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Class 12 Chemistry Revision Notes Chapter 11 Alcohols, Phenols and Ethers 

• Structure of alcohols:
• Preparation of alcohols:
  a) From alkene
  
  b) From esters
  
  c) From aldehydes and ketones
  
  d) From carboxylic acids
  
• Structure of phenols:

• Preparation of phenols:

a) From benzene

b) From chlorobenzene

c) From cumene

d) From aniline

• Physical properties of alcohols and phenols:

a) Boiling points: Boiling points of alcohols and phenols are higher in comparison to other classes of compounds. This is because the –OH group in alcohols and phenols is involved in intermolecular hydrogen bonding.

The boiling points of alcohols and phenols increase with increase in the number of carbon atoms. This is because of increase in van der Waals forces with increase in surface area.
In alcohols, the boiling points decrease with increase of branching in carbon chain. This is because of decrease in van der Waals forces with decrease in surface area.
b) Solubility: Solubility of alcohols and phenols are soluble in water due to their ability to form hydrogen bonds with water molecules. The solubility of alcohols decreases with increase in size of alkyl/aryl (hydrophobic) groups.

• Chemical properties of alcohols:

I. Reactions involving cleavage of O–H bond: Alcohols react as nucleophiles:
a) Reaction with metals

b) Esterification reaction



II. Reactions of alcohols involving cleavage of carbon – oxygen (C–O) bond:
a) Reaction with hydrogen halides

b) Reaction with phosphorus trihalides

c) Dehydration reaction
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d). Oxidation reaction
i)
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• Chemical properties of phenols:

I. Reactions involving cleavage of O–H bond: Alcohols react as nucleophiles:
a) Reaction with metals

b) Esterification reaction



II. Other chemical reactions of phenols:

III. Acidic nature of phenol and alcohol:
a). Phenol > H2O > Primary alcohol > Secondary alcohol > Tertiary alcohol.
The acidic character of alcohols is due to the polar nature of O–H bond. Alkyl group is an electron-releasing group (–CH3, –C2H5) or it has electron releasing inductive effect (+I effect). Due to +I effect of alkyl groups, the electron density on oxygen increases. This decreases the polarity of O-H bond. And hence the acid strength decreases.

b) Phenol is more acidic than alcohol: In phenol, the hydroxyl group is directly attached to the sp2hybridised carbon of benzene ring which acts as an electron withdrawing group whereas in alcohols, the hydroxyl group is attached to the alkyl group which have electron releasing inductive effect. In phenol, the hydroxyl group is directly attached to the sp2hybridised carbon of benzene ring whereas in alcohols, the hydroxyl group is attached to the sp3hybridised carbon of the alkyl group. The sp2hybridised carbon has higher electronegativity than sp3hybridised carbon. Thus, the polarity of O–H bond of phenols is higher than those of alcohols. Hence, the ionisation of phenols is higher than that of alcohols.
The ionisation of an alcohol and a phenol takes place as follows:

In alkoxide ion, the negative charge is localised on oxygen while in phenoxide ion, the charge is delocalised.

The delocalisation of negative charge makes phenoxide ion more stable and favours the ionisation of phenol. Although there is also charge delocalisation in phenol, its resonance structures have charge separation due to which the phenol molecule is less stable than phenoxide ion.

c) In substituted phenols, the presence of electron withdrawing groups such as nitro group enhances the acidic strength of phenol. On the other hand, electron releasing groups, such as alkyl groups, in general, decreases the acid strength. It is because electron withdrawing groups lead to effective delocalisation of negative charge in phenoxide ion.

• Differentiate between organic compounds:

• Structure of ethers:
  
• Preparation of ethers:

a) From alcohols

b) From alkyl halide and sodium alkoxide

Here, the alkyl halide should be primary and alkoxide should be tertiary. In case of aromatic ether, the aromatic part should be with phenoxide ion.

• Physical properties of ethers

a) Miscibility: Miscibility of ethers with water resembles those of alcohols of the same molecular mass. This is due to the fact that just like alcohols, oxygen of ether can also form hydrogen bonds with water molecule.

b) Boiling points:
Ethers have much lower boiling points than alcohols. This is due to the presence of hydrogen bonding in alcohols. Hydrogen bonding is absent in ethers.

• Chemical properties of ethers:

a) Cleavage of C–O bond in ethers:
R-O-R’ + HX → R-X + R’OH
Excess
The order of reactivity of hydrogen halides is as follows: HI >HBr>HCl
Alkyl halide formed is always the lower alkyl group. But if a tertiary alkyl group is present, the alkyl halide is always tertiary. In case of phenolic ethers, the cleavage occurs with the formation of phenol and alkyl halide.
b) Electrophilic substitution reaction in aromatic ethers:

The electrophilic substitution reaction of aromatic ether involves the following reaction:

• Other conversion reactions:

a) Phenol to salicyldehyde

b) Phenol to benzene diazonium chloride

Class 12 Chemistry Revision Notes Chapter 10 Haloalkanes and Haloarenes

• Nature of C-X bond in alkyl halides: X is more electronegative than carbon. So, the C-X bond is polarized with C having a partial positive charge and X having a partial negative charge.
• Preparation of haloalkanes:

a) 

b) 

c) 

d) Halogen Exchange Method:

• Preparation of haloarenes:

a) By elecrophilic substitution reaction:

b) Sandmeyer’s reaction:

c) Gattermann reaction:

d) From Diazonium Chloride:

e). Balz – Schiemann reaction:

• Physical properties of haloalkanes:

a) Solubility

1. Although haloalkanes are polar in nature, yet they are practically very slightly soluble in water.
2. In order for a haloalkane to dissolve in water, energy is required to overcome the attractions between the haloalkane molecules and break the hydrogen bonds between water molecules.
3. However Haloalkanes are not able to form hydrogen bonds with water and therefore, less energy is released when new attractions are set up between the haloalkane and the water molecules because these are not as strong as the original hydrogen bonds in water molecules.
4. As a result, solubility of haloalkanes in water is low.

b) Density

1. Simple fluoro and chloroalkanes are lighter than water while bromides and polychlorodevrivatives are heavier than water.
2. With the increase in number of carbon atoms, the densities go on increasing. With the increase in number of halogen atoms, the densities go on increasing. The densities increase in the order: Fluoride < chloride < bromide < iodide
3. The density also increases with increasing number and atomic mass of the halogen.

c) Boiling Points

1. Molecules of organic halogen compounds are generally polar.
2. Due to the polarity as well as higher molecular mass as compared to the parent hydrocarbon, the intermolecular forces of attraction (dipole – dipole and van der Waals) between the molecules are stronger in halogen derivatives of alkanes.
3. As a result melting and boiling points of chlorides, bromides and iodides are considerably higher than those of the parent hydrocarbon of comparable molecular mass.
4. For the same alkyl group the boiling points of alkyl chlorides, bromides and iodides follow the order RI >RBr>RCl> RF where R is an alkyl group. This is because with the increase in the size of the halogen, the magnitude of van der Waals force increase.
5. In general, the boiling points of chloro, bromo and iodo compounds increase with increase in the number of halogen atoms.
6. For the same halogen atom, the boiling points of haloalkanes increase with increase in the size of alkyl groups.
7. For isomeric alkyl halides, the boiling points decrease with branching. This is because branching of the chain makes the molecule more compact and, therefore, decrease the surface area. Due to decrease in surface area, the magnitude of van der Waals forces of attraction decreases and consequently, the boiling points of the branched chain compound is less than those of the straight chain compounds.

• Physical Properties of Haloarenes:

a. These are generally colourless liquids or crystalline solids.

b. These are heavier than water.

c. Melting and boiling points of haloarenes
i. Melting and boiling points of haloarenes are nearly the same as those of alkyl halides containing the same number of carbon atoms.
ii. The boiling points of monohalogen derivatives of benzene are in the order: iodo>bromo>chloro>fluoro
iii. For the same halogen atom, the melting and boiling points increase as the size of the aryl group increases.
iv. The melting point of para isomer is quite higher than that of ortho or meta isomers. This is due to the fast that is has symmetrical structure and therefore, its molecules can easily pack loosely in the crystal lattice. As a result intermolecular forces of attraction are stronger and therefore, greater energy is required to break its lattice and it melts at higher temperature.

• Chemical properties of haloalkanes:

Nucleophilic substitution reaction:

Mechanism of Nucleophilic Substitution Reaction:

SN1 Mechanism

1. First order reaction.
2. Rate = k [RX] [Nu]
3. Racemic mixture
4. One step reaction
5. Order: CH3X < 10< 20< 30

SN2 Mechanism

1. Second order reaction
2. Rate = k [RX]
3. Inversion of configuration
4. Two step reaction
5. Order: CH3X > 10> 20> 30

• Elimination reaction: Dehydrohalogentaion(– elimination): When a haloalkane with β-hydrogen atom is heated with alcoholic solution of potassium hydroxide, there is elimination of hydrogen atom from β-carbon and a halogen atom from the α-carbon atom. As a result, an alkene is formed as a product. Zaitsev rule (also pronounced as Saytzeff) is followed.It states that “In dehydrohalogenation reactions, the preferred product is that alkene which has the greater number of alkyl groups attached to the doubly bonded carbon atoms.”
• Reaction with metals:

a) Reaction with Magnesium

b) Wurtz reaction

• Chemical properties of haloarenes:

a) Dow’s Process

b) With halogens

c) With conc. nitric and sulphuric acid

d) On heating with conc. sulphuric acid

e) With methyl chloride

f) With acetyl chloride

g) Fittig reaction: 

h) Wurtz – fittig reaction: 

i) Other conversions:

Coordination Compounds Class 12 Notes Chemistry Chapter 9

• Co-ordination compounds:

1. A coordination compound contains a central metal atom or ion surrounded by number of oppositely charged ions or neutral molecules. These ions or molecules re bonded to the metal atom or ion by a coordinate bond.
2. Example: 
3. They do not dissociate into simple ions when dissolved in water.

• Double salt

1. When two salts in stoichiometric ratio are crystallised together from their saturated solution they are called double salts
2. Example:  (Mohr’s salt)
3. They dissociate into simple ions when dissolved in water.

• Coordination entity:

1. A coordination entity constitutes a central metal atom or ion bonded to a fixed number of ions or molecules.
2. Example: In – represents coordination entity.

