and properties of matter (atoms and molecules).

The electromagnetic radiation is passed onto a sample matter and the response is observed and recorded. A plot of the response as a function of wavelength is referred to as a spectrum.

Spectroscopy is the modern technique which has many merits (advantages) over the classical techniques.

• Accuracy of these techniques is very high as compared to classical techniques.
• These techniques are quick (i.e. not time consuming).
• There is no wastage of sample. The sample used in investigation remains unchanged and can be reused for other studies, except in case of mass spectrometry, where the sample is destroyed.

Types of spectroscopic techniques

• Spectroscopy methods can be categorized depending on the types of radiation and the interaction between the radiation with matter.
• On the basis of nature of the interaction between the energy and the material, spectroscopy is classified as absorption, emission and scattering spectroscopy.
• On the basis of radiation involved in the interaction, spectroscopy can be radiowave, microwave, infrared, ultraviolet-visible, x-ray, andgamma ray spectroscopy.

There are many different types of spectroscopy, but the most common types used for chemical analysis are as follows:

• Nuclear magnetic resonance (NMR) spectroscopy
• Infrared (IR) spectroscopy
• Mass spectrometry
• Ultraviolet and visible spectroscopy
• Atomic spectroscopy
• Raman spectroscopy

Nuclear magnetic resonance (NMR) spectroscopy

If a nucleus like proton is placed in an external magnetic field, the magnetic moment of proton will be oriented either with or against the external magnetic field.

Out of two orientations, the one aligned with applied field is more stable ( i.e. associated with smaller energy) than the one aligned against the applied field (i.e. associated with higher energy). The difference in energy of two orientations is denoted by ∆E. Thus, if we desire to flip the nucleus from lower energy state to higher energy state, an amount of ∆E will have to absorbed by the nucleus.

In organic chemistry, we are more interested in the protons as the nucleus, as hydrogen is constituent of almost every organic compound. The particular branch of NMR spectroscopy where the nucleus is proton is called proton magnetic resonance (PMR) spectroscopy.

A PMR spectrum can be recorded by placing the substance containing hydrogen nucleus in magnetic field of different strength and noting the magnetic field strength at which the absorption of energy corresponding to flipping of proton from lower to higher energy state takes place.

Information that can be obtained from NMR spectrum (spectroscopy):

The following informations regarding structure of the molecule can be obtained from NMR spectroscopy:

• The number of signal tells us how many different types of protons are present in molecule.
• The position of signal tells us about the electronic environment of proton (shielding and deshielding of protons).
• The intensities of signals tells us how many protons of each kind are present.
• The splitting of signal into several peaks tells us about the environment of a proton with respect to other nearby protons.

Equivalent and non-equivalent protons:

Equivalent protons: Protons having the same environment in a molecule absorb the same energy (magnetic field) and give one signal(peak) in the NMR spectrum, such protons are called equivalent protons.

Non-equivalent protons: Protons having the different environment in a molecule absorb the different energy (magnetic field) and give different signals(peaks) in the NMR spectrum, such protons are called non-equivalent protons.

From the number of signals, we can tell how many different types of protons are present in the molecule.

For example:

Shielding and deshielding of protons:

When a compound is placed in a magnetic field, the electrons around the protons also generate a magnetic field called “induced magnetic field”. The induced magnetic field either oppose or supports the applied magnetic field.

• If the induced field opposes the applied field, the effective field strength experienced by the protons decreases. The proton is said to be shielded in this case. The shielded proton absorbs upfield in NMR spectrum as a greater applied field strength is required for the excitation of protons to higher level.
• If the induced field supports(rainforces) the applied field, the effective field strength experienced by the protons increases. The proton is said to be deshielded in this case. The deshielded proton absorbs downfield in NMR spectrum as a lower applied field strength is sufficient for the excitation of protons to higher level.

As a result of shielding and deshielding of protons, there is a shift in the position of the NMR signal as compared with standard substance(i.e. TMS), which is called chemical shift.

Chemical shift:

The shift in the position of PMR signal compared with a standard substance (i.e. TMS) as a result of shielding and deshielding by electrons is known as chemical shift.

