CHEMISTRY_QUICK_NOTES.pdf

Chemistry Quick Revision Notes 1 BOM SERIES 176/200 MCQs matched with MDCAT 2012.
146/200 MCQS Matched with ETEA 2019 with BOM. www.BANKOFMCQS.com
Quick revision Notes
By BOM
1 STOICHIOMETRY
Quantitative analysis
 Glucose → diabetes
 Steroids+stimultants → athletes
 Cholesterol → blood of diabetes
Dozen 12, Gross 144, Mole NA
Mole
 Atomic mass → element
 Molecular mass → compound
 Formula mass → ionic compound
No of moles
 Mass/molar mass
 No of particles/NA
 No of formula units/NA
 Volume /molar volume (22.4dm3)
Neutralization reaction
HCl + NaOH → NaCl + H2O
STP: 00C or 273 K, 1 atm
1 dm3 = 1 litre= 1000 cm3
% of an element = (total mass of an
element/total mass of compound)x 100
Sum of all percentages must equal to 100
Actual yield is less than theoretical
because of
 Reversibility
 Side reaction
 Mechanical loss
 Human error
 Reaction condition (conc, T and P)
 % yield = (actual yield/theoretical
yield) x 100
2 ATOMIC STRUCTURES
 Candescent → light by heat
 Charcoal → desired colour → proper
time
 Sodium → gold or yellow colour
 Zinc → smoke effect
John Dalton laws of chemical
combination
 Law of cons of matter
 Law of definite proportion
 Law of multiple proportion
 Gold Stein → cathode rays
 G.J Stoney → electrons
Cathode rays properties
 Hittorf →1869→straight
 Crooke→1870→mv
 Pierre+Thomson→1895-
97→negative
 Thomson →1897 → electric field
deflection
 Cathode rays→ high reducing effect
 For electron, e/m = 1.602 x 10-19 /
9.11 x 10-31 = 1.7588 x 1011
 for proton, e/m = 1.602 x 10 -
19/1.6726 x 10-27 = 9.54 x 10 7
masses

Electron → 9.11 x 10-31
 proton →1.6726 x 10-27

neutron → 1.6749 x 10-27
 h = 6.6262 x 10 -34 Js
 k = 9 x 109
 c = 2.99 x 108 m/s
 R = 1.097 x 107
Bohr model
 rn =
 En = -
 f =
 wave number = f/c =
for Z=1 = R
rays → wavelength (nm)
gamma →>0.006
x-rays → 0.006 -8
ultra violet→8-380
visible→380-760
infrared→760-10 lac
microwaves→10 lac-30 crore
radio waves→>30 crore
spectral lines
 Lyman →ultravoilet→n=1
 Bamer→visible → n=2
 Paschan→near IR → n=3
 Brackett→mid IR→n=4
 Pfund→far IR → n=5
Bohr’s theory defect
 Can’t explain multiplicity of spectral
lines
 Can’t explain magnetic field effect
(Zeeman effect)
 Can’t explain electric field affect
(stark effect)
 Against Heisenberger’s principle
X rays production method
 Roentgen method
 Coolidge method
 Betatron metod
X rays properties
 Neutral
 Reflected + refracted+diffracted
 Penetrate
X ray has 3 types
Moseley’s law
Underroot v = a (Z-b)
a → proportionality constant
b → screening constant
 S→ spherical
 P→ dumbbell
 D→ sausage
 F→ complicated
 Parallel spin ↑↑
 Antiparallel or pair up spins ↑↓
 Total nodes = angular +radial nodes =
n-1
 Angular nodes = l
 Radial nodes = (n-1) – l
 Aufbau principle → building up
 Pauli’s principle → no 2 es have same
quantum numbers set
 Hund’s rule→ degenerate orbitals
Principle N K→1,l→2 2n2
Azimuthal L 0 to n-1 2(2l+1)
Magnetic M -l --0--+1 2l+1
Spin s
3 THEORIES OF COVALENT BONF AND
SHAPES OF MOLECLES
Cystine(sulphur containing) → hairs
Disulphide bond(alpha covalent bond)
stronger than hydrogen bond

Lone pair – lone pair > lone pair-bond
pair > bond pair – bond pair
Shapes of molecules
BeCl2 Linear 1800C
BF3 triangular planar 1200C
SnCl2 Angular structure Less than
1200C
CH4 Tetrahedral
structure
109.50C
NH3 Triangular
structure
107.50C
H2O Angular structure 104.50C
Sigma bond
s-s overlap → H-H
s-p overlap → H-Cl
p-p overlap → Cl-Cl
Hybridization
Hybridiz
ation
e.g. Angl
e
S
chara
cter
P
chara
cter
Sp3 Meth
ane
109.
50C
25 % 75 %
Sp2 Ethan
e
1200
C
33 % 66 %
sp Acety
lene
1800
C
50 % 50 %
Lower energy → bonding molecular
orbital → sigma and pi
Higher energy → anti bonding → sigma
star and pi star
Bond order = (pair of bonding molecular
orbital – pair of anti-bonding molecular
orbital) / 2
Bond order of
H2 → 1
He2 → 0
O2 → 2
N2 → 3
→ greater the bond order, greater the
bond dissociation energy, smaller the
bond length
Stability of molecule
Stable → nb > na
Unstable → nb < na
Energy need to break one mole H2 → 436
KJ/mol
Energy contributed by
H → 36.21 x 10 -23 KJ/mol
Cl → 19.73 x 10 -23 KJ/mol
Expected energy for HCl →55.94x 10 -23
KJ/mol
Experimentally energy for HCl → 72.39 x
10 -23 KJ/mol
→ greater the charge difference, greater
is the bond dissociation energy
Radius of carbon
C → 77 pm (c-c → 154 9m)
Cl → 99 pm (Cl-Cl→198pm)
H → 37 pm
H-Cl → 136 (expected)
H-Cl → 127 pm (experimentally→
because its hetronuclear molecule)
Energy difference trend
HF > HCl = HBr > HI
Ionic character, if E.N difference is
Less than 0.9 → non polar
Between 0.9-1.7 → polar to some ionic
character
Greater than 1.7 → ionic
Dipole movement (vector )
Si unit Cm other unit is debye
1D = 3.34 x 10-30 Cm
Measure by electric condenser
% of ionic character = (μobserver / μionic) x
100
Dipole movement fr
Water→ 1.84 D
CO2 → zero
CO → has dipole
dependence
Physical properties → intermolecular
forces (Vander waal)
Chemical properties → intramolecular
forces
Reaction between
Covalent→ slow
Ionic → fast
Mixing of silver nitrate and sodium
chloride solutions produces a white ppt
of silver chloride
Ionic → non-directional
Covalent → direction
4. GASES
Gas density is 10-3 times less than solid
and liquid
Density of
Oxygen gas →0.00142 g cm-3
Oxygen liquid →1.149 g cm-3
Oxygen solid →1.426 g cm-3
Liquid natural gas → vehicles
Liquid oxygen → patients
Liquid nitrogen → dermatologists
Liquefaction of air → to obtain gases
Liquid Helium below 2.17 K →zero
viscosity and climbing the wall of vessel
Kinetic molecular theory of gas
Small size of gas molecule
Negligible attractive forces
K.E is directly proportional to absolute
Temperature
No effect of force of gravity on gas
molecule
Manometer → measuring pressure
Barometer → measuring atmospheric
pressure
Torriccellian barometer → common
barometer
1 atm at 273.15K = 76 cm Hg = 760 mm
of Hg = 760 torr = 101325 Pa = 10125
Nm-1 = 14.7psi
Boyle’s law Charles
law
Avogadro’s
law
V α 1/P V α T V α n
PV = K V/T = K V/n = K
P1V1 = P2 V2 V1/T1 =
V2 /T2
V1/n1 = V2
/n2
Curve line
graph
Straight
line
Straight line
New volume at x0C = y + (y/273) x
Y → volume at 00C
Ideal as equation
PV = nRT
P1V1 / T1 = P2 V2 / T2
When V=22.4 dm3 and T=247K then
R = 0.0821 dm3 atm mole-1 K-1
R = 62.4 dm3 mm of Hg mole-1 K-1
R = 62.4 dm3 torr mole-1 K-1
R = 62400 cm3 atm mole-1 K-1
For SI unit. When P=101325 N-2,
V=0.0224m3 anf T= 273K then
R = 8.1343 Nm mole-1 K-1

