temperatures [32]. At room temperature oxygen reacts by a straightforward association reaction mechanism with the dimer anion, as determined from the measured
product ion concentration as a function of the reaction time, and the negative dependence of the reaction rate with temperature:
Au 2 þ O 2 !
k Au 2 O 2
ð1Þ
As mentioned above no adsorption of CO is observed at around room temperature.
When both reactants, O 2 and CO, are introduced into the ion trap, the reaction kinetics of Au 2
À changes drastically, as seen by the offset in the Au 2
À signal. This
offset increases when the partial pressure of CO is augmented. In addition, at
temperatures below 200 K the intermediate with the stoichiometry Au 2 (CO)O 2
À
could be isolated (Figure 17.2A-b). The ion stoichiometry clearly shows that CO
and O 2 are able to coadsorb onto an Au 2
À dimer. From the kinetics of all observable ions, Au 2
À , Au 2 O 2
À , and Au 2 (CO)O 2
À , measured under a multitude of different reaction conditions the catalytic conversion of CO to CO 2 could unambiguously
be detected. The reaction mechanism that fulfills all the prerequisites and fits all
kinetic data measured under all the different reaction conditions could be described by the reaction equations below:
Au 2
À þ O 2 !
k1 Au 2 O 2
À
ð2aÞ
Au 2 O 2
À þ CO T
k2
k3
Au 2 (CO)O 2
À
ð2bÞ
Au 2 (CO)O 2
À þ CO !
k4 Au 2
À þ 2CO 2
ð2cÞ
By studying the kinetics as a function of concentration and temperature (Figure
17.2B) some important findings were obtained. (1) Undoubtedly O 2 adsorption
precedes CO adsorption in the catalytic reaction. This is, furthermore, supported
by the reaction kinetics of Au 2
À when only O 2 or CO is present in the trap: The
adsorption of O 2 is faster by at least a factor of 10 than the adsorption of CO molecules [39]. (2) These experiments revealed that varying the oxygen partial pressure
inside the trap mainly affects the pseudo-first order rate constant k 1 , such that k 1
increases linearly with increasing p(O 2 ) and, to a smaller extent, the rate constant
k 4 , which was attributed to the higher abundance of Au 2 (CO)O 2
À . (3) An increase
in the carbon monoxide partial pressure results, at room temperature, in an increasing back pseudo-first order reaction rate k 2 (Eq. (2)), i.e. the Au 2
À offset is
increasing. Resolving the kinetics at lower temperature, 190 K, reveals a clearly
positive dependence of k 4 (Eq. (2c)) on p(CO), also correlated to the Au 2
À offset.
(4) Finally, with decreasing temperature the final equilibrium concentration of
Au 2 (CO)O 2
À increases, whereas the Au 2 O 2
À concentration decreases (Figure
17.2B). This is consistent with the decreasing rate constants k 1 , k 2 , and k 3 with increasing temperature and this negative temperature dependence is indicative of
barrierless reaction steps, explained in the framework of the Lindemann theory
[39, 40]. Most importantly, however, k 4 increases with increasing temperature.
Hence, an activation barrier is involved in the corresponding reaction (2c).
17.3 Chemical Reactions and Catalytic Processes on Free and Supported Clusters 557
product ion concentration as a function of the reaction time, and the negative dependence of the reaction rate with temperature:
Au 2 þ O 2 !
k Au 2 O 2
ð1Þ
As mentioned above no adsorption of CO is observed at around room temperature.
When both reactants, O 2 and CO, are introduced into the ion trap, the reaction kinetics of Au 2
À changes drastically, as seen by the offset in the Au 2
À signal. This
offset increases when the partial pressure of CO is augmented. In addition, at
temperatures below 200 K the intermediate with the stoichiometry Au 2 (CO)O 2
À
could be isolated (Figure 17.2A-b). The ion stoichiometry clearly shows that CO
and O 2 are able to coadsorb onto an Au 2
À dimer. From the kinetics of all observable ions, Au 2
À , Au 2 O 2
À , and Au 2 (CO)O 2
À , measured under a multitude of different reaction conditions the catalytic conversion of CO to CO 2 could unambiguously
be detected. The reaction mechanism that fulfills all the prerequisites and fits all
kinetic data measured under all the different reaction conditions could be described by the reaction equations below:
Au 2
À þ O 2 !
k1 Au 2 O 2
À
ð2aÞ
Au 2 O 2
À þ CO T
k2
k3
Au 2 (CO)O 2
À
ð2bÞ
Au 2 (CO)O 2
À þ CO !
k4 Au 2
À þ 2CO 2
ð2cÞ
By studying the kinetics as a function of concentration and temperature (Figure
17.2B) some important findings were obtained. (1) Undoubtedly O 2 adsorption
precedes CO adsorption in the catalytic reaction. This is, furthermore, supported
by the reaction kinetics of Au 2
À when only O 2 or CO is present in the trap: The
adsorption of O 2 is faster by at least a factor of 10 than the adsorption of CO molecules [39]. (2) These experiments revealed that varying the oxygen partial pressure
inside the trap mainly affects the pseudo-first order rate constant k 1 , such that k 1
increases linearly with increasing p(O 2 ) and, to a smaller extent, the rate constant
k 4 , which was attributed to the higher abundance of Au 2 (CO)O 2
À . (3) An increase
in the carbon monoxide partial pressure results, at room temperature, in an increasing back pseudo-first order reaction rate k 2 (Eq. (2)), i.e. the Au 2
À offset is
increasing. Resolving the kinetics at lower temperature, 190 K, reveals a clearly
positive dependence of k 4 (Eq. (2c)) on p(CO), also correlated to the Au 2
À offset.
(4) Finally, with decreasing temperature the final equilibrium concentration of
Au 2 (CO)O 2
À increases, whereas the Au 2 O 2
À concentration decreases (Figure
17.2B). This is consistent with the decreasing rate constants k 1 , k 2 , and k 3 with increasing temperature and this negative temperature dependence is indicative of
barrierless reaction steps, explained in the framework of the Lindemann theory
[39, 40]. Most importantly, however, k 4 increases with increasing temperature.
Hence, an activation barrier is involved in the corresponding reaction (2c).
17.3 Chemical Reactions and Catalytic Processes on Free and Supported Clusters 557