Electrochemical chlorine evolution at rutile oxide (110) surfaces.
ABSTRACT Based on density functional theory (DFT) calculations we study the electrochemical chlorine evolution reaction on rutile (110) oxide surfaces. First we construct the Pourbaix surface diagram for IrO(2) and RuO(2), and from this we find the chlorine evolution reaction intermediates and identify the lowest overpotential at which all elementary reaction steps in the chlorine evolution reaction are downhill in free energy. This condition is then used as a measure for catalytic activity. Linear scaling relations between the binding energies of the intermediates and the oxygen binding energies at cus-sites are established for MO(2) (M being Ir, Ru, Pt, Ti). The linear relations form the basis for constructing a generalized surface phase diagram where two parameters, the potential and the binding energy of oxygen, are needed to determine the surface composition. We calculate the catalytic activity as function of the oxygen binding energy, giving rise to a Sabatier volcano. By combining the surface phase diagram and the volcano describing the catalytic activity, we find that the reaction mechanism differs depending on catalyst material. The flexibility in reaction path means that the chlorine evolution activity is high for a wide range of oxygen binding energies. We find that the required overpotential for chlorine evolution is lower than the overpotential necessary for oxygen evolution.
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ABSTRACT: In this paper, density functional theory (DFT) calculations are performed to analyze the electrochemical water-splitting process producing molecular oxygen (O2) and hydrogen (H2). We investigate the trends in the electro-catalytic properties of (1 1 0) surfaces of three rutile-type oxides (RuO2, IrO2, and TiO2). The two first of these oxide anodes show lower O2-evolving over-potentials than metal anodes, due to weak O binding but strong hydroxyl (HO∗) binding on the surface. Furthermore, the binding energies of O, HO, and HOO on the (1 1 0) surfaces fulfill universal linear relations similar to those found on metal surfaces.Journal of Electroanalytical Chemistry. 01/2007; 607:83-89.
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ABSTRACT: The catalytic oxidation of hydrogen chloride to chlorine (Deacon reaction) is an attractive means to recover Cl2 from HCl-containing waste streams in the chemical industry. Recently, RuO2 supported on TiO2 rutile has been industrially implemented for large-scale chlorine manufacture. However, detailed understanding of the reaction mechanism over claimed catalysts has not been acquired. Our tests in a fixed-bed reactor at ambient pressure concluded that RuO2 powder is highly active for Cl2 production. Characterization of the fresh and used RuO2 samples by ex situ XRD and XPS revealed no appreciable alteration of the bulk structure and limited chlorination of the catalyst. Ab initio thermodynamics predicted that the initial state of the RuO2(110) surface is partially over-oxidized, while the surface after reaction contains both oxygen and chlorine. DFT described the Deacon reaction with a Mars-van Krevelen type mechanism consisting of five steps: hydrogen abstraction from HCl, recombination of atomic chlJournal of Catalysis. 01/2008; 255(1):29-39.
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ABSTRACT: Based on an extensive set of density functional theory calculations it is shown that for a class of catalytic reactions there is a universal, reactant independent relation between the reaction activation energy and the stability of reaction intermediates. This leads directly to a universal relationship between adsorption energies and catalytic activity, which is used to pinpoint what it is that determines the best catalyst for a given reaction. The universality principle rationalizes a number of known facts about catalysts and points to new ways of improving them.Journal of Catalysis. 01/2002;
Electrochemical chlorine evolution at rutile oxide (110) surfacesw
Heine A. Hansen,aIsabela C. Man,bFelix Studt,abFrank Abild-Pedersen,ab
Thomas Bligaardacand Jan Rossmeisl*a
Received 24th August 2009, Accepted 7th October 2009
First published as an Advance Article on the web 11th November 2009
Based on density functional theory (DFT) calculations we study the electrochemical chlorine
evolution reaction on rutile (110) oxide surfaces. First we construct the Pourbaix surface diagram
for IrO2and RuO2, and from this we find the chlorine evolution reaction intermediates and
identify the lowest overpotential at which all elementary reaction steps in the chlorine evolution
reaction are downhill in free energy. This condition is then used as a measure for catalytic
activity. Linear scaling relations between the binding energies of the intermediates and the oxygen
binding energies at cus-sites are established for MO2(M being Ir, Ru, Pt, Ti). The linear relations
form the basis for constructing a generalized surface phase diagram where two parameters, the
potential and the binding energy of oxygen, are needed to determine the surface composition.
We calculate the catalytic activity as function of the oxygen binding energy, giving rise to a
Sabatier volcano. By combining the surface phase diagram and the volcano describing the
catalytic activity, we find that the reaction mechanism differs depending on catalyst material.
