## Tamilnadu Samacheer Kalvi 12th Chemistry Notes Chapter 9 Electro Chemistry Notes

Electro chemistry : The branch of chemistry that deals with the study of electrical energy transport and the inter conversion of electrical and chemical energy is called electro chemistry.

Ohm’s law : At a constant temperature, the current flowing through the cell (I) is directly proportional to the voltage across the cell (V)
I ∝ V
(OR)
I = V/R
(OR) V = IR

Resistivity: Resistance of an electrolytic solution is directly proportional to the length (1) and inversely proportional to the cross sectional area (A)
R ∝ l/A

Conductivity: The reciprocal of the resistance gives the conductance of an electrolytic solution.
C = 1/R

Specific conductivity: The reciprocal of the specific resistance
(OR)
The conductance of a cube of an electrolytic solution of unit dimensions are called specific conductance.
$$\kappa=\frac{1}{\rho} \cdot \frac{l}{A}$$

Molar conductivity (∧m): It is defined as the conductance of a solution containing one mole of the electrolyte dissolved in it.
$$\Lambda_{\mathrm{m}}=\frac{\kappa \times 10^{-3}}{\mathrm{M}} \mathrm{sm}^{-1} \mathrm{~m}^{3} \mathrm{~mol}^{-1} \text {(or) } \mathrm{mho} \mathrm{m}^{3} \mathrm{~mol}^{-1} \text {. }$$

Equivalent conductance (∧): It is defined as the conductance of an electrolyte solution containing one gram equivalent of the electrolyte.
$$\Lambda=\frac{\kappa \times 10^{-3}}{N} \mathrm{sm}^{-1} \mathrm{~m}^{3}$$
(gram equivalent)-1 (or) mhom3 (gram equivalent)-1.

Kohlraush’s law: At infinite dilution, the limiting molar conductivity of an electrolyte is equal to the sum of the limiting of molar conductivities of its constituent ions.

Application of Kohlraush’s law: Kohlraush’s law used as,

• Calculation of molar conductance at infinite dilution of a weak electrolyte.
• Calculation of degree of dissociation of weak electrolytes.
• Calculation of solubility of a sparingly soluble salts.

Electro chemical cell: It is a device which inter converts chemical into electrical energy and vice versa.

Galvanic Cell (Voltaic cell): It is a device in which a spontaneous chemical reaction generates an electric current i.e., it converts chemical energy into electrical energy. It is commonly known as a battery.

Electrolytic cell: It is a device in which an electric current from an external source drives a non-spontaneous reaction i.e., it converts electrical energy into chemical energy.

Electromotive force: The force that pushes the electrons away from the anode and pulls them towards cathode is called the electromotive force (emf) or the cell potential.

Standard Hydrogen Electrode (SHE): It is used as the reference electrode. It has been assigned an orbitary emf of exactly zero volt. It consists of a platinum electrode in contact with 1M HC1 solution and 1 atm hydrogen gas. The hydrogen gas is bubbled through the solution at 25°C.

Electrode potential (E): Electromotive force of a cell in which the electrode on the left is a standard hydrogen electrode and the electrode on the right is the electrode in question.

Standard electrode potential (E°): The value of the standard emf of a cell in which molecular hydrogen under standard pressure is oxidised to solvated protons at the left hand electrode.

Nernst equation: It is the one which relates the cell potential and the concentration of the species involved in an electrochemical reaction.

Electrolysis: The process of chemical decomposition of an electrolyte in solution or molten state by the passage of electric current is called electrolysis.

Faraday’s first law of electrolysis: The mass of the substance (m) liberated at an electrode during electrolysis is directly proportional to the quantity of charge (Q) passed through the cell.
m ∝ Q (OR) m ∝ It (OR) m = Z It.

Faraday’s second law of electrolysis: When the same quantity of charge is passed through the solutions of different electrolytes, the amount of substances liberated at the respective electrodes are directly proportional to their electrochemical equivalents.

Electrochemical equivalent: It is defined as the amount of substance deposited or liberated at the electrode by a charge of one coulomb (one ampere current passing for one second).

Battery: It is a device that produces electrons through electrochemical reactions, and contains positive and negative terminals. A battery consists of one or more electro chemical cells, which transform stored chemical energy directly into electrical energy.

Fuel cell: The galvanic cell in which the energy of combustion of fuels is directly converted into electrical energy is called the fuel cell.

Corrosion: The spontaneous destruction of metals due to their interaction with environment
is called corrosion.

Protection of metals from corrosion: Following methods helps to protect metals from corrosion,

• Coating metal surface by paint.
• Galvanizing – by coating with another metal such as zinc.
• Cathodic protection

1. Variation of molar conductivity with concentration.

 Concentration (M) Molar conductance ( x10‘3Sm2 mol-1) NaCl KCl HCl 0.1 10.674 12.896 39.132 0.01 11.851 14.127 41.20 0.0001 12.374 14.695 42.136

2. ∧°m values for various compounds.

 Electrolyte ∧°m at 298 K Difference KCl NaCl 149.86 126.45 23.41 KBr NaBr 151.92 128.51 23.41 knO3 NaNO3 114.96 121.55 23.41
 Electrolyte ∧°m at 298 in K Difference KBr KCl 151.92 149.86 2.06 NaBr KCl 128.51 126.45 2.06 LiBr LiCl 117.09 115.03 2.06

Samacheer Kalvi 12th Chemistry Notes

## Tamilnadu Samacheer Kalvi 12th Chemistry Notes Chapter 8 Ionic Equilibrium Notes

Acid: The term ‘acid’ is derived from the latin word ‘acidus’ meaning sour, which turns the blue litmus into red.