• Central atom or ion:

1. In a coordination entity, the atom/ion to which a fixed number of ions/groups are bound in a definite geometrical arrangement around it, is called the central atom or ion.
2. Example: In ,  is the central metal ion.

• Ligands:

1. A molecule, ion or group that is bonded to the metal atom or ion in a complex or coordination compound by a coordinate bond is called ligand.
2. It may be neutral, positively or negatively charged.
3. Examples: etc.

• Donor atom:

1. An atom of the ligand attached directly to the metal is called the donor atom.
2. Example: In the complex ,CN is a donor atom.

• Coordination number:

1. The coordination number (CN) of a metal ion in a complex can be defined as the number of ligand donor atoms to which the metal is directly bonded.
2. Example: In the complex , the coordination number of Fe is 6.

• Coordination sphere:

1. The central atom/ion and the ligands attached to it are enclosed in square bracket and are collectively termed as the coordination sphere.
2. Example: In the complex is the coordination sphere.

• Counter ions:

1. The ions present outside the coordination sphere are called counter ions.
2. Example: In the complex , K+ is the counter ion.

• Coordination polyhedron:

1. The spatial arrangement of the ligand atoms which are directly attached to the central atom/ ion defines a coordination polyhedron about the central atom.
2. The most common coordination polyhedra are octahedral, square planar and tetrahedral.
3. Examples: is square planar,  is tetrahedral while [Cu(NH3)6]3+ is octahedral.

• Charge on the complex ion: The charge on the complex ion is equal to the algebraic sum of the charges on all the ligands coordinated to the central metal ion.
• Denticity: The number of ligating (linking) atoms present in ligand is called denticity.
• Unidentate ligands:

1. The ligands whose only one donor atom is bonded to metal atom are called unidentate ligands.
2. Examples: 

• Didentate ligands:

1. The ligands which contain two donor atoms or ions through which they are bonded to the metal ion.
2. Examples: Ethylene diamine () has two nitrogen atoms, oxalate ion  has two oxygen atoms which can bind with the metal atom.

• Polydentate ligand:

1. When several donor atoms are present in a single ligand, the ligand is called polydentate ligand.
2. Examples: In , the ligand is said to be polydentate and Ethylenediaminetetraacetate ion  is an important hexadentate ligand. It can bind through two nitrogen and four oxygen atoms to a central metal ion.

• Chelate:

1. An inorganic metal complex in which there is a close ring of atoms caused by attachment of a ligand to a metal atom at two points.
2. An example is the complex ion formed between ethylene diamine and cupric ion, .

• Ambidentate ligand:

1. Ligands which can ligate (link) through two different atoms present in it are called ambidentate ligand.
2. Example: and . Here, can link through N as well as O while  can link through S as well as N atom.

• Werner’s coordination theory:

1. Werner was able to explain the nature of bonding in complexes.
2. The postulates of Werner’s theory are:

a). Metal shows two different kinds of valencies: primary valence and secondary valence.

b). The ions/ groups bound by secondary linkages to the metal have characteristic spatial arrangements corresponding to different coordination numbers.

c). The most common geometrical shapes in coordination compounds are octahedral, square planar and tetrahedral.

• Primary valence

1. This valence is normally ionisable.
2. It is equal to positive charge on central metal atom.
3. These valencies are satisfied by negatively charged ions.
4. Example: In , the primary valency is three. It is equal to oxidation state of central metal ion.

• Secondary valence

1. This valence is non – ionisable.
2. The secondary valency equals the number of ligand atoms coordinated to the metal. It is also called coordination number of the metal.
3. It is commonly satisfied by neutral and negatively charged, sometimes by positively charged ligands.

• Oxidation number of central atom: The oxidation number of the central atom in a complex is defined as the charge it would carry if all the ligands are removed along with the electron pairs that are shared with the central atom.
• Homoleptic complexes: Those complexes in which metal or ion is coordinate bonded to only one kind of donor atoms. For example: 
• Heteroleptic complexes: Those complexes in which metal or ion is coordinate bonded to more than one kind of donor atoms. For example: 
• Isomers: Two or more compounds which have same chemical formula but different arrangement of atoms are called isomers.
• Types of isomerism:

a). Linkage isomerism

b). Solvate isomerism or hydrate isomerism

c). Ionisation isomerism

d). Coordination isomerism

1. Structural isomerism
2. Stereoisomerism

a). Geometrical isomerism

b). Optical isomerism

• Structural isomerism:

1. It arises due to the difference in structures of coordination compounds.
2. Structural isomerism, or constitutional isomerism, is a form of isomerism in which molecules with the same molecular formula have atoms bonded together in different orders.

• Ionisation isomerism:

1. It arises when the counter ion in a complex salt is itself a potential ligand and can displace a ligand which can then become the counter ion.
2. Example: 

• Solvate isomerism:

1. It is isomerism in which solvent is involved as ligand.
2. If solvent is water it is called hydrate isomerism, e.g.,  and .

• Linkage isomerism:

1. It arises in a coordination compound containing ambidentate ligand.
2. In the isomerism, a ligand can form linkage with metal through different atoms.
3. Example:  and .

• Coordination isomerism:

1. This type of isomerism arises from the interchange of ligands between cationic and anionic entities of different metal ions present in a complex.
2. Example:  and .

• Stereoisomerism: This type of isomerism arises because of different spatial arrangement.
• Geometrical isomerism: It arises in heteroleptic complexes due to different possible geometrical arrangements of ligands.
• Optical isomerism: Optical isomers are those isomers which are non-superimposable mirror images.
• Valence bond theory:

1. According to this theory, the metal atom or ion under the influence of ligands can use its (n-1)d, ns, np or ns, np, nd orbitals for hybridisation to yield a set of equivalent orbitals of definite geometry such as octahedral, tetrahedral, and square planar.
2. These hybridised orbitals are allowed to overlap with ligand orbitals that can donate electron pairs for bonding.

Coordination Number	Type of hybridisation	Shape of hybrid
4		Tetrahedral
4		Square planar
5		Trigonalbipyramidal
6	(nd orbitals are involved – outer orbital complex or high spin or spin free complex)	Octahedral
6	d orbitals are involved –inner orbital or low spin or spin paired complex)	Octahedral

• Magnetic properties of coordination compounds:

A coordination compound is paramagnetic in nature if it has unpaired electrons and diamagnetic if all the electrons in the coordination compound are paired.

Magnetic moment  where n is number of unpaired electrons.

• Crystal Field Theory:

1. It assumes the ligands to be point charges and there is electrostatic force of attraction between ligands and metal atom or ion.
2. It is theoretical assumption.

• Crystal field splitting in octahedral coordination complexes:

• Crystal field splitting in tetrahedral coordination complexes:

• For the same metal, the same ligands and metal-ligand distances, the difference in energy between eg and t2g level is 
• Metal carbonyls:

1. Metal carbonyls are homoleptic complexes in which carbon monoxide (CO) acts as the ligand.
2. Example: 
3. The metal-carbon bond in metal carbonyls possess both s and p character.
4. The M–C  bond is formed by the donation of lone pair of electrons from the carbonyl carbon into a vacant orbital of the metal.
5. The M–C bond is formed by the donation of a pair of electrons from a filled d orbital of metal into the vacant antibonding orbital of carbon monoxide.
6. The metal to ligand bonding creates a synergic effect which strengthens the bond between CO and the metal.

Class 12 Chemistry Quick Revision Notes Chapter 8 The d and f Block Elements

The d -Block elements:

1. The elements lying in the middle of periodic table belonging to groups 3 to 12 are known as d – block elements.
2. Their general electronic configuration is  where (n – 1) stands for penultimate (last but one) shell.

• Transition element:

1. A transition element is defined as the one which has incompletely filled d orbitals in its ground state or in any one of its oxidation states.
2. Zinc, cadmium, mercury are not regarded as transition metals due to completely filled d – orbital.

• The f-Block elements: The elements constituting the f -block are those in which the 4 f and 5 f orbitals are progressively filled in the latter two long periods.
• Lanthanoids: The 14 elements immediately following lanthanum, i.e., Cerium (58) to Lutetium (71) are called lanthanoids. They belong to first inner transition series. Lanthanum (57) has similar properties. Therefore, it is studied along with lanthanoids.
• Actinoids: The 14 elements immediately following actinium (89), with atomic numbers 90 (Thorium) to 103 (Lawrencium) are called actinoids. They belong to second inner transition series. Actinium (89) has similar properties. Therefore, it is studied along with actinoids.
• Four transition series:

1. 3d – transition series. The transition elements with atomic number 21(Sc) to 30(Zn) and having incomplete 3d orbitals is called the first transition series.
2. 4d – transition series. It consists of elements with atomic number 39(Y) to 48 (Cd) and having incomplete 4d orbitals. It is called second transition series.
3. 5d – transition series. It consists of elements with atomic number 57(La), 72(Hf) to 80(Hg) having incomplete 5d orbitals. It is called third transition series.
4. 6d – transition series. It consists of elements with atomic number 89(Ac), 104(Rf) to 112(Uub) having incomplete 6d orbitals. It is called fourth transition series.

• General Characteristics of transition elements:

a) Metallic character: All transition elements are metallic in nature, i.e. they have strong metallic bonds. This is because of presence of unpaired electrons. This gives rise to properties like high density, high enthalpies of atomization, and high melting and boiling points.
b) Atomic radii: The atomic radii decrease from Sc to Cr because the effective nuclear charge increases. The atomic size of Fe, Co, Ni is almost same because the attraction due to increase in nuclear charge is cancelled by the repulsion because of increase in shielding effect. Cu and Zn have bigger size because the shielding effect increases and electron electron repulsions repulsion increases.
c) Lanthanoid Contraction: The steady decrease in the atomic and ionic radii of the transition metals as the atomic number increases. This is because of filling of 4f orbitals before the 5d orbitals. This contraction is size is quite regular. This is called lanthanoid contraction. It is because of lanthanoid contraction that the atomic radii of the second row of transition elements are almost similar to those of the third row of transition elements.
d) Ionisation enthalpy: There is slight and irregular variation in ionization energies of transition metals due to irregular variation of atomic size. The I.E. of 5d transition series is higher than 3d and 4d transition series because of Lanthanoid Contraction.
e) Oxidation state: Transition metals show variable oxidation states due to tendency of (n-1)d as well as ns electrons to take part in bond formation.
f) Magnetic properties: Most of transition metals are paramagnetic in nature as a result of which they give coloured compounds and it is all due to presence of unpaired electrons. It increase s from Sc to Cr and then decreases because number of unpaired and then decrease because number of unpaired electrons increases from Sc to Cr and then decreases. They are rarely diamagnetic.
g) Catalytic properties: Most of transition metals are used as catalyst because of (i) presence of incomplete or empty d – orbitals, (ii) large surface area, (iii) varuable oxidation state, (iv) ability to form complexes, e.g., Fe, Ni, V2O3, Pt, Mo, Co and used as catalyst.
h) Formation of coloured compounds: They form coloured ions due to presence of incompletely filled d – orbitals and unpaired electrons, they can undergo d – d transition by absorbing colour from visible region and radiating complementary colour.
i) Formation of complexes: Transition metals form complexes due to (i) presence of vacant d – orbitals of suitable energy (ii) smaller size (iii) higher charge on cations.
j) Interstitial compounds: Transition metals have voids or interstitials in which C, H, N, B etc. can fit into resulting in formation of interstitial compounds. They are non – stoichiometric, i.e., their composition is not fixed, e.g., steel. They are harder and less malleable and ductile.
k) Alloys formation: They form alloys due to similar ionic size. Metals can replace each other in crystal lattice, e.g., brass, bronze, steel etc.