The flexibility in reaction path means that the chlorine evolution activity is high for a wide range
of oxygen binding energies. We find that the required overpotential for chlorine evolution is
lower than the overpotential necessary for oxygen evolution.
Chlorine is an essential product for the global chemical
industry – approximately 50% of the total turnover of the
chemical industry depends on chlorine and caustic soda.1
Chlorine production by chlor-alkali processes is one of the
largest current technological applications of electrochemistry.2
The most active anode catalysts are usually based on RuO2,
however, RuO2 is barely stable at the high potentials.
Therefore RuO2 is mixed with IrO2 and additives such as
TiO2and SnO2, in order to improve the stability. The most
commonly used electrocatalyst in industrial chlorine processes
is the so-called Dimensionally Stable Anodes (DSA
named according to its improved lifetime compared with
earlier used electrocatalysts.3
The equilibrium potential for Cl2evolution is 1.36 V at
room temperature and standard conditions, which is slightly
larger than the equilibrium potential for oxygen evolution,
which is 1.23 V under the same conditions. This means that
under chlorine evolution the simultaneous evolution of oxygen
tends to occur as a parasitic side reaction, especially at high
current densities. However, depending on the employed
catalyst, oxygen evolution usually requires a somewhat larger
s) which is
overpotential than chlorine evolution. Unfortunately, RuO2is
known to be a good catalyst for oxygen evolution as well as for
chlorine evolution. This suggests an overlap of the activity
volcanoes for the two reactions, and it has in fact been
suggested that high catalytic activity for chlorine evolution is
fundamentally linked with high oxygen evolution activity.2
Anodic chlorine evolution at oxide electrodes, and especially
chlorine evolution on RuO2, has been widely studied
experimentally. A variety of different reaction mechanisms
have been suggested based on indirect experimental quantities
such as Tafel slopes and reaction orders.4Among the possible
reactions are the Volmer–Tafel reaction5
2* + 2Cl?- 2Cl* + 2e?- 2* + Cl2+ 2e?, (1)
the Volmer–Heyvrosky reaction6
* +2Cl?- Cl* + e?+ Cl?- * +Cl2+ 2e?, (2)
and the Khrishtalik reaction7,8
* +2Cl?- Cl* + e?+ Cl?- (Cl*)++ 2e?+ Cl?
- * + Cl2+ 2e?. (3)
Here, * is an active site, which may be a surface oxygen or a
Very little is known about the reaction mechanism for
chlorine evolution and about the atomic-scale structure of
the surface, which depends strongly on catalyst material,
electrostatic potential, and electrolyte. Recent developments
within density functional theory analysis of electrochemical
reactions have opened up the possibilities to study these
reactions at the atomic scale.9In particular, investigations of
fuel cell catalysis such as oxygen reduction10–13and methanol
aDepartment of Physics, Building 307, Center for Atomic-scale
Materials Design, Technical University of Denmark,
DK-2800 Kgs. Lyngby, Denmark
bComputational.Materials Design ApS, Fysikvej-Building 307-DTU,
DK-2800 Kgs. Lyngby, Denmark
cMaterials Sciences Division, Lawrence Berkeley National
Laboratory, Berkeley, CA94720, USA
w Electronic supplementary information (ESI) available: Details
regarding construction of the phase diagrams; Tables S1–S5, and
linear relations. See DOI: 10.1039/b917459a
This journal is ? c the Owner Societies 2010 Phys. Chem. Chem. Phys., 2010, 12, 283–290 | 283
PAPERwww.rsc.org/pccp | Physical Chemistry Chemical Physics
mechanisms and surface composition under electrocatalytic
reaction conditions. In a previous study the oxygen evolution
reaction was investigated.18There it was established that it is
possible to describe the trends in the oxygen evolving activity
using one single descriptor: the adsorption energy of O-atoms
on the surfaces. RuO2was determined to be the most active
rutile (110) surface, which is in good agreement with the
experiments. Recently, theoretical studies have addressed
so-called Deacon process.19,20However, in spite of the
significant importance of electrochemical chlorine evolution,
detailed electronic structure studies of this process have not
appeared in the literature.