R = 8.1343 J mole-1 K-1
R = 1.98722 cal mole-1 K-1
1 cal = 4.184 J
1dm3 of H2 weights 0.0899 g
1dm3 of O2 weights 1.4384 g
Compressibility factor = Z = PV/nRT = 1
(for ideal gas)
For ideal gas plot for PV/nRT against P is
parallel to x-axis
Deviation from laws occurs at Low T and
high P
One molecule is 300 times of its diameter
far from another molecule
Graph of PV/nRT against P
at 00C
 H2 → increase
 N2 → little decrease then increase
 CO2 → little more decrease then
increase
at 1000C
 H2 → increase
 N2 → little more decrease from that of
170C then increase
 CO2 → very more decrease from that
of 170C decrease then increase
Ideal gas equation → PV = nRT
Real gas equation → (P+a/V2) (V-b) = RT
→ n=1
(P+an2/V2) (V-b) = RT → n=1
Plasma
 99% universe is made of plasma
 1.5 million ball of plasma in sun
 Macroscopically neutral
 Electrically conductive
 Color of plasma depends upon gas
used
5. LIQUIDS
Evaporation→ all temperature
Intramoleular forces → within molecule
covalent bond
coordinate covalent bond
Intermoleular forces → among molecules
→ Van der Waal’s forces
hydrogen bonding → H-F,O,N
dipole-dipole interaction → polar
London dispersion Forces(dipole-
induced dipole interaction) → polar +
non polar → short range forces
Covalent >Hydrogen > dipole diploe → London
dispersion
CH3F → polar →dipole-dipole → B.P 194.7 K
C2H6 → non-polar → London dispersion →B.P 184.5 K
CHCl3 → polar → dipole-dipole → 76.8 0C
CCl4 → non-polar → London-dispersion →61.2 0C
Properties of Hydrogen Bonding
 20 times weaker than covalent bond
 HF is weaker acid than HCl,HBr,HI
 Ethyl alcohol is miscible in water in
all proportion
 Sticky action of glue and honey is due
t H-bonding
 Ice has less density than water
B.P
B.P α size Straight > branch
Neon → -245.9 0C
Xenon → -107.1 0C
Fourine → gas → -
188.1 0C
Iodine → solid →
184.4 0C
Butane → 272.5 0C
2-Methyl
propane→261.3 0C
Evaporationat constant temperature,
evaporation continues at the same rate
Factors affecting evaporation
Surface area → directly
Temperature → directly
IMF → inversely
Evaporation of water is slower than ether
Vapour pressure
Pressure exerted by vapours when it is in
equilibrium with liquid
Factors affecting vapour pressure
IMF → inversely
Water is 24 mm of Hg at 250C
Ether is 537 mm of Hg at 250C
Temperature → directly
Water at 250C is 24 mm of Hg
Measurement of vapour pressure
Volumetric method
Manometric method → V.P = P + Δh
Boiling point
At B.P. K.E of molecule become maximum
Temperature remains constant at B.P
Factors affecting boiling point
IMF → directly
External pressure → varies with external
pressure
Place Pressure
(atm)
B.P(K)
Sea leve 1 373
Muree hills 0.921 371
Mount
Everst
0.425 345
Application of B.P
Pressure cooking, an example of
increased pressure
More pressure, more heat absorbed, cook
quickly
Vacuum distillation, an example of
reduced pressure
At atmospheric pressure, B.P of
CS2 → 46.30 0C
CL4 → 76.50 0C
C2H5OH → 78.26 0C
C6H6 → 80.15 0C
H2O →100 0C
CH3COOH → 118.50 0C
Normal pressure → glycerine boils at
2900C
Reduced pressure → glycerine boils at
210 0C
Advantage of vacuum distillation
Reduce time for distillation
Avoid thermal decomposition
Molar heat of fusion, ΔHf → for ice is 6 K
J/mol
Molar heat of vaporization, ΔHvap →for
water is 40.7 KJ/mol
Molar heat of sublimation, ΔHsub →for one
mole iodine 62.3 KJ/mol
Polar molecules have higher values of
molar heat, such as H2O, SO2 and NH3.
Liquid surface tension
Amount of energy required to expand
thesurface of liquid by unit area. Unit is
Jm-2 or Nm-1.
Surface tension decrease with increase in
temperature.
Measurement of surface tension
Torsion method
Capillary method
Drop method or stalagmometer meter
Surface tension = γl = (nw dl /nl dw) x γw

Viscosity
Viscosity near sides → high
Velocity near sides → low
Centre → high velocity
Centre → low viscosity
Factors affecting viscosity
Molecular size → directly
Molecular shape →irregular > regular
shape
IMF → directly
Temperature → inversely
Viscosity = ηl = (dl tl /dw tw) x ηw
Molten sulphur at 1400C with ring shape
S8 is less viscous than long chain Sn
molecule at 1900C.
Viscosity of water is more than alcohol
due to stronger hydrogen bonding.
SI unit =kg/ms
1 poise = 0.1 kg/ms
Water= 1 centipoise at 250C
Liquid crystal
Optical properties like crystalline solids
Surface tension and viscosity like liquid
Anisotropic optical behavior of solids
Stearin between 52.10C and 62.60C
Cholesteryl benzoate between 1450C and
1790C
H2O → 2H + O ΔH = 920 K J /mol
H2O(l) → H2O(g) ΔH =40.7 K J /mol
6. SOLIDS.
Sodium iodide → goiter
Table salt → food
Potassium permanganate → disinfectant
Amorphous solids are made by silicates
fusing with
Borax oxide
Aluminium oxide
Phosphorous pentaoxide
3D arrangement of particles in solid →
lattice
Types of solids
Crystalline / true solids
Amorphous solids/ super cooled liquids
Properties of crystalline solids
Symmetry
Plane of symmetry >1
Axis of symmetry >1
Centre of symmetry = 1
Geometric shape
Melting point
Cleavage plane
Habit of crystal (shape of growth)
Crystal growth (10% urea→NaCl →
needle growth)
Anisotropy
Isomorphism
NaCl3 - MgO Cubic structure
ZnO - CdS Hexagonal
KNO3 - NaNO3
- CaCO3
Rhombohedral
Polymorphism
KNO3 - AgNO3 Rhombohedral +
orthorhombic
CaCO3 Trignol +
orthorhombic
Allotropy
Suphur Rhombohedral and
monoclinic
Oxygen O2 and O3
Carbon Diamond, graphite and
bucky balls
Tin Grey tine cubic and white
tin tetragonal
Transition temperature
Sulphur 95.5 0C
Tin 13.2 0C
KON3 128.5 0C
Six parameters = 3 edges +e angles
Cubic lattices
Simple cubic lattice → p type
Body center cubic → I type
Face center cubic → f type
In NaCl , Na is attracted by 6 Cl and Cl is
attracvted by 6 Na
Types of crystalline solids
Ionic
crystal
Long range,
Never exist in liquid or
gas, Soluble in polar,
NaCl, MgO ,NaBr
Metallic
crystal
Malleable → sheets,
Ductile → wires,
Only few are soft,
Copper, iron, aluminium,
sodium, silver
Covalent
crystal
isoluble in polar
Diamond, carborundum,
silicon carbide
Molecular
crystal
Tightly packed patters
Soft.
May be:
Polar
→ sugar and ice
Non-polar
→ solidified noble gas,
CO2, S, P and I
→ polar molecular crystals have high
boiling point than non-polar.
→ ice structure → regular tetrahedron
Molecular crystals are soft and have low
melting points.
Ionic, covalent and metallic crystals are
hard and have high m.ps.
7. CHEMICAL EQUILIBRIUM
Kc is independent to initial
concentrations.
Kc change with change in temperature.
Kp = Kc(RT)Δn
Kp = Kx(P)Δn = Kx(RT/V)Δn
Kp = Kn(P/N)Δn
Reaction Quotient Qc
Qc = Kc → equilibrium
Qc > Kc → reverse reaction
Qc < Kc → forward reaction
Prediction of chemical reaction
Kc very large → reaction almost complete
Kc very small →small product is formed
Kc neither small or large→ appreciable
quantities
Optimum temperature for the formation
of
NO2 → 3000 0C
NH3 → 450 0C
SO3 → 400-5000C and 1.5-1.7 atm
Solubility product (Ksp)
→ product of molar concentrations of its
ions in the saturated solution
→ usually very small at room
temperature
→Temperature dependent
Ionic product
→ product of molar concentrations of its
ions in the saturated solution at
particular solution.
Ionic product = Ksp →saturated solution
in equilibrium with excess solid
Ionic product > Ksp →super saturated
Ionic product < Ksp → unsaturated
solution

Common ion effect
NH4Cl suppresses NH4OH
HCl suppresses NaCl
9. CHEMICAL KINETICS
Rate = k P[A]a [B]b→ here K is velocity or
rate constant
If a and b = 1 then rate=k
K has fixed value for reaction under given
T and P
When conc doubles and rate doubles →
first order
When conc doubles and rate quadruples
→ second order
Isolation method → to find order
Factors affecting rate of reaction
Nature of reactants
Conc of reactant
Particle size of solid reacting gases
Temperature
Catalyst
Theories of Reaction rate
Collision theory of reaction rate →
collision with proper orientation and
proper activation energy
Transitions
State theory
1032molecules per litre per second at
standard conditions.
Oder is experimentally determined
parameter
Under same STP,the decomposition rate
of
H →4.4 x10-3 mol/dm3-hour
N2O5 → 9.4 x 105 mol/dm3-hour
Transition state
Impossible to isolate it
Molecular weight
Intermediate distance
A definite enthalpy
Loss structure
Ability to rotate and vibrate
Greater energy than reactant and product
Catalyst for hydrogen peroxide
decomposition is MnO2
Catalyst for SO3 is NO2
Catalyst for NH3 is Fe2O3 or Mn2 O
Catalyst for C2H4 + H2 → C2H6 is Ni
Ptyalin →→ starch to sugar
Pepsin → →protein to simple molecules
13 S AND P BLOCK ELEMENTS
 Cs → superoxide
 Normal oxide; X2O + H2O → 2XOH
 Peroxide; X2O2 + 2H2O → 2XOH +
H2O2
 Superoxide; 2XO2 + 2H2O →2XOH +
H2O2 + O2
 Normal oxide; X2O +2HCl → 2XCl
+H2O
 Peroxide; X2O2 +2HCl → 2XCl
+H2O2
 Superoxide; 2XO2 + 2HCl →2XCl
+H2O2 + O2
 Polarizing power is directly
proportional to decomposition.
 Polarizing power is directly
proportional to charge and inversely
to radius.
 Hydroxides stability decreases
downward
 Sulphates stability increases down
the group.
 carbonates stabilityincreases down
the group.
 All carbonates of alkaline earth
metals are insoluble in neutral
medium while all dissolves in solids
and decomposes at red heat.
 CaSO4 is sufficiently soluble in water.
 Strontium and barium sulphates are
almost insoluble.
 Radius → Increases(F<Cl<Br<I)
 E.N → Decreases (F>Cl>Br>I)
 E.A → Decreases (Cl>Br>F>I)
 M.P and B.P → Increases (F>Cl>Br>I)
 Bond enthalpy; (Cl>Br>F>I)
 Bond enthalpy for hydrogen halides;
(HF>HCl>HBr>HI)
 Oxidizing agent; F2 > Cl2 > Br2 > I2
 Reducing agent; I > Br > Cl > F
 HF is weaker acid than HCL, HBR and
HI.
 Br reduces sulphur in sulphuric acid
from +6 oxidation state to +4.
 I reduce sulphur in suphuric acid
from +6 to -2 oxidation state.
Simple oxide Per oxide
Be → -
Mg → -
Ca → -
Sr → →
Ba → →
Ra → -
 SO2 is colourless.
 PCl3 → colourless fuming liquid
 PCl5 → straw colour solid
 S2Cl2 → orange, foul smelling liquid
 Metal oxides → Basic →Ionic
 Non-Metal oxides → Acidic →
Covalent
 Li → normal oxide
 Na → normal and peroxide
 K → peroxide and superoxide
 Rb →superoxide
 Li → Red
 Na → Yellow
 K → Lilac
 Rb → Red
 Cs → Blue/violet
 Magnesium sources
 Sea water
 Underground brines
 Mineral dolomite
 Magnesite (MgCO3)
 Calcium sources
 Sea shell (CaCO3)
 Gypsum (CaSO4.2H2O)
 CO2 → acidic
 Si O2 → acidic
 Ge O2 → amphoteric
 Sn O2 → amphoteric
 Pb O2 → amphoteric
 C O → neutral
 Sn O → amphoteric
 Pb O → amphoteric
 SiO2 → White
 PbO2 → Brown
 CO2 → Gas
 Si O2 → Solid

