Co

7.17 Ammonia Synthesis: State of the Bellwether ReactionA Hellman, Chalmers University of Technology, Gothenburg, SwedenK Honkala, University of Jyvaskyla, Jyvaskyla, FinlandS Dahl, Technical University of Denmark, Kgs. Lyngby, DenmarkCH Christensen, DONG Energy, Fredericia, DenmarkJK Nørskov, SUNCAT Center for Interface Science and Catalysis, SLAC National Accelerator Laboratory, Menlo Park, CA, USA;Stanford University, Stanford, CA, USA

ã 2013 Elsevier Ltd. All rights reserved.

7.17.1 Introduction 4597.17.2 The Ammonia Synthesis Process 4607.17.3 Reaction Mechanism, Active Sites, and Reaction Kinetics 4627.17.4 Micro-Kinetic Models of Catalytic Ammonia Synthesis 4657.17.4.1 Activity Trends 4687.17.5 Discovery of New Catalysts for Ammonia Synthesis 4707.17.6 New Perspectives in Ammonia Synthesis 4717.17.7 Conclusion 473References 473

7.17.1 Introduction

It has been argued that ammonia synthesis from the elements is

among themost important discoveries of the last century. “What

is the most important invention of the twentieth century? Aero-

planes, nuclear energy, space flight, television, and computers

will be the most common answers. Yet none of these can match

the synthesis of ammonia from its elements.”1 The industrial

Haber–Bosch process enables the mass production of fertilizers

that is needed to feed the continuously growing population of

the Earth (Figure 1). Ammonia production currently consumes

more than 1% of all energy produced globally.2 This is note-

worthy because the energy efficiency of ammonia synthesis has

been improved dramatically since its discovery. Today, the

majority of all ammonia produced in the world is used directly

in the fertilizer industry, but new applications also are emerging,

such as the use of ammonia in the removal of nitrogen oxides

from exhaust gases by selective catalytic reduction (SCR). Still,

the volume of such applications is limited and, within the

foreseeable future, ammonia production is expected to follow

population growth unless new breakthrough applications are

developed, for instance, the use of ammonia as an energy

carrier.3

Any method for ammonia synthesis must address the neces-

sary activation of the extraordinarily strong triple bond in the N2

molecule. For instance, N2 activation in the gas phase requires

extreme conditions. However, such a method for nitrogen fixa-

tion was actually commercialized during the early twentieth

century, but its energy consumption was enormous.4 Nature

relies on an alternative route, where enzymes catalyze the step-

wise hydrogenation of the nitrogen molecule, which allows for

nitrogen fixation at ambient conditions.5,6 Mankind, however,

relies exclusively on the Haber–Bosch process for the supply of

ammonia not covered by the enzyme process. Today, the energy

consumption for this process is remarkably low compared with

othermeans of ammonia production, and the technology is well

suited for large-scale manufacturing.7,8

mprehensive Inorganic Chemistry II http://dx.doi.org/10.1016/B978-0-08-09777

The fundamentals of the technology behind the Haber–

Bosch process were originally developed by Fritz Haber9 and

Carl Bosch10 almost a century ago. First, Haber determined the

thermodynamic equilibrium of ammonia under atmospheric

pressure and high temperature, 1000 �C, with the help of an

iron (Fe) catalyst.11 Even though the amount of ammonia

produced was small, its success showed that ammonia could

be produced directly from gaseous N2 and H2. Later, Haber

suggested more optimal operational conditions: pressures

around 150–200 atm and temperatures around 500 �C. Fol-lowing this development, Bosch solved several technical prob-

lems related to operating a chemical plant at these extreme

conditions and, within his team, created the technology

required to commercialize ammonia production. The scientific

community soon recognized their pioneering work. In 1918,

Fritz Haber received The Nobel Prize in Chemistry “for the

synthesis of ammonia from its elements,” and in 1931, Carl

Bosch was awarded The Nobel Prize in Chemistry “in recogni-

tion of his contributions to the invention and development of

chemical high pressure methods.” Hence, the developments in

physical chemistry and chemical engineering in the early part

of the last century played a crucial role in the development of

the Haber–Bosch process, and vice versa.

During the period 1909–1912, Alwin Mittasch12 con-

ducted, for the first time ever, a large-scale screening experi-

ment to find a substitute to Haber’s more exotic osmium- and

uranium-based catalysts preferentially employed at that time.