Base: Base tastes bitter and turns the red litmus to blue.

Arrhenius Concept:

• An acid is a substance that dissociates to give hydrogen ions in water.
• A base is a substance that dissociates to give hydroxyl ions in water.

Limitations of Arrhenius Concept:

• Arrhenius theory does not explain the behaviour of acids and bases in non aqueous ‘ solvents such as acetone, Tetrahydrofuran etc.
• This theory does not account for the basicity of the substances like ammonia (NH3) which do not possess hydroxyl group.

Lowry – Bronsted Theory :

• An acid is defined as a substance that has a tendency to donate a proton to another substance.
• A base is a substance that has a tendency to accept a proton form other substance.

Limitations of Lowry – Bronsted Theory :
Substances like BF3, AlCl3 etc., that do not donate protons are known to behave as acids.

Lewis concept:

• An acid is a species that accepts an electron pair (Lewis acid)
• Base is a species that donates an electron pair. (Lewis base)

Strong acid : A strong acid is the one that is almost completely dissociated in water.

Weak acid : A weak acid is only partially dissociated in water.

Auto ionisation of water : The pure water itself has a little tendency to dissociate.
i.e., one water molecule donates a proton to an another water molecule. This is known as auto ionisation of water.

pH scale : The term pH is derived from the French word ‘Purissance de hydrogene’ meaning, the power of hydrogen. It is defined as the negative logarithm of base 10 of the molar concentration of the hydronium ions present in the solution.

Ostwald’s dilution law : Ostwald’s dilution law relates the dissociation constant of the weak acid (Ka) with its degree of dissociation (α) and the concentration (c).

Degree of dissociation (α): It is the fraction of the total number of moles of a substance that dissociates at equilibrium.

Common ion effect: The dissociation of the weak acid is suppressed in the presence of a salt containing an ion common to the weak electrolyte. It is called the Common ion effect.

Buffer solution : Buffer is a solution which consists of a mixture of a weak acid and its conjugate base (or) a weak base and its conjugate acid. This buffer solution resists drastic changes in its pH upon addition of a small quantities of acids (or) bases, and this ability is called buffer action.

Types of Buffer Solution : There are two types of buffer solutions.

1. Acidic buffer solution : a solution containing a weak acid and its salt.
Example : solution containing acetic acid and sodium, acetate
2. Basic buffer solution : a solution containing a weak base and its salt.
Example : Solution containing NH4OH and NH4Cl

Buffer index : Buffer index ‘β’ is defined as the number of gram equivalents of acid or base added to one litre of the buffer solution to change its pH by unity.$$\beta=\frac{\mathrm{dB}}{\mathrm{d}(\mathrm{pH})}$$

Here,
dB = number of gram equivalents of acid / base added to one litre of buffer solution.
d(pH) = The change in the pH after the addition of acid / base.

Salt Hydrolysis : Salts completely dissociate in aqueous solutions to give their constituent ions. The ions so produced are hydrated in water. In certain cases, the cation, anion or both react with water and the reaction is called salt hydrolysis.

Solubility Product: It is defined as the product of the molar concentration of the constituent ions, each raised to the power of its stoichiometric co-efficient in a balanced equilibrium equation.

Molar solubility : The maximum number of moles of solute that can be dissolved in one litre of the solution.

Difference between Lewis acids and Lewis base:

 Lewis acids ‘            Lewis bases Electron deficient molecules such as BF3, AlCl3, BeF2 etc… Molecules with one (or) more lone pairs of electrons. NH3, H2O, R-O-H,R-O-R, R – NH2 All metal ions (or) atoms Examples: Fe2+ ,Fe3+ ,Cr3+ ,Cu2+ etc… All anions F–,Cl– ,CN– , SCN–, SO42- etc… Molecules that contain a polar double bond Examples : SO2, CO2, SO3 etc… Molecules that contain carbon – carbon multiple bond Examples: CH2 = CH2, CH = CH etc… Molecules in which the central atom can expand its octet due to the availability of empty d – orbitals Example: SiF4, SF4, FeCl3 etc.. All metal oxides CaO, MgO, Na2O etc… Carbonium ion (CH3)3 c+ Carbanion CH3–

Kw values at different temperatures are given in the following table.

 Temperature (°C) Kw 0 1.14 xl0-15 10 2.95 x 1o-15 25 1.00 x 10-14 40 2.71x 10-14 50 5.30 x 10-14

Samacheer Kalvi 12th Chemistry Notes

## Tamilnadu Samacheer Kalvi 12th Chemistry Notes Chapter 7 Chemical Kinetics Notes

Chemical kinetics – The word kinetics is derived from the Greek word “kinesis” meaning movement. Chemical kinetics is the study of the rate and the mechanism of chemical reactions, proceeding under given conditions of temperature, pressure, concentration etc.

Rate of chemical reaction – A rate is a change in a particular variable per unit time. In a chemical reaction, the change in the concentration of the species involved in a chemical reaction per unit time gives the rate of a reaction.
Rate = $$\frac{-[\text { Change in the concentration of the reactants }]}{[\text { Change in time }]}$$

Unit of the rate of a reaction:
Unit of rate = $$\frac{\text { unit of concentration }}{\text { unit of time }}$$

Concentration is expressed in number of moles per litre and time is expressed in seconds and therefore the unit of the rate of a reaction is mol L-1s-1

Average rate – It is obtained by dividing the change in concentration of any of the reactant or product by the time taken for the change.