The commonly used scale for expressing the chemical shift is δ-scale(ppm). Tetramethylsilane(TMS) is taken as internal reference. The position of TMS signal is taken as 0 ppm. Next scale is τ and is given as τ=10- δ.

From the chemical shift, the electronic environment of proton can be determined. Protons with different environment (non equivalent protons) have different chemical shift values.

Chemical shift value also tells us about the type of proton i.e. aliphatic, aromatic, alcoholic, carboxylic, etc.

Q) Why do we choose tetramethylsilane i.e. (CH₃)₄Si as a standard substance for recording chemical shift?

Spin-spin coupling (Splitting of NMR signals) :

• Spin-spin coupling in NMR spectroscopy is the effect of one nucleus’s magnetic field on other nuclei within the molecule, causing splitting of the NMR signals.
• We assume that one type of protons give rise to one signal(peak) in PMR (NMR) spectrum but in actual practice it is not so. For example:

• This compound contains two types of protons, therefore we expect to observe two signals but actually five signals are observed.
• This phenomenon of splitting of a peak into several peaks is called splitting of signals.
• Splitting of signals is due to spin-spin coupling between the neighboring protons. This can be explained as:

• Consider the absorption by one of the proton ‘b’. Magnetic field produced by neighboring proton ‘a’ may have two possible orientations with respect to applied magnetic field.

Case I : The magnetic field produced by the proton is aligned with the applied magnetic field.

Case II : The magnetic field produced by the proton is aligned against the applied magnetic field.

In the first case, the proton ‘b’ experience some increased magnetic field strength and absorb at a lower applied field (downfield).

In the second case, the proton ‘b’ experience some decreased magnetic field strength and absorb at a higher applied field (upfield).

Thus, one peak will be split into two peaks (doublet).

• Now, consider the absorption by the proton ‘a’. Magnetic field produced by two neighboring protons ‘b’ may have three possible orientations with respect to applied magnetic field.

Case I : The magnetic field produced by two protons is aligned with the applied magnetic field.

Case II : The magnetic field produced by two protons is aligned against the applied magnetic field.

Case III : One proton is aligned with the applied field and the other proton is aligned against the applied magnetic field.

In the first case, the proton ‘a’ experience some increased magnetic field strength and absorb at a lower applied field (downfield).

In the second case, the proton ‘a’ experience some decreased magnetic field strength and absorb at a higher applied field (upfield).

In the third case, the position of the signal will not change.

Thus, one peak will be split into three peaks (triplet).

Mass spectrometry

• In mass spectrometry, the molecules are bombarded with a stream of high energy electrons. The energetic electrons knock out generally one most loosely bound electron from the molecule. This process produces molecular ions or radical cations.
• The molecular ions being energetic is further fragmented to produce smaller ions called daughter ions or fragmented ions.

Each ion has certain m/e ratio i.e. ratio of mass to charge of the ions.

These ions are accelerated by electric field and the ions with particular m/e ratio are detected and recorded by mass spectrometer.

A graph between relative abundance (intensity) and m/e values of the ions is called mass spectrum.

Applications of Mass Spectrometry:

Mass spectrometry is an efficient method to elucidate the chemical composition of a sample or molecule. More recently, it has been used to classify biological products, in particular proteins and protein complexes, in a number of species. Usually, mass spectrometers can be used to classify unknown substances by molecular weight measurement, to measure known compounds, and to determine the structure and chemical properties of molecules.

• Due to its capability to distinguish between substances, Mass spectrometry is used to determine unknown substances.
• To identify the isotopes of a substance.
• In analytical laboratories that study the chemical, physical and biological properties of substances. It is favored over several other analytical techniques as it has less background interference since it is performed in a vacuum.

Infrared (IR) spectroscopy or Vibrational spectroscopy

• IR spectroscopy detects the absorption of light by a compound, in the IR region of the electromagnetic spectrum.
• The molecular vibrations are of two types – (i) stretching and (ii) bending vibration. A stretching vibration causes change in the interatomic distance while the bending vibration causes the change in bond angle.
• On absorption of light(radiation), the molecules are excited from lower vibrational levels to higher ones.
• Molecules will absorb such frequencies as are needed to excite molecules from lower vibrational level to permitted higher energy levels.
• Every bond and every functional group has a specific absorption frequency to excite the molecule to higher vibrational level.
• A graph between the amount of absorbance (or transmittance) of IR light against the frequency (or wavelength) of this light is called infrared (IR) spectrum.