The aim of the present study is to analyze the surface
structure and the activity trends underlying electrochemical
chlorine evolution over rutile oxides. We start by analyzing
IrO2and RuO2and we construct surface phase diagrams of
the rutile (110) surfaces. This allows us to derive plausible
mechanisms of the electrochemical chlorine evolution based
on the reaction intermediates. We determine the lowest
potential where Cl evolution is possible. Applying adsorption
energy correlations, we can determine a reduced set of key
energetic descriptors for the surface reactions involved, and
generalize the analysis of IrO2and RuO2to a trend study
where all the material dependence is included in a single
descriptor, in this case the oxygen binding energy.
have deepened the insight into reaction
All electronic structure calculations have been carried out
using density functional theory (DFT), with the RPBE
functional for exchange and correlation.21A periodically
repeated 4-layer slab is chosen for the rutile (110) surfaces of
RuO2, IrO2, TiO2, and PtO2. A vacuum layer of 16 A˚is used
to separate the slab from its periodically repeated images. We
use a 2?1 surface unit cell and 4?4?1 Monkhorst–Pack
Kohn–Sham equations are solved using a plane wave basis
with a cutoff of 350 eV for the eigenstates, and the densities are
described using a cutoff corresponding to 500 eV. Vanderbilt
ultrasoft pseudopotentials are used to deal with the ion
cores.23A Fermi smearing of 0.1 eV is used, and energies are
extrapolated to an electronic temperature of 0 K. The two
bottom layers of the slab are fixed in their bulk structure,
while the two top layers as well as possible adsorbates on it
are relaxed until the sum of the absolute forces is less than
0.05 eV A˚?1. All calculations are performed using the Dacapo
and ASE simulation package.24
The surface of the unit cell contains two bridge and two cus
sites, which means that the total coverage at each type site
varies between 50% and 100%. We consider all relevant
combinations of adsorption site and adsorbates. We find that
adsorbates bind stronger at bridge sites than on cus sites and
bridge sites are therefore occupied with oxygen for a
large range of conditions. We therefore focus on cus sites
throughout this paper. Mixed phases where different kinds of
adsorbates are mutually present at the cus sites may exist,
however, we find that they are in general only stable in very
narrow windows of conditions.
We consider the adsorption of Clc, OHcand Ocat a
adsorbed at two cus sites. The adsorption energy of chlorine
is calculated using:
c, as well as the formation of O2ccand Cl(Oc)2
DE(Clc) = E(Clc) ? E(c) ?1
For oxygen the energy is calculated relative to water
DE(Oc) = E(Oc) ? E(c) ? E(H2O) + E(H2)
and for ClOcwe apply the combined reference energy states
The adsorption energy of O2ccis defined with reference to
water and hydrogen
2E(Cl2)? E(H2O)+E(H2). (6)
DE (O2cc) = E(O2cc) ? E(2c) ? 2E(H2O) + 2E(H2),
and the adsorption energy of Cl(Oc)2is defined by
DE(Cl(Oc)2) = E(Cl(Oc)2) ? E(2c) ? 2E(H2O)
+ 2E(H2). ?1
The changes in the interaction between the liquid electrolyte
and the surface upon adsorption of molecules are expected to
be small as long as all hydrogen bonds are saturated. It has
previously been shown that the O and OH adsorption energies
at the cus site is changed by less than 0.05 eV by the presence
of water at the surface on RuO2.18These interactions are
therefore neglected in the present study.
Furthermore, the effect of the local field in the Helmholtz
layer is not accounted for. Previously, it has been shown that
for metal surfaces the effect of the field is negligible for
adsorbates with small dipole moments perpendicular to the
surface.25For RuO2we find that applying a homogeneous
external field up to ?0.53 V A˚?1, corresponding to a 1.6 V
potential drop across a 3 A˚thick Helmholtz layer, changes the
relative adsorption energies by less than 0.11 eV.
The above simplifications are expected to be independent on
the catalyst material, and therefore the resulting trends in
catalytic activity should only be weakly affected by them.
Variations in the adsorption energy of e.g. oxygen on the
(110) surfaces are several eV between e.g. IrO2and TiO2, while
differences in water interaction and field effects are at least an
order of magnitude smaller, and therefore vanish on the
adsorption energy scale.
Surface phase diagram
There are four parameters determining the surface composition:
the potential, the pH, the concentration of Cl?, and the
electrode material. Only the latter is directly available in the
simulations, and the other three parameters can be included
analytically as described below. By applying the computa-
tional standard hydrogen electrode,9it is possible to construct
surface Pourbaix diagrams, and identify the most stable
structure of the catalyst surface at a range of potentials and
pH values.26At conditions where oxygen and chlorine
284 | Phys. Chem. Chem. Phys., 2010, 12, 283–290 This journal is ? c the Owner Societies 2010
evolution are negligible, the structure of the catalyst surface is
determined by the equilibrium with water, protons and
chloride ions. The oxidation of water may lead to the
formation of OHcor Octhrough
H2O(l) +c" HOc+ H+(aq) + e?" Oc+ 2H+(aq) + 2e?
Chloride ions may be exchanged with the surface via
Cl?(aq) +c" Clc+ e?