Instantaneous rate of a reaction – The rate of a reaction at a particular moment of time is called the Instantaneous rate of a reaction.

Rate law – The expression in which reaction rate is given in term of molar concentration of the reactants with each term raised to some power, which may or may not be same as the stoichiometric coefficient of the reacting species in a balanced chemical equation.
For a general reaction, xA+yB → products

The rate law for the above reaction is generally expressed as Rate = k[A]m [B]n
Where k is proportionality constant which is called rate constant. The values of m and n represent the reaction order with respect to A and B respectively.

Order of a reaction – The sum of exponents of the concentration of the reactants in the rate law expression.

Differences between rate and rate constant of a reaction

 Rate of a reaction Rate constant of a reaction 1. It represents the speed at which the reactants are converted into products at any instant. It is a proportionality constant 2. It is measured as decrease in the concentration of the reactants or increase in the concentration of products. It is equal to the rate of reaction, when the concentration of each of the reactants in unity 3. It depends on the initial concentration of reactants. It does not depend on the initial concentration of reactants.

Elementary reaction – Each and every single step in a reaction mechanism is called an elementary reaction.

Molecularity – The total number of reactant species that are involved in an elementary step is called molecularity of that particular step.

Differences between order and Molecularity

 Order of a reaction Molecularity of a reaction 1. It is the sum of the powers of concentration terms involved in the experimentally determined rate law. It is the total number of reactant species that are involved in an elementary step. 2. It can be zero (or) fractional (or) integer. It is always a whole number, cannot be zero or a fractional number. 3. It is assigned for a overall reaction. It is assigned for each elementary step of mechanism.

Examples for the first order reaction

(i) Decomposition of dinitrogen pentoxide

(ii) Decomposition of thionylchloride;

(iii) Decomposition of the H202 in aqueous solution;

(iv) Isomerisation of cyclopropane to propene.

Pseudo first order reaction – A second order reaction can be altered to a first order reaction by taking one of the reactant in large excess. Such reaction is called pseudo first order reaction.

Examples for a zero order reaction
(i) Photochemical reaction between H2 and Cl2

(ii) Decomposition of N2O on hot platinum surface

(iii) lodination of acetone in acid medium is zero order with respect to iodine.

Half life period of a reaction – The half life period of a reaction is defined as the time required for the reactant concentration to reach one half its initial value.

Collision theory – According to this theory, chemical reactions occur as a result of collisions between the reacting molecules.

Arrhenius equation

k= $$\mathrm{Ae}^{-\left(\frac{\mathrm{E}_{a}}{\mathrm{RT}}\right)}$$
A = Arrhenius factor (frequency factor)
R = Gas constant
k = Rate constant
Ea = Activation energy
T Absolute temperature (in K)

Factors affecting the reaction rate are,

1. Nature and state of the reactant
2. Concentration of the reactant
3. Surface area of the reactant
4. Temperature of the reaction
5. Presence of a catalyst

Samacheer Kalvi 12th Chemistry Notes

## Tamilnadu Samacheer Kalvi 12th Chemistry Notes Chapter 6 Solid State Notes

Solids – The substances having definite shape mass and volume are called solids.

General characteristics of solids –

• Solids have definite volume and shape.
• Solids are rigid and incompressible
• Solids have strong cohesive forces.
• Solids have short inter atomic, ionic or molecular distances.
• Their constituents ( atoms , ions or molecules) have fixed positions and can only oscillate about their mean positions

Types of solids – They are two types.

1. Crystalline solids
2. Amorphous solids

Crystalline solids – In a crystalline solid, the particles are arranged in a regular and repetitive three dimensional arrangement.

Amorphous solids – In a amorphous solid, the particles are arranged in an irregular and non- repetitive three dimensional arrangement.

Isotropy – Isotropy means uniformity in all directions. In solid state isotropy means having identical values of physical properties such as refractive index electrical conductance etc., in all directions.

Anisotropy – Anisotropy is the property which depends on the direction of measurement.

Ionic solids – The structural units of an ionic crystal are cations and anions. They are bound together by strong electrostatic attractive forces. To maximize the attractive force, cations are surrounded by as many anions as possible and vice versa. Ex., NaCl crystal

Characteristics of ionic crystals

• Ionic solids have high melting points.
• These solids do not conduct electricity, because the ions are fixed in their lattice positions.
• They do conduct electricity in molten state (or) when dissolved in water because, the ions are free to move in the molten state or solution.
• They are hard as only strong external force can change the relative positions of ions.

Covalent solids – In covalent solids the constituents are bound together in a three dimensional network entirely by covalent bonds. Ex., Diamond

Molecular solids – In molecular solids, the constituents are neutral molecules. They are held together by weak vander waals forces. Ex., Naphthalene, solid CO2, urea

Metallic solids – In metallic solids, the lattice points are occupied by positive metal ions and a cloud of electrons pervades the space. Ex., all kind of metals (Cu, Ni, Fe).

Crystal lattice – Crystalline solid is characterised by a definite orientation of atoms, ions or molecules, relative to one another in a three dimensional pattern. The regular arrangement of these species throughout the crystal is called a crystal lattice.