Functional group region and finger print region in IR spectrum:

• Two important areas (reasons) in a IR spectrum are the 4000-1450 cm^-1 and 1450-500 cm^-1
• The higher frequency region (4000-1450 cm^-1) is called functional group region and the lower frequency region (1450-500 cm^-1) is called finger print region.

• In the fingerprint region, the spectra usually consist of bending vibrations within the molecule.The pattern of peaks is more complicated. The fingerprint region is important because each different compound produces its own unique pattern of peaks (like a fingerprint) in this region.
• In the functional group region, the spectra usually consist of stretching vibrations within the molecule. This region contains relatively few peaks. For example:

1. The peak (band) at 3200-3600 cm^-1 is due to O-H bond stretching of alcohols and phenols.
2. The peak (band) at 1690-1760 cm^-1 is due to C=O bond stretching of aldehydes, ketones, carboxylic acids and esters.

 Applications of IR-spectroscopy:

Infrared spectroscopy is widely used in industry as well as in research. It is a simple and reliable technique for measurement, quality control and dynamic measurement. It is also employed in forensic analysis in civil and criminal analysis.

Some of the major applications of IR spectroscopy are as follows:

1. Identification of functional group and structure elucidation:

Entire IR region is divided into group frequency region and fingerprint region. Range of group frequency is 4000-1500 cm^-1 while that of finger print region is 1500-400 cm^-1.

In group frequency region, the peaks corresponding to different functional groups can be observed. According to corresponding peaks, functional group can be determined.

Each atom of the molecule is connected by bond and each bond requires different IR region so characteristic peaks are observed. This region of IR spectrum is called as finger print region of the molecule. It can be determined by characteristic peaks.

2. Identification of substances:

IR spectroscopy is used to establish whether a given sample of an organic substance is identical with another or not. This is because large number of absorption bands is observed in the IR spectra of organic molecules and the probability that any two compounds will produce identical spectra is almost zero. So if two compounds have identical IR spectra then both of them must be samples of the same substances.

IR spectra of two enatiomeric compound are identical. So IR spectroscopy fails to distinguish between enantiomers.

For example, an IR spectrum of benzaldehyde is observed as follows.

C-H stretching of aromatic ring-3080 cm^-1

C-H stretching of aldehyde-2860 cm^-1 and 2775 cm^-1

C=O stretching of an aromatic aldehyde-1700 cm^-1

C=C stretching of an aromatic ring-1595 cm^-1

C-H bending-745 cm^-1 and 685 cm^-1

No other compound than benzaldehyde produces same IR spectra as shown above.

3. Studying the progress of the reaction

Progress of chemical reaction can be determined by examining the small portion of the reaction mixture withdrawn from time to time. The rate of disappearance of a characteristic absorption band of the reactant group and/or the rate of appearance of the characteristic absorption band of the product group due to formation of product is observed.

4. Detection of impurities

IR spectrum of the test sample to be determined is compared with the standard compound. If any additional peaks are observed in the IR spectrum, then it is due to impurities present in the compound.

UV – Visible spectroscopy or electronic spectroscopy

• The Principle of UV-Visible Spectroscopy is based on the absorption of ultraviolet light or visible light by chemical compounds. When the matter absorbs the light, it undergoes excitation and de-excitation, resulting in the production of a spectrum.
• When matter absorbs ultraviolet radiation, the electrons present in it undergo excitation. This causes them to jump from a ground state (an energy state with a relatively small amount of energy associated with it) to an excited state (an energy state with a relatively large amount of energy associated with it).
• It is important to note that the molecule does not absorb just any radiation. It absorbs only that radiation which possess appropriate energy required to electronic transition.
• The absorption of radiation is observed (detected) with the help of spectrophotometer.
• The graph between the amount of radiation absorbed by the sample(absorbance) and the wavelength of the radiation is called absorption spectrum.

Different types of electronic transitions:

When a molecule absorbs the radiations, the electrons are excited to higher levels. The electrons involved could be σ electron (occupying σ molecular orbital) or π electron (occupying π molecular orbital) or n electron (non-bonding). In the diagram below, σ, π and n electrons have been indicated in a molecule of aldehyde (RCHO).