Chap No. 14 D and F BLOCK
ELEMENTS
 D-shell; (n-1)d1-10 ns0,1,2,3
 Binding energy increase upto VIB and
then decreases.
 [Cr(NH3) 6](NO3) 3
→Hexaamminechromium (lll) nitrate
 K2 [PtCl6] → Potassium
hexachoroplatinate (lV)
 [CO(NH3) 3(NH2)3] →
Trinitrotriaminecobalt (lll)
 [CO(NH3)4(Cl2)]Cl →
Dichlorotetraaminecobalt (lll)
chloride
 Na3[Co(NO2)6] → sodium
hexachlorocobaltater (lll)
 Na3[Fe(CN)6] → sodium
hexacyanoferrate (lll)
 K3 [Fe(CN)6] → potassium
hexacyanoferrate (lll)
 Na2 [Fe(N)(CN)5] → Sodium
pentacyanonitrosylferrate (lll)
 [CO(en)2 (Cl)2] →Dichloro-Bi-
ethylenediamine cobalt (ll)
 Weakly attracted → paramagnetic
 Weakly repelled → diamagnetic
 Paramagnetic → one or more
unpaired electron
 Diamagnetic → pared electrons

Ferromagnetic → five unpaired
electrons like Fe+3 and Mn+2
 Fe → Haber process → Ammonia
synthesis
 V2O5 → Contact process → H2SO4
synthesis
 TiCl4 → polymerization of ethene to
polythene
 Ni, Pt and Pd → hydrogenation of
unsaturated hydrocarbons.
 Cu → oxidation of ethanol to
acetaldehyde.
CrO Cr2O3 CrO3
+2 +3 +6
Basic Amphoteric Acidic
Ionic Ionize to
some extent
Covalent
Chromous
salt
Chromic
compound
Oxidation Stable Reduction
 [Cu(H2O)6]+2 → blue solution
 [Cu(H2O)4(OH)2] → blue ppt
 Medium → neutral
 [Cu(H2O)6]+2 → blue solution
 [Cu(H2O)4(OH)2] → blue ppt
 6→ octahedral
 4-→tetrahedral and square planar
 5→triognal bipyramidral and square
pyramidal
 [CuCl4]2- and [CoCl] 2- → tetrahedral
ion
 Cisplatin {Pt(NH3)2Cl2} →square
planar
 Gold → Cu = 20-25 %, Au=70-75
% → 18 carat
 Brass → Cu = 60-80 %, Zn=20-40
%
 Bronze → Cu =75-90 %, Sn=10-25
%
 Steel → Fe=90-95 %, C=0.1-2 %
 Varieties of iron
 Malleable or Wrought iron → 0.1 –
0.25 %
 Steel → 0.25-2 %
 Cast or pig iron → 2-3 %
 +2 → Ferrous → pale green
 Manganese:
 Pyrolusite → magnet
 Rhodochrosite
 Franklinite
 Psilomelane
 Manganite
 Iron ores
 Red haematite → Fe2O3
 Brown haematite or limonite → 2F O2
O3 .H2 O
 Magnetite → Fe3 O4
 Copper Ores
 Malachite → CuCO3Cu(OH)2
 Azurite → 2 CuCO3Cu(OH)2
 Chalcocite → Cu2S
 Copper pyrite → CuFeS2
 Sc3+ → Colourless
 Ti3+→ Purple
 Ti4+→ Colourless
 Cr3+→ Blue
 Mn2+→ Green
 Fe3+→ Yellow
 Co2+→ Blue
 Ni2+→ Green
 Cu2+→ Blue
 Zn2+→ Colourless
 → 33-43-23-22-22
 →p-bgy-bgb
 V+2 → Violet
 V+3 → Green
 V+4 → Blue
 V+5 → Colorless
 K2Cr2O4 → yellow solution.
 K2Cr2O7 → orange solution.
 Cr+3 → green
 +3 → manganic compound
 +2 →manganous compounds
 C +3 stable and Mn+2 stable
 7 → 6 (dark green solution)→acidic
 7→6→ 4 (dark brown precipitate) →
basic
 Iron (II) salts → pale green
 Iron (III) salts → yellow or brown
 +3 → Ferric →yellow or yellow
brown
 +2 very easily oxidize to ferric ion.
 Cobalt (II) nitrate → red
 K2Cr2O7 → orange
 KCrO4 → yellow
 Nickel (II) chloride → green
 KMnO4 → red
 White titanium oxide → white paint.
 C 21(in nature) Specific gravity 7.2.
manganese 12 and 7.4
 The solution turns from dark purple
to faint pink colour at equivalence
point.
 MnO4
- + C2O4
-2+ → Mn+2 + CO2
 KMnO4 + Fe+2 → Mn+2 + Fe+3
 S2O8
2- + 2Cl- →2SO4
-2 + I2
 S2O8
2- + 2FE2+ → 2SO4
-2 + 2Fe+3
 2Fe+3 + 2I- →2Fe+2 + I2
 2Fe+3 + 2I- →2Fe+2 + I2
 S2O8
2- + 2FE2+ → 2SO4
-2 + 2Fe+3
 [Fe(H2O)6]+2 + H2O →
[Fe(H2O)5(OH)]+1 + H3O
 [Fe(H2O)6]+3 + H2O →
[Fe(H2O)5(OH)]+2 + H3O
 [Fe(H2O)6]+3 + 3OH-
→[Fe(H2O)3(OH)3]0 + 3H2O
 [Fe(H2O)6]+2 + 3OH- →[no reaction
 [Fe(H2O)6]+2 + 2NH3 →
[Fe(H2O)4(OH)2]+ 2NH4
+

 [Fe(H2O)6]+3 + 3NH3 →
[Fe(H2O)3(OH)3]+ 3NH4
+
 [Fe(H2O)4(OH)2] → Orange
 [Fe(H2O)3(OH)3] → Brown
 Fe+2 + CO3
2- → FeCO3
 2[Fe(H2O)6]+3 + 3CO3
2- →
[Fe(H2O)3(OH)3]+3CO2 + 3H2O
 If you add thiocyanate ions, SCN- to a
solution containing iron (III) ions,
you will get an intense red solution
containing the ion [Fe(SCN(H2O)5]2+.
 [Fe(H2O)6]3+ + SCN- → [Fe(SCN)
(H2O)5]2+
 +1 → Diamagnetic → colourless →
Cu2O, CuCl, CuBr
 +2 → cupric compound → Coloured →
CuO, CuF2, CuCl2, CuCO3, CuSO4.
 +3 → found in oxides → KCuO2, a
blue black solid.
 Yttrium barium copper oxide
(YBa2Cu3O7) consists of both Cu(II)
and Cu(III) centres.
 [Cu(H2O)6]+2 + 2OH-
→[Cu(H2O)4(OH)2] + 2H2O
 [Cu(H2O)6]+2 + 2NH3 →
[Cu(H2O)4(OH)2] + 2NH4
-
 [Cu(H2O)4(OH)2] +4NH3 + 2OH- +
2H2O
 Cu+3 + CO3
-2 →CuCO3
 CHAP NO.15 ORGANIC COMPOUNDS
 Organic chemistry deals with carbon
based compounds.
 The world organic means life or
living.
 In 1928, friedrich Wohler
synthesized urea from ammonium
cyanate.
 Coal occurs in rock strata in layers
called coral beds.
 Wood → Peat → Coal
 Coal is major source of aromatic
compounds.
 Petroleum means crude oil
 Petroleum is also called mineral oil/
crude oil/ liquid gold.
 Petroleum includes only crude oil in
strict sense.
 Petroleum includes both crude oil
and natural gas in common sense.
 Coal gas is also called town gas.
 Coal is converted to petroleum by
Fischer-Tropsch process.
 Some products of Biotechnology:
 reactions are slow because they
involves breaking of some bonds and
formation of new bonds.
 Many farmers in USA grow maize for
ethanol.
 The first fullerenes was discovered in
1885 by Herold Kroto, James Heath,
Seam O’Brion, Robert Curl and
Richard Smalley.
 In 2010, fullerenes were also
discovered in outer space.
 Extract sodium test or Lassaign’s
solution(L.S) is prepared by heating
the substance with sodium metal in
fusion tube till tube become hot and
after cooling its filtered.
 O.C + CuO →CO2
 CO2 + Ca(OH)2 → CaCO3 milky colour
 O.C + CuO →H2O
 H2O + CuSO4 CuSO4.5H2O blue colour
 L.S + NaOH + FeSO4 boil-cool +FeCl3 +
HCl/H2SO4 prussian blue or green
colour
 L.S +acetic acid + lead acetate black
ppt of lead sulphide
 Detection of Halogens in organic
compounds L.S + conc nitric acid
+silver nitrate solution →
 White ppt soluble in (NH3 )OH
→Chlorine
 Yellow ppr slightly soluble in (NH3
)OH →Bromine
 Deep yellow ppt insoluble in(NH3
)OH →Iodine
 Detection of Oxygen in organic
compounds
 Can’t be test directly
 Tests for oxygen containing
functional group
 Formation of water in nitrogen
atmosphere
 Combustion analysis
 Characteristics of homologous series:
 Each series have its own formula.
 Members of series have same
chemical properties.
 Series members have same method
of preparation.
 Physical properties increase with
increase in molecular mass.
not organic are :
 Carbon containing alloys
 Simple oxides of carbon
 Allotropes of carbon
 Metal carbonates
 Bicarbonates
 Carbonyls
 Cyanides
 Cyanates
 Sulfides
Organic compound sources
 Fossil remains
 Petroleum
 Natural gas
 Coal exists in different
 forms like:
 Lignite (low C %)
 Sub- Bituminous coal
 Bituminous coal
 Anthracite (high C %)
 Benzylpencilin → an antibiotic
 Insulin → A hormone
 Polyhydroxybutyrate → A
Biodegradable thermoplastics
 Renin → an enzyme Chemosensory
protein(CSP)
 Destructive distillation of coal:
 Coke → a reducing agent
 Coal tar → used for fertilizer making
 Coal gas
 Ammoniacal liquor → H+CO
Life molecules includes:
 Proteins
 Nucleic acids
 Enzymes
 Fats
 Lipids etc
 Quinonee → antimalarial
 Aspirine → cardiac disease and pain
killer
 Borneol → anti-inflammatory
 Benzyle benzoate → scabicide
 Galantamine hydrobromide →
alzheimer’s disease
 Some fullerenes are:
 C20 →C60 →C70 →C76 →C84
 Smallest is C20 but most common I
sC60.
 Alkane → CnH2n + 2
 Cycloalkanes → CnH2n
 Alkenes → CnH2n