Unit Cell – A basic repeating structural unit of a crystalline solid is called a unit cell.

Primitive and non-primitive unit cell:

• A unit cell that contains only one lattice points is called a primitive unit cell, which is made up from the lattice points at each of the comers.
• In case of non-primitive unit cells there are additional lattice points, either on a face of the unit cell or with in the unit cell.

Seven classes of primitive crystal systems :

 System Crystallographic axes Crystallographic angles (i) Cubic a = b = c α = β = γ = 90° (ii) Rhombohedral a = b = c α = β = γ ≠ 90° (iii) Hexagonal a = b ≠ c α – β = 90°, γ  = 120° (iv) Tetragonal a ≠ b ≠ c α = β = γ = 90° (v) Orthorhombic A ≠ b ≠ c α = β = γ = 90° (vi) Monoclinic a ≠ b ≠ c α = β = 90°, β ≠ 90° (vii) Triclinic a ≠ b ≠ c α ≠ β ≠ γ

Packing efficiency – There is some free space between the spheres of a single layer and the spheres of successive layers. The percentage of total volume occupied by these constituent spheres gives the packing efficiency of an arrangement.

Schottky defect: – Schottky defect arises due to the missing of equal number of cations and anions from the crystal lattice. This effect does not change the stoichiometry of the crystal. Example :NaCl.

Frenkel defect: – Frenkel defect arises due to the dislocation of ions from its crystal lattice. The ion which is missing from the lattice point occupies an interstitial position.
Example: AgBr

Metal excess defect: – Metal excess defect arises due to the presence of more number of metal ions as compared to anions.Ex., Alkali halides (NaCl, KCl).

Metal deficiency defect: – Metal deficiency defect arises due to the presence of less number of cations than the anions. This defect is observed in a crystal in which, the cations have variable oxidation states. Example: FeO

Impurity defect: – This defect arises due to cation vacancies. In some ionic compounds presence of impurity produce some defects called impurity defects.
Ex, (i) presence of Cd++ ion as impurity in AgCl.
(ii) presence of Sr++ ion as impurity in AgCl.

Samacheer Kalvi 12th Chemistry Notes

## Tamilnadu Samacheer Kalvi 12th Chemistry Notes Chapter 5 Coordination Chemistry Notes

Double Salts:

• Addition compounds are formed when stoichiometric amounts of two or more stable compounds combine together.
• The Compounds which exist only in the solid state but dissociate into their constituent simple ions when dissolved in water. Such addition compounds are called double salts.
• For example, Mohr’s Salt – FeSO4.(NH4)2SO4.6H2O Potash alum – K2SO4.Al2(SO4)3.24 H2O

Coordination Compounds :

• The transition metals have a tendency to form complexes. The name is derived from the Latin words “complexus” and “coordinate” which mean “hold” and “to arrange” respectively.
• The compounds which retain their identity even in the solution, and have properties entirely different from those of the constituent simple ions are called coordination compounds.
• For example, [Cu(NH3)4]SO4 is a typical coordination compound.

Coordination Entity :

Coordination entity is an ion or a neutral molecule, composed of a central atom, usually a metal and the array of other atoms or groups of atoms (ligands) that are attached to it. In the formula the coordination entity is enclosed in sequare brackets. For example, in potassium ferrocyanide, K4[Fe(CN)6], the coordination entity is [Fe(CN)6]4-

Central metal atom / ion :
The central atom/ion is the one that occupies the central position in a coordination entity and binds other atoms or groups of atoms (ligands) to itself, through a coordinate covalent bond. For example, in K4[Fe(CN)6] the central metal ion is Fe2+.

Ligands :
The ligands are the atoms or groups of atoms bound to the central atom/ion. The atom in a ligand that is bound directly to the central metal atom is known as a donor atom. For example, in K4[Fe(CN)6] the ligand is CN ion.

Coordination sphere :
The complex ion of the coordination compound containing the central metal atom/ion and the ligands attached to it, is collectively called coordination sphere and are usually enclosed in square brackets with the net charge. For example, the coordination compound K4[Fe(CN)6] contains the complex ion [Fe(CN)6]4- and is referred as the coordination sphere.

Coordination polyhedron :
The three dimensional spacial arrangement of ligand atoms/ ions that are directly attached to the central atom is known as the coordination polyhedron. For example, in K4[Fe(CN)6] the coordination polyhedra is octahedral. The coordination polyhedra of [Ni(CO)4] is tetrahedral.

Coordination number :
The number of ligand donor atoms bonded to a central metal ion in a complex is called the coordination number of the metal. For example, in K4[Fe(CN)6] the coordination number of Fe2+ is 6.

Oxidation state :
The oxidation state of a central atom in a coordination entity is defined as the charge it would bear if all the ligands were removed along with the electron pairs that were shared with the central atom. For example, in the coordination entity [Fe(CN)6]4- the oxidation state of iron is represented as (II).

Types of complexes :
The coordination compounds can be classified into the following types based on,
(i) the net charge of the complex ion
(ii) kinds of ligands present in the coordination entity.

• Cationic complex : Carries a net positive charge. E.g: [Ag(NF3)2]
• Anionic complex : Carries a net negative charge. E.g: [Ag(CN)2)]
• Neutral complex : Bears no net charge. E.g: [Ni(CO)4].