• Preparation of Potassium dichromate (): It is prepared by fusion of chromate ore (FeCr2O4) with sodium carbonate in excess of air.
  
  
  
• Effect of pH on chromate and dichromate ions: The chromates and dichromates are inter-convertible in aqueous solution depending upon pH of the solution. The oxidation state of chromium in chromate and dichromate is the same.
  
  
• Potassium dichromate acts as a strong oxidizing agent in acidic medium:
  
• Preparation of Potassium permanganate (KMnO4):

a) Potassium permanganate is prepared by fusion of MnO4 with alkali metal hydroxide (KOH) in presence of O2 or oxidising agent like KNO3. It produces dark green K2MnO4 which undergoes oxidation as well as reduction in neutral or acidic solution to give permanganate.


b) Commercially, it is prepared by the alkaline oxidative fusion of MnO2 followed by the electrolytic oxidation of manganate (Vl).


c) In laboratory, Mn²+ salt can be oxidized by peroxodisulphate ion to permanganate ion.
In acidic medium:
In neutral or faintly basic medium: 

• Properties of Lanthanoids:

1. +3 oxidation state is most common along with +2 and +4.
2. Except Promethium, they are non – radioactive.
3. The magnetic properties of lanthanoids are less complex than actinoids.

• Properties of Actinoids:

1. Actinoids also show higher oxidation states such as +4, +5, +6 and +7.
2. They are radioactive.
3. The magnetic properties of the actinoids are more complex than those of the lanthanoids.
4. They are more reactive.

• Mischmetall

1. It is a well-known alloy which consists of a lanthanoid metal and iron and traces of S, C, Ca and Al.
2. A good deal of mischmetall is used in Mg-based alloy to produce bullets, shell and lighter flint.

Class 12 Chemistry Quick Revision Notes Chapter 7 The P-Block Elements

• The p-Block elements: Elements belonging to groups 13 to 18 of the periodic table are called p-block elements.
• General electronic configuration of p-block elements: The p-block elements are characterized by the ns2np1-6 valence shell electronic configuration.
• Representative elements: Elements belonging to the s and p-blocks in the periodic table are called the representative elements or main group elements.
• Inert pair effect: The tendency of ns2 electron pair to participate in bond formation decreases with the increase in atomic size. Within a group the higher oxidation state becomes less stable with respect to the lower oxidation state as the atomic number increases. This trend is called ‘inert pair effect’. In other words, the energy required to unpair the electrons is more than energy released in the formation of two additional bonds.

GROUP 15 ELEMENTS

• Nitrogen family: The elements of group 15 – nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb) and bismuth (Bi) belong to configuration is .
• Atomic and ionic radii:

1. Covalent and ionic radii increase down the group.
2. There is appreciable increase in covalent radii from N to P.
3. There is small increase from As to Bi due to presence of completely filled d or f orbitals in heavy elements.

• Ionisation energy:

1. It goes on decreasing down the group due to increase in atomic size.
2. Group 15 elements have higher ionisation energy than group 14 elements due to smaller size of group 15 elements.
3. Group 15 elements have higher ionization energy than group 16 elements because they have stable electronic configuration i.e., half-filled p-orbitals.

• Allotropy: All elements of Group 15 except nitrogen show allotropy.
• Catenation:

1. Nitrogen shows catenation to some extent due to triple bond but phosphorus shows catenation to maximum extent.
2. The tendency to show catenation decreases down the group.

• Oxidation states:

1. The common oxidation states are +3, +5 and –3.
2. The tendency to show –3 oxidation state decreases down the group because of decrease in electronegativity by the increase in atomic size.
3. The stability of +5 oxidation state decreases whereas stability of +3 oxidation state increases due to inert pair effect.
4. Nitrogen shows oxidation states from –3 to +5.
5. Nitrogen and phosphorus with oxidation states from +1 to +4 undergo oxidation as well as reduction in acidic medium. This process is called disproportionation.
   

• Reactivity towards hydrogen:

1. All group 15 elements from trihydrides, .
2. It belongs to  hybridisation.
3. The stability of hydrides decreases down the group due to decrease in bond dissociation energy down the group.
   

• Boiling point:
  

1. Boiling point increases with increase in size due to increase in van der Waals forces.
2. Boiling point of NH3 is more because of hydrogen bonding.

• Bond angle:
  

1. Electronegativity of N is highest. Therefore, the lone pairs will be towards nitrogen and hence more repulsion between bond pairs. Therefore bond angle is the highest. After nitrogen, the electronegativity decreases down the group.
2. Basicity decreases as NH3> PH3> AsH3> SbH3< BiH3. This is because the lone pair of electrons are concentrated more on nitrogen and hence the basicity will be maximum in the case of NH3. It will decrease down the group as the electronegativity decreases down the group. The reducing power of hydrides increases down the group due to decrease in bond dissociation energy down the group.

• Reactivity towards oxygen:

1. All group 15 elements from trioxides () and pentoxides ().
2. Acidic character of oxides decreases and basicity increases down the group. This is because the size of nitrogen is very small.
3. It has a strong positive field in a very small area. Therefore, it attracts the electrons of water O-H bond to itself and release H+ ions easily.
4. As we move down the group, the atomic size increases and so, the acidic character of oxide decreases and basicity increases down the group.

• Reactivity towards halogen:Group 15 elements form trihalides and pentahalides.

1. Trihalides: These are covalent compounds and become ionic down the group with hybridisation, pyramidal shape.
2. Pentahalidesa). They are lewis acids because of the presence of vacant d – orbitals.b). They possess  hybridisation and hence possess trigonalbirpyamidal shape.
3. PCl5 is ionic in solid state and exist as 
4. In PCl5, there are three equatorial bonds and two axial bonds. The axial bonds are longer than equatorial bonds because of greater repulsion from equatorial bonds.
5. Nitrogen does not form pentahalides due to absence of d– orbitals.

• Reactivity towards metals: All elements react with metals to form binary compounds in –3 oxidation state.
• Anomalous behaviour of nitrogen: The behaviour of nitrogen differs from rest of the elements.

Reasons:

i. It has a small size.
ii. It does not have d – orbitals
iii. It has high electronegativity
iv. It has high ionization enthalpy

• Dinitrogen:

a) Preparation:

b) Physical Properties:

i) It is a colourless, odourless, tasteless and non – toxic gas.
ii) It is chemically un-reactive at ordinary temperature due to triple bond in N ≡ N which has high bond dissociation energy.

• Ammonia:

1. Ammonia molecule is trigonal pyramidal with nitrogen atom at the apex.
2. It has 3 bond pairs and 1 lone pair.
3. N is hybridised.
4. Preparation:

Haber’s process:

Pressure 20010 Pa Temperature 773 K Catalyst is FeO with small amounts of  and 

• Nitric Acid:

Ostwald Process: The NO thus formed is recycled and the aqueous  can be concentrated by distillation upto ~ 68% by mass. Further concentration to 98% can be achieved by dehydration with concentrated . Nitric acid is strong oxidizing agent in the concentrated as well as in the dilute state.

Phosphorus:

a) It shows the property of catenation to maximum extent due to most stable P – P bond.

b) It has many allotropes, the important ones are:
i) White phosphorus
ii) Red phosphorus
iii) Black phosphorus

• White phosphorus:

1. Discrete tetrahedral P4 molecules
2. Very reactive
3. Glows in dark
4. Translucent waxy solid
5. Soluble in but insoluble in water
6. It has low ignition temperature, therefore, kept under water

• Red phosphorus

1. Polymeric structure consisting of chains of P4 units linked together
2. Less reactive than white phosphorus
3. Does not glow in dark
4. Has an iron grey lustre
5. Insoluble in water as well as 

• Black phosphorus

1. Exists in two forms –  black phosphorus and  black phosphorus
2. Very less reactive
3. Has an opaque monoclinic or rhombohedral crystals

• Phosphine

1. It is highly poisonous, colourless gas and has a smell of rotten fish.
2. Preparation
   
   
   

• Chlorides of Phosphorous:

a) Phosphorus Trichloride
i) It is a colourless oily liquid.

ii) Preparation

iii) With water,
It gets hydrolysed in the presence of moisture.

iv) Pyramidal shape, sp3 hybridisation

v) With acetic acid

vi). With alcohol

b) Phosphorus pentachloride

1. Yellowish white powder.
2. Trigonalbipyramidal shape, sp3dhybridisation .
3. Preparation
4. 
   
5. With water
6. 
   
7. With acetic acid
8. 
9. With alcohol
10. With metals

GROUP 16 ELEMENTS

• Oxidation states:

1. They show -2, +2, +4, +6 oxidation states.
2. Oxygen does not show +6 oxidation state due to absence of d – orbitals.
3. Po does not show +6 oxidation state due to inert pair effect.
4. The stability of -2 oxidation state decreases down the group due to increase in atomic size and decrease in electronegativity.
5. Oxygen shows -2 oxidation state in general except in and 
6. Thus, the stability of +6 oxidation state decreases and +4 oxidation state increases due to inert pair effect.

