Berzelius  is  credited  with  origina3ng  the  chemical  terms  "catalysis",  "polymer",  "isomer"  and  "allotrope"  

Berzelius  is  credited  with  iden3fying  the  chemical  elements  silicon,  selenium,  thorium,  and  cerium.  Students  working  in  Berzelius's  laboratory  also  discovered  lithium,  and  vanadium  

Jöns  Jacob  Berzelius  Born  20  August  1779  Väversunda,  Östergötland,  Sweden    Died  7  August  1848  (aged  68)  Stockholm,  Sweden      

Catalysis  

Gold oxidation catalyst Photcatalyst in the form of nanoflower

Es3mates  are  that  90%  of  all  commercially  produced  chemical  products  involve  catalysts  at  some  stage  in  the  process  of  their  manufacture.[1]  In  2005,  cataly3c  processes  generated  about  $900  billion  in  products  worldwide.[2]    

1.  Recognizing  the  Best  in  Innova3on:  Breakthrough  Catalyst".  R&D  Magazine,  September  2005,  pg  20.  2.   hNp://www.climatetechnology.gov/library/2005/tech-­‐op3ons/tor2005-­‐143.pdf  

Anything  that  increases  the  rate  of  a  process  is  a  "catalyst",  a  term  derived  from  Greek  καταλύειν,  meaning  "to  unite"    The  phrase  catalyzed  processes  was  coined  by  Jöns  Jakob  Berzelius  in  1836  to  describe  reac3ons  that  are  accelerated  by  substances  that  remain  unchanged  a`er  the  reac3on.      Humphry  Davy  discovered  the  use  of  pla5num  in  catalysis.    Probably  the  most  important  metal  in  catalysis.  Wilhelm  Ostwald  at  Leipzig  University  started  a  systema3c  inves3ga3on  into  reac3ons  that  were  catalyzed  by  the  presence  of  acids  and  bases;  Ostwald  was  awarded  the  1909  Nobel  Prize  in  Chemistry.    Other  recent    Noble  prices  in  Chemistry  for  Catalysis:    2011    for  palladium-­‐catalyzed  cross  couplings  in  organic  synthesis,  2007    for  chemical  processes  on  solid  surfaces,  2005  for  the  development  of  the  metathesis  method  in  organic  synthesis,  2001  for  chirally  catalysed  oxida3on  and  reduc3on  reac3ons  

1.  Catalysts  work  by  providing  an  (alterna3ve)  mechanism  involving  a  different  transi3on  state  and  lower  ac3va3on  energy.  Consequently,  more  molecular  collisions  have  the  energy  needed  to  reach  the  transi3on  state.  Hence,  catalysts  can  enable  reac3ons  that  would  otherwise  be  blocked  or  slowed  by  a  kine3c  barrier.  The  catalyst  may  increase  reac3on  rate  or  selec3vity,  or  enable  the  reac3on  at  lower  temperatures.      2.  In  the  catalyzed  elementary  reac3on,  catalysts  do  not  change  the  extent  of  a  reac3on:  they  have  no  effect  on  the  chemical  equilibrium  of  a  reac3on  because  the  rate  of  both  the  forward  and  the  reverse  reac3on  are  both  affected.    3.  The  produc3vity  of  a  catalyst  can  be  described  by  the  turn  over  number  (or  TON)  and  the  cataly3c  ac3vity  by  the  turn  over  frequency  (TOF),  which  is  the  TON  per  3me  unit.      4.  The  catalyst  stabilizes  the  transi3on  state  more  than  it  stabilizes  the  star3ng  material.  It  decreases  the  kine3c  barrier  by  decreasing  the  difference  in  energy  between  star3ng  material  and  transi3on  state.  

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Principles of Catalysis

• A catalyst opens a new pathway with a lower activation barrier for reaction to follow. • The Gibbs Energy of the reaction is unchanged. • There are no stable intermediates in the catalytic pathway.