Isomerism :
It is the phenomenon in which more than one coordination compounds having the same molecular formula have different physical and chemical properties due to different arrangement of ligands around the central metal atom.

Structural isomers :
The coordination compounds with same formula, but have different connections among their constituent atoms are called structural isomers or constitutional isomers.

This type of isomers arises when an ambidentate ligand is bonded to the central metal atom/ion through either of its two different donor atoms.

Coordination isomers :
This type of isomers arises in the coordination compounds having both cation and anion as complex ions. The interchange of one or more ligands between the cationic and the anionic coordination entities result in different isomers.

Ionisation isomers :
This type of isomers arises when an ionisable counter ion itself can act as a ligand. The exchange of such counter ions with one or more ligands in the coordination entity will results in ionisation isomers.

Solvate isomers :
The exchange of free solvent molecules such as water, ammonia, alcohol etc. in the crystal lattice with a ligand in the coordination entity will give different isomers. These type of isomers are called solvate isomers.

Stereoisomers :
The stereoisomers of a coordination compound have the same chemical formula and connectivity between the central metal atom and the ligands. But they differ in the spatial arrangement of ligands in three dimensional space.

Geometrical isomers :
Geometrical isomerism exists in heteroleptic complexes due to different possible three dimensional spatial arrangements of the ligands around the central metal atom.

Optical isomerism :
Coordination compounds which possess chairality exhibit optical isomerism similar to organic compounds. The pair of two optically active isomers which are mirror images of each other are called enantiomers. Their solutions rotate the plane of the plane polarised light either clockwise or anti clockwise and the corresponding isomers are called’d’ (dextrorotatory) and T (levorotatory) forms respectively.

Homoleptic complex :
The central metal ion / atom is coordinated to only one kind of ligands is called a homoleptic complex. E.g: [Co(NH3)6]3+, [Fe(H2O)6]2+.

Heteroleptic complex :
The central metal ion/atom is coordinated to more than one kind of ligands is called a heteroleptic complex. E.g:[Co(NH3)5Cl]2+, [Pt(NH3)Cl2].
Crystal field splitting power of various ligands :

The above series is known as spectrochemical series.

Strong field ligands : The ligands present on the right side of the spectrochemical series such as carbonyl causes relatively larger crystal field splitting and are called strong field ligands.

Weak field ligands : The ligands present on the left side of the spectrochemical series such as iodide causes relatively smaller crystal fields splitting and are called weak fields ligands

Colour of the complex: Most of the transition metal complexes are coloured. A substance exhibits colour when it absorbs the light of a particular wavelength in the visible region and transmit the rest of the visible light. When this transmitted light enters our eye, our brain recognises its colour. The colour of the transmitted light is given by the complementary colour of the absorbed light.

Metallic carbonyls : Metal carbonyls are the transition metal complexes of carbon monoxide, containing metal-carbon bond. In these complexes CO molecule acts as a neutral ligand.

Synergic effect : In metal carbonyls, electron density moves from ligands to metal through sigma bonding and from metal to ligand through Pi-bonding. This is called synergic effect.

Labile complexes : In some cases, complexes can undergo rapid ligand substitution, such complexes are called labile complexes.

Inert complexes : In some cases, complexes undergo ligand substitution very slowly (or sometimes no substitution). Such complexes are called inert complexes.

Stability constant: The stability of a coordination complex is a measure of its resistance to the replacement of one ligand by another. The stability of a complex refer to the degree of association between two species involved in an equilibrium.

C is platin : It is a square planar coordination complex cis – [Pt(NH3)2Cl2], in which two similar ligands are in adjacent positions. It is a platinum based anticancer drug.

Coordination number : The number of ligand donor atoms bonded to central metal ion in a complex is called the coordination number of the metal.

Naming the central metal:

 Element Name of the metal in Cation complex Anionic complex Cr Chromium Chromate Zn Zinc Zincate A1 Aluminium Aluminate Fe Iron Ferrate Cu Copper Cuperate Co Cobalt Cobaltate Pb Lead Plumbate Ag Silver Argentate Sn Tin Stannate Au Gold Aurate Pt Platinum Platinate

List of absorbed wavelength and their complementary colour

 Wave length(λ) of absorbed light (Å) Wave number(v) of the absorbed light (cm-1) Colour of absorbed light Observed Colour 4000 25000 Violet Yellow 4750 21053 Blue Orange 5100 19608 Green Red 5700 17544 Yellow Violet 5900 16949 Orange Blue 6500 15385 Red Green

Samacheer Kalvi 12th Chemistry Notes

## Tamilnadu Samacheer Kalvi 12th Chemistry Notes Chapter 4 Transition and Inner Transition Elements Notes

Position of d-block elements: Group – 3 to group – 12 of the modem periodic table is the position of d-block elements.

Electronic configuration: [Noble gas] (n – l)d1-10 ns1-2

Metallic behaviours: Except Cu, Ag, Au all are hard. Most of them are hexagonal close packed, cubic close packed (or) body centred cubic which are the characteristics of true metals.

Variation (or) trend in properties of d-block elements:

 Property Across the period Down the group 1. Melting point First increases and reaches maximum value then decreases Decreases 2. Atomic and ionic size Generally decreases, but in 3d series, atomic radius decreases Sc to V and then upto Cu, remains the same. Generally increases 3. Ionization enthalpy Increases Decreases 4. Oxidation state Except Sc, all exhibit variable oxidation states generally increase and then decrease. Decreases

Standard electrode potentials: It is the value of the standard emf of a cell in which molecular hydrogen under standard pressure ( latm) and temperature (273K) is oxidised to solvated protons at the electrode.