 Alkynes → CnH2n - 2
 Alkyl → CnH2n + 1
CHAP# 16 HYDROCARBONS
 First four alkanes are colourless
gases.
 C5– C17 are colourless liquids.
 Onward from C17, alkanes are wax
like soft solids.
 Alkanes are none-polar so insoluble
in water.
 Alkanes are soluble in CCl4 and C6 H6.
 B.P of alkanes increases with
molecular weight.
 Straight chain alkane have high B.P
than isomeric branch chain
 Melting point of alkane increase with
molecular weight.
 There is no regularity in the melting
point of alkane with the no of carbon
atoms in molecule.
 The specific gravity of alkanes
normally increase with molecular
weight.
 Viscosity of alkanes increase with
increase in no of carbon atoms.
 Cyclopropane and cyclobutane are
gases while rest of cycloalkanes are
liquids.
 Melting point and boiling point of
cycloalkanes increase with increase
in no of carbon atoms.
 Cyclopropane and cyclobutane have
greater angle strain and hence
undergoes ring opening reactions.
 Cycleproapne undergo ring opening
reaction with H2/Ni and HBr to give
open chain products.
 Cyclobutane undergo ring opening
reaction under severs conditions.
 Greater no of alkyl groups attached,
greater is the stability of alkenes.
isobutylene is more stable than
1.butene
 Dehydrahion of alcohol at 1700 C in
presence of sulphuric acid gives
alkene.
 Dehydrohalogenation of alkyl halide
in presence of alcoholic solution of
KOH/NaOH gives alkene and alkyl
halide.
 The hydrogenation of alkenes is
industrially used for the conversion
of vegetable oils into ghee.
 Hydrogenation of alkene is done at
200-250o C.
 Alkenes react with H2SO4 to produce
hydrogen sulphates, which on
hydrolysis yields alcohol at 100oC.
 The bromination of alkenes provides
a useful test for the presence of
double bond.
 The colour of bromine rapidly
discharge as the colourless dibromo
compound is formed.
 Alkene react with hypohalous acids
(X-OH) to form halohydrins.
 Halohydrins are organic compound
having hydroxyl group and halogen
at adjacent carbon atom.
 Alkenes react with oxygen in the
presence of silver catalyst at
temperature 3000C to epoxides.
 Epoxides on acid hydrolysis produce
glycol.
 When ozone is passed through alkene
in presence of inert solvent like CCl4
to form ozonide.
 Ozonide are being explosive cannot
be isolated.
 Ozonide on treatment with Zn and
water cleavage at position of double
bond to form carbonyl compounds.
 Ozonolysis is done for locating the
position of the double bond in
unknown alkene.
 The process by which simple
molecules chemically join together to
form large molecules with high
molecular weight, is called
polymerization.
 Polyethylene are also known as
polyethene.
 The temperature for polymerization
of ethane is 100-3000C with pressure
of 1K-2K atm.
 Nicol prism made of calcite, CaCO3 act
as polarizer.
 Optical isomerism is type of
isomerism in which the isomer differ
in their interaction towards plane
polarized light.
 Two mirror images of single
compound that cannot be
superimpose are called enantiomers
of each other
 Ethyne, propyne and butyne are
gases.
 C5 – C12 alkynes are liquids and
higher are solids at room
temperature and pressure.
 All alkynes are odourless and
colourless except acetyls.
 Acetylene has garlic like odour.
 Alkanes are insoluble in water.
 Alkynes are slightly soluble in water.
 The boiling point of alkynes increase
with increase in no of carbon atoms.
 The melting point of alkynes did not
show regular pattern
 Alkenes are more denser than alkane
and alkene
 Alkynes are less reactive than alkene.
 Terminal alkyne and acetylene are
acidic in nature.
 If acetylene or terminal alkyne is
treated with solution of sodium
amide (NaNH2) in liquid ammonia,
sodium acetylide is obtained.
 Acetylene and 1-alkyne react with
ammoniacal solutions of cuprous
chloride and silver nitrate to form
acetylides and alkynides of these
metals.
 Copper and cilver acetylides are
highly explosives in dry conditions.
They are decomposed by acids such
as HNO3 to regenerate acetylene.
 None-terminal alkynes can be
distinguisher from terminal alkynes
by Cu2Cl2 and NH4 OH or Ag(NO3 )2
and NH4 OH
 Addition reactions of alkynes are:
 Hydrogenation → alkenes
 Reduction by dissolving metal →
alkynides/ acetylides
 Hydrohalogenation → haloalkane
 Hydration → carbonyl compounds
 Halogination → tetrahaloalkane
 Ozonolysis → ozonides →ketones →
carboxylic acids
 Electrophilic aromatic substitution
reaction:
 Nitration → nitrobenzene
 Sulphonation →benzene sulphonic
acid
 Halogenation → halobenzene
 Friedel-crafts’s acylation → alkyl
benzene

 Friedel craft’s acylation → aromatic
ketones
 Hydrogenation of alkynes leads to
alkene and onward to form alkane
the reaction is stopped by poisoning
Pd catalyst with BaSO4 + quinolone
(lindlar’s catalyst).
 1-alkynes and terminal alkynes react
with metals in liquid ammonia to
form salts like alkynides or
acetylides.
 Alkynes react with ozone to form
ozonide.
 Ozonide may be decomposed by
water to give ketones.
 Ketones are foxidized by H2O2 to
form carbonyl compounds.
 Bezene is colourless liquid at room
temperature and pressure.
 Benzene has pecular smell and
burning tastes.
 The specific gravity of benzene
is0.8788.
 M=benzene melts at 5.50C and boils
at 80.20C.
 Benzene is highly inflammable.
 The representation of real structure
as a weighted average of two or more
contributing structures is called
resonance
 The hybridization of C in benzene is
sp2
 Benezene have 6 CC and6 CH sigma
bonds.
 The resonance energy of benzene is
152 kJ/ mol.
 *(320-208)due to unusual stability,
benzene does not give addition
reactions like those of alkenes.
 Benzene prefers to undergo
electrophilic substitution reactions
rather than additions reactions.
 Benzene is less reactive than alkene.
 Benzene react with hydrogen in the
presence of Ni or Pt catalyst at 1500C,
under high pressure to form
cycolohexane.
 Benzene react with chlorine or
bromine in the presence of ultralight
to form hexachloride.
 Benzene reacts with concentrated
nitric acid in the presence of
concentrated sulphuric acid at 60oC
to form nitrobenzene.
 An electrophile NO2+ is produced by
reaction of H2 SO4 and HNO3.
 Benzene react with concentrated H2
SO4 at 1200C or fuming H2 SO4 at
room temperature to give benzene
sulphonic acid.
 Fuming sulphuric acid is
concentrated sulphuric acid in which
SO3 has been dissolved.
 Treatment of benzene with n-propyl
chloride gives isopropyl benzene
rather than the expected n-propyl
benzene.
 Benzene reacts with alkyl halides in
the presence of AlCl3 to form alkyl
benzenes.
 Benzene reacts with acid halides in
the presence of a lewis acid catalyst
(AlCl3) to give aromatic ketones.
 Effects of substitution of benzene:
 Ortho/para directing groups are
activators except halogens.
 Meta directing groups are
deactivators.
 When phenol is nitrated, the reaction
yield only the n-nitrophenol and
pnitrophenol in ratio of 53% and
47%.
 Using nitrated mixture (conc HNO3 +
conc H2 SO4), benzene can be nitrated
at 600C to form nitrobenzene.
 Dinitrobenzene is obtained if
reaction is carried at 1000C.
 Trinitrobenzebe is obtained by using
mixture of fuming nitric acid and
sulphuric acid at 1000C.
 Trinitrotoluene → explosives
 C-C→1.540A
 C=C→1.340A
 C≡C→1.190A
 C-H→1.090A
 C→E.N→2.5
 H→E.N→2.1
 Cl2 is taken in excess→ CCl4.
 Cl2 is taken in limited→ CH3Cl.
 methane→ 890.95 kJ / mol.
 ethane →1559 kJ /mol.
 sufficient O2 gives CO2
 limited O2 gives CO
 very limited O2 gives C
 1-butene →2719 kJ/mol
 Cis-2-butene →2712 kJ/mol
 Trans-2-butene →2707 kJ/mol
 Isobutylene → 2703 kJ/mol
 *(BSTI)
 in 1,3-Butadiene
 C-C→ 1.48oA not 1.54 oA.
 C=C→1.37oA not 1.33 oA
 RIGHT OR clockwise →
dextrorotatory or +
 left or anticlockwise → Levorotatory
ior –
 Copper acetylide → red ppt *(CAR)
 Silver acetylide → white ppt *(SAW)
 Toluene → CH3
 Phenol →OH
 Alinine →NH2
 Nezoic acid →COOH
 O-xylene, m-xylene, p-xylene → 2 CH3
 Catechol, resorcinol, hydroquinonee
→ 2OH
 Mesitylene → 3 CH3
 Durene → 4 CH3
 Nepthalane → 2 benzene
 Anthracene → 3 benzene
Ortho/para directing substituents:
 OR
 OH
 NH2
 NR2
 NHR
 CH3
 C2 H5
 C6 H5
 halogens
 *(R groups)
Meta directing substituents:
 CR
 COR
 CH
 COH
 CN
 NO2
 SO3 H