In all, some 3000 catalyst compositions were tested in about

20 000 small-scale experiments. Eventually, he arrived at an

Fe-based catalyst, which actually has a composition very simi-

lar to that of the catalyst used industrially today.12 Although in

the same experiments ruthenium (Ru) was found to be an

interesting catalyst candidate,13 it was not until the 1970s14,15

that Ru was widely recognized as the best elementary-metal

catalyst for ammonia synthesis.16,17 Today, it is known that, at

low temperature and close to the thermodynamic equilibrium,

Ru exhibits considerably higher activity than Fe.16,18–21

4-4.00725-7 459

6

5

4

3

2

1

Wor

ld p

opul

atio

n (b

illio

ns)

Con

sum

ptio

n of

nitr

ogen

fert

ilize

r(m

egat

ons

of n

itrog

en)

01900 1925 1950

Year1975

100

80

60

40

20

0

Figure 1 The sudden growth in the global consumption of nitrogenfertilizer during the twentieth century was caused by the increase in theworld population. The strong correlation between the two indicates thatthe Haber–Bosch process is of paramount importance for increasingthe population. Reproduced from V. Smil, Nature 1999, 400, 415, withpermission Copyright © 1999 Scientific American, Inc. All rightsreserved.

460 Ammonia Synthesis: State of the Bellwether Reaction

However, because of the higher price of Ru and a shorter

catalyst lifetime, the dominance of iron-based catalysts has

only recently been challenged by promoted Ru catalysts22–24

in a few installations.

A simple view of heterogeneous catalysis is that the catalyst

provides a surface onto which the reactants temporarily adsorb

and, during this process, some intramolecular bonds are weak-

ened sufficiently for old bonds to be broken and new bonds to

be created. The bonds between the reaction intermediates and

the catalyst must be weak enough so that the products can be

released. In the Haber–Bosch process, reactant gas molecules N2

and H2 are dissociated and adsorbed onto well-dispersed, nano-

sized metal particles, and ammonia is formed by the stepwise

hydrogenation of nitrogen atoms on the catalyst surface. The

catalytic cycle is closed by the desorption of ammonia. The rela-

tive simplicity and importance of ammonia synthesis have made

it ‘the textbook example’ or ’the bellwether reaction‘ in heteroge-

neous catalysis. As such, it is commonly used as a test reaction to

develop new concepts and ideas (see also Chapter 7.01).25–29

The ultimate goal of catalysis research is to design and tune

the activity, selectivity, and lifetime of catalysts by controlling

their structural properties at the atomic level. The identification

of concepts to achieve this goal is still one of the key issues of

research in catalysis. A variety of different strategies have been

put forward, most of which are based on structure–reactivity

relationships, taking different aspects of these exceedingly

complex systems into account.

In the historical development of surface science and the

goal of atomic understanding of heterogeneous catalysis,

ammonia synthesis also has played a central role.25,30 This

development has allowed the design of new ammonia catalysts

based on atomistic insight20 and, lately, made it possible to

understand and quantitatively describe ammonia synthesis

from first-principle calculations.31

7.17.2 The Ammonia Synthesis Process

In the Haber–Bosch process, ammonia is synthesized from a gas

containing H2 and N2 at approximately stoichiometric ratio.

N2 þ 3H2 P2NH3 � 99:22kJ mole�1

Importantly, ammonia synthesis is an equilibrium-limited

exothermic reaction at relevant reaction conditions, and it is

required that the synthesis gas must be free of any reactive

contaminants that poison the catalyst. It is mainly these

requirements that determine the optimal process design of an

ammonia synthesis plant and an ammonia reactor.

A modern implementation of the Haber–Bosch process

provides a highly efficient method for producing ammonia in

large quantities. There are many slightly different process

designs,32 all of which are based on synthesis gas production

from a carbonaceous feedstock, steam, and air, followed by

ammonia synthesis in a recycle loop. A standard process

scheme utilizing natural gas as feedstock is shown schemati-

cally in Figure 2.

The first step in the process is desulphurization of the

natural gas feed in order to avoid sulfur poisoning of the

downstream catalysts. The second step is steam reforming,

where the natural gas is reacted with steam to produce an

equilibrium mixture of H2, CO, CO2, and CH4. This reaction

is endothermic, so it is carried out in tubular reactors placed in

a heated furnace in order to supply the heat for the reaction

and to maximize the equilibrium content of the desired prod-

ucts H2 and CO. Following this, air is added to supply the N2

required for ammonia production, while the O2 from the air

converts the remaining CH4 in an exothermic reaction that

increases the temperature and the H2 and CO content further.