Magnetic properties: Most of the compounds of transition elements are paramagnetic.

Magnetic moment: μ = $$g \sqrt{S(S+1)} \mu_{B}$$
where g = 2, S is the total spin quantum number of unpaired electrons and pB is Bohr Magneton.

Catalytic property: Most of the transition metals and their compounds act as catalyst.

Alloy formation: An alloy is formed by blending a metal with one or more other metals.

Formation of interstitial compounds: It is a compound that is formed when small atoms like hydrogen, boron, carbon or nitrogen are trapped in the interstitial holes in a metal lattice.

Formation of complex: Transition elements have a tendency to form complex compounds with a species that has an ability to donate an electron pair to form a coordinate covalent bond.

K2Cr2O7: Potassium dichromate is an orange red crystalline solid.

Structure of dichromate ion:

Potassium permanganate -KMnO4: It is a dark purple crystals.

Structure of permanganate ion:

Baeyer’s reagent: Cold, dilute alkaline KMnO4

f-block elements: 1. Lanthanoids 2. Actinoids

Lanthanoids and Actinoids differences:

 Lanthanoids Actinoids Electronic configuration [Xe] 4f1-14 5d0-1 6s2 Electronic configuration [Rn] 5f1-14 6d0-27s2 Common oxidation state is +3 Common oxidation state is +4 Differentiating electron eneters 4f orbital Differentiating electron eneters in 5f orbital Binding energy of 4f orbitals are higher Binding energy of 5f orbitals are lower They show less tendency to form complexes They show greater tendency to form complexes Most of the lanthanoids are colourless Most of the actinoids are coloured. E.g : U3+ (red), U4+ (green), UO22+(yellow) They do not form oxo cations They do form oxo cations such as UO22+, NpO22+ etc. Besides +3 oxidation states lanthanoids show +2 and +4 oxidation states in few cases. Besides +3 oxidation states actinoids show higher oxidation states such as +4, +5, +6 and +7.

Samacheer Kalvi 12th Chemistry Notes

## Tamilnadu Samacheer Kalvi 12th Chemistry Notes Chapter 3 p-Block Elements-II Notes

Group-15 elements: About 78% of earth atmosphere contains dinitrogen gas. The other elements arsenic, antimony and bismuth are present as sulphides and are not very abundant.

Nitrogen:

• Nitrogen, the principle gas of atmosphere (78% by volume) is separated industrially from liquid air by fractional distillation.
• Nitrogen gas is rather inert.
• It can form complexes with metal.

Uses of nitrogen:

• Nitrogen is used for the manufacture of ammonia, nitric acid and calcium cyanamide etc.
• Liquid nitrogen is used for producing low temperature required in cryosurgery, and in biological preservation.

Ammonia:

• Ammonia is formed by the hydrolysis of urea.
• It is a pungent smelling gas and it is lighter than air.
• It is highly associated through strong hydrogen bonding.
• Ammonia act as a reducing agent.

Nitric acid:

• Nitric acid is prepared by heating equal amounts of potassium or sodium nitrate with concentrated H2SO4.
• Nitric acid act as an oxidising agent.
• Some metals such as Al, Fe, Co, Ni and Cr are rendered passive in concentrated nitric acid due to the formation of a layer of their oxides on the metal surface.

Uses of nitric acid:

• Nitric acid is used as a oxidising agent and in the preparation of aquaregia.
• Salts of nitric acid are used in photography (AgNO3) and gunpowder for firearms. (NaNO3)

Allotopic forms of phosphorus: Three forms namely,

• White phosphorous
• Red phosphorous
• Black phosphorous

Uses of phosphorous:

• The red phosphorus is used in the match boxes.
• It is also used for the production of certain alloys such as phosphor bronze.

Phosphine:

• Phosphine is the most important hydride of phosphorous.
• It is prepared by action of sodium hydroxide with white phosphorous.
• It is poisonous gas with rotten fish smell.

Uses of Phosphine: Phosphine is used for producing smoke screen as it gives large smoke. In a ship, a pierced container with a mixture of calcium carbide and calcium phosphide, liberates phosphine and acetylene when thrown into sea. The liberated phosphine catches fire and ignites acetylene. These burning gases serves as a signal to the approaching ships. This is known as Holmes signal.

Phosphorous trichloride:

• When a slow stream of chlorine is passed over white phosphorous, PCl3 is formed.
• PCl3 is used as a chlorinating agent and for the preparation of H3PO3.

Phosphorous petachloride:

• hen PCl3 is treated with excess chlorine, PCl5 is obtained.
• PCl5 is chorinating agent and is useful for replacing hydroxyl groups by chlorine atom.

Groups-16 elements:

• Elements belonging group-16 are called chalgogens or ore forming elements.
• First element oxygen, the most abundant element, exists in both as dioxygen in air and in combined form as oxides.
• The other elements of this groups are, Sulphur, Selenium, Tellurium and Pollonium.

Oxygen:

• The atmosphere and water contain 23% and 83% by mass of oxygen respectively.
• The decomposition of hydrogen peroxide in the presence of catalyst produce oxygen.
• Under ordinary condition oxygen exists as a diatomic gas.
• Oxygen exists in two allotropic forms namely dioxygen (O2) and ozone (O3).