 Cyclohexene evolves 120 kJ /mol
 1,3-cyclohexadiene gives 232 kJ/ mol
 1,3,5-cyclohexartiene give 208 kL/
mol

 The main types of reactions of
benzene are:
 Addition reactions
 Electrophilic substitution reactions
 Oxidation reactions
 Electrophilic aromatic substitution
reaction:
 Nitration → nitrobenzene
 Sulphonation →benzene sulphonic
acid
 Halogenation → halobenzene
 Friedel-crafts’s acylation → alkyl
benzene
 Friedel craft’s acylation → aromatic
ketones
 Alkane
→sp3→tetrahedral→PARAFFIN
 Alkene→sp2→planar→OLEFIN
 Alkyne→ sp1→linear→ACETYLENE
 Chapter No.17 Alkyl halide
 Methyl and ethyl halides are gases at
room temperature.
 Alkyl halide upto C18 are colourless
liquids.
 Alkyl halides are water insoluble.
 Alkyl halides have high boiling point
than corresponding alkane.
 For a given alkyl group, the boiling
point increase with increase of size of
halogen atom.
 For given halogen atom the boiling
point increase with increasing size of
alkyl group.
 Reaction of halogen acids with
alcohol gives alkyl halide and water.
 By the action of phosphorous
trihalides on alcohol, alkyl halides are
obtained.
 Phosphorous trihalides are produced
in situ by the action of red
phosphorus on halogen.
 By the action of thionyl chloride on
alcohol, alkyl halides are produced
along with HCl.
 Pyridine being base absorbs HCl after
its production.
 Alkanes react with halogens in
presence of uv light or at 4000C to
yield alkyl group.
 Order of strength of C-X bonds,
 → C-F>C-Cl>C-Br>C-I
 Order of reactivity of alkyl halide,
 → R-F<R-Cl<R-Br<R-I
 Greater the no of alkyl groups,
greater is the stability of the
carbocation.
 Tertiary carbocation is more stable
than secondary and primary.
 Base has a species that have affinity
for proton.
 Nucleophile has the ability to form
bond with carbon atom.
 A base attack hydrogen atom in the
elimination reaction.
 A nucleophile attacks carbon atom in
the substitution reactions.
 Tertiary alkyl halide →(SN1).
 Primary alkyl halides → (SN2).
 Unimolecular substitution (SN1→
tertiary alkyl halide →polar solvent.
 Bimolecular substitution (SN2) →
primary alkyl halide→ none-polar
solvent.
 polar →secondary alkyl halide →SN1.
 none-polar → SN2.
 Elimination reaction takes place in
the presence of base.
 E1 → tertiary alkyl halide.
 E2 → primary alkyl halide.
 E1 → double step
 E2 → single step
 A stronger base will favor in
elimination.
 A stronger nucleophile will favor
substitution.
 Ethoxide is strong base.
 Anion of thioalcohol (C2H5S-) is
strong nucleophile.
 Crowding within molecules of
substrate also generally favors
elimination over substitution
reaction.
 Alkyl groups stabilizes alkene more
than the substitution product.
 All those organic compound that
contain at least one carbon metal
bond are called organometallic
compound.
 Alkyl or aryl magnesium halides are
commonly known as Grignard
Reagents.
 The general formula of Grignard
reagents are R-Mg-X.
 Grignard reagents are prepared by
action of alkyl or aryl halide on
freshly prepared magnesium metal in
the presence of anhydrous or dry
ether.
 Grignard reagents cannot be isolated,
therefore, it’s ethereal solution is
directly used in the synthetic
reactions.
 Increasing size of alkyl or aryl group
make the formation of Grignard
reagents difficult. → I > Br > Cl
 Alkyl or aryl magnesium fluorides are
not known.
 akylbromides are most suitable for
preparation of Grignard reagents
because alkyliodides are expensive.
 Characteristics reactions of Grignard
reagents are
 nucleophilic substitution and
nucleophilic addition reactions
 Formaldehyde on reaction with
Grignard reagents gives primary
alcohol.
 Higher aldehyde on reaction with
Grignard reagents gives secondary
alcohol.
 Ketones on reaction with Grignard
reagents gives tertiary alcohol.
 Grignard reagents react with esters
to form carbonyl compounds.
 Grignard reagents react with ethyl
formate gives secondary alcohol at
the end.
 Grignard reagents on reaction with
ethyl acetate to give tertiary alcohol.
 Grignard reagents reacts with CO2
and forms carboxylic acids.
 Hemoglobin have iron while plants
chlorophyll has Magnesium.
 Amines are important nitrogen
containing organic compounds.
 Amines are derivatives of NH3 in
which one or more hydrogen group is
replaced by one or more similar or
different alkyl groups.
 The functional group of amine may
be: NH2 NH N

 On basis of number of alkyl groups
directly bonded to nitrogen atom,
amines may be primary, secondary
and tertiary amines.
 Amines in common system are
named as “Alkyl amine”.
 In IUPAC naming, alkynes are named
as substituent attached to alkane and
named as “n-Aminoalkane”.
 Lower molecular weight amines are
generally gases or lower boiling
liquids at room temperature.
 Amines has ammonia like smell.
 Amines have high boiling point than
alkane due to hydrogen bonding.
 All primary secondary and tertiary
amines have hydrogen bonding with
water molecules.
 Only primary and secondary amine
are able to form hydrogen bonding
among their molecule.
 Molecules of tertiary amine can’t
form hydrogen bonding so its boiling
point is lower than other.
 Boiling point trend: ter < sec < pri
 Amines have trigonal pyramidal
shape.
 Amines are basic in nature due to
lone pain of electron on nitrogen.
 Amine react with acids to form salts.
 Basidity is directly proportional to no
of alkyl groups.
 Order of basidity: R3N > R2NH > RNH2
> NH3
 When R-X is heated with alcoholic
NH3, it yields a mixture of primary,
secondary and tertiary amines and
quaternary ammonium salt.
 The reaction of R-X with alcoholic
NH3 is also called alkylation of
ammonia.
 Primary amines are prepared by the
reduction of nitro alkanes(R-NO2) in
the presence of “Pt/Pd/Ni or lithium
aluminum hydride (LiAlH4) in ether.
 When nitriles or alkyl cyanides(R-
CN) are reduced they yield the
corresponding primary amines.
 Primary amines are obtained when
simple amides are reduced by lithium
aluminum hydride in water.
 On the basis of lone pair on nitrogen,
Amines act as nucleophilic reagent.
 The lone pair of nitrogen in amine is
available to the electron deficient
reagents called electrophiles.
 When primary amines are treated
with alkyl halides, they produce a
mixture of secondary, tertiary amines
and quaternary ammonium salts.
 Primary amine react with aldehyde
and ketones yielding condensation
products called imines.
 Imines are also called Schiff’s bases.
 Primary amines react with acid
chloride or acid anhydride to
produce Nsubstituent amides.
 Secondary amines react with acid
chloride to N,N-disubtituted amides.
 Tertiary amine have no directly
attached hydrogen therefore they do
not react with acid chloride to
produce amides.
 When primary aliphatic amines are
treated with nitrous acid, they yield
highly unstable diazonium salt.
 Nitrous acid being unstable acid is
prepared in situ by the reaction of
NaNO2 and dil HCl.
 CHAP# 18 (A) ALCOHOLS
 Dihydric alcohol (diols) are usually
called glycols because of sweet taste.
 Lower alcohols are colourless, toxic
liquids.
 Alcohols have characteristics sweet
smell.
 Boiling point of alcohol is higher than
alkane due to hydrogen bonding.
 Boiling point of alcohol increase
regularly with the increase in the
number of carbon atoms.
 Lower alcohols C1-C4 are completely
soluble in water in all proportions.
 Angle in water is 109.50C and Angle
in alcohol is 1090C.
 Order of acidity of alcohols: pri > sec
> ter
 R-OH + NaOH → No Reaction.
 Alkenes react with concentrated
sulphuric acid to produce alkyl
hydrogen sulphates, which on
hydrolysis yields alcohols.
 Alcohols can be prepared by
hydrolysis of alkyl halides by means
of water or an aqueous alkali.
 Reaction with RMGX
 Formaldehy→primary alcohol.
 Aldehyde react → secondary alcohol.
 Ketones → tertiary alcohols.
 Reduction of aldehyde and ketone to
alcohol is done at 2000C and 10atm.
 Formate esters on reaction with
Grignard reagent secondary alcohol
while other esters form tertiary
alcohol.
 Both carboxylic acid and esters can
be reduced to primary alcohols with
Li Al H4.
 Carboxylic acid cannot be reduced
with H2/Ni or Na+C2H5 -OH
 The reactions of alcohol may be
substitution reaction or elimination
reactions.
 Alcohols react with halogen acids to
form corresponding alkyl halides.
 Order of reactivity: HI > HBr > HCl
→→ Ter > sec > pri (alcohols)
 HCL react only in the presence of
catalyst (anhydrous ZnCl2).
 Lucas test: in lucas test, alcohols are
treated with a solution of HCl and
ZnCl2 to form alkyl halides.
 Tertiary alcohols → immediately.
 Secondary alcohols → slower.
 Primary alcohols →slowly.
 Alcohols react with thionyl chloride
to form alkyl chlorides.
 Phosphorus Trihalides also form
alkyl halides with alcohols.
 Alcohols when treated with
concentrated sulphuric acid at 1700C
undergo dehydration to form
alkenes.
 Alcohols react with carboxylic acid to
form esters (RCOOR). This process is
called Esterification.
 In Esterification, concentrated H2SO4
is used as catalyst.
 The reaction of Esterification is
reversible