• Ionisation enthalpy:

1. Ionisation enthalpy of elements of group 16 is lower than group 15 due to half-filled p-orbitals in group 15 which is more stable.
2. However, ionization enthalpy decreases down the group.

• Electron gain enthalpy:

1. Oxygen has less negative electron gain enthalpy than S because of small size of O.
2. From S to Po electron gain enthalpy becomes less negative to Po because of increase in atomic size.

• Melting and boiling point:

1. It increases with increase in atomic number.
2. Oxygen has much lower melting and boiling points than sulphur because oxygen is diatomic ( ) and sulphur is octatomic ().

• Reactivity with hydrogen:

1. All group 16 elements form hydrides.
2. They possess bent shape.
3. Bond angle: 

• Acidic nature: 
  This is because the H-E bond length increases down the group. Therefore, the bond dissociation enthalpy decreases down the group.
• Thermal stability: 
  This is because the H-E bond length increases down the group. Therefore, the bond dissociation enthalpy decreases down the group.
• Reducing character: 
  This is because the H-E bond length increases down the group. Therefore, the bond dissociation enthalpy decreases down the group.
• Reactivity with oxygen:  and 

1. Reducing character of dioxides decreases down the group because oxygen has a strong positive field which attracts the hydroxyl group and removal of becomes easy.
2. Acidity also decreases down the group.
3. is a gas whereas SeO2 is solid. This is because  has a chain polymeric structure whereas SO2 forms discrete units.

• Reactivity with halogens: EX2, EX4 and EX6

1. The stability of halides decreases in the order 
2. This is because E-X bond length increases with increase in size.
3. Among hexa halides, fluorides are the most stable because of steric reasons.
4. Dihalides are  hybridised and so, are tetrahedral in shape.
5. Hexafluorides are only stable halides which are gaseous and have  hybridisation and octahedral structure.
6.  is a liquid while H2S is a gas. This is because strong hydrogen bonding is present in water. This is due to small size and high electronegativity of O.

Oxygen:

The compounds of oxygen and other elements are called oxides.

• Oxides: The compounds of oxygen and other elements are called oxides.
• Types of oxides:

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1. Acidic oxides: Non- metallic oxides are usually acidic in nature.
2. Basic oxides: Metallic oxides are mostly basic in nature. Basic oxides dissolve in water forming bases e.g.,
3. Amphoteric oxides: They show characteristics of both acidic as well as basic oxides. 
   
4. Neutral oxides: These oxides are neither acidic nor basic. Example: Co, NO and N2O

• Ozone:

1. Preparation: It is prepared by passing silent electric discharge through pure and dry oxygen 10 – 15 % oxygen is converted to ozone.
   
2. Structure of Ozone: Ozone has angular structure. Both O = O bonds are of equal bond length due to resonance.

• Sulphur:

1. Sulphur exhibits allotropy:
   1. Yellow Rhombic ( – sulphur)
   2. Monoclinic (- sulphur)
2. 
3. At 369 K both forms are stable. It is called transition temperature.
4. Both of them have S8 molecules.
5. The ring is puckered and has a crown shape.
6. Another allotrope of sulphur – cyclo  ring adopts a chair form.
7. S2is formed at high temperature (1000 K).
8. It is paramagnetic because of 2 unpaired electrons present in anti bonding* orbitals like O₂.

Sulphuric acid:

By contact process

1. Preparation:
2. Exothermic reaction and therefore low temperature and high pressure are favourable.
3. It is dibasic acid or diprotic acid.
4. It is a strong dehydrating agent.
5. It is a moderately strong oxidizing agent.

GROUP 17 ELEMENTS

• Atomic and ionic radii: Halogens have the smallest atomic radii in their respective periods because of maximum effective nuclear charge.
• Ionisation enthalpy: They have very high ionization enthalpy because of small size as compared to other groups.
• Electron gain enthalpy:

1. Halogens have maximum negative electron gain enthalpy because these elements have only one electron less than stable noble gas configuration.
2. Electron gain enthalpy becomes less negative down the group because atomic size increases down the group.

• Electronegativity:

1. These elements are highly electronegative and electronegativity decreases down the group.
2. They have high effective nuclear charge.

• Bond dissociation enthalpy:

1. Bond dissociation enthalpy follows the order: 
2. This is because as the size increases bond length increases.
3. Bond dissociation enthalpy of Cl2 is more than F2 because there are large electronic repulsions of lone pairs present in F2.

• Colour: All halogens are coloured because of absorption of radiations in visible region which results in the excitation of outer electrons to higher energy levels.
• Oxidising power:

1. All halogens are strong oxidisingagents because they have a strong tendency to accept electrons.
2. Order of oxidizing power is: 

• Reactivity with Hydrogen:

1. Acidic strength: HF <HCl<HBr< HI
2. Stability: HF >HCl>HBr> HI. This is because of decrease in bond dissociation enthalpy.
3. Boiling point: HCl<HBr< HI < HF. HF has strong intermolecular H bonding. As the size increases van der Waals forces increases and hence boiling point increases.
4. % Ionic character: HF >HCl>HBr> HI Dipole moment: HF >HCl>HBr> HI. Electronegativity decreases down the group.
5. Reducing power: HF <HCl<HBr< HI

• Reactivity with metals:

1. Halogens react with metals to form halides.
2. Ionic character: MF >MCl>MBr> MI. The halides in higher oxidation state will be more covalent than the one in the lower oxidation state.

• Interhalogen compounds:

Reactivity of halogens towards other halogens:

1. Binary compounds of two different halogen atoms of general formula X  are called interhalogen compounds where n = 1, 3, 5, or 7. All these are covalent compounds.
2. Interhalogen compounds are more reactive than halogens because X-X is a more polar bond than X-X bond.
3. All are diamagnetic.
4. Their melting point is little higher than halogens.
5. XX’ (CIF, BrF, BrCl, ICl, IBr, IF) (Linear shape) XX’₃ () (Bent T- shape) XX’₅ –, (square pyramidal shape) XX’₇ – (Pentagonal bipyramidal shape)

• Oxoacids of halogens:

1. Fluorine forms only one oxoacid HOF (Fluoric (I) acid or hypofluorous acid) due to high electronegativity.
2. Acid strength: 
3. Reason:
4. Acid strength: HOF >HOCl>HOBr> HOI. This is because Fluorine is most electronegative.

GROUP 18 ELEMENTS:

• Ionisation enthalpy:

1. They have very high ionization enthalpy because of completely filled orbitals.
2. Ionisation enthalpy decreases down the group because of increase in size.

• Atomic radii: Increases down the group because the number of shells increases down the group.

• Electron gain enthalpy: They have large electron gain enthalpy because of stable electronic configuration.
• Melting and boiling point: It has low melting and boiling point due to the presence of only weak dispersion forces.
• Shapes:

 is linear,  is square planar and  is distorted octahedral.  is known but no true compound of He Ne and Arare known.

• Compounds of Xe and F:
• 
• 
• 
• 

 and  are powerful fluorinating agents.

• Compounds of Xe and O:

General Principles and Processes of Isolation of Elements class 12 Notes Chemistry

• Minerals: The naturally occurring chemical substances in the earth’s crust which are obtained by mining are known as minerals.
• Metals may or may not be extracted profitably from them.
• Ores: The rocky materials which contain sufficient quantity of mineral so that the metal can be extracted profitably or economically are known as ores.
• Gangue: The earthy or undesirable materials present in ore are known as gangue.
• Metallurgy: The entire scientific and technological process used for isolation of the metal from its ores is known as metallurgy.
• Chief Ores and Methods of Extraction of Some Common Metals:

Sodium metal

1. Occurrence: Rock salt (NaCl), Feldspar ()
2. Extraction method: Electrolysis of fused NaCl or NaCl/ 
3. Inference: Sodium is highly reactive and hence, it reacts with water.

Copper metal

1. Occurrence: Copper pyrites (), Malachite (), Cuprite () Copper glance ()
2. Extraction method: Roasting of sulphide partially and reduction.
3. Inference: It is self-reduction in a specially designed converter. Sulphuric acid leaching is also employed.

Aluminum metal

1. Occurrence: Bauxite:( where 0<x<1), Cryolite (),Kaolinite ()
2. Extraction method: Electrolysis of Al2O3 dissolved in molten cryolite or in Na3AlCl6
3. Inference: A good source of electricity is needed in the extraction of Al

Zinc metal

1. Occurrence: Zinc blende or Sphalerite (ZnS), Zincite (ZnO), Calamine ()
2. Extraction method: Roasting and then reduction with carbon.
3. Inference: The metal may be purified by fractional distillation.

Lead metal

1. Occurrence: Galena (PbS)
2. Extraction: Roasting of the sulphide ore and then reduction of the oxide.
3. Inference: Sulphide ore is concentrated by froth floatation process.

Silver metal

1. Occurrence: Argentite ()
2. Extraction method: Sodium cyanide leaching of the sulphide ore and finally replacement of Ag by Zn.
3. Inference: It involves complex formation and displacement.

Gold metal

1. Occurrence: Native, small amounts in many ores such as those of copper and silver
2. Extraction method: Cyanide leaching, same as in case of silver
3. Inference: Gold reacts with cyanide to form complex

Iron metal

1. Occurrence: Haematite (), Magnetite (), Siderite (),Iron pyrites ()
2. Extraction method: Reduction with the help of CO and coke in blast furnace.
3. Inference: Limestone is added as flux which removes  as calcium silicate (slag) floats over molten iron and prevents its oxidation. Temperatures approaching 2170 K is required.

• Steps of metallurgy:
  1. Concentration of ore
  2. Conversion of concentrated ore to oxide
  3. Reduction of oxide to metal
  4. Refining of metal
• Concentration of ore: The process of removal unwanted materials like sand, clay, rocks etc from the ore is known as concentration, ore – dressing or benefaction. It involves several steps which depend upon physical properties of metal compound and impurity (gangue). The type of metal, available facilities and environmental factors are also taken into consideration.
• Hydraulic washing (or gravity separation): It is based on difference in densities of ore and gangue particles. Ore is washed with a stream of water under pressure so that lighter impurities are washed away whereas heavy ores are left behind.
• Magnetic separation: This method is based on the difference in magnetic and non – magnetic properties of two components of ore (pure and impure). This method is used to remove tungsten ore particles from cassiterite (). It is also used to concentrate magnetite (), chromite () and pyrolusite () from unwanted gangue.
• Froth floatation process: It is based on the principle that sulphide ores are preferentially wetted by the pine oil or fatty acids or xanthates etc., whereas the gangue particles are wetted by the water.

