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CHEMISTRY  
 
Unit  4:  General  Principles  of  Chemistry  I  –  Rates,  Equilibria  and  Further  Organic  Chemistry  
1) How  far?  –  Entropy  
a) Demonstrate  an  understanding  that,  since  endothermic  reactions  can  occur  spontaneously  
at  room  temperature,  enthalpy  changes  alone  do  not  control  whether  reactions  occur  
i) First  law  of  thermodynamics  –  the  total  amount  of  energy  always  stays  the  same  
ii) Second   law   of   thermodynamics   –   energy   spontaneously   tends   to   flow   from   being  
concentrated  in  one  place  to  becoming  diffused  and  spread  out  
(1) ∆Stotal  >  0  for  a  spontaneous  change  at  the  specified  temperature  
(2) Endothermic   reactions   can   occur   spontaneously   –   feasibility   is   determined   by  
enthalpy  and  entropy  change    
(3) Determines  whether  a  physical  or  chemical  change  is  likely  to  happen  at  a  particular  
temperature  
(4) Determines  whether  redox  reactions  will  take  place  
(5) Determines  the  position  of  equilibrium  
iii) …  This  is  because  energy  and  matter  tend  to  spread  out  or  disperse  
iv) Entropy,  S  –  the  measure  of  how  spread  out  energy  is  /  JK-­‐1mol-­‐1  
(1) An  increase  in  disorder  results  in  a  higher  entropy  
(2) Entropy   is   related   to   the   number   of   ways   the   particles   within   a   substance   can   be  
arranged  and  how  the  quanta  of  energy  can  be  distributed  within  a  substance  
 
b) Demonstrate   an   understanding   of   entropy   in   terms   of   the   random   dispersal   of   molecules  
and  of  energy  quanta  between  molecules  
i) Quanta  –  discrete  quantities  of  energy  that  atoms  absorb  
(1) Atoms  rotate,  vibrate  and  translate  (spin,  shake  and  move)  requiring  energy  (quanta)  
(2) Molecules  have  distinct,  rotational  and  vibrational  energy  levels    
ii) At   any   moment   in   time,   a   molecule   will   have   a   set   amount   of   rotational   and/or  
vibrational  energy  depending  upon  how  many  quanta  of  energy  it  has  absorbed  
iii) Molecules  in  contact  with  one  another  can  transfer  quanta  between  them  
iv) The  more  energy  quanta  a  substance  has,  the  higher  the  entropy  
 

c) Demonstrate  an  understanding  that  the  entropy  of  a  substance  increases  with  temperature,  
that  entropy  increases  as  solid  →  liquid  →  gas  and  that  perfect  crystals  at  zero  kelvin  have  
zero  entropy  

i) A   solid   is   much   more   ordered   (less   disordered)   than   a   liquid,   which   in   turn   is   more  
ordered  than  a  gas  
(1) Gaseous  water  has  a  greater  entropy  than  liquid  water,  which  has  a  greater  entropy  
than  ice  
(2) S  solid  <  S  liquid  <  S  gas  
ii) As  temperature  increases,  so  does  disorder  ie
...
 The  entropy  of  ethane  is  greater  than  that  of  methane  

(2) Eg2
...
  In   a   crystalline   solid,   can   be   rearranged   into   many  
possible  disordered  arrangements,  eg  in  a  solution,  so  the  probability  of  disorder  is  greater  
than  order    
i) A   solid   is   much   more   ordered   (less   disordered)   than   a   liquid,   which   in   turn   is   more  
ordered  than  a  gas  
ii) One   possible   ordered   arrangement   in   a   crystalline   solid   can   be   rearranged   into   many  
possible   disordered   arrangements   in   a   solution   –   the   probability   of   disorder   is   greater  
than  order  in  solution  (relative  to  gas)  
iii) Therefore,  gaseous  water  has  a  greater  entropy  than  liquid  water,  which  has  a  greater  
entropy  than  ice  
iv) S  solid  <  S  liquid  <  S  gas  
 
f) Interpret  the  natural  direction  of  change  as  being  in  the  direction  of  increasing  total  entropy  
(positive  entropy  change),  eg
...
8H2O  
with  solid  ammonium  chloride    
(1) Ba(OH)2
...
 transfer  of  heat/energy  
 
-­‐   ∆H/   T  
j) Use   the   expression   ∆Ssurroundings   =  
to   calculate   the   entropy   change   in   the  
surroundings  and  hence  ∆Stotal  

i) ∆Ssurroundings  =  

!∆!
!