 Using strong oxidizing agent such as
“Na2Cr2O7 + H2SO4” or “KMnO4 +
H2SO4”, alcohols can oxidized to
carbonyl compounds and finally to
acids.
 Primary alcohols are first oxidized to
aldehydes and then to acids.
 Secondary alcohols are first oxidized
to ketone and then to carboxylic
acids
 Tertiary alcohols are stable to
oxidation under normal conditions.
 Ethylene glycol when treated with
acidic KMnO4 or K2Cr2O7 results in
the formation of formic acid.
 Ethylene glycol when treated with
periodic acid (HIO4) or lead tetra
acetate ((C2H5COO)4 Pb) , ethylene
glycol gives formaldehyde.
 The sulphur analogues of alcohols are
called Thiols and are called Thiols or
alkyl hydrogen sulphdes or
Mercaptans.
 The functional group of thiols is –SH.
 Thiols react with insoluble salts and
hence named is Mercaptans.
 Methanthiol is gas while ethanthiol
and higher members are colourless,
volatile liquids at STP.
 methanthiol and ethanthiol are
added to natural gas in minute
amounts to make gas leakage
detectable by smell.
 Thiols have lower boiling point than
alcohol due to lack of hydrogen
bonding.
 Thiols are insoluble in water.
 CHAP# 18 (A) PHENOL AND
ETHER
 The world phenol is used for specific
compound “hydroxyl benzene”.
 Phenols are usually named as
derivatives of the parent phenol
(C2H5OH).
 The C-O-H angle in phenol is 109.50.
 The C-O-H angle in methanol is
108.50.
 In phenol the cix carbon atoms are
sp2 hybridized and internal angle is
1200.
 The C-O bond in phenol is slightly
shorter than that of methanol.
 The C-O bond length in phenol is
1.360A and in methanol is 1.42o.
 Phenol are colourless liquids or low
melting crystalline solids at room
temperature.
 Phenols have characteristics odour.
 The vapours of phenol is itself toxic.
 The boiling point of phenol is slightly
higher than that of alcohol due to
strong hydrogen bonding.
 Phenol are more soluble than alcohol
in water.
 Above 65oC, phenol and water are
completely soluble.
 The liquid phenol containing 5% of
water is known as carbolic acid. 421.
 Carbolic acid is used as disinfectant
and germicide.
 Acidity: Carboxylic acid > water >
phenol > alcohol
Pka Ka
Carboxylic acid 5 10-5
Water 7 10-7
Phenol 10 10-10
Alcohol 46-18 10-16 -
01-18
 Being acidic, phenol react with NaOH
or Na metal to form salt (Ar-ONa).
 Phenol can be prepared from bezene
sulphonic acid.
 The sodium phenoxide is treated
with dilute HCl to form phenol.
 Chlorobenzene is hydrolysed with
aqueous NaOH at high temperature
and pressure to form phenol. this
process was developed by Dow
company of USA in 1928.
 Cumene is also called isopropyl
benzene.
 A solution of benzenediazonium
chloride is warmed on a water bath
at 500C
 Benzenediazonium chloride is
prepared from aniline.
 Phenol exhibit two types of reactions:
 Reaction due to hydroxyl group
 Reaction due to aromatic ring.
 The presence of OH group in phenol
increase the reactivity of phenol.
 Phenol react with bromine water or
aqueous bromine to give ppt of 2,4,6
tribromophenol.
 Chlorine react with phenol and
forms:
 O-Bromophenol 15%
 P-Bromophenol 85%
 With dilute HNO3, phenol reacts to
form ortho and para nitrophenool.
 2,4,6-Trinitrophenol is also known as
Picric acid.
 Phenol being acidic in nature, react
with sodium metal to form salt with
the release of H2 gas.
 Phenol undergo oxidation with air
(O2) or chromic acid (CrO3) to form
pBenzoquinonee.
 Compounds which contain a hydroxyl
group in side chain attached to an
aromatic ring are not phenols, they
are called aromatic alcohols.
 The first person to demonstrate
ether’s use as anesthetic was Dr.
Morton in 1896.
 The common home disinfectant is
chlorine bleach.
 Chlorine bleach, a 5% solution of
sodium hypochlorite.
 Antiseptics are antimicrobial
substances that are applied to living
tissues to reduce possibility of
infection.
 kill bacteria→bactericidal.
 kill fungi → fungicidal.
 kill bacteria spores →sporicidal.
 kill viruses → virucidal.
 CHAP# 19 CARBONYL COMPOUNDS
1: ALDEHYDES AND KETONES
 Common system naming of Aldehyde:
their names derived from carboxylic
acid containing same carbon atom
but “ic acid” is replaced by aldehyde.
 In IUPAC naming of aldehyde “e” of
alkane is replaced by “al”.
 In common naming system, ketone
are named as “alkyl ketone”.
 If two same groups are attached to
carbonyl carbon, it’s called
symmetrical ketone.
 If two different groups are attached
to carbon, it’s called unsymmetrical
carbon.

 Formaldehyde is gas at room
temperature while other aldehyde
are colourless liquids.
 Acetone, the simplest ketone is liquid
is room temperature with pleasant
odour
 All the members of ketones are
colourless liquids except acetone.
 Lower members of aldehyde and
ketones upto C4 are water soluble.
 Their solubility decreases as the size
of the molecules increase.
 The most soluble in water is
fomaldehyde.
 Carbonyl compounds do not form
hydrogen bonding with each other.
 Carbonyl compounds form hydrogen
bonding with water molecule due to
oxygen.
 Order of boiling point: alkane/ether <
aldehyde/ketone < Alcohol
 The boiling point of aldehyde and
ketone increase with increase in the
molecular weight. Ethanol >
Methanal.
 The carbon and oxygen of carbonyl
group are sp2 hybridized.
 The length of CO single bond is
1.430A.
 The length of CO double bond is
1.230A.
 Ozone react vigorously with alkene
and form ozonide which is unstable.
 Ozonide is reduced directly to
aldehydes and ketones by zinc and
water. This reaction is called
ozonolysis.
 Water adds to alkene in presence of
mercuric sulphate and sulphuric acid
to form enol which is unstable,.
 The enol intermediate undergoes
arrangement to form aldehydes and
ketones depending on starting alkyne
used.
 Friedel-Crafts acylation of aromatic
gives aromatic ketones.
 When benzene is treated in the
presence of Lewis acid, AlCl3 with
acid halide, an aromatic ketone is
produced.
 The aldehydes and ketones
undergoes addition reactions as
compared to alkenes.
 The presence of base increase the
nucleophilic character of the reagent.
 The presence of acid increase the
electrophilic character of the
carbonyl carbon atom inducing more
positive charge on it and thus
enhances its ability to be attacked by
weak nucleophiles.
 Carbonyl compounds are weak Lewis
bases which can be protonated.
 In addition reaction of carbonyl atom
its geometry changes from trigonal to
tetrahedral as its changes from sp2 to
sp3.
 An acid catalyzed reaction will take
place with weak nucleophile.
 A base catalyzed addition reaction
will take place with strong
nucleophile. 505. Ketones are less
reactive than aldehyde
 Ketone is less reactive than aldehyde
due to Steric Hindrance and
Electronic effect
 An alkyl group neutralize positive
charge on carbonyl atom decreasing
its reactivity towards nucleophile.
 Aldehydes and ketones are reduced
to saturated hydrocarbons by:
 Clemmenson Reduction: by using
Zinc amalgam and conc HCl Wolf-
Kishner method: by using hydrazine.
 Aldehydes and ketones are reduced
to alkanes in the presence of Zinc
amalgam and HCl as reducing agent.
 When aldehyde or ketone is treated
with hydrazine (NH2NH2),
ahydrazone is obtained.
 A hydrazine on heating with KOH in
boiling with ethylene glycol gives
corresponding alkanes.
 In Clemmenson-reducation and Wolf
Kishner reducation, alkane is
produced at the end.
 Aldehydes and ketones are easily
reduced to primary and secondary
alcohols respectively by using metal
hydrides as reducing agents.
 The most common metals hydrides
are Lithium aluminum hydride
(LiAlH4) and sodium borohydride
(NaBH4).
 Reduction of aldehydes and ketones
by using hydrocyanic acid is done in
basic medium.
 Acetophenonee react with
hydrocyanic acid to form
acetophenonee cyanohydrin.
 Aldehydes react with ammonia to
form solid aldehyde ammonia.
 Some important ammonia
derivatives are:
 Alkyl amine R-(NH2)
 Hydroxyl amine (NH2OH)
 Hydrazine (NH2NH2)
 Phenyl hydrazine (C6H5NHNH2)
 Primary amine react with aldehyde
and ketone to form unstable
compound which losses water to
form product with CH double bond,
called imines.
 Aldehydes and ketones form oxime
on reaction with hydroxyl amine.
 Aldehyde and ketone react with
hydrazine to form hydrazine.
 Alcohols are weak nucleophile, an
acid catalyst (H2SO4) is used.
 Hemiacetal contain both alcohol and
ether is functional group.
 Acetal have two ether functional
group.
 Aldehydes are more easily oxidized
than ketones.
 The hydrogen atom attached to
carbonyl group in aldehyde is
oxidized to OH group. (RCHO →
RCOOH)
 Aldehydes can be oxidized by much
milder oxidizing agent such as:
 Tollen’s Reagent
 Fehling’s Solution
 Benedict’s Solution
 The Tollen’s Reagent is ammonium
silver nitrate (2Ag (NH3)2 OH).
 Tollen’s reagent reaction is also
called mirror test.
 Ammonioum carboxylate is formed
by the reaction of aldehyde with
ammonical silver nitrate.

 The fehling’s solution is 2Cu(OH)2 +
NaOH.
 If aldehyde react with Fehling’s
solution, the deep blue colour of
cupric ion is reduced to ret ppt of
cupric oxide.
 Oxidation by Fehling solution is used
widely for the estimation of glucose
in blood and urine.
 Ketones having hydrogen attached to
alkyl group or also called alpha
carbon can be oxidized in the
presence of K2Cr2O7 / H2SO4, KMnO4 /
H2 SO4 / conc HNO3 etc which
involves breaking C-C bondin case of
unsymmetrical ketone, the carbonyl
group remain smaller alkyl group
 Aliphatic carboxylic acids are also
commonly called fatty acids because
esters of several higher members are
fats.
 dicarboxylic acids → dioic acids
(IUPAC).
 aliphatic acids C1-C10 are liquids with
distinctive odours.
 Higher members of acid homologous
series are wax-like solids.
 Acetic acid → 4-5% of vinegar
 Butyric acid i→rancid butter.
 Anhydrous ethanoic acid freezes at
170C to form a solid which look like
ice.It is, therefore also known as
glacial acetic acid.
 Carboxylic acid are more polar than
alcohol.
 Solubility of carboxylic acid in water
decreases as their relative molecular
mass increases.
 Structural features formic acid.
 CO double bond is 120 pm.
 CO single bond is 134 pm.
 angles of H-C-O in carboxylic acids is
1110A.
 angles of H-C=O in carboxylic acids is
1240A.
 angle of O-C=C in carboxylic acids is
1250A.
 The hybridization of hydroxyl oxygen
in carboxylic acid is sp2.
 Lone pair from hydroxyl oxygen
makes the carbonyl group less
electrophilic than that of aldehyde
and ketone.
 Any electron withdrawing
substituent increase the acidity of
acid.
 Any electron donating group
decrease the acidity of the acid.
 Reaction of carbon dioxide with
Grignard reagent is known as
carboxylation of Grignard reagent.
(RMgX+CO2→acid)
 The reaction that provide an
extension to length of carbon chain is
reaction of CO2 with R-Mg-X.
 Compounds having cyanide (–CN)
group are called alkyl nitriles or alkyl
cyanides.
 The carbon nitrogen triple bond in
alkyl nitriles can be hydrolyzed to
carboxylic acid in aqueous acid
medium.
 Primary alcohol can be oxidized to
carboxylic acids by oxidizing agents
like acidified potassium
permanganate or potassium
dichromate etc.
 on oxidation, Primary alcohol
→aldehyde →carboxylic acid.
 Oxidation of aldehydes in the
presence of oxidizing agents like
KMnO4, K2Cr2O7 or Ag2O gives
carboxylic acid with the same
number of carbon atoms.
 Aromatic carboxylic acids can be
prepared by the oxidation of aliphatic
side chain (alky group) present on
the benzene ring, with oxidizing
agent like KMnO4, K2Cr2O7.
 any side chain is converted to
carboxyl group.
 In oxidation of alkyl benzene, the
methyl group is oxidized not the
aromatic ring, this show the striking
stability of aromatic rings towards
oxidizing agents.
 Carbon atom of carboxylic acid is less
positive than that of aldehyde and
ketones so it did not undergoes
addition or condensation reactions
like that of aldehyde and ketone.
 The OH donate electron to CO in
carboxylic acid and hence reduce its
partial positive charge so it I not
attacked by nucleophiles as
compared to aldehyde and ketone.
 The most reactive derivatives of
carboxylic acid derivatives are alkyl
halide.
 Acyl chloride→ most common + less
expensive than bromides and iodides.
 Alkyl chlorides can be prepared by
the reaction of acids with thionyl
chlorides or phosphorus
pentachloride (PCl5).
 Acid anhydrides are derived formic
aids by removing water from two
carboxylic acid molecules.
 Naming of acid anhydrides: the name
of acid of carboxylic is replaced by
anhydride like carboxylic acid →
carboxylic anhydride.
 important and commercially
available anhydride are acetic
anhydride or ethanoic anhydride.
 The dehydrating agent is P2O5.
 While naming ester the R part of OR
is named first and then followed by
the name of the acids, where by “ic
acid“ is replaced by “ate”.
 When a carboxylic acid and alcohol
are heated in the presence of acid
catalyst, equilibrium is established
with the formation of ester and
water.
 Esterification is reaction of an acid
with alcohol.
 Ethyl acetate is important ester
which can be prepared by the
reaction of acetic acid with ethanol.
 Esters can also be prepared by the
reaction of an alcohol with acid
halides or acid anhydride.
 Amides are named by replacing “ic
acid” corresponding acid by word
“amide”.
 Amides can be prepared by the
reaction of ammonia with carboxylic
acid to form first ammonium salts
which on heating produces acid
amides.
 Amides can also be prepared by the
reaction of ammonia with ester or
acetyl chloride.
 Order of acid derivatives towards
nucleophile are:
 Acylhalide > acid anhydride > ester >
amide >Nitrile *(Cl-O-OR-NH2-CN)
 Acylhalide, acid anhydride, ester,
amide and nitrile on hydrolysis yield
corresponding carboxylic acid.
 Acylhalide and acid anhydride on
reaction with alcohol yield ester.

 Acylhalide, acid anhydride and ester
on ammonolysis yield an amide.
 Carboxylic acid can be reduced to the
corresponding alcohols using lithium
tetrahydrido aluminate in dry ethoxy
ethane.
 The removal of carbon dioxide from a
carboxylic acid takes place is known
as decarboxylation.
 Decarboxylation of carboxylic acid
takes place when its sodium salt is
heated with soda lime to form
alkanes.
 Soda lime is dry mixture of caustic
soda, NaOH and quick lime CaO.
 Ascorbic acid occur naturally in fruit,
used as preservatives.
 Ascorbic acid inhibits fungal growth
but allow bacterial growth.
 Benzoic acid and sodium benzoate
have inhibitory effect on the growth
of yeast.
 The tartness in lemon is due to
carboxylic acid.
 Oranges have citric acid.
 Acetic acid present in vinegar is
responsible for giving sour taste.
 Malic acid found in unripe fruit gives
these fruit a sour or tart taste.
 Acyl group react with benzene in the
presence of lewis acid to form
aromatic ketones.
 On hydrolysis, anhydrides form
corresponding carboxylic acid.
 Hydrolysis of esters is called
saponification.
 Hydrolysis of esters is used to make
soaps from fats.
 Ester can be reduced to primary
alcohol in the presence of reducing
agent in ether which is used as
solvent.
 Ester reacts with two equivalent of
Grignard Reagent to give tertiary
alcohol.
 Ester react with R-Mg-X to form
ketone, ketone react with another R-
Mg-X to form tertiary alcohols.
 Amides on hydrolysis form the
corresponding carboxylic acids. This
reaction is slow and requires acid or
base as catalyst.
 Amides can be reduced to primary
amine in presence of aluminum
hydride.
 Alkyl nitriles or simply nitriles are
also considered as derivatives of
carboxylic acid.
 Alkyl nitriles can be obtained from
carboxylic acids, though they do not
contain acyl group.
 On boiling with a dilute minerals acid
or dilute alkali, nitriles are
hydrolyzed forming carboxylic acids.
 Alkyl cyanide when treated with a
reducing agent such as sodium and
ethanol or lithium aluminium
hydride (lithium tetra
hydrioaluminate(III) in
ethoxyethane, itriles are primary
reduced to amines.
 Nitriles on reaction with Grignard
reagent produce Ketones
Physics Formulas
Motion and Force
 v = s/t a = v/t
 vf = vi +at s = vit + ½ at2
 2as = vf2 – vi2 S = vave x t
 Vave =( vi + vf )/2
 g = 9.8 ms-2 = 32 ft-2
 F = ma →a = v/t → P = mv = P = F t
 Impulse; J = F x t = ∆P
 J = ∆P
 Law of conservation of momentum; ∆p = 0
 Elastic collision in one dimension; [v1 + v2] = [v1’+
v2’]
 Magnitude of projectile velocity; Vf = √(v_fx^2+〖
v_fy^2〗^ )
 Height of projectile; H = vi2sin2θ/2g
 Time of flight; T = 2 vi sinθ/g
 Time of summit or time to reach to highest point; T
= vi sinθ/g
 Range; R = vi2 sin 2θ/g
 Rmax = vi2/g
 R = Rmax at 450
 Work and Energy
 W = Fd cosθ
 Power; p=W/t or p =Fv
 1 watt = Js-1
 1 hp = 746 watts
 K.E = ½ mv2
 P.E = mgh
 Efficiency = output/input = W x D/P x d
 Circular motion
 Absolute potential energy =Fr = - GmMe/Re (-
because work is done against gravity)
 Gravitational potential = E/m = GMe/Re
 For escape velocity compare K.E with Absolute
potential energy; vesc = √(〖2GM〗_e/r_e ) → vesc
= √(〖2gr〗_e )
 G = 6.67 x 10-11 Nm2kg-2
 Re = 6.4 x 106 m
 Me = 6 x 1024 kg
 Vesc = 11.2 x 103 ms-1
 Wh = K.E + fh → (Wh = loss in potential energy)
 Loss in P.E = Gain inn K.E + work done against
friction
 E = mc2 →(c= 3 x 108 ms-1)
 Rotational and circular motion
 Angular velocity; ω = ∆θ/∆t
 Angular acceleration; α = ∆ω/∆t → a = α x r
 v = r ω
 Fc = mv2/r
 ac = -(v2/r)
 Centrifugal force= mv2/r
 F sin θ = mv2/r
 F cos θ = mg
 Tan θ = v2/gr
 Τorque = r F = rma = rm (rα) =( r2m)α = I α
 Moment of inertia; I = mr2
 Ring or thin walled cylinder inertia(I) = MR2

 Disc or solid cylinder inertia = ½ MR2
 Disc inertia = ½ M (R22 + R12 )
 Solid sphere inertia = 2/5 MR2
 Solid rod or meter stick inertia = 1/12 Ml2
 Rectangular plate inertia = 1/12 M (a2+b2)
 Angular momentum = L = r x p = r mv = rmrω
=r2mω = Iω
 L = rmv → L/t = rmv/t = rma = rF = τ
 L/t = τ
 Linear kinetic energy = ½ mv2
 Rotational kinetic energy = ½ Iω2
 Velocity of hoop = v = √gh
 Velocity of disc = v = √(4/3 gh)
 Critical velocity = v = 7.9 km2
 The orbital velocity = v =√(〖GM〗_e/r)
 Lift at rest → T =w
 Lift moving downward → T = w – ma
 Lift moving upward → T = w + ma
 Lift falling freely = T mg-ma = 0
 Frequency for artificial satellite → f = 1/2π √(g/r)
 Oscillation
 Frequency → f=1/T
 Angular frequency → ω = 2πf
 Time period → T = 2π/ω
 Velocity of projection → vy = ω√(r^2-x^2 )
 Simple pendulum time period → T = 2π √(L/g)
 Simple pendulum potential energy = ½ kx2
 Simple pendulum kinetic energy = ½ kx02 -½ kx2
 Total energy of simple pendulum = ½ kx02
 Resonance frequency = Fn = nf1
 Phase → θ =ω t
 Waves
 Transverse wave speed → v=√(T x L )/M or
v=√(T )/m
 Longitudinal waves speed → v=√(E )/ρ
 Phase change→ 2π = λ
 Phase difference → δ = 2π/λ
 Speed of sound by newton → v = √((ρ_m gh)/ρ) =
281 ms-1
 Laplace correction → v = √((〖γρ〗_m gh)/ρ) = 332
ms-1
 Chap No.11 ELECTROSTATICS
 1 e = 1.602 x 10-19 C
 Q = ne
 Coulomb’s Law; F = k (q1 q2)/r2
 K = 1/4πεo
 K = 9.0 x 109 N m2 C-2
 εo = 8.85 x 10 -12 C2 N-1 m-2
 εr = ε/ε0
 Fmed = (F vac)/εr
 E = F/q = V/d = K q/r2
 Ф = E A cos θ = N m2 C-1
 Ф = Q/ε0
 E due to sheet of charge; E = ς/2ε
 E due to charge palates; E = ς/ε
 V = W/Q = U/Q Volt = Joule / Coulomb
 Electric potential energy; U = K Qq/r
 Electric potential; V = W/Q = Fr/Q = K Q/r
 Potential Gradient = E = - ΔV/Δr
 1 eV =1.602 x 10-19 C x 1V → (1 eV = 1.602 x
10-19 J)
 C = Q/V = C V-1 = farad
 Charge density; ς = Q/A
 Cvac = Q/V = (ε0 A)/d = (ε0 εr A)/d
 εr = Cmed / Vvac
 Capacitors In Series;
 Q = Q1 = Q2 =Q3
 V =V1 + V2 + V3
 1/Ce = 1/C1 + 1/C2 + 1/C3
 Capacitors In Parallel;
 Q = Q1 = Q2 = Q3
 V = V1 + V2 +V3,
 Ce = C1 + C2 + C3
 Electric dipole; P = q d
 Energy = U = UV/2= CV2/( 2) = 1/2 (A ε0 εr )/d
(Ed)2
 Energy density; μ=U/Ad=1/2 εo εr E2
 Maximum charge on capacitor = C x e.m.f
 q/q0 = 63.2 % →for charging
 q/q0 = 36.7 % →for discharging
 q = q0 (1-e-t/RC ) →for charging
 q = q0 e-t/RC →for discharging
 CURRENT ELECTRICITY
 Current, I = Q/t → C s-1 = A
 Drift velocity order = 10-5 m/s.
 V = IR
 Tan θ = I/V = 1/R
 Resistance, R = V/I → 1Ω = 1V/1A
 R = ρ L/A → Ω.m
 Conductance, G = 1/R → Siemen(S) or mho
 Conductivity, ς = 1/ρ =L/RA →mho/m or S/m
 Pure metals R inc with T inc.
 Electrolytes and insulators, R dec with T inc.
 ΔR = αR0 T → RT = R0 (1+αT)
 Temperature co-efficient of Resistance, α = RT –
R0/R0T → K-1
 Resistivity, ρ T = ρ 0 (1+αT) OR α = ρ T – ρ 0/ ρ
0T → K-1
 Electromotive Force, ε = W/q → 1 volt = 1
joule/coulomb
 Open circuit, I = 0 so V= ε
 Terminal Voltage, Vt = ε - Ir
 Power, P = W/t = VI → 1 Watt = 1V x 1A
 1 kWh = 1 unit of electrical energy
 1 J = 1W x 1s
 Maximum output power, (Pout)max = ε2 /4r = ε2
/4R
 Thermo emf, ε = αT + ½ βT2
 KCL, ƩI = 0
 KVL, Ʃε = ƩV = ƩIR
 KCL based on L.O.C.O.CHARGE
 KVL based on L.O.C.O.ENERGY
 Wheatstone Bridge, X = PQ/R
 Potentiometer, ε2 /ε1 = I2 /I1
 Tan θ = I/V = 1/R
 ELECTROMAGNETISM
 Force on current carrying wire, F=BIL sin θ.
 Magnetic field or magnetic induction, B = F/IL →1
tesla =1 NA-1 m-1 = 1 Wb m-2

 1 T = 104 G
 Magnetic Flux, Ф = B A cos θ → 1 Wb = 1 N m A-
1.
 Ampere’s Law, B ∝ I/r = μ0 (I/2πr) OR ƩB.ΔL =
μ0 I
 Bnet = B1 + B2
 Magnetic field due to current carrying solenoid, B =
μ0 n I → n=N/L
 Motion of charge particle in uniform magnetic field,
F=q v B sin θ
 Centripetal Force = Magnetic force → mv2/r = qvB
 Time period of charge particle in B, T = 2πm/qB
 Frequency of charge particle in B, f = qB/2πm
 Velocity selector, FE = FM → qE = qvB → v
= E/B
 Torque on current carrying coil, τ = NBIA cos θ
 Ρestoring torque, τ = C θ
 Galvanometer, NBIA cos θ = C θ → I = Cθ/NAB
→ I ∝ θ
 Conversion of galvanometer into ammeter, small R
connected in parallel
 Conversion of galvanometer into voltmeter, large
R in series are connected
 Ammeter, Rs = Rg Ig / (I – Ig) → Ideal ammeter
→ 0 R
 Voltmeter, Rh = (V/I¬g) – Rg → Ideal
voltmeter → infinite R
 ELECTROMAGNETIC INDUCTION
 Faraday’s Law, ε ∝ N (ΔФ/Δt) → ε = N (ΔФ/Δt )
 Lenz Law, ε = –N (ΔФ/Δt )
 Flux motional emf, ε = Blv sin θ
 Rate of work done, W= Bilv
 Rate of production of electrical energy, energy =ε I
 W = energy → Bilv = εI → ε = Blv
 Power, P = F v
 ε = L ΔI/Δt or ε = N ΔФ/Δt → LI = NФ
 Self-Inductance, L = NФ /I
 ε = M ΔI/Δt or ε = N ΔФ/Δt → MI = NФ
 Mutually inductance, M = NФ /I
 F = 1/T
 Induced emf, ε = NAB cosωt or NAB ω sinωt
 ε = εmax sin ωt
 Back emf, V = ε + IR
 Ns / Np = Vs / Vp = Ip /Is
 DAWN OF MODERN PHYSICS
 E = m0 c2
 L= L0 √((1=v2)/c2)
 T = t0 √((1=v2)/c2)
 M = m0 √((1=v2)/c2)
 λmax T = 0.2898 x 10-2 m k (Wein’s displacement
law)
 E = ς T4 (Steffan-Bolts Law)
 ς = 5.67 x 10-8 Wm-1 K-4
 E = n h f
 K.Emax = e V0
 K.Emax = h f – Ф
 H f0 = Ф = hc/λ
 K.Emax = hf - Hf0
 Hf = K.E +hf’
 P= E/c
 Δλ =E/(m0 c) 1-cos⁡θ
 1/f' = 1/f + E/(m0 c) 1-cos⁡θ
 Ephoton = Eelectron + Epositron
 Photon rest mass energy = 2m0c2 = 1.02 MeV
 h/fc = mve- + mve+
 λ = h/p = h/mv
 Δp = h/λ and Δx = λ
 (Δp)(Δx) = h
 (ΔE)(Δt) = h
 ATOMIC SPECTRA
 1/(λ ) = R ( 1/(P2 ) - 1/(n2 ) )
 R =E0 / hc
 R == 1.097 x 107m-1.
 mvr = nh/2π.
 h = planks constant = 6.6256 x 10-34 j s.
 E = hf = En – Ep
 rn = (n2 h2)/(4 π k m e2 )
 En = - ( 2 π2 2 k m e4 )/(n2 h2)
 En =-E0/(n2 ) = 2.17 x 10-18 j/ n2 = +13.6 ev/ n2
 rn = n2 r1 → r1 = 0.53 0A.
 1 0A = 10- m
 2πr=nλ
 eV → hfmax = hc/λmin
 λmin = hc/eV
 excited state for 10-8 s.
 metastable state for 10-3 s
 NUCLEAR PHYSICS
 Nuclear size is of the order of 10-14 m.
 The mass of the nucleus is of the order of 10-27 kg.
 ½ mv2 = Vq
 Bqv = mv2/r
 Bqv = mv2/r → m = Bqr/v
 ½ mv2 = Vq → v2 = 2Vq/m
 So m = qr2B2/2V
 Δm = Zmp + Nmn – M(A,Z)
 The binding energy in MeV is 931 x Δm.
 The binding energy per nucleon = Eb/A.
 0n1 → 1H1 + -1β0 + antineutrino 12 MIN
 ΔN/Δt =-λN
 R =- ΔN/Δt =λN
 N= N0e-λt
 1 Bq = 1 decay per second
 1 Ci = 3.70 x 1010 decay/s
 λT ½ = 0.693
 The charge on u,t and c, in term of electron is +2/3e.
 The charge on s,t and b in term of electron is -1/3e.
 proton =2U→D.

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