The reaction mixture is then cooled and followed by a two-step

water–gas shift process, where CO is reacted with steam and

transformed into H2 and CO2. Since all oxygen-containing

molecules poison the ammonia synthesis catalyst, CO2 is sub-

sequently removed by absorption and a final CO/CO2 cleanup

is carried out, for example, by the methanation reaction.

Before sending the synthesis gas into the ammonia reactor,

it is compressed and, at a point before it reaches the ammonia

catalyst, water is removed by condensation. As mentioned

above, ammonia synthesis is an equilibrium-limited exother-

mic reaction, and the ammonia product is favored by low

temperature and high pressure. The high reaction temperature

necessary for obtaining sufficiently high reaction rate, together

with constraints on the synthesis pressure, excludes full con-

version of the synthesis gas.32 This means that most of the N2

and H2 passes unreacted through the reactor, and must be

recycled in the loop after being cleaned for ammonia. Since

the synthesis gas that is fed into the loop contains some inert

gases (CH4 and Ar), a purge is also needed in the loop to avoid

too large a build-up of these diluents.

Optimal heat management is important for obtaining an

energy-efficient ammonia synthesis process. The coupling

between the exothermic process steps (secondary reforming,

water–gas shift, ammonia synthesis, and cooling steps) and the

endothermic process steps (steam reforming water evaporation

and heating steps) is achieved with suitable heat exchangers

and proper steam generation and usage. This integration of

Ammonia Synthesis: State of the Bellwether Reaction 461

the processes, together with the high-pressure conditions

(100–200 bar) for the Haber–Bosch process, means that very

large plants are most efficient and that the huge global ammo-

nia production is predominantly carried out in a few hundred

centralized plants.

When designing the high-pressure ammonia reactor, an

important criterion is to keep the reactor and catalyst volume

small. As the temperature in an adiabatic reactor will increase as

ammonia is synthesized by 14–18 �C for each percent of ammo-

nia produced,2,34 somekind of cooled reactormust be employed

to reach a reasonable ammonia concentration. To illustrate the

reactor design process, the so-called optimal operation lines are

an essential aid. Using only a single catalyst, it is well known that

theminimumcatalyst volumewill beobtained in a reactorwhere

the optimal operating line is followed as closely as possible. The

Natural gas

Air

400–500 �C

1200 �C

1000 �C

800 �C

Desulphu-rization

Primaryreforming tubular

Secondaryreforming Shift Methanato

350 �C

400 �C

300

325

200 �C

220 �C

H2O

CO2

Figure 2 Process scheme of ammonia synthesis. (Courtesy of Haldor Tops

350 450

% N

H3

550Temperature (�C)

Figure 3 The schematic picture of an internally cooled reactor and of an adiand operational lines are shown in the ammonia concentration versus tempeparticular reactor is also shown. Reprinted with permission from Jacobsen, CNørskov, J. K., J. Catal. 2002, 205, 382–387.

optimal operating line, occasionally also called the maximum

rate line, illustrates the temperatures at which themaximum rate

is reached at given pressure and gas composition. Different syn-

thesis loop configurations have been proposed to make the

actual operation line follow the optimal operating line as closely

as possible. Figure 3 shows how it is done in an internally cooled

reactor with a countercurrent flow of synthesis gas in the cooling

tubes or in an indirectly cooled reactor, where the gas is cooled

by heat exchange between the individual beds. The resulting

operating lines are also shown. These two reactor designs have

been extensively used in the ammonia synthesis industry.

The Tennessee Valley Authority reactor that was common from

1930 to 1965 typifies the internally cooled reactor. Today, most

new industrial ammonia synthesis reactors are adiabatic two- or

three-bed radial-flow arrangements with indirect cooling.

450–500 �C

rPurge

synthesis

�C

300 �C

NH3

KMRKM

100 �C

0 �C

�C

oe AS.)