Collectors are added to enhance the non-wettability of the mineral particles.
Froth stabilizers such as cresols, aniline etc., are added to stabilize the froth.
If two sulphide ores are present, it is possible to separate the two sulphide ores by adjusting proportion of oil to water or by adding depressants.
For example, in the case of an ore containing ZnS and PbS, the depressant used is NaCN. It selectively prevents ZnS from coming to froth but allows PbS to come with the froth.

• Leaching (Chemical separation): It is a process in which ore is treated with suitable solvent which dissolves the ore but not the impurities.
• Purification of Bauxite by leaching ( Baeyer’s process):

a) Step 1: 
b) Step 2: 
c) Step 3: 
d) Concentration of Gold and Silver Ores by Leaching:


Where M=AgorAu

• Conversion of ore into oxide: It is easier to reduce oxide than sulphide or carbonate ore. Therefore, the given ore should be converted into oxide by any one of the following method:

1. roasting
2. calcination

• Roasting:

1. It is a process in which ore is heated in a regular supply of air at a temperature below melting point of the metal so as to convert the given ore into oxide ore.
2. Sulphide ores are converted into oxide by roasting
3. It is also used to remove impurities as volatile oxides
4. example – 

• Calcination
  1. It is a process of heating ore in limited supply of air so as to convert carbonate ores into oxides.
  2. Carbonate ores are converted into oxide by roasting
  3. It is also used to remove moisture and volatile impurities.
  4. Example – 
• Reduction of oxide to metal: The process of converting metal oxide into metal is called reduction. It needs a suitable reducing agent depending upon the reactivity or reducing power of metal. The common reducing agents used are carbon or carbon monoxide or any other metals like Al, Mg etc.
• Thermodynamic principles of metallurgy: Some basic concepts of thermodynamics help in understanding the conditions of temperature and selecting suitable reducing agent in metallurgical processes:

1. Gibbs free energy change at any temperature is given by ΔG = ΔH – TΔS where ΔG is free energy change, ΔH is enthalpy change and ΔS is entropy change.
2. The relationship between and K is = –2.303 RT log K where K is equilibrium constant. R = 8.314 JK^-1 mol^-1, T is temperature in Kelvin.
3. A negative   means +ve value of K i.e., products are formed more than the reactants. The reaction will proceed in forward direction.
4. If ΔS is +ve, on increasing temperature the value of  increases so that  >  and  will become negative.

• Coupled reactions: If reactants and products of two reactions are put together in a system and the net ΔG of two possible reactions is –ve the overall reaction will take place. These reactions are called coupled reactions.
• Ellingham diagrams: The plots betweenof formation of oxides of elements vs. temperature are called Ellingham diagrams. It provides a sound idea about selecting a reducing agent in reduction of oxides.Such diagrams help in predicting the feasibility of a thermal reduction of an ore.  must be negative at a given temperature for a reaction to be feasible.
• Limitations of Ellingham Diagrams: It does not take kinetics of reduction into consideration, i.e., how fast reduction will take place cannot be determined.
• Reduction of iron oxide in blast furnace: Reduction of oxides takes place in different zones.
  
  

1. At 500 – 800 K (lower temperature range in blast furnace)
2. At 900 – 1500 K (higher temperature range in blast furnace)
3. Limestone decomposes to CaO and CO2
4. Silica (impurity) reacts with CaO to form calcium silicate which forms slag. It floats over molten iron and prevents oxidation of iron.
   

• Types of iron:
  1. Pig iron: The iron obtained from blast furnace is called pig iron. It is impure from of iron contains 4% carbon and small amount of S,.P, Si and Mn. It can be casted into variety of shapes.
  2. Cast iron: It is made by melting pig iron with scrap iron and coke using hot air blast. It contains about 3% of carbon content. It is extremely hard and brittle.
  3. Wrought iron: It is the purest form of commercial iron. It is also called malleable iron. It is prepared by oxidative refining of pig iron in reverberatory furnace lined with haematite which oxidises carbon to carbon monoxide.
     

The substance which reacts with impurity to form slag is called flux e.g. limestone is flux.





The metal is removed and freed from slag by passing through rollers.

• Electrolytic Reduction (Hall – Heroult Process): Purified bauxite ore is mixed with cryolite () or CaF2 which lowers its melting point and increases electrical conductivity. Molten mixture is electrolysed using a number of graphite rods as anode and carbon lining as cathode.

The graphite anode is useful for reduction of metal oxide to metal.


At cathode: 
At anode: 

Graphite rods get burnt forming CO and CO2. The aluminium thus obtained is refined electrolytically using impure Al as anode, pure Al as cathode and molten cryolite as electrolyte.
At anode: 
At cathode: 

• Electrolysis of molten NaCl:
  
  At cathode: 
  At anode: 

Thus sodium metal is obtained at cathode and  (g) is liberated at anode.

• Refining: It is the process of converting an impure metal into pure metal depending upon the nature of metal.
• Distillation: It is the process used to purify those metals which have low boiling points, e.g., zinc, mercury, sodium, potassium. Impure metal is heated so as to convert it into vapours which changes into pure metal on condensation and is obtained as distillate.
• Liquation: Those metals which have impurities whose melting points are higher than metal can be purified by this method. In this method, Sn metal can be purified. Tin containing iron as impurities heated on the top of sloping furnace. Tin melts and flows down the sloping surface where iron is left behind and pure tin is obtained.

• Electrolytic refining: In this method, impure metal is taken as anode, pure metal is taken as cathode, and a soluble salt of metal is used as electrolyte. When electric current is passed, impure metal forms metal ions which are discharged at cathode forming pure metal.
  At anode:
  At cathode: 
• Zone refining: It is based on the principle that impurities are more soluble in the melt than in the solid state of the metal. The impure metal is heated with the help of circular heaters at one end of the rod of impure metal. The molten zone moves forward along with the heater with impurities and reaches the other end and is discarded. Pure metal crystallizes out of the melt. The process is repeated several times and heater is moved in the same direction. It is used for purifying semiconductors like B, Ge, Si, Ga and In.
• Vapour phase refining: Nickel is purified by Mond’s process. Nickel, when heated in stream of carbon monoxide forms volatile Ni(CO)4 which on further subjecting to higher temperature decomposes to give pure metal.
  
• Van-Arkel method: It is used to get ultra pure metals. Zr and Ti are purified by this process. Zr or Ti are heated in iodine vapours at about 870 K to form volatile ZrI4 or TiI4 which are heated over tungsten filament at 1800K to give pure Zr or Ti.
  
  
• Chromatographic method: It is based on the principle of separation or purification by chromatography which is based on differential adsorption on an adsorbent. In column chromatography,  is used as adsorbent. The mixture to be separated is taken in suitable solvent and applied on the column. They are then eluted out with suitable solvent (eluent). The weakly adsorbed component is eluted first. This method is suitable for such elements which are available only in minute quantities and the impurities are not very much different in their chemical behaviour from the element to be purified.

Class 12 Chemistry Quick Revision Notes Chapter 5 Surface 

• Adsorption:
  (i) The accumulation of molecular species at the surface rather than in the bulk of a solid or liquid is termed as adsorption.
  (ii) It is a surface phenomenon.
  (iii) The concentration of adsorbate increases only at the surface of the adsorbent.
• Adsorbate: It is the substance which is being adsorbed on the surface of another substance.
• Adsorbent: It is the substance present in bulk, on the surface of which adsorption is taking place.
• Desorption: It is the process of removing an adsorbed substance from a surface on which it is adsorbed.
• Absorption:
  (i) It is the phenomenon in which a substance is uniformly distributed throughout the bulk of the solid.
  (ii) It is a bulk phenomenon.
  (iii) The concentration is uniform throughout the bulk of solid.
• Sorption: When adsorption and absorption take place simultaneously, it is called sorption.
• Enthalpy or heat of adsorption: Since, adsorption occurs with release in energy, i.e., it is exothermic in nature. The enthalpy change for the adsorption of one mole of an adsorbate on the surface of adsorbent is called enthalpy or heat of adsorption.
• Types of adsorption: There are different types of adsorption namely,

1. Physical adsorption
2. Chemical adsorption

• Physical adsorption
  (i) If the adsorbate is held on a surface of adsorbent by weak van der Waals’ forces, the adsorption is called physical adsorption or physisorption.
  (ii) It is non-specific.
  (iii) It is reversible.
  (iv) The amount of gas depends upon nature of gas, i.e., easily liquefiable gases like NH₃, CO₂, gas adsorbed to greater extent than H₂ and He. Higher the critical temperature of gas, more will be the extent of adsorption.
  (v) The extent of adsorption increases with increase in surface area, e.g. porous and finely divided metals are good adsorbents.
  (vi) There are weak van der Waals’ forces of attraction between adsorbate and adsorbent.
  (vii) It has low enthalpy of adsorption (20 – 40 kJ mol-1).
  (viii) Low temperature is favourable.
  (ix) No appreciable activation energy is needed.
  (x) It forms multimolecular layers.
• Chemical adsorption or chemisorption:

(i) If the forces holding the adsorbate are as strong as in chemical bonds, the adsorption process is known as chemical adsorption of chemisorption.
(ii) It is highly specific.
(iii) It is irreversible.
(iv) The amount of gas adsorbed is not related to critical temperature of the gas.
(v) It also increases with increase in surface area.
(vi) There is strong force of attraction similar to chemical bond.
(vii) It has enthalpy heat of adsorption (180 – 240 kJ mol-1).
(viii) High temperature is favourable.
(ix) High activation energy is sometimes needed.
(x) It forms unimolecular layers.

• Factors affecting adsorption of gases on solids:

a. Nature of adsorbate: Physical adsorption is non-specific in nature and therefore every gas gets adsorbed on the surface of any solid to a lesser or greater extent. However, easily liquefiable gases like NH₃,HCl, CO₂, etc. which have higher critical temperatures are absorbed to greater extent whereas H₂, O₂, N₂ etc. are adsorbed to lesser extent. The chemical adsorption being highly specific, therefore, a gas gets adsorbed on specific solid only if it enters into chemical combination with it.
b. Nature of adsorbent: Activated carbon, metal oxides like aluminum oxide, silica gel and clay are commonly used adsorbents. They have their specific adsorption properties depending upon pores.
c. Specific area of the adsorbent: The greater the specific area, more will be the extent of adsorption. That is why porous or finely divided forms of adsorbents adsorb larger quantities of adsorbate. The pores should be large enough to allow the gas molecules to enter.
d. Pressure of the gas: Physical adsorption increases with increase in pressure.