Uses of oxygen:

• Oxygen is one of the essential component for the survival of living organisms.
• It is used in welding (oxyacetylene welding)
• Liquid oxygen is used as fuel in rockets etc…

Allotrophic forms of sulphur: The crystallline allotropic forms are,

• Rhombic sulphur
• Monoclinic suphur

Amorphous allotropic forms are,

• Plastic sulphur
• Milk of sulphur
• Colloidal sulphur

Sulphur dioxide:

• Sulphur dioxide is prepared in the laboratory treating a metal or metal sulphite with sulphuric acid.
• Sulphur dioxide gas is formed in volcanic eruptions.
• It is an acidic oxide.
• It can be liquefied at 2.5 atomospheric pressure and 288 K temperature.

Uses of Sulphur dioxide:

• Sulphur dioxide is used in bleaching hair, silk, wool etc…
• It can be used for disinfecting crops and plants in agriculture.

Sulphuric acid:

• Sulphuric acid can be manufactured by lead chamber process or contact process.
• Pure sulphuric acid is a colourless,viscous liquid.
• It is highly reactive and it can act as dehydrating agent.

Uses of Sulphuric acid:

• Sulphuric acid is used in the manufacture of fertilisers, ammonium sulphate and super phosphates and other chemicals such as hydrochloric acid, nitric acid etc…
• It is used as a drying agent and also used in the preparation of pigments, explosives etc..

Group-17 elements:

• The halogens are present in combined form as they are highly reactive.
• They are also called as salt producers.
• The elements present in group-17 are, Fluorine, Chlorine, Bromine, Iodine and Astatine.

Chlorine:

• The main source of chlorine is sodiumchloride from sea water.
• Chlorine is highly reactive.
• Chlorine is manufactured by the electrolysis of brine in electrolytic process or by oxidation of HCl by air in Deacon’s process.
• It is a greenish yellow gas with pungent irritating odour.
• Bleaching powder is produced by passing chlorine gas through dry slaked time.

Uses of Chlorine:

• Purification of drinking water.
• Bleaching of cotton textiles, paper and rayon.
• It is used in extraction of gold and platinum.

Hydrochloric acid:

• It is prepared by the action,of sodium chloride and concentrated sulphuric acid.
• When three parts of cone. HCl and one part of cone. HNO3 are mixed aquaregia (Royal water) is obtained.

Uses of Hydrochloric acid:

• Hydrochloric acid is used for the manufacture of chlorine, ammonium chloride, glucose from com starch etc.,
• It is used in the extraction of glue from bone and also for purification of bone black.

Hydrogen halides:

• Direct combination is a useful means of preparing hydrogen chloride. The reaction between hydrogen and fluorine is violent while the reaction between hydrogen and bromine or hydrogen and iodine are reversible and don’t produce pure forms.
• The hydrogen halides are extremely soluble in water due to the ionisation.
• Solution of hydrogen halides are acidic and known as hydrohalic acids.

Interhalogen compounds: Each halogen combines with other halogen to form a series of compounds called interhalogen compounds.

Group-18 elements:

• All the noble gases occur in the atmosphere.
• Elements belongs to group-18 are Helium, Neon, Argon, Krypton, Xenon and Radon.
• The first ionisation energy decreases from helium to radon.
• They are extremely stable and have a small tendency to gain or lose electrons.
• Noble gases are monoatomic.

Uses of Helium:

• Helium and oxygen mixture is used by divers in place of air oxygen mixture. This prevents the painful dangerous condition called bends.
• Helium is used to provide inert atmosphere in electric arc welding of metals.
• Helium has lowest boiling point hence used in cryogenics (low temperature science).
• It is much less denser than air and hence used for filling air balloons.

Uses of Neon: Neon is used in advertisement as neon sign and the brilliant red glow is caused by passing electric current through neon gas under low pressure.

Uses of Argon: Argon prevents the oxidation of hot filament and prolongs the life in filament bulbs

Uses of Krypton: Krypton is used in fluorescent bulbs, flash bulbs etc… Lamps filed with krypton are used in airports as approaching lights as they can penetrate through dense fog.

Uses of Xenon: Xenon is used in fluorescent bulbs, flash bulbs and lasers. Xenon emits an intense light in discharge tubes instantly. Due to this it is used in high speed electronic flash bulbs used by photographers.

Uses of Radon: Radon is radioactive and used as a source of gamma rays.Radon gas is sealed as small capsules and implanted in the body to destroy malignant i.e. cancer growth

Hydrogen halide

 Property Order Reactivity of hydrogen Decreases from fluorine to iodine Stability Decreases from HF to HI Volatility of the hydrides HF < HI < HBr < HCl Thermal stability HF > HI > HBr > HCl Boiling point HCl < HBr < HI Acid strength Increases from HF to HI

Samacheer Kalvi 12th Chemistry Notes

## Tamilnadu Samacheer Kalvi 12th Chemistry Notes Chapter 1 Metallurgy Notes

Metallurgy: Metallurgy relates to the science and technology of metals. Metallurgy helps to study the various steps involved in the extraction of metal and the chemical principles behind these processes.

Mineral: A naturally occurring substance obtained by mining which contains the metal in free state or in the form of compounds like oxides, sulphides etc. is called a mineral.

Ore: Minerals that contains a high percentage of metal, from which it can be extracted conveniently and economically are called ores.