ClOcmay be formed by first having Ocadsorbed on the
surface followed by Cl?adsorption on top of Oc. At potentials
where evolution of Cl2 or O2 is appreciable, the surface
structure is, however, controlled by the steady-state reaction.
At standard conditions (zero pH), H+(aq) + e?is in
2H2(g) at zero potential vs. the standard
hydrogen electrode. At finite pH and potential the chemical
potential of a proton and an electron is:
m(H+(aq)) + m(e?) =1
2mH2(g)? eUSHE+ kBTln(10)pH.
Similarly, Cl?(aq) is in equilibrium with1
standard conditions at the potential of standard chlorine
2Cl2 + e?under
2Cl2+ e?@USHE= 1.36 V.(12)
For an arbitrary potential and activity we therefore obtain
m(Cl?(aq)) ? m(e?) =1
2mCl2(g)? e(USHE? 1.36 V) + kBTlnaCl?.
Eqns (11) and (13) allow us to calculate the free energies of Oc,
OHc, Clc, and ClOcadsorbed at a surface site in the
electrochemical environment, based on calculations of the
gas-phase molecules rather than the solvated ions.
The free energy of adsorption for a surface with an
adsorbate at USHE= 0 V is given by
DG = DE + DZPE ? TDS + DGref,
where DZPE is the change in zero point energy, T is the
temperature, DS is the change in entropy upon absorption,
and DE is the DFT-calculated adsorption energy. The
zero-point energy contribution and the entropy for the
adsorbed species are obtained from harmonic vibrational
analysis and from tables of thermodynamic properties in the
case of gas-phase species. The numbers for DZPE and ?TDS
are listed in the ESIw (in Table S1). The correction DGrefis
1.36 eV for Clc, ClOcand Cl(Oc)2and zero for HOcand Oc,
and is related to the reversible potentials of the chlorine and
hydrogen electrodes, respectively (see ESI for detailsw).
To obtain a measure of the activity we apply a simplified
estimate: the chlorine evolution reaction is considered possible
if and only if all the involved reaction steps are neutral or
downhill in free energy. For a given reaction we can determine
the lowest potential for which this is the case. Due to the
significant challenges in treating reaction barriers for electro-
chemical processes, we do not include reaction barriers in the
present study, and can therefore not directly compare e.g. the
relative rates of the Volmer–Tafel and the Volmer–Heyvrosky
reactions. Our approach can thus be viewed as a ‘‘lower-bound
overpotential analysis’’ of the chlorine evolution activity.
Since barriers of surface reactions27as well as barriers for
proton transfer reactions28are known to often be linearly
dependent on the reaction energy, we expect that the trends
are conserved even when barriers are included.
Results and discussion
Surface phase diagram for IrO2
Fig. 1 shows the interesting part of the phase diagram of
IrO2. At pH = 7 the surface sites are covered by OH and O at
most potentials. At low potential, the surface is covered by OH
groups (not shown). Increasing the potential oxidizes OH to O
first at the bridge sites and then at the cus sites. Eventually
formation of OOH becomes thermodynamically favored.
When this happens, we expect oxygen evolution to become
appreciable,18and the surface structure is then determined by
the kinetics of the steady-state evolution of oxygen. The
formation of chlorine adsorbates directly at the cus sites
requires pH o ?3. Formation of Cl at the bridge sites requires
even lower pH.
We would expect that for a good catalyst the formation of
the Cl intermediate has DG B 0 eV near 1.36 V and that
there are free sites available for the formation of this inter-
mediate. A mechanism involving Cl adsorbed directly at an Ir
cation, does not fulfill any of these requirements. Instead we
see from the phase diagram that a ClOcintermediate is
thermodynamically favored for U 4 1.5 V in the pH range
from 0 to 3.
H+and H2O at 298.15 K and aCl? = 1. The regions where we expect
chlorine or oxygen evolution to become significant have been marked.
candbdenote cus site and bridge sites, respectively. The adsorbate
phases are shown in the insets. Ir atoms are cyan, O atoms are red,
H atoms are white and Cl atoms are green.
Surface phase diagram for IrO2(110) in equilibrium with Cl?,
This journal is ? c the Owner Societies 2010Phys. Chem. Chem. Phys., 2010, 12, 283–290 | 285
This suggests the following sequence of intermediates on
Oc+ 2Cl?(aq) - ClOc+ Cl?(aq) + e?
- Oc+ Cl2(g) + 2e?
as both steps have |DG| = 0.14 eV at U = 1.36 V, and a
significant amount of Ocsites exist at U 4 1.36 V. The
reaction is written here as a Volmer–Heyvrosky mechanism.