• Adsorption isotherm:
• The variation in the amount of gas adsorbed by the adsorbent with pressure at constant temperature can be expressed by means of a curve is termed as adsorption isotherm.
• Freundlich Adsorption isotherm: The relationship between  and pressure of the gas at constant temperature is called adsorption isotherm and is given by fracxm=kp1/n(n>1)

Where x- mass of the gas adsorbed on mass m of the adsorbent and the gas at a particular temperature k and n depends upon the nature of gas

• The solidfirst increases with increase in pressure at low pressure but becomes independent of pressure at high pressure.

Taking logarithm on both sides, we get,

• If we plot a graph between logand log P, we get a straight line.

The slope of the line isand intercept will be equal to log k.

• Catalyst: These are substances which alter the rate of a chemical reaction and themselves remain chemically and quantitatively unchanged after the reactionand the phenomenon is known as catalysis.

• Promoters: These are the substances which increase the activity of catalyst. Example – Mo is promoter whereas Fe is catalyst in Haber’s Process.
  
• Catalytic poisons (Inhibitors): These are the substances which decrease the activity of catalyst. Example -Arsenic acts as catalytic poison in the manufacture of sulphuric acid by ‘contact process.’
• Types of catalysis:

There are two types of catalysis namely,

1. Homogeneous catalysis: When the catalyst and the reactants are in the same phase, this kind of catalytic process is known as homogeneous catalysis.
2. Heterogeneous catalysis: When the catalyst and the reactants are in different phases, the catalytic process is said to be heterogeneous catalysis.
3. Activity of catalyst: It is the ability of a catalyst to increase the rate of a chemical reaction.
4. Selectivity of catalyst: It is the ability of catalyst to direct a reaction to yield a particular product (excluding others).

For example: CO and H₂ react to form different products in presence of different catalysts as follows:

1. 
2. 
3. 

• Shape – selective catalysis: It is the catalysis which depends upon the pore structure of the catalyst and molecular size of reactant and product molecules. Example – Zeolites are shape – selective catalysts due to their honey- comb structure.
• Enzymes: These are complex nitrogenous organic compounds which are produced by living plants and animals. They are actually protein molecules of high molecular mass. They are biochemical catalysts
• Steps of enzyme catalysis:

(i) Binding of enzyme to substrate to form an activated complex.
(ii) Decomposition of the activated complex to form product.

• Characteristics of enzyme catalysis:

(i) They are highly efficient. One molecule of an enzyme can transform 106 molecules of reactants per minute.
(ii) They are highly specific in nature. Example – Urease catalysis hydrolysis of urea only.
(iii) They are active at optimum temperature (298 – 310 K). The rate of enzyme catalysed reaction becomes maximum at a definite temperature called the optimum temperature.
(iv) They are highly active at a specific pH called optimum pH.
(v) Enzymatic activity can be increased in presence of coenzymes which can be called as promoters.
(vi) Activators are generally metal ions Na⁺, Co²⁺ and Cu²⁺ etc. They weakly bind to enzyme and increase its activity.
(vii) Influence of inhibitors (poison): Enzymes can also be inhibited or poisoned by the presence of certain substances.

• True solution:

(i) It is homogeneous.
(ii) The diameter of the particles is less than 1 nm.
(iii) It passes through filter paper.
(iv) Its particles cannot be seen under a microscope.

• Colloids:

(i) It appears to be homogeneous but is actually heterogeneous.
(ii) The diameter of the particles is 1 nm to 1000 nm.
(iii) It passes through ordinary filter paper but not through ultra-filters.
(iv) Its particles can be seen by a powerful microscope due to scattering of light.

• Suspension:

(i) It is heterogeneous.
(ii) The diameter of the particles are larger than 1000 nm.
(iii) It does not pass through filter paper.
(iv) Its particles can be seen even with naked eye.

• Dispersed phase: It is the substance which is dispersed as very fine particles.
• Dispersion medium: It is the substance present in larger quantity.
• Classification of colloids on the basis of the physical state of dispersed phase and dispersion medium:

Name	Dispersed phase	Dispersed medium	Examples
Solid sol	solid	Solid	Coloured gem stones
Sol	Solid	Liquid	Paints
Aerosol	Solid	Gas	Smoke, dust
Gel	Liquid	Solid	Cheese, jellies
Emulsion	Liquid	Liquid	Hair cream, milk
Aerosol	Liquid	Gas	Mist, fog, cloud
Solid sol	Gas	Solid	Foam rubber, pumice stone
Foam	Gas	Liquid	Whipped cream

• Classification of colloids on the basis of nature of interaction between dispersed phase and dispersion medium, the colloids are classified into two types namely,

1. Lyophobic sols
2. Lyophilic sols

• Lyophobic sols:

(i) These colloids are liquid hating.
(ii) In these colloids the particles of dispersed phase have no affinity for the dispersion medium.
(iii) They are not stable.
(iv) They can be prepared by mixing substances directly.
(v) They need stabilizing agents for their preservation.
(vi) They are irreversible sols.

• Lyophilic sols:

(i) These colloids are liquid loving.
(ii) In these colloids, the particles of dispersed phase have great affinity for the dispersion medium.
(iii) They are stable.
(iv) They cannot be prepared by mixing substances directly. They are prepared only by special methods.
(v) They do not need stabilizing agents for their preservation.
(vi) They are reversible sols.

• Classification of colloids on the basis of types of particles of the dispersed phase:

There are three types of colloids based on the type of dispersed phase, namely,

1. Multimolecular colloids: The colloids in which the colloidal particles consist of aggregates of atoms or small molecules. The diameter of the colloidal particle formed is less than 1 nm.
2. Macromolecular colloids: These are the colloids in which the dispersed particles are themselves large molecules (usually polymers). Since these molecules have dimensions comparable to those of colloids particles, their dispersions are called macromolecular colloids, e.g., proteins, starch and cellulose form macromolecular colloids.
3. Associated colloids (Micelles): Those colloids which behave as normal, strong electrolytes at low concentrations, but show colloidal properties at higherconcentrations due to the formation of aggregated particles of colloidal dimensions. Such substances are also referred to as associated colloids.

• Kraft Temperature (Tk):Micelles are formed only above a certain temperature called Kraft temperature.
• Critical Micelle Concentration (CMC): Micelles are formed only above a particular concentration called critical micelle concentration.
• Soaps: These are are sodium or potassium salts of higher fatty acids e.g., sodium stearate CH₃(CH₂)₁₆COO-Na⁺
• Methods of preparation of colloids:

1. Chemical methods: Colloids can be prepared by chemical reactions leading to the formation of molecules. These molecules aggregate leading to formation of sols.
2. Electrical disintegration or Bredig’s Arc method: In this method, electric arc is struck between electrodes of the metal immersed in the dispersion medium. The intense heat produced vaporizes the metal which then condenses to form particles of colloidal size.
3. Peptization: It is the process of converting a precipitate into colloidal sol by shaking it with dispersion medium in the presence of a small amount of electrolyte. The electrolyte used for this purpose is called peptizing agent.

• Purification of colloids:

1. Dialysis: It is a process of removing a dissolved substance from a colloidal solution by means of diffusion through a suitable membrane.
2. Electro dialysis. The process of dialysis is quite slow. It can be made faster by applying an electric field if the dissolved substance in the impure colloidal solution is only an electrolyte.
3. Ultrafiltration: It is the process of separating the colloidal particles from the solvent and soluble solutes present in the colloidal solution by specially prepared filters, which are permeable to all substances except the colloidal particles.
4. Ultracentrifugation: In this process, the colloidal solution is taken in a tube which is placed in ultracentrifuge. On rotating the tube at very high speed, the colloidal particles settle down at the bottom of the tube and the impurities remain in solution. The settled particles are mixed with dispersion medium to regenerate the sol.

• Properties of colloids:Positively charged colloidal particles:(i) These include hydrated metallic oxides such as Fe₂O₃.H₂O, Cr₂O₃.H₂O, Al₂O₃.H₂O
  (ii) Basic dye stuff like malachite green, methylene blue sols.
  (iii) Example – Haemoglobin (blood).Negatively charged colloidal particles:(i) Metallic sulphides like As₂S₃, Sb₂S₃ sols.
  (ii) Acid dye stuff like eosin, methyl orange, Congo red sols.
  (iii) Examples – Starch sol, gum, gelatin, clay, charcoal, egg albumin, etc.
  1. Colour: The colour of colloidal solution depends upon the wavelength of light scattered by the colloidal particles which in turn depends upon the nature and size of particles. The colour also depends upon the manner in which light is received by the observer.
  2. Brownian movement: Colloidal particles move in zig – zag path. This type of motion is due to colliding molecules of dispersion medium constantly with colloidal particles.
  3. Colligative properties: The values of colligative properties (osmotic pressure, lowering in vapour pressure, depression in freezing point and elevation in boiling point) are of small order as compared to values shown by true solutions at the same concentrations.
  4. Tyndall effect: The scattering of a beam of light by colloidal particles is called Tyndall effect. The bright cone of light is called the Tyndall cone.
  5. Charge on colloidal particles: Colloidal particles always carry an electric charge. The nature of this charge is the same on all the particles in a given colloidal solution and may be either positive or negative.
  6. Helmholtz electrical double layer: When the colloidal particles acquire negative or positive charge by selective adsorption of one of the ions, it attracts counter ions from the medium forming a second layer. The combination of these two layers of opposite charges around colloidal particles is called Helmholtz electrical double layer.
  7. Electrokinetic potential or zeta potential: The potential difference between the fixed layer and the diffused layer of opposite charges is called electrokinetic potential or zeta potential.
  8. Electrophoresis: The movement of colloidal particles under an applied electric potential is called electrophoresis.
  9. Coagulation or precipitation: The process of settling of colloidal particles as precipitate is called coagulation.
• Hardy – Schulze rules:

i) Oppositely charged ions are effective for coagulation.
ii) The coagulating power of electrolyte increases with increase in charge on the ions used for coagulation. Examples – Al³⁺> Ba²⁺> Na⁺ for negatively charged colloids. Fe (CN)₆]^4->>>Cl^– for positively charged colloids.