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“A catalyst accelerates a chemical reaction without appearing in any of the products. An equilibrium is equilibrated faster, but the position of the equilibrium will not be changed”

The world market for catalysts is estimated to be more than $ 2x109 and the total value of chemicals produced by means of catalysis exceeds $ 1500x109

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Heterogeneous Catalysis

Typical  cataly5c  materials    The  chemical  nature  of  catalysts  is  as  diverse  as  catalysis  itself,  although  some  generaliza3ons  can  be  made.      a.  Proton  acids  are  probably  the  most  widely  used  catalysts,  especially  for  the  

many  reac3ons  involving  water,  including  hydrolysis  and  its  reverse.      b.  Mul3func3onal  solids  o`en  are  cataly3cally  ac3ve,  e.g.  zeolites,  alumina,  

higher-­‐order  oxides,  graphi3c  carbon,  nanopar3cles,  and  facets  of  bulk  materials.  

   c.      Transi3on  metals  are  o`en  used  to  catalyze  redox  reac3ons.  Many  cataly3c      processes,  especially  those  used  in  organic  synthesis,  require  so  called  "late  transi3on  metals",  which  include  palladium,  pla5num,  gold,  ruthenium,  rhodium,  and  iridium.    Chemical  species  that  improve  cataly3c  ac3vity,  without  themselves  being  ac3ve,  are  called  co-­‐catalysts  or  promoters.  

Temperature  and  the  Rate  Constant  •  The  rates  of  chemical  reac3ons  are  sensi3ve  to  temperature:  most  reac3ons  slow  down  at  lower  temperatures  and  speed  up  at  higher  temperatures.    –  This  temperature  dependence  is  contained  in  the  rate  constant,  k.  Rate  =  k  [A]  –  Increasing  the  value  of  k  increases  the  rate  of  the  reac3on.  For  many  reac3ons,  every  increase  in  temperature  by  10°C  doubles  the  reac3on  rate.    •  The  temperature  dependence  of  k  is  given  by  the  Arrhenius  equa5on:        

)/exp( RTEAk a−=

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Collision  Theory    •  Collision  theory  views  the  reac3on  rate  as  the  result  of  par3cles  colliding  with  a  certain  frequency  and  minimum  energy.  –  Par3cles  must  collide  in  order  to  react,  but  most  collisions  do  not  result  in  a  reac3on,  either  because  the  par3cles  do  not  hit  each  other  hard  enough,  or  they  are  turned  the  wrong  way,  etc.  –  As  the  number  of  colliding  reactants  increases,  the  chances  of  two  reactants  colliding  also  increases.  Thus,  increasing  concentra3on  increases  the  rate  of  the  reac3on.  –  Anything  that  increases  the  number  of  effec3ve  collisions  increases  the  rate.  

For  a  general  reac3on    A  +  BC  →  AB  +  C    as  A  and  BC  collide,  their  electron  clouds  repel  each  other.  The  energy  needed  to  overcome  this  repulsion  comes  from  the  kine3c  energy  of  the  par3cles,  and  is  converted  to  the  poten3al  energy  of  the  A-­‐-­‐-­‐B-­‐-­‐-­‐C  complex.  

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Factors  that  Influence  Effec:ve  Collisions    •  Not  every  collision  between  reactant  molecules  leads  to  the  forma3on  of  a  product  molecule.  The  number  of  effec5ve  collisions,  which  actually  lead  to  the  forma3on  of  a  product  molecule,  depends  on  three  factors:  –  the  exponen5al  factor,  f  —  the  frac:on  with  enough  energy  to  react  (related  to  the  ac3va3on  energy).  –  the  collision  frequency,  Z  —the  number  of  collisions  per  unit  of  3me.  –  the  orienta5on  factor,  p  —the  frac:on  of  collisions  with  the  correct  orienta3on.  