However, as we only consider the stability of the adsorbed
intermediate, we cannot compare the relative rates of the
Tafel, Heyvrosky and Khrishtalik steps.
Surface phase diagram for RuO2
The phase diagram for RuO2(110) turns out to be a bit more
complicated (see Fig. 2). At pH = 7, the surface is dominated
by species formed by the oxidation of water. At low potential,
only the bridge sites are covered by OH. When the potential is
increased, OH is formed at the cus sites, before OH is oxidized
to O. We find that oxygen association at the cus sites
is exothermic by 0.71 eV for the fully O-covered surface. The
association barrier is only 0.18 eV, while desorption of O2ccis
endothermic by 1.16 eV. O2ccwill therefore most likely be
present at the surface rather than Oc. Oxygen evolution could
happen by further oxidation of the surface
H2O + O2cc- O2c+c+ H2O
- O2c+ OHc+ H++ e?. (17)
The stability of the O2c+ OHcstructure relative to H2O and
H+is indicated in Fig. 3. Desorption of O2from this surface
has DG = ?0.1 eV, however, so when O2c+ OHcstarts to
form, we expect oxygen evolution to become important.
Additional barriers could exist, but we will not go further into
the details of oxygen evolution.
On RuO2–chlorine species formed at pH o 1.3, however,
oxygen is still the most stable adsorbate near U = 1.36 V. We
find that the 2Ob+ Oc+ OClcintermediate is metastable
relative to a 2Ob+ Cl(Oc)2structure. The latter structure
forms at U 4 1.5 V, and we expect this to be the intermediate
O2cc+ 2Cl?(aq) - Cl(Oc)2+ Cl?(aq) + e?
- Oc+ Cl2(g) + 2e?. (18)
Adsorption of oxygen at the cus sites leads to an increase in
the work function. This is consistent with negatively charged
O2ccadsorbates. Subsequent formation of Cl(Oc)2leads to a
decrease in the work function. This is consistent with Cl(Oc)2
being more positively charged than O2cc, as has been suggested
for the Khrishtalik mechanism.
We note in passing that the formation of O2ccand Cl(Oc)2
depend on the presence of pairs of Ru cus sites at the surface,
and it may not be relevant for alloys of e.g. TiO2and RuO2.
Neglecting the formation of O2ccand Cl(Oc)2and considering
the IrO2reaction path eqn (15), we find that at U = 1.36 V
Cl?, H+and H2O at 298.15 K and aCl? = 1. The regions where we
expect chlorine or oxygen evolution to become significant have been
marked.candbdenote cus site and bridge sites, respectively.
Surface phase diagram for RuO2(110) in equilibrium with
0.59DE(Oc) ?2.26 eV (’ – vacant neighboring cus-sites, m – Cl
neighbor, K – O neighbor), the adsorption energy of ClO (red) at cus:
DE(ClOc) = 0.52DE(Oc) + 0.62 eV (
– O neighbor, – average adsorption energy of ClO for the fully
covered surface vs. average adsorption energy for O for fully covered
DE(Clc) = ?0.48DE(Oc) + 0.68 eV (
neighboring cus sites,
—Cl atop O vs. O with O neighbors,
atop O vs. O with ClO neighbors,
atop O for fully covered surface on vs. average adsorption energy of O
for fully covered surface), the adsorption energy of O2ccvs. the average
adsorption energy of Oc(yellow): DE(O2cc) = 0.94DE(Oc) + 1.96 eV,
and the adsorption energy of Cl(Oc)2vs. the average adsorption energy
of Oc(green): DE(Cl(Oc)2 = 0.56DE(Oc) + 2.51 eV. The mean
absolute error of the fits are below 0.21 eV.
The adsorption energies of chlorine at cus (black): DE(Clc) =
– vacant neighboring cus sites,
– Cl atop O vs. O with vacant
– average adsorption energy of Cl
286 | Phys. Chem. Chem. Phys., 2010, 12, 283–290 This journal is ? c the Owner Societies 2010
each step has |DG| = 0.05–0.12 eV, depending on whether
there is Ocor OClcadsorbed at the other cus site in the (2 ? 1)
unit cell. The trend in the change of the function upon
formation of Ocand ClOcis similar to the change of the work
function upon formation of O2ccand Cl(Oc)2.