Allowed and forbidden transitioins:

Molecules absorb only such radiations which have appropriate energy to excite the electrons to allowed higher level. The probability of a particular electronic transition has found to depend upon the value of extinction coefficient (€) and certain other factors. According to these factors, transitions can be divided into two categories:
(i) Allowed transitions
(ii) Forbidden transitions

(i) Allowed transitions – These are the transitions having extinction coefficient(€)= 10⁴   or more. π→ π* transitions fulfill these requirements.

For example, in 1,3-butadiene which shows absorption at 217nm has  € value 21000 represents an allowed transition.

 (ii) Forbidden transitions – These are the transitions having extinction coefficient(€) less than 10⁴ . n→ π* transitions are forbidden transitions.

For example, transition of saturated aldehydes which shows weak absorption near 290nm has  € value 100 represents a forbidden transition.

Chromophore and Auxochrome :

Chromophore:

The chromophore was previously defined as a functional group which gives (imparts) the colour to the compound. For example- nitro group is a chromophore because its presence in a compound gives yellow colour to the compound. But these days the term chromophore is defined as any group which absorbs electromagnetic radiation in a visible or UV region, which may or may not impart any colour to the compound.

Some of the important chromophores are :

Auxochrome :

It is a group which itself does not act as a chromophore but when attached to a chromophore, it shifts the absorption towards longer wavelength along with an increase in the intensity of absorption.

Some commonly known auxochromic groups are: -OH, -NH₂, -OR, -NHR, and –NR₂. For example: When the auxochrome –NH₂ group is attached to benzene ring, its absorption changes from λ_max=225 nm (ɛ_max 203) to λ_max=280 (ε_max1430)
All auxochromes have one or more non-bonding pairs of electrons. If an auxochrome is attached to a chromophore, it helps in extending the conjugation by sharing of non-bonding pair of electrons as shown below.

Bathochromic and Hypsochromic shift:

Bathochromic or red shift:

Shift of an absorption of light towards higher wavelength (lower energy) or towards red portion of spectrum is known as bathochromic shift or red shift.

This effect may be produced by changing the polarity of solvent. It is also produced if two or more chromophores are present in conjugation.

For example, 1,3-butadiene shows absorption at 217nm.

Hypsochromic or blue shift:

Shift of an absorption of light towards lower wavelength (higher energy) or towards blue portion of spectrum is known as hypsochromic shift or blue shift.

This effect may be produced by the removal of conjugation (auxochrome) or by changing the polarity of solvent.

For example, protonation of aniline causes a blue shift from 280 nm to 203 nm because the aniline ion formed by protonation has no electron pair(i.e. conjugation is removed).

Hyperchromic and Hypochromic effect:

Increase in intensity of absorption is called hyperchromic effect while the decrease in intensity of absorption is called hypochromic shift.

Introduction of auxochromes (conjugation) into the system causes hyperchromic effect and removal of conjugation causes hypochromic effect.

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Triphenyl phosphonium alkylide (simply phosphorous ylide) is called witting reagent. When carbonyl compound (aldehyde or ketone) is treated with an ylide, olefin (alkene) is formed. This reaction is called witting reaction.

Mechanism of witting reaction

Question- Answer from witting reaction

Q) Give the reaction and mechanism when 2-hexanone is treated with witting reagent.

When 2-hexanone is treated with witting reagent, 2-methylhex-1-ene is formed. This reaction is called witting reaction. The reaction involved and mechanism is given below.

Mechanism:

Q) Give two different methods for witting synthesis of 2-methyl-1-hexene.

2-methyl-1-hexene can be prepared using witting synthesis by following two methods:

Method-I :

Method-II :

Q) Write product and mechanism of the reaction:

This reaction is witting reaction. The reaction involved and mechanism is given below:

Mechanism:

References

• Morrison, R.T. , Boyd, R.N., Organic Chemistry, Sixth edition, Prentice-Hall of India Pvt. Ltd., 2008.
• March, j., Advanced Organic Chemistry, Fourth edition, Wiley Eastern Ltd. India, 2005.
• https://chemicalnote.com/category/organic-chemistry/name-reactions/

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