Gangue: Ores are associated with non-metallic impurities, rocky materials and siliceous matter which are collectively known as gangue.

Roasting: Roasting is the method, usually applied for the conversion of sulphide ores into their oxides. In this method, the concentrated ore is oxidized by heating it with excess of oxygen in a suitable furnace below the melting point of the metal.

Calcination: It is the process in which the concentrated ore is strongly heated in the absence of air. Here ore is converted into their oxides.

Smelting: It is the process of reducing the roasting metallic oxide to metal in molten condition.

Blistered copper: The solidified copper obtained after the reduction of copper matte has blistered appearance. This is called blistered copper.

Ellingham diagram: The graphical representation of variation of the standard Gibbs free energy of reaction for the formation of various metal oxides with temperature is called Ellingham diagram.

Distillation: In this method, the impure metal is heated to evaporate and the vapours are condensed to get pure metal.

Liquation: This method is employed to remove the impurities with high melting points from metals having relatively low melting points such as tin.

Electrolytic refining: The crude metal is refined by electrolysis. It is carried out in an electrolytic cell containing aqueous solution of the salts of the metal of interest. The rods of impure metal are used as anode and thin strips of pure metal are used as cathode.

Application of Aluminium (Al): Aluminium is the most abundant metal and is a good conductor of electricity and heat. It also resists corrosion. The following are some of its applications:

• Many heat exchangers/sinks and our day to day cooking vessels are made of aluminium.
• It is used as wraps (aluminium foils) and is used in packing materials for food items.
• Aluminium is not very strong, However, its alloys with copper, manganese, magnesium and silicon are light weight and strong and they are used in design of aeroplanes and other forms of transport.
• As Aluminium shows high resistance to corrosion, it is used in the design of chemical reactors, medical equipments, refrigeration units and gas pipelines.
• Aluminium is a good electrical conductor and cheap, hence used in electrical overhead and electric cables with steel core for strength.

Application of Zinc (Zn):

• Metallic zinc is used in galvanising metals such as iron and steel structures to protect them from rusting and corrosion.
• Zinc is also used to produce die-castings in the automobile, electrical and hardware industries.
• Zinc oxide is used in the manufacture of many products such as paints, rubber, cosmetics, pharmaceuticals, plastics, inks, batteries, textiles and electrical equipment. Zinc sulphide is used in making luminous paints, fluorescent lights and x-ray screens.
• Brass, an alloy of zinc is used in water valves and communication equipment, as it is highly resistant to corrosion.

Application of Iron (Fe):

• Iron is one of the most useful metals and its alloys are used everywhere including bridges, electricity pylons, bicycle chains, cutting tools and rifle barrels.
• Cast iron is used to make pipes, valves, pump stoves, etc.
• Magnets can be made of iron and its alloys and compounds.
• An important alloy of iron is stainless steel, and it is very resistant to corrosion. It is used in architecture, bearings, cutlery, surgical instruments and jewellery. Nickel steel is used for making cables, automobiles and aeroplane parts. Chrome steels are used for manufacturing cutting tools and crushing machines.

Application of Copper (Cu):

• Copper is the first metal used by the humans and extended use of its alloy bronze resulted in a new era, ’ Bronze age ’.
• Copper is used for making coins and ornaments along with gold and other metals.
• Copper and its alloys are used for making wires, water pipes and other electrical parts.

Application of Gold (Au):

• Gold, one of the expensive and precious metals. It is used for coinage, and has been used as standard for monetary systems in some countries.
• It is used extensively in j ewellery in its alloy form with copper. It is also used in electroplating to cover other metals with a thin layer of gold which are used in watches, artificial limb joints, cheap jewellery, dental fillings and electrical connectors.
• Gold nanoparticles are also used for increasing the efficiency of solar cells and also used as catalysts.

Samacheer Kalvi 12th Chemistry Notes

## Tamilnadu Samacheer Kalvi 12th Computer Science Notes Chapter 6 Control Structures Notes

Control structure:
A program statement that causes a jump of control from one part of the program to another is called control structure or control statement.

Control Structures in Python:
There are three important control structures namely

• Sequential
• Alternative or Branching
• Iterative or Looping

Sequential Statement:
A sequential statement is composed of a sequence of statements which are executed one after another. A code to print your name, address and phone number is an example of sequential statement.

Alternative or branching statements in Python:
The types of alternative or branching statements provided by Python are:

• Simple if statement
• if..else statement
• if..elif statement

i) Simple if statement
Simple if is the simplest of all decision making statements. Condition should be in the form of relational or logical expression.

Syntax:
if:<condition>:
statements-block1

(ii) if..else statement
The if ..else statement provides control to check the true block as well as the false block. Following is the syntax of ‘if..else’ statement.

Syntax:
if:<condition>:
statements-block 1
else:
statements-block 2

iii) if..else statement
The if ..else statement provides control to check the true block as well as the false block. Following is the syntax of ‘if..else’ statement.

Syntax:
if:<condition>:
statements-block 1
else:
statements-block 2

Iteration or Looping constructs :

• Iteration or loop are used in situation when the user need to execute a block of code several of times or till the condition is satisfied.
• A loop statement allows to execute a statement or group of statements multiple times.

Looping constructs in Python:
Python provides two types of looping constructs:

• while loop
• for loop
1. While loop:
Syntax:
while< condition > :
statements block 1
[else:
statements block 2]

Samacheer Kalvi 12th Computer Science Notes