Scaling relations. As mentioned above, it is possible to
construct the surface phase diagram and reaction intermediate
as function of pH, potential and the Cl?concentration for a
given material. The aim is now to generalize the analysis,
not studying a single or a few oxide surfaces but rather
determining a descriptor which will be a continuous material
variable. The starting point of our analysis is to establish
correlations between adsorption energies of intermediates on
various (110) rutile oxide surfaces. Such relations can be useful
in establishing simplified models describing the surface activity
and composition, and can be suitable for subsequent screening
purposes.29In Fig. 3 the adsorption energies of Clcand ClOc
as defined above are plotted against the O binding at the
cus-site with the same environment at the surface. The plot
clearly shows that the Cl and O adsorption energies are
linearly correlated. Such linear energy relations between
adsorption energies of hydrogenated species (CHx, OH, SH,
and NHx) and the adsorption of the corresponding unhydro-
genated atoms: (C, O, S, and N) have previously been shown
for transition metals30,31and transition metal compounds
including oxides.18,32The scaling of Cl with respect to O is
very similar to the scaling of OH with respect to O. This
reflects the fact that Cl has a valency of one like the oxygen
atom in OH. ClOcthus also scales as OH (and similar to
HOOc). The present results suggest that the oxygen adsorption
energy is a general measure (a so-called ‘‘descriptor’’) for the
reactivity of oxides which has also been suggested for the case
of cations in oxides by Pankratiev.33
Adsorption of Cl atop Ocis determined by:
DE = E(ClOc) ? E(Oc) ?1
2E(Cl2) = DE(ClOc) ? DE(Oc)
The linear scaling relations established above makes it possible
to analyze the reaction, not only for a specific metal oxide
surface, but for potential metal oxide catalyst surfaces with
continuously varying reactivity as measured by the adsorption
energy of oxygen at the cus-site. The obtained reactivity curves
will then be continuous in the oxygen adsorption energy,
whereas specific oxides (e.g. RuO2, IrO2, PtO2, and TiO2) will
show up as discrete points. The descriptor approach provides a
fast overview of the ‘‘phase-space’’ of materials, but leaves the
problem of how to find specific materials with the desired
descriptor properties unanswered.
Generalized phase diagram
Since the binding energy of all intermediates at the cus sites
scales directly with DE(Oc), it is possible to construct a
generalized phase diagram showing the most stable phase at
potential U as function of the material-dependent descriptor,
We choose the electrolytic conditions such that when
increasing the potential the most stable form of chlorine goes
directly from Cl?to Cl2, which means that the pH value
should be between ?1 and 3. HCl(aq) is more stable than Cl?
at pH values below B?1, whereas HClO(aq) becomes stable
at pH values higher than B3. We keep the electrolyte pH and
Cl?concentration fixed (pH = 0, aCl? = 1) and investigate the
surface phase diagram as a function of DE(Oc) and potential.
This approach is not a limitation of the method, since other
electrolyte conditions can be treated just by changing the free
This is shown in Fig. 4. In the limit of weak binding, oxygen
association becomes exothermic and barrierless, so phases like
ClOcand Cl(Oc)2cannot form. From the linear relations we
O2cc- O2(g) + 2c
has |DG| o 0 for DE(Oc) 4 2.97 eV. We therefore chose to
consider only OHcand Clcfor DE(Oc) 4 2.97 eV. The free
energies of OHcand Clcare within 0.01–0.27 eV depending on
the oxide, and we expect some coexistence in these regions of
phase space. The range of DE(Oc) for some rutile oxides is seen
in Fig. 5. For IrO2and RuO2the line at lowest DE(Oc) marks
the adsorption energy with free neighboring cus sites, and the
line at highest DE(Oc) marks the adsorption energy with Oc
neighbors as calculated in the (2 ? 1) unit cell. For PtO2and
TiO2the line at weakest binding marks the binding energy at
high Occoverage. The variation in adsorption energy with
coverage may be seen as an uncertainty arising from neglecting
adsorbate–adsorbate interactions. For the considered oxides
DE(Oc) is more affected by changing the oxide than changing
the Occoverage. The change of the most stable adsorbate
when the potential is increased is qualitatively reproduced for
Chlorine evolution activity. Firstly, we investigate the
mechanism involving ClOc
Oc+ 2Cl?(aq) - ClOc+ Cl?(aq) + e?
- Oc+ Cl2(g) + 2e?. (21)
of potential, U, and the surface reactivity descriptor, DE(Oc). Metal
ions are blue, O atoms red, hydrogen atoms white, and chlorine atoms
are green. The regions in the figure are determined by the most stable
surface configuration at the given potential. The phase borders are
defined by the equilibrium point of the reactions. So for example, the
border between the surface with Ocon the surface and the surface
with ClOcis defined by: Oc+ Cl?(aq) " ClOc+ e?, DG(Oc) ?
DG(ClOc) ? eUSHE= 0.