350 450

% N

H3

550Temperature (�C)

abatic three-bed reactor, cooled by two heat exchangers. The equilibriumrature plot. The corresponding operation lines for that reaction in that. J. H.; Dahl, S.; Boisen, A.; Clausen, B. S.; Topsøe, H.; Logadottir, A.;

00.5N2+ 1.5H2

0.5N2-* + 1.5H2

oten

tial e

nerg

y/kJ

mol

-1

NH2-* + H-*

NH3-*

NH3

17 Fe~21

~33

50

46

~41

259

106

-100

-200

462 Ammonia Synthesis: State of the Bellwether Reaction

In such reactors, the ammonia concentration and the tempera-

ture vary in the two (or three) beds, as shown in Figure 3. The

position of the optimal operation lines depends on the type of

catalyst; however, the differences are small and the lines are

generally positioned 30–60 �C below the equilibrium curve.

The absolute reaction rate is much more dependent on the

catalyst of choice and, in Section 7.17.5, we will discuss how

the reactor volume can also be minimized by choosing different

catalysts in different parts of the reactor.

N-* + 3H-*

P

NH-* + 2H-*-300

Figure 4 The reaction energy diagram for ammonia synthesis on aniron surface as inferred from surface-science measurements. Energiesare in kJ mol�1. Reprinted with permission from Hinrichsen, O. Catal.Today 1999, 53, 177–188.

7.17.3 Reaction Mechanism, Active Sites, andReaction Kinetics

In heterogeneously catalyzed reactions, like ammonia synthesis,

the catalytic action comes from the interaction of reactants and

reaction intermediates with the active surface sites of a solid

catalyst. The direct link between activity and the number of active

sites and the output of products froma catalyst has been a spur to

establish an understanding of surface reactions at the atomic

level.35 However, it has taken more than half a century for the

surface-science community to develop techniques that can rou-

tinely probe the atomicprocesses of interest.36,37 Throughout the

1960s to the early 1990s, thousands ofwell-defined single-crystal

surface structures were studied with various surface-science

techniques, such as low-energy electron diffraction, Auger elec-

tron spectroscopy, and x-ray photoelectron spectroscopy (XPS).

These studies were followed by experimental and theoretical

investigations of the nature of the chemical bonding of adsor-

bates at surfaces and interfaces. Detailed electronic structures of

surface chemical bonds were obtained by spectroscopic tech-

niques, such as ultraviolet and XPS, and x-ray emission spectros-

copy. First-principles calculations helped to rationalize the

experimental findings and to provide a consistent framework

for understanding surface chemical bonding.38,39 For example,

the d-bandmodel correlates the binding strengthof an adsorbate

to a transition-metal surface and the local d-electron density of

the binding site.28 Along with the experimental data, the d-band

model provides answers to fundamental questions such as why

bulk gold is noble while small gold clusters are not,40,41 andwhy

surface defect sites (steps and kinks) are more chemically active

than sites of higher coordination on flat surfaces.42

Long before surface-science experiments43–48 provided direct

observations of the reaction intermediates, the mechanism of

ammonia synthesis was basically resolved. Already as early as

1933, Emmett et al.49,50 found strong experimental evidence that

N2 dissociation over Fe catalysts is the rate-limiting step for

ammonia synthesis under industrially relevant conditions.

About 30 years later, surface-science studies revealed a detailed

picture of the N2 dissociation process.45–48,51,52 Following the

identification of the rate-determining step, the connection

between ammonia synthesis at elevated temperatures and pres-

sures and ultra-high vacuum surface-science measurements on

iron-based catalysts (single-crystal surfaces)was established both

experimentally45–48,53–55 and theoretically (Figure 4).25,56–59

The first surface-science studies on ammonia synthesis

focused on N2 dissociation and subsequent ammonia forma-

tion over various Fe crystal surfaces at low pressures. Based

on these results, it was established that the dissociation reac-

tion is structure sensitive and the open Fe(111) and Fe(211)

surfaces are the most active.61 This was attributed to the access

to the so-called C7 active site (surface atoms with seven

nearest neighbors, Figure 5). Activity measurements at high

pressure on similar single-crystal surfaces showed that Fe

crystal faces with C7 sites have significantly higher catalytic

activity for ammonia synthesis than other Fe crystal faces

(Figure 5).54 These findings, together with the investigations

of elementary reaction steps45,62 and the effects of catalyst

support and additives,55,62 were important in establishing

an understanding of industrial ammonia synthesis based on

surface science.

First-principles studies63,64 show that the adsorption energy

of N on Fe(100) is higher compared to the more close-packed

(111) and (110) surfaces. Especially, the adsorbed phase c

(2�2)-N/Fe(100) is particularly stable, indicating that this

phase will form under high pressure of nitrogen. Consequently,

at higher coverage of N, the surfaces (111) and (110) might very

well reconstruct to form a stable c(2�2)-N/Fe(100) overlayer.