• Types of emulsions:
  1. Water dispersed in oil: When water is the dispersed phase and oil is the dispersion medium. E.g. butter
  2. Oil dispersed in water: When oil is the dispersed phase and water is the dispersion medium. E.g. milk
• Emulsification: It is the process of stabilizing an emulsion by means of an emulsifier.
• Emulsifying agent or emulsifier: These are the substances which are added to stabilize the emulsions. Examples – soaps, gum
• Demulsification: It is the process of breaking an emulsion into its constituent liquidsby freezing, boiling, centrifugation or some chemical methods.

Class 12 Chemistry Revision Notes Chapter 4 Chemical Kinetics

• Chemical kinetics: It is the branch of chemistry that deals with the study of reaction rates and their mechanisms.
• Rate of reaction: It is the change in concentration of reactant (or product) in unit time.
• The unit of rate of reaction is mol L-1s-1.
• A + B → C + D
  Rate of disappearance of 
  where d[A] is small change in conc. of ‘A’ and dt is small interval of time
  Rate of disappearance of 
  Where d[B] is small change in conc. of ‘B’ and dt is small interval of time
  Rate of appearance of 
  Where d[C] is small change in conc. of ‘C’ and dt is small interval of time
  Rate of appearance of 
  Where d[D] is small change in conc. of ‘D’ and dt is small interval of time
  Rate 
• Rate law or rate equation: It is the expression which relates the rate of reaction with concentration of the reactants. The constant of proportionality ‘k’ is known as rate constant.
• Average rate: It is the rate of reaction measured over a long time interval.
• Average rate 
  where is Δx change in concentration and Δt is large interval of time.
• Instantaneous rate: It is the rate of reaction when the average rate is taken over a particular moment of time. Instantaneous rate .
  where dx is small change in conc. and dt is the smallest interval of time.
  It is the expression which relates the rate of reaction with concentration of the reactants.
• Rate constant: When the concentration of reactants is unity, then the rate of reaction is known as rate constant. It is also called specific reaction rate.
  The constant of proportionality ‘k’ is known as rate constant.
• Molecularity of a reaction: The total number of atoms, ions or molecules of the reactants involved in the reaction is termed as its molecularity. It is always in whole number and is never more than three. It cannot be zero.
• Order of a reaction: The sum of the exponents (power) of the concentration of reactants in the rate law is termed as order of the reaction. It can be in fraction. It can be zero also.
  If rate law expression for a reaction is Rate = k [A]x [B]y
  Then its order of reaction = x + y
• Order cannot be determined with a given balanced chemical equation. It can be experimentally determined.
• Integrated rate law for zero order reaction: R → P
  
  
  If we plot a graph between concentration of R vs time t, the graph is a straight line with slope equal to -k and intercept is equal to [Ro].
• Half- life of a reaction: The time taken for a reaction, when half of the starting material has reacted is called half- life of a reaction.
• For zero order reaction, the half-life time is 
• For first order reaction, the half-life time is , where ‘k’ is rate constant.
  It is independent of initial concentration for first order reaction.
• Rate law for first order reaction: RP
  
  where ‘k’ is rate constant or specific reaction rate, [Ro] is initial molar conc., [R] is final molar conc. after time ‘t’.
  
  where ‘a’ is initial conc. reacted in time ‘t’ final conc., after time ‘t’ is (a – x).
• If we plot a graph between ln[R] with time, we get a straight line whose slope = – k and intercept ln[Ro].
• To calculate rate constant for first order gas phase reaction of the type
  A(g)B(g) + C(g)
  
  Where pi is initial pressure of A, pt is total pressure of gaseous mixture containing A , B, C
  t = A + B + C
• Pseudo first order reaction: The reaction which is bimolecular but order is one is called pseudo first order reaction. This happens when one of the reactants is in large excess.Example – Acidic hydrolysis of ester (ethyl acetate).
  
• Activation energy (Ea): It is extra energy which must be possessed by reactant molecules so that collision between reactant molecules is effective and leads to the formation of product molecules.
• Arrhenius equation of reaction rate: It gives the relation between rate of reaction and temperature.
  
  
  
  where k = rate constant, A = frequency factor, Ea = energy of activation R = gas constant, T = temperature in Kelvin,
  
• Probability factor or Steric factor
  
  Where ZAB represents the collision frequency of reactants, A and B, represents the fraction of molecules with energies equal to or greater than Ea and P is called the probability or steric factor.
• Mechanism of reaction: It is the sequence of elementary processes leading to the overall stoichiometry of a chemical reaction.
• Activated complex: It is an unstable intermediate formed between reacting molecules. Since, it is highly unstable and it readily changes into product.
• Rate determining step: It is the slowest step in the reaction mechanism.
• The number of collisions per second per unit volume of the reaction mixture is known as collision frequency (Z).

Class 12 Chemistry Quick Revision Notes Chapter 3 Electrochemistry

• Oxidation: It is defined as a loss of electrons while reduction is defined as a gain of electrons.
  In a redox reaction, both oxidation and reduction reaction takes place simultaneously.
• Direct redox reaction: In a direct redox reaction, both oxidation and reduction reactions take place in the same vessel. Chemical energy is converted to heat energy in a direct redox reaction.
• Indirect redox reaction: In indirect redox reactions, oxidation and reduction take place in different vessels.
  In an indirect redox reaction, chemical energy is converted into electrical energy. The device which converts chemical energy into electrical energy is known as an electrochemical cell.
• In an electrochemical cell:
  1. The half-cell in which oxidation takes place is known as oxidation half-cell
  2. The half-cell in which reduction takes place is known as reduction half-cell.
  3. Oxidation takes place at anode which is negatively charged and reduction takes place at cathode which is positively charged.
  4. Transfer of electrons takes place from anode to cathode while electric current flows in the opposite direction.
  5. An electrode is made by dipping the metal plate into the electrolytic solution of its soluble salt.
  6. A salt bridge is a U shaped tube containing an inert electrolyte in agar-agar and gelatine.
• Salt bridge: A salt bridge maintains electrical neutrality and allows the flow of electric current by completing the electrical circuit.
• Representation of an electrochemical cell:
  1. Anode is written on the left while the cathode is written on the right.
  2. Anode represents the oxidation half-cell and is written as: Metal/Metal ion (Concentration)
  3. Cathode represents the reduction half-cell and is written as: Metal ion (Concentration)/Metal
  4. Salt bridge is indicated by placing double vertical lines between the anode and the cathode
  5. Electrode potential is the potential difference that develops between the electrode and its electrolyte. The separation of charges at the equilibrium state results in the potential difference between the metal and the solution of its ions. It is the measure of tendency of an electrode in the half cell to lose or gain electrons.
  6. Standard electrode potential: When the concentration of all the species involved in a half cell is unity, then the electrode potential is known as standard electrode potential. It is denoted as EΘ.
• According to the present convention, standard reduction potentials are now called standard electrode potential.
• Types of electrode potential: There are 2 types of electrode potentials namely,
  1. Oxidation potential
  2. Reduction potential
• Oxidation potential: It is the tendency of an electrode to lose electrons or get oxidized.
• Reduction potential: It is the tendency of an electrode to gain electrons or get reduced.
  Oxidation potential is the reverse of reduction potential.
• The electrode having a higher reduction potential have higher tendency to gain electrons and so it acts as a cathode whereas the electrode having a lower reduction potential acts as an anode.
• The standard electrode potential of an electrode cannot be measured in isolation.
• According to convention, the Standard Hydrogen Electrode is taken as a reference electrode and it is assigned a zero potential at all temperatures.
• Reference electrode: Standard calomel electrode can also be used as a reference electrode
• SHE: Standard hydrogen electrode consists of a platinum wire sealed in a glass tube and carrying a platinum foil at one end. The electrode is placed in a beaker containing an aqueous solution of an acid having 1 Molar concentration of hydrogen ions. Hydrogen gas at 1 bar pressure is continuously bubbled through the solution at 298 K. The oxidation or reduction takes place at the Platinum foil. The standard hydrogen electrode can act as both anode and cathode.
• If the standard hydrogen electrode acts as an anode:
  
  If the standard hydrogen electrode acts as a cathode:
  
• In the electrochemical series, various elements are arranged as per their standard reduction potential values.
• A substance with higher reduction potential value means that it has a higher tendency to get reduced. So, it acts as a good oxidising agent.
• A substance with lower reduction potential value means that it has a higher tendency to get oxidised. So, it acts as a good reducing agent.
• The electrode with higher reduction potential acts as a cathode while the electrode with a lower reduction potential acts as an anode.
• The potential difference between the 2 electrodes of a galvanic cell is called cell potential and is measured in Volts.
• The cell potential is the difference between the reduction potential of cathode and anode.
  E cell = E cathode – E anode
• Cell potential is called the electromotive force of the cell (EMF) when no current is drawn through the cell.
• Nernst studied the variation of electrode potential of an electrode with temperature and concentration of electrolyte.
• Nernst formulated a relationship between standard electrode potential EΘ and electrode potential E.
  
  
• Electrode potential increases with increase in the concentration of the electrolyte and decrease in temperature.
• Nernst equation when applied to a cell, it helps in calculating the cell potential.
  
• At equilibrium, cell potential Ecteell becomes zero.
• Relationship between equilibrium constant Kc and standard cell potential EΘcell:
  
• Work done by an electrochemical cell is equal to the decrease in Gibbs energy
  
• The substances which allow the passage of electricity through them are known as conductors.
• Every conducting material offers some obstruction to the flow of electricity which is called resistance. It is denoted by R and is measured in ohm.
• The resistance of any object is directly proportional to its length l and inversely proportional to its area of cross section A.
  