The most stable surface at pH = 0 and aCl? = 1 as a function
This journal is ? c the Owner Societies 2010 Phys. Chem. Chem. Phys., 2010, 12, 283–290 | 287
The potential at which all steps are neutral or downhill is:
U = Ueq+ |DG(ClOc) ? DG(Oc)|/e,
where Ueqis the equilibrium potential for chlorine evolution,
in this case 1.36 VSHE. Secondly, we investigate the mechanism
O2cc+ 2Cl?(aq) - Cl(Oc)2+ Cl?(aq) + e?
- Oc+ Cl2(g) + 2e?. (23)
The potential at which all steps are neutral or downhill is:
U = Ueq+ |DG(Cl(Oc)2) ? DG(O2cc)|/e.
Thirdly we consider a mechanism involving Clcadsorbed
directly at the metal cus site.
2Cl?(aq) +c- Cl?(aq) + Clc+ e?
- Cl2(g) +c+ 2e?. (25)
This mechanism could be relevant for oxides with weaker
adsorption energy at the cus site than RuO2. However, our
calculations suggest that this mechanism will be somewhat
poisoned by OHcformation at the cus sites. The potential
where all steps are neutral or downhill is:
U = Ueq+ |DG(Clc)|/e, (26)
RuO2and PtO2have |DG(Clc)| o 0.05 eV for high and low
coverage of Clc, respectively, and could in principle work as
good catalysts following this path; however, for RuO2we find
the cus sites to be blocked by O2cc.
Since the different chlorine evolution potentials all are
functions of DE(Oc), the potentials can be plotted directly on
the phase diagram as shown in Fig. 5. This is similar to the
Sabatier activity volcano curves known from heterogeneous
catalysis.34To have the surface phase diagram in the same plot
as the potential volcano directly assures that the activity
volcano and the stable surface configuration agree. In other
words, the different activity plots are relevant in different areas
of the phase diagram, which are easily obtained by looking at
Fig. 5. The thick black line marks the volcanoes, where the
mechanism involves one of the most stable surfaces as an
intermediate at the potential where all steps are neutral or
downhill in free energy. We note that the surface composition
during oxygen and chlorine evolution is not determined by
equilibrium, but rather by a steady state. However, it seems
equilibrium is one of the intermediates during the reaction.
Fig. 5 also shows the activity of IrO2and RuO2based on Oc
adsorption energy, with the error bars corresponding to the
variation of the Ocadsorption energy when going from low to
If the accuracy of the linear relations are taken into account,
the three investigated mechanisms form a single volcano with a
broad plateau for DE(Oc) from 1.5 eV to 3.5 eV. Deviations
from the linear relations could be important in this area. The
agreement between the detailed analysis for IrO2and RuO2
and the linear relations is therefore surprisingly good.
We find RuO2to be at the top of the volcano, whereas IrO2
binds Cl on top Octoo weakly. TiO2does not show up on this
activity scale. To our knowledge, only a few studies of the
relative activity of rutile oxides have been carried out. Kuhn
and Mortimer found IrO2and RuO2to have similar activities
and to be more active than TiO2. Mixtures of TiO2with Ir and
Ru are more active than mixtures of TiO2with Cr, Co, or Pt.35
Arikado et al. found the overpotential to increase in the order
RuO2o Ti/PtO2o IrO2.36Kelly et al. found the specific
activity of Ru sites at RuxTi1?xO2to be 45% more active than
the Ir sites at IrxTi1?xO2.4,37We note the discrepancy between
the relative activity of RuO2, IrO2, and PtO2could be because
different preparation methods may lead to different surface
roughness factors and different concentrations of residual
chlorine in thermally prepared oxides.4The high activity of
RuO2and IrO2relative to TiO2is in agreement with experi-
ments. The rutile crystal structure of PtO2is not the most
stable structure for PtO2. It is however possible that some
PtO2may be found in the rutile crystal phase if PtO2is mixed
with oxides that do form the rutile crystal phase.
For comparison the potential for oxygen evolution is also
shown in Fig. 5. It is seen that the potential for chlorine
evolution is lower than the potential for oxygen evolution in
spite of the lower equilibrium potential for oxygen evolution.
This is the reason why electrochemical chlorine evolution is
possible. It is also seen that a good oxygen evolution catalyst is
also a good chlorine evolution catalyst. A comparison of the
experimental potentials for OER and ClER has suggested that
the selectivity of oxides does not depend appreciably on the
catalyst material.2The potential of chlorine evolution changes
with the potential of oxygen evolution with a slope of 1.
Interestingly, one of the biggest outliers in the comparison
above was a Pt/MnO2catalyst in acid where the potential for
oxygen evolution was 0.3 eV higher than the potential of
chlorine evolution. MnO2 has an oxygen binding energy
around 3.2 eV.38Based on Fig. 5 we would therefore expect
the potential for chlorine evolution to be 0.4 V lower than the
potential for oxygen evolution. We note that at other pH
paths involving ClOc, Cl(Oc)2and Clc(from left to right). The domains
of the most stable surface structure as function of potential and
oxygen binding energy is marked by gray. To be truly active, the
intermediate should form at sites that are stable, as this makes the
active site abundant. The solid black line shows the combined Sabatier
volcano taking into account the stability of the active sites for a given
mechanism. The Sabatier volcano for oxygen evolution18(dashed blue
line) shows OER always requires a higher potential than ClER. The
activity of IrO2and RuO2are indicated with error bars derived from
the variation of the Ocadsorption energy with varying Occoverage.
Sabatier volcanoes (black dotted) for the considered reaction
288 | Phys. Chem. Chem. Phys., 2010, 12, 283–290 This journal is ? c the Owner Societies 2010
values the competition between chlorine and oxygen evolution
will change; see Fig. 1 and Fig. 2.
Single-crystal experiments on RuO2show that the (110)
surface is less active for Cl2evolution than the (101) and the
(320) surfaces.8,39On polycrystalline RuO2,40mixed RuO2+
TiO2,41and RuO2(320)39the activity depends on pH, whereas
the activity of RuO2(110) is independent on pH. The variation
of activity with pH has been explained by the reaction
OHc" Oc+ H+(aq) + e?
determining the availability of active Ocsites.7,39,40,42This
clearly requires Ocand OHcto be near equilibrium at the
reaction conditions for Cl2evolution. Since the bridge sites
and cus-sites on the (110) surface of rutile oxides fulfill the
same scaling relations between O and OH adsorption as
perovskites,32it is reasonable to assume the scaling relations
are identical for all rutile oxide surfaces. In this case Fig. 5
applies for any rutile oxide surface, but with the oxygen
adsorption energy depending on the specific surface facet. It
has been argued that the binding energy on the stepped (320)
surface is stronger than on the (110) surface.39If it is assumed
the O2ccand Cl(Oc)2intermediates form at the (320) surface as
well, Fig. 5 shows that as DE(Oc) is decreased from ca. 2.6 eV,
the OHc–O2ccequilibrium shifts to higher potential, which
leads to increased blocking of the active sites by OHcat a fixed
overpotential. We note the overpotential at constant current is
found to be 80 meV lower on the (320) surface than on
the (110) surface,39and thus within the vertical error bars
indicated in Fig. 5.
Based on DFT calculations, we have established linear scaling
relations between Cl, ClO, and O adsorption energies at the
cus-sites of rutile oxides. These linear energy relations enable
the construction of a generalized surface phase diagram where
potential and binding of oxygen are the descriptors determining
the surface composition. By applying an electrochemical–
thermodynamic approach we can make the first simple
theoretical analysis of the electrocatalytic chlorine evolution
reaction based on the free energies of the reaction inter-
mediates. A lower-bound to the overpotential required for
driving the reaction is thereby determined as function of
the oxygen adsorption energy. This approach is an electro-
chemical analogue to the Sabatier analysis used in hetero-
geneous catalysis. Combining the surface phase diagram and
the Sabatier volcano, one obtains a qualified suggestion for the
surface structure during reaction condition. The analysis
shows that ClO or Cl(Oc)2will form spontaneously on the
cus-sites of IrO2 and RuO2 at the potential required for
chlorine evolution. This indicates that the Cl2evolution occurs
through these intermediates on IrO2and RuO2. The potential
necessary for Cl2evolution is always smaller than the potential
for oxygen evolution for oxides exhibiting certain oxygen
adsorption energies. A simple explanation is that the oxide
evolution reaction involves three intermediates, and since the
bindings of these intermediates to the catalyst are linearly
related, there cannot be found a material that binds all of them
to the surface with exactly the right binding strength. This is
the reason for the high overpotential even at the top of the
oxygen evolution volcano. The chlorine evolution reaction, on
the other hand, involves only a single intermediate, and a
material that optimizes this bond could in principle exist. This
difference in overpotential is consistent with experiments2and
rationalizes experimental findings.
The authors thank Prof. M. T. M. Koper for useful discus-
sions. The Center for Atomic-scale Materials Design is funded
by the Lundbeck Foundation. This work was supported by the
Danish Center for Scientific Computing through Grant No.
HDW-1103-06, the European Commission (Marie Curie
Research Training Network MRTNCT-2006-032474) and
The Danish Council for Strategic Research though the
HyCycle Center (No. 2104-07-0041).
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