Furthermore, the dissociation of N2 on Fe(111) has, essentially,

two different pathways: one with a low energy barrier but with a

large entropic barrier, and the other with a high energy barrier.64

The low-energy-barrier path includes several precursor states,

which results in a complicated dissociation path. Under normal

ammonia synthesis conditions, the low-barrier, high-entropy

path will dominate, but at the highest temperatures, the high-

barrier process may become more effective.

After studying ammonia synthesis over iron, the attention

turned to ruthenium, which is a more active catalyst than iron,

especially at high ammonia concentrations. This gave rise to

industrial interest in ruthenium-based ammonia catalysts in

the 1990s.

In the first attempts to measure the N2 dissociation rate on

different Ru surface facets, the process was found to be struc-

ture insensitive.65 As with the case of iron, it was possible to

model the activity of technical catalysts at high-pressure con-

ditions based on the surface-science results,60 so there seemed

to be some consistency. However, there was one problem;

experimentally, the activation energy for the dissociation reac-

tion was found to be around 0.40 eV on the Ru(0001) surface,

which disagrees with a calculated barrier of 1.36 eV.66 The

answer to this discrepancy was found in a joint experimental

and theoretical effort,67 where dissociative adsorption of N2

Ammonia Synthesis: State of the Bellwether Reaction 463

over a closed packed Ru(0001) surface was revisited. Measure-

ments were done on both clean and Au-passivated surfaces. In

the latter case, small amounts of Au (1–2% on the surface)

were deposited on the surface, where they preferentially deco-

rate reactive step sites. The results show that the presence of Au

reduces the N2 dissociation rate almost completely, demon-

strating the high structure sensitivity of N2 dissociation on Ru.

Based on the measurements, it was concluded that the reaction

takes place solely at the step sites, with a dissociation barrier of

0.4 eV. The rate on the planar terrace was measured to be at

least nine orders of magnitude smaller than that at the steps at

500 K. The first ab initio calculations supported this and they

gave a corresponding 1.5-eV barrier difference. Laser-assisted

associative desorption of N2 from Ru(0001) also confirmed

the reactivity difference between step and terrace atoms.68

14

12

10

8

6

4

2

0(111) (211) (100)

Surface orientation

Mol

NH

3/cm

2 s

´10

-9

(210) (110)

T = 400 �C

20 atm 3:1 H2:N2

Fe

C4

C5

C7

C7

Fe

Figure 5 Produced ammonia over different Fe facets (left) and the correspoLi, Y. Proc. Natl. Acad. Sci. 2011, 108, 917–924.

-10

-11

-12

-13

0.0015 0.0020 0.0025

1–2% Au on Ru(0001)Ea= 1.3 ± 0.2 eV

Ru(0001)Ea= 0.4 ± 0.1 eV

0.00301/T (1/K)

log

S0

0.0035

-14

-1

0DE [e

V/N

2] 1

2

Figure 6 (Left) An Arrhenius plot of measured thermal sticking coefficients o0.01–0.02 mL of gold. (Right) Results from density-functional calculations comcurve shows the adsorption and transition-state energies for the dissociation onThe energy zero is taken to be the energy of the N2 molecule in the gas phase. ReLarsen, J. H.; Chorkendorff, I.; Tornqvist, E.; Nørskov, J. K. Phys. Rev. Lett. 19

Figure 6(a) highlights the difference in sticking coefficient

observed on these surfaces, and Figure 6(b) compares the

calculated N2 dissociation barriers over stepped and flat sur-

faces. The barrier on a step site is very close to the experimen-

tally observed value on bare Ru, which strongly supports the

interpretation that steps (or sites with similar characteristics)

dominate the N2 dissociation.

The conclusion was that, to efficiently break the strong

triple bond, Ru steps should expose a threefold hollow site

and a bridge site close together and form an active site for

dissociation. In this configuration, atoms of a nitrogen mole-

cule do not simultaneously bind to the same Ru surface atom

at the transition state during dissociation (see the structure at

the bottom of Figure 6(b)). This has later been confirmed by

other studies.69,70 According to the terminology developed by

The first layer The second layer The third layer

(111)

C4

C6

C4 C5

C7C8

C8

Fe(210)

(211)

Fe(100)