  Where ρ is called specific resistance or resistivity.
• The SI unit of specific resistivity is ohm metre.
• The inverse of resistance is known as conductance, G
• Unit of conductance is ohm-1 or mho. It is also expressed in Siemens denoted by S.
• The inverse of resistivity is known as conductivity. It is represented by the symbol .
• The SI unit of conductivity is Sm-1. But it is also expressed in Scm-1.
• Conductivity = Conductance × Cell constant
• For measuring the resistance of an ionic solution, there are 2 problems:
  1. Firstly, passing direct current changes the composition of the solution
  2. Secondly, a solution cannot be connected to the bridge like a metallic wire or a solid conductor.
• Conductivity cell: The problem of measuring the resistance of an ionic solution can be resolved by using a source of alternating current and the second problem is resolved by using a specially designed vessel called conductivity cell.
• A conductivity cell consists of 2 Pt electrodes coated with Pt black. They have area of cross section A and are separated by a distance ‘l’. Resistance of such a column of solution is given by the equation:
• 
  Where  is called cell constant and is denoted by the symbol 
• Molar conductivity of a solution: It is defined as the conducting power of all the ions produced by dissolving 1 mole of an electrolyte in solution.
  Molar conductivity 
  Where = Conductivity and M is the molarity Unit of Molar conductivity is Scm2 mol-1
• Equivalent conductivity: It is the conductivity of all the ions produced by dissolving one gram equivalent of an electrolyte in solution. Unit of equivalent conductivity is S cm2 (g equiv) -1
  Equivalent conductivity: 
• Kohlrausch’s Law of independent migration of ions: According to this law, molar conductivity of an electrolyte, at infinite dilution, can be expressed as the sum of individual contributions from its individual ions.
• If the limiting molar conductivity of the cations is denoted by and that of the anions by , then the limiting molar conductivity of electrolyte is:
  Molar conductivity, 
  Where v+ and v- are the number of cations and anions per formula of electrolyte
• Degree of dissociation: It is ratio of molar conductivity at a specific concentration ‘c’ to the molar conductivity at infinite dilution. It is denoted by.
  
• Dissociation constant:  WhereKa is acid dissociation constant, ‘c’ is concentration of electrolyte, α is degree of ionization.
• Faraday constant: It is equal to charge on 1 mol of electrons. It is equal to 96487 C mol-1 or approximately equal to 96500 C mol-1.
• Faraday’s first law of electrolysis: The amount of substance deposited during electrolysis is directly proportional to quantity of electricity passed.
• Faraday’s second law of electrolysis: If same charge is passed through different electrolytes, the mass of substance deposited will be proportional to their equivalent weights.
• Products of electrolysis: The products of electrolysis depend upon
• The nature of electrolyte being electrolyzed and the nature of electrodes. If electrode is inert like platinum or gold, they do not take part in chemical reaction i.e. they neither lose nor gain electrons. If the electrodes are reactive then they will take part in chemical reaction and products will be different as compared to inert electrodes.
• The electrode potentials of oxidizing and reducing species. Some of the electrochemical processes although feasible but slow in their rates at lower voltage, these require extra voltage, i.e. over voltage at which these processes will take place. The products of electrolysis also differ in molten state and aqueous solution of electrolyte.
• Primary cells: A primary cell is a cell in which electrical energy is produced by the reaction occurring in the cell, e.g. Daniel cell, dry cell, mercury cell. It cannot be recharged.
• Dry Cell:
  At anode 
  At cathode 
  The net reaction: 
• Mercury Cell: The electrolyte is a paste of KOH and ZnO.
  At Anode: 
  At cathode: 
  The net reaction: 
• Secondary cells: Those cells which are used for storing electricity, e.g., lead storage battery, nickel – cadmium cell. They can be recharged.
• Lead storage battery:
  Anode: 
  Cathode: 
  The overall cell reaction consisting of cathode and anode reactions is:
  
  On recharging the battery, the reaction is reversed.
• Nickel cadmium cell: It is another type of secondary cell which has longer life than lead storage cell but more expensive to manufacture.
  The overall reaction during discharge is
  
• Fuel cells:
  At Anode: 
  At cathode: 
• Overall reaction:
  
• Corrosion:
  Oxidation: 
  Reduction: 
• Galvanization: It is a process of coating zinc over iron so as to protect it from rusting.
• Cathodic protection: Instead of coating more reactive metal on iron, the use of such metal is made as sacrificial anode.

Class 12 Chemistry Quick Revision Notes

The difference in boiling points of solution  and pure solvent  is called elevation in boiling point 

• Solutions: Solutions are the homogeneous mixtures of two or more than two components.
• Binary solution: A solution having two components is called a binary solution.
• Components of a binary solution.
  It includes solute and solvent.
  1. When the solvent is in solid state, solution is called solid solution.
  2. When the solvent is in liquid state, solution is called liquid solution.
  3. When the solvent is in gaseous state, solution is called gaseous solution.
• Concentration: It is the amount of solute in given amount of solution.
• Mass by volume percentage (w/v): Mass of the solute dissolved in 100 mL of solution.
• Molality (m) is the number of moles of solute present in 1kg of solvent.
• Molarity (M) is the number of moles of solute present in 1L of solution.
• Normality is the number of gram equivalent of solute dissolved per litre of solution.
• Solubility: It is the maximum amount that can be dissolved in a specified amount of solvent at a specified temperature.
• Saturated solution: It is a solution in which no more solute can be dissolved at the same temperature and pressure.
• In a nearly saturated solution if dissolution process is an endothermic process, solubility increases with increase in temperature.
• In a nearly saturated solution if dissolution process is an exothermic process, solubility decreases with increase in temperature.
• Henry’s Law: It states “at a constant temperature the solubility of gas in a liquid is directly proportional to the pressure of gas”. In other words, “the partial pressure of gas in vapour phase is proportional to the mole fraction of the gas in the solution”.
• When a non-volatile solute is dissolved in a volatile solvent, the vapour pressure of solution is less than that of pure solvent.
• Raoult’s law: It states that “for a solution of volatile liquids the partial vapour pressure of each component in the solution is directly proportional to its mole fraction”.
• Using Dalton’s law of partial pressure the total pressure of solution is calculated.
  
• Comparison of Raoult’ law and Henry’s law: It is observed that the partial pressure of volatile component or gas is directly proportional to its mole fraction in solution. In case of Henry’s Law the proportionality constant is KH and it is different from p10 which is partial pressure of pure component. Raoult’s Law becomes a special case of Henry’s Law when KH becomes equal to p10 in Henry’s law.
• Classification of liquid–liquid solutions: It can be classified into ideal and non-ideal solutions on basis of Raoult’s Law.
• Ideal solutions:
  1. The solutions that obey Raoult’s Law over the entire range of concentrations are known as ideal solutions.
  2. 
  3. The intermolecular attractive forces between solute molecules and solvent molecules are nearly equal to those present between solute and solvent molecules i.e. A-A and B-B interactions are nearly equal to those between A-B.
• Non-ideal solutions:
  1. When a solution does not obey Raoult’s Law over the entire range of concentration, then it is called non-ideal solution.
  2. 
  3. The intermolecular attractive forces between solute molecules and solvent molecules are not equal to those present between solute and solvent molecules i.e. A-A and B-B interactions are not equal to those between A-B
• Types of non- ideal solutions:
  1. Non ideal solution showing positive deviation
  2. Non ideal solution showing negative deviation
• Non ideal solution showing positive deviation
  1. The vapour pressure of a solution is higher than that predicted by Raoult’s Law.
  2. The intermolecular attractive forces between solute-solvent molecules are weaker than those between solute-solute and solvent-solvent molecules i.e., A-B < A-A and B-B interactions.
• Non ideal solution showing negative deviation
  1. The vapour pressure of a solution is lower than that predicted by Raoult’s Law.
  2. The intermolecular attractive forces between solute-solvent molecules are stronger than those between solute-solute and solvent-solvent molecules i.e. A-B > A-A and B-B interactions.
• Azeotopes: These are binary mixtures having same composition in liquid and vapour phase and boil at constant temperature. Liquids forming azeotrope cannot be separated by fractional distillation.
• Types of azeotropes: There are two types of azeotropes namely,
  1. Minimum boiling azeotrope
  2. Maximum boiling azeotrope
• The solutions which show a large positive deviation from Raoult’s law form minimum boiling azeotrope at a specific composition.
• The solutions that show large negative deviation from Raoult’s law form maximum boiling azeotrope at a specific composition.
• Colligative properties: The properties of solution which depends on only the number of solute particles but not on the nature of solute are called colligative properties.
• Types of colligative properties: There are four colligative properties namely,
  1. Relative lowering of vapour pressure
  2. Elevation of boiling point
  3. Depression of freezing point
  4. Osmotic pressure
• Relative lowering of vapour pressure: The difference in the vapour pressure of pure solvent  and solution  represents lowering in vapour pressure.
• Relative lowering of vapour pressure: Dividing lowering in vapour pressure by vapour pressure of pure solvent is called relative lowering of vapour pressure 
• Relative lowering of vapour pressure is directly proportional to mole fraction of solute. Hence it is a colligative property.
• Elevation of boiling point: 
• For a dilute solution elevation of boiling point is directly proportional to molal concentration of the solute in solution. Hence it is a colligative property.
• Depression of freezing point: The lowering of vapour pressure ofsolution causes a lowering of freezing point compared to that of pure solvent.The difference in freezing point of the pure solvent  and solution is called the depression in freezing point.
• For a dilute solution depression in freezing point is a colligative property because it is directly proportional to molal concentration of solute.
• Osmosis: The phenomenon of flow of solvent molecules through a semi permeable membrane from pure solvent to solution is called osmosis.
• Osmotic pressure: The excess pressure that must be applied to solution to prevent the passage of solvent into solution through a semipermeable membrane is called osmotic pressure.
• Osmotic pressure is a colligative property as it depends on the number of solute particles and not on their identity.
• For a dilute solution, osmotic pressure () is directly proportional to the molarity (C) of the solution i.e. = CRT
• Osmotic pressure can also be used to determine the molar mass of solute using the equation 
• Isotonic solution: Two solutions having same osmotic pressure at a given temperature are called isotonic solution.
• Hypertonic solution: If a solution has more osmotic pressure than other solution it is called hypertonic solution.
• Hypotonic solution: If a solution has less osmotic pressure than other solution it is called hypotonic solution.
• Reverse osmosis: The process of movement of solvent through a semipermeable membrane from the solution to the pure solvent by applyingexcess pressure on the solution side is called reverse osmosis.
• Colligative properties help in calculation of molar mass of solutes.
• Abnormal molar mass: Molar mass that is either lower or higher than expected or normalmolar mass is called as abnormal molar mass.
• Van’t Hoff factor: Van’t Hoff factor (i)accounts for the extent of dissociation or association.
• Value of i is less than unity in case solute undergo association and the value of i is greater than unity in case solute undergo dissociation.
• Inclusion of van’t Hoff factor modifies the equations for colligative properties as: