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HSC Chemistry Preparation Tips Part - I
1. Solid State
 Classification of Solids


 Classification of crystalline
   solids
   Unit cells, two and three dimensional
    lattices and number of atoms per unit
    cell and
   Number of Atoms in the unit cell :
    Unit cell      scc bcc   fcc   hcp

    Number of      1   2     4     3
    atoms
    Packing efficiency :

    scc     bcc      fcc    hcp
    52.4%   68 %     74%    74%
   Relation between radius (r) of an
    atom and edge length (a) of cubic
    unit cell :
    scc       bcc           fcc
          a         3              a
    r         r         a   r
          2         4             2 2
 Packing in solids
 Density of unit cell
   ∙   Density of the crystal :
               z   M
       d       3
           a       NA
   Packing in voids of ionic solids

   Defects in crystal structure

   Electrical properties

   Magnetic properties
2. Solutions & Colligative
       Properties
  Types of solution
 Concentration of solution of solids
   in liquids
                           W
                         n
  Number ofnmoles,        M mol
  Molarity = V mol dm-3 (or M)
  Normality = gra m e q. gram eq.dm -3
                   V
 Solubility of gases in liquids
 Solid solution
 Colligative properties and
  molecular masses
 Lowering of vapour pressure
n
   Molality = W t. of solven t   in kg
    mol kg -1 (or m)

   Raoult’s law: Psoln = x1 P0
     P0        P   W2 M 1
          P0       W1M 2
   Elevation of boiling point

        Tb = Kbm and Tf = kf m

                    W2   1000
       Tb = Kb
                    W1   M2
       Depression of freezing point
                     W2     1000
          Tf = Kf     W1   M2


       Osmotic pressure
                                        2
                                                C2
         At constant temperature :             C1
                                        1

                                            2
                                                 T2
         At constant concentration :            T1
                                            1
 van’t Hoff equation :   = CRT

                          n
                              RT
                          V
                          W RT
                           MV
   van’t Hoff factor (i)

       Tb (o b )       Tf (o b )             M th
                                      ob
       Tb (th)         Tf (th)        th
                                             M ob


              1        M th        M ob    (For dissociation)

          n        1        M th           (M ob    M th )
1      M oh    M th (For association)
   n 1          M ob     (M ob   M th )

    i 1
         (For dissociation)
    n 1

     n
            (1     i) (For association)
    n 1
 Abnormal molecular mass


 Van’t Hoff factor
3. Chemical Thermodynamics
       & Energetics
   Basic concepts in thermodynamics


   Nature of Heat and Work


   Internal Energy
 W = - P (V2 – V1) = -P V
  (For expansion)
 W = P (V2 – V1) – P V
  (For compression)
                           V2
 Wmax = - 2.303 nRT log10
                               V1

                               P1
   Wmax = - 2.303 nRT log10   P2
 First Law of Thermodynamics
 Enthalpy
     U=q+W
    H = U + PV
    H= U+P V
    H = U + nRT
 Enthalpy of physical changes
 Thermo chemistry
 Spontaneous processes
  (Irreversible processes)
 Gibbs free energy
    S   q rev    H
           T      T
    G = H – TS
    G= H-T S
         G0 = - 2.303 RT log10K
         G = 0, the system is at
        equilibrium
         G < 0, the process is
        spontaneous
         G > 0, the process is non-
        spontaneous.
   Third law of thermodynamics.
4. Electrochemistry
 Redox Reaction


 Conductance in electronic
  solutions
Potential difference (V)
 Resistance (R)
                        Electric current (I)

                                       1       1
 Electrical conductance (G)
                                       R
   or S

                         a
 Resistivity ( )   R          cm
                         1

   Cell constant =       cm-1(or m– 1)
                      a
                          C e ll C o n s tan t
   Conductivity (k)
                           R e s is tan ce

   Molar conductivity (∧m) = (k in Ω-1
                                     k
   m-1 and C in mol m-3) OR         C
k   1000
   ∧m
              C
    (k in Ω-1 m-1 and C in mol m-3)
                                   0       0
   Kohlarausch’s law : ∧0
                                       m
   Degree of dissociation   ( )
                                       0
                                                   2
                                                   m   C
   Dissociation constant (ka)                 (
                                           0       0       m)
 Electrochemical cells

 Electrolytic cells

 Galvanic or voltaic cells

 Electrode potentials and cells potential
0                0                             0
   E cell           E red   (cathode)             E red   (anode)


                             0            0.0592                   n
   EM n       / M
                         E Mn                           log10 [M       ]
                                 /M
                                               n

                   0             0 .0 5 9 2              [P r o d u c ts]
    E c e ll     E c e ll                     lo g1 0
                                     n                 [R e a c ta n ts]
0               0
                G            n   FE cell
          ∆G = nFEcell
          ∆G0 = - RTln K

          0          0.0592
       E cell                log10 K
                         n


 For spontaneous cell reaction :
        Ecell > 0; ∆G < 0
5. Chemical Kinetics
    Rate of reaction
     For a reaction, aA + bB           cC + dD
         1   [A ]   1   [B]   1   [C]     1    [D ]
         a    t     b    t    c    t     d      t

      Average rate =
     Rate law : Rate = k [A]a          [B]b
 Dependence of rate on reactant
     concentration
          2 .3 0 3       [A ]0
      k
            t  log0 [A ]    t

      (for first order reaction)

                 0.693
     t ½ = k (for first order
      reaction)
[A ]0      [A ]t
         k=
                   t
          (For zero order reaction)

                [A ]0
         t½=             (For zero order reaction
                 2k

       Molecularity of elementary reactions
       Collision theory and activation energy
 Temperature dependence of
  reaction rates (Arrhenius equation)
  K = Ae– Ea/RT (Arrhenius equation)

                            Ea
    Log10k = log10A – 2 .3 0 3R T

    Log10
 Effect of catalyst on rates of
   reactions
    k2       E a ( T2   T1)
    k1   2 .3 0 3R      T1    T2
6. General & Processes
of Isolation of Elements
 Concentration of an ore
 Oxidation – reduction
 Refining of crude metal
 Extraction of Zinc


 Extraction of Iron


 Extraction of Aluminium


 Extraction of Copper
7. p–Block Elements
   Group 15 elements

   Group 16 elements

   Group 17 elements

   Group 18 elements
   Reference electrodes

   Common types of cells

   Fuel cells

   Electrochemical series

   Corrosion
Thank You

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HSC Chemistry Preparation Tips Part - I