Ac:va:on  Energy,  Ea  •  The  height  of  the  barrier  is  the  ac5va5on  energy,  Ea,  and  the  configura:on  of  the  atoms  at  the  maximum  poten3al  energy  is  the  transi5on  state  or  ac:vated  complex  (++).  –  If  the  reactant  par3cles  collide  with  an  energy  less  than  Ea,  they  bounce  apart.  –  If  the  collision  energy  is  greater  than  Ea  (and  orienta3on  is  right),  there  is  enough  energy  to  overcome  the  repulsions,  and  they  react.  –  In  the  transi3on  state,  the  reactant  bonds  are  in  the  process  of  breaking,  and  the  product  bonds  are  in  the  process  of  forming.  –  The  higher  Ea  is,  the  slower  the  reacTon  will  be  

Winger and Polanyi’s representation of Arrhenius model of activation barriers to reactions

The  Frequency  Factor  •  The  exponen5al  factor,  f,  is  the  frac:on  of  collisions  with  enough  energy  to  react:        where  R  is  the  gas  constant  8.314  J  K-­‐1  mol-­‐1.    •  At  higher  temperatures,  the  distribu3on  of  collision  energies  broadens  and  shi`s  to  higher  energies,  enlarging  the  frac3on  of  collisions  with  energy  greater  than  Ea.  This  makes  f  a  larger  number.  

)/exp( RTEf a−=

Reac:on  Rate  and  Temperature    •  The  collision  frequency,  Z,  is  the  number  of  collisions  which  occur  in  a  given  unit  of  3me.    •  For  a  gas  at  room  temperature  and  a  pressure  of  1  atmosphere,  each  molecule  undergoes  about  109  collisions  per  second,  or  1  collision  every  10-­‐9  s.    –  If  every  collision  resulted  in  a  reac3on,  every  gas  phase  reac3on  would  be  over  in  10-­‐9  s.  Most  reac3ons  are  obviously  much  slower  than  this.    –  For  a  reac3on  where  Ea  is  75  kJ/mol,  at  298  K  f=  7  x  10-­‐14  only  7  collisions  in  100  trillion  are  energe3c  enough  to  cause  a  reac3on  to  occur!    •  The  collision  frequency  is  directly  propor3onal  to  the  concentra3on  of  the  reactants.  

Molecular  Orienta:on    •  Not  all  collisions  with  energy  greater  than  Ea  lead  to  a  reac3on:  the  molecules  have  to  be  facing  each  other  the  right  way  when  they  hit  each  other.  •  The  frac3on  of  collisions  having  the  right  orienta3on  is  called  the  orienta5on  factor,  p.  

The  Arrhenius  EquaTon  •  All  of  these  factors  can  be  combined  into  a  single  equa3on:        •  p  and  Z  are  o]en  combined  into  a  frequency  factor,  A  (A  =  pZ);  in  this  form,  this  equaTon  is  known  as  the  Arrhenius  equa5on  (Svante  Arrhenius,  1889):        •  Rearranging  the  Arrhenius  equa3on,    we  can  obtain  the  form  of  an  equa3on    of  a  line:      

)/exp( RTEppZfk a−==

)/exp( RTEAk a−=

RTEAk a /lnln −=

Rate-­‐Determining  Steps  •  Usually  one  step  in  a  mechanism  is  much  slower  than  the  other  steps,  and  acts  as  a  “boNleneck”  for  the  reac3on;  the  rate  of  this  step  limits  how  fast  the  overall  reac3on  can  occur,  and  is  known  as  the  rate  determining  step.      •  The  rate  law  for  the  rate-­‐determining  step  represents  the  rate  law  for  the  overall  reacTon.  

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Overcoming unfavorable thermodynamic (water splitting)

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H2O à H2 + ½ O2 ΔG +286 kJ/mol, 2.3 eV, T = 3000 °C

HP and HT electrolysis Solar thermal (Almeria, Spain) Hyrosol-2 (100kW)

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In  an  eight-­‐atom  cluster,  all  of  the  atoms  are  on  the  surface.  However,  the  dispersion,  D,  defined  as  the  number  of  surface  atoms  divided  by  the  total  number  of  atoms  in  the  cluster,  declines  rapidly  with  increasing  cluster  size.    

Supporting catalysts Dispersion - nano

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Fritz Haber's successful synthesis of ammonia in 1909, capturing nitrogen from the air, brought him fame and wealth. In 1911, he moved to Berlin to head the Kaiser Wilhelm Institute for Physical Chemistry and Electrochemistry. In Berlin, he became friend with Albert Einstein.

The making of Ammonia

N2 + 3 H2 à 2 NH3 Nothing happens in the system without a catalyst as T raised until 1000oC or higher. Above this temperature some H2 molecules are dissociated to H atoms. H2 à 2 H (atoms) For example at 1430oC with p(H2) = 150 atm., the partial pressure of H atoms is ca. 0.1 % only above 3000 oC where N2 molecules dissociates to N atoms and ammonia can be synthesized in reasonable quantities. N2 à 2 N (atoms) The role of the catalyst in ammonia synthesis is that of making the reaction go sufficiently fast (by facilitating the dissociation of molecular nitrogen) so that significant rates are obtained.

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+

+

Gas phase reaction

Catalytic reaction

Energy profiles for the series of reaction steps to make ammonia from N2 and H2 by both homogeneous gas-phase and iron-catalyzed reactions. The role of the catalyst in decreasing the energy barrier to reaction can be seen (vales are in kJ mol-1)

N N N N

N

N

N N N N N

N N

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Kinetic definition of catalysis

Paul J. Crutzen Born: 3 December 1933, Amsterdam, the Netherlands. The Nobel Prize in Chemistry 1995

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As  you  can  also  see  in  the  figure  The  catalyst  has  not  changed  the  thermodynamics,    ΔH  and  therefore  ΔG  and  Kp  are  unchanged,  it  only  affected  the  transi3on  state.  

Hydrogenating organic compounds in the presence of finely disintegrated metals

The Nobel Prize in Chemistry 1912

Born: 5 November 1854, Carcassonne, France Died: 14 August 1941, Toulouse, France

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Rh/Al2O3

50 nm x 50 nm

5  nm

a

bAu/TiO2

Examples of Catalysts

PtRu/CeO2

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The active sites: acidity in zeolites

There is one acid hydrogen for every tetrahedrally bonded aluminium. These active sites are distributed uniformly throughout the bulk and are bridging hydroxyl groups. These are the classic Bronsted acid sites, the intrinsic strength of which is a function both of the particular local environment and also the Si/Al ratio.

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The active sites: bi-functional catalysts

Example: the Pt/Al2O3 catalysts used in the hydroprocessing of petrochemicals, the metal serves to dissociate H2, while the acid support serves to catalyze the build-up of vital carbonium ion intermediates.

H:H methyl cyclo-propane

2-butene butane

a

b

cxy

z

Pt

Al3+ O2-

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The concept of “active site” is therefore very wide. Some examples of adsorbed surface complex are showed, you can observe how reactants interacts with the catalysts surface depending on the nature and distribution of the active sites.

A. NH3 (Lewis base) coordinately linked to Al+3 ions (Lewis acid) on Al2O3 surface. B and C. Linear and bridge adsorption of CO on Pt. D and E. Dissociative adsorption on Pt of H2 or ethane. F. Dissociative adsorption of N2 on Fe. G. Heterolytic dissociative adsorption of H2 on the ZnO surface. H. Adsorbed complex with charge transference. I. Adsorption of isobutene on silica alumina where the acid surface proton (σ-OH) was transferred to the isobutene. J and K. Possibilities of ethylene adsorption on Pt. L. Adsorption of O2 on metal oxides with charge transference. M. Dissociative adsorption of O2. N. Heterolitic dissociative adsorption of propylene on ZnO.

Hydrogen Production (by “Steam Reforming” and “Water-gas Shift”)

CO clean-up (by “Methanation”)

Ammonia synthesis (by “Haber-Bosch”)

CH4

H2O

Heat

CO2

H2 with CO and CO2 impurities

Pure H2

N2

CH4 and H2O

NH3

1 % of the World’s energy production

à Fertilizer à Food for 2-3 billion people

(160 million tons/year)