HSC Chemistry Preparation Tips Part - I
Upcoming SlideShare
Loading in...5
×
 

HSC Chemistry Preparation Tips Part - I

on

  • 5,850 views

Important tips for HSC Chemistry 2013.

Important tips for HSC Chemistry 2013.

Statistics

Views

Total Views
5,850
Views on SlideShare
5,846
Embed Views
4

Actions

Likes
1
Downloads
64
Comments
0

2 Embeds 4

http://pinterest.com 3
http://www.pinterest.com 1

Accessibility

Categories

Upload Details

Uploaded via as Microsoft PowerPoint

Usage Rights

© All Rights Reserved

Report content

Flagged as inappropriate Flag as inappropriate
Flag as inappropriate

Select your reason for flagging this presentation as inappropriate.

Cancel
  • Full Name Full Name Comment goes here.
    Are you sure you want to
    Your message goes here
    Processing…
Post Comment
Edit your comment

HSC Chemistry Preparation Tips Part - I HSC Chemistry Preparation Tips Part - I Presentation Transcript

  • 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% View slide
  •  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 View slide
  •  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 & Processesof 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