  • 2. 1. Solid State  Classification of Solids  Classification of crystalline solids
  • 3. Unit cells, two and three dimensional lattices and number of atoms per unit cell and  Number of Atoms in the unit cell : Unit cell scc bcc fcc hcp Number of 1 2 4 3 atoms
  • 4. Packing efficiency : scc bcc fcc hcp 52.4% 68 % 74% 74%
  • 5. Relation between radius (r) of an atom and edge length (a) of cubic unit cell : scc bcc fcc a 3 a r r a r 2 4 2 2
  • 6.  Packing in solids  Density of unit cell ∙ Density of the crystal : z M d 3 a NA
  • 7. Packing in voids of ionic solids  Defects in crystal structure  Electrical properties  Magnetic properties
  • 8. 2. Solutions & Colligative Properties  Types of solution  Concentration of solution of solids in liquids W n  Number ofnmoles, M mol  Molarity = V mol dm-3 (or M)  Normality = gra m e q. gram eq.dm -3 V
  • 9.  Solubility of gases in liquids  Solid solution  Colligative properties and molecular masses  Lowering of vapour pressure
  • 10. n  Molality = W t. of solven t in kg mol kg -1 (or m)  Raoult’s law: Psoln = x1 P0 P0 P W2 M 1 P0 W1M 2
  • 11. Elevation of boiling point  Tb = Kbm and Tf = kf m W2 1000  Tb = Kb W1 M2
  • 12. Depression of freezing point W2 1000  Tf = Kf W1 M2  Osmotic pressure 2 C2  At constant temperature : C1 1 2 T2  At constant concentration : T1 1
  • 13.  van’t Hoff equation : = CRT n RT V W RT MV
  • 14. van’t Hoff factor (i) Tb (o b ) Tf (o b ) M th ob Tb (th) Tf (th) th M ob 1 M th M ob (For dissociation)  n 1 M th (M ob M th )
  • 15. 1 M oh M th (For association)  n 1 M ob (M ob M th ) i 1  (For dissociation) n 1  n (1 i) (For association) n 1
  • 16.  Abnormal molecular mass  Van’t Hoff factor
  • 17. 3. Chemical Thermodynamics & Energetics  Basic concepts in thermodynamics  Nature of Heat and Work  Internal Energy
  • 18.  W = - P (V2 – V1) = -P V (For expansion)  W = P (V2 – V1) – P V (For compression) V2  Wmax = - 2.303 nRT log10 V1 P1  Wmax = - 2.303 nRT log10 P2
  • 19.  First Law of Thermodynamics  Enthalpy  U=q+W  H = U + PV  H= U+P V  H = U + nRT
  • 20.  Enthalpy of physical changes  Thermo chemistry  Spontaneous processes (Irreversible processes)  Gibbs free energy  S q rev H T T  G = H – TS
  • 21. G= H-T S  G0 = - 2.303 RT log10K  G = 0, the system is at equilibrium  G < 0, the process is spontaneous  G > 0, the process is non- spontaneous.  Third law of thermodynamics.
  • 22. 4. Electrochemistry  Redox Reaction  Conductance in electronic solutions
  • 23. Potential difference (V)  Resistance (R) Electric current (I) 1 1  Electrical conductance (G) R or S a  Resistivity ( ) R cm 1
  • 24.   Cell constant = cm-1(or m– 1) a C e ll C o n s tan t  Conductivity (k) R e s is tan ce  Molar conductivity (∧m) = (k in Ω-1 k  m-1 and C in mol m-3) OR C
  • 25. k 1000  ∧m C (k in Ω-1 m-1 and C in mol m-3) 0 0  Kohlarausch’s law : ∧0 m  Degree of dissociation ( ) 0 2 m C  Dissociation constant (ka) ( 0 0 m)
  • 26.  Electrochemical cells  Electrolytic cells  Galvanic or voltaic cells  Electrode potentials and cells potential
  • 27. 0 0 0  E cell E red (cathode) E red (anode) 0 0.0592 n  EM n / M E Mn log10 [M ] /M n 0 0 .0 5 9 2 [P r o d u c ts] E c e ll E c e ll lo g1 0  n [R e a c ta n ts]
  • 28. 0 0  G n FE cell  ∆G = nFEcell  ∆G0 = - RTln K 0 0.0592  E cell log10 K n  For spontaneous cell reaction : Ecell > 0; ∆G < 0
  • 29. 5. Chemical Kinetics  Rate of reaction  For a reaction, aA + bB cC + dD 1 [A ] 1 [B] 1 [C] 1 [D ] a t b t c t d t Average rate =  Rate law : Rate = k [A]a [B]b
  • 30.  Dependence of rate on reactant concentration 2 .3 0 3 [A ]0 k  t log0 [A ] t (for first order reaction) 0.693  t ½ = k (for first order reaction)
  • 31. [A ]0 [A ]t  k= t (For zero order reaction) [A ]0  t½= (For zero order reaction 2k  Molecularity of elementary reactions  Collision theory and activation energy
  • 32.  Temperature dependence of reaction rates (Arrhenius equation)  K = Ae– Ea/RT (Arrhenius equation) Ea  Log10k = log10A – 2 .3 0 3R T  Log10
  • 33.  Effect of catalyst on rates of reactions k2 E a ( T2 T1) k1 2 .3 0 3R T1 T2
  • 34. 6. General & Processes of Isolation of Elements  Concentration of an ore  Oxidation – reduction  Refining of crude metal
  • 35.  Extraction of Zinc  Extraction of Iron  Extraction of Aluminium  Extraction of Copper
  • 36. 7. p–Block Elements  Group 15 elements  Group 16 elements  Group 17 elements  Group 18 elements
  • 37. Reference electrodes  Common types of cells  Fuel cells  Electrochemical series  Corrosion