 

(1) Remember  the  negative  sign  
(2) Remember  to  multiply  –ΔH  by  1000  as  it  is  usually  in  kJmol-­‐1  not  Jmol-­‐1  
(3) Remember  that  ‘T’  must  always  be  given  in  Kelvin  (°C  +  273)  
ii) When   an   exothermic   reaction   takes   place,   heat   energy   is   transferred   to   the   surrounding  
air,  causing  an  increase  in  disorder  of  the  air  molecules  
(1) The  Maxwell  Boltzmann  distribution  of  energies  shows  a  greater  range  of  energy  at  a  
higher  temperature,  so  the  molecules  are  more  random  or  disordered  
(2) ∆Ssurroundings  will  always  be  positive  for  exothermic  reactions  
(3) ∆Ssurroundings  will  always  be  negative  for  endothermic  reactions  
iii) If  the  surroundings  are  hot,  the  entropy  increase  is  small  because  the  molecules  have  a  
high  entropy  already  as  they  are  already  in  chaotic  motion  
 
k)  Demonstrate   an   understanding   that   the   feasibility   of   a   reaction   depends   on   the   balance  
between  ∆Ssystem   and  ∆Ssurroundings,  and  that  at  higher  temperatures  the  magnitude  of  
∆Ssurroundings   decreases  and  its  contribution  to   ∆Stotal  is  less
...
 
!∆!

∆Ssystem  

∆Ssurroundings  =  

Positive  
Positive  

Positive    
Negative  
(endothermic  reaction)  

Negative  

!

 

Positive  
(exothermic  reaction)  

∆Stotal  
Always  >  0,  therefore  always  feasible  
If    ∆Ssystem  >  -­‐  ∆Ssurroundings  
If  ‘T’  is  increased,  ∆Ssurroundings  
decreases  in  magnitude,  it  will  become  
less  negative  and  less  significant  
If    ∆Ssurroundings  >  -­‐  ∆Ssystem  
If  ‘T’  is  decreased,  ∆Ssurroundings  
increase  in  magnitude,  it  will  become  
more  positive  and  more  significant  
Always  <  0,  therefore  never  feasible  

Negative  
Negative  
 
l) Demonstrate  an  understanding  of  and  distinguish  between  the  concepts  of  thermodynamic  
stability  and  kinetic  inertness  
i) If  ∆Stotal  >  0,   the  reactants  are  thermodynamically  unstable  relative  to  the  products  –  
the  reaction  is  thermodynamically  feasible  at  the  specified  temperature  
(1) If  ∆Stotal   <  0,  the  reactants  are  thermodynamically  stable  relative  to  the  products  –    
the  reverse  reaction  is  thermodynamically  spontaneous    
(2) The   more   positive   the   total   change   in   entropy   value,   the   more   the   position   of  
equilibrium  will  lie  to  the  right  

(3) Endothermic  reactions  can  happen  only  if  ∆Ssystem   is  positive  –  more  likely  to  take  
place  at  higher  temperatures  as  the  magnitude  of  ∆Ssurroundings  decreases  
(4) Exothermic   reactions   are   thermodynamically   favourable   even   when   the   entropy   of  
the   system   is   negative,   so   long   as   the   entropy   change   of  the   surroundings   outweighs  
the  entropy  change  of  the  system  
ii) Kinetic  inertness  
(1) A  small  activation  energy  means  that  the  reactants  are  kinetically  unstable  relative  to  
the  products    
(2) A   large   activation   energy   means   that   the   reactants   are   kinetically   stable/inert  
relative  to  the  products  
(3) Although   ∆Stotal   >   0,   there   could   be   a   very   high   activation   energy   that   prevents   it  
from  actually  taking  place  
 
m)  Calculate  ∆Ssystem   and  ∆Ssurroundings   for  the  reactions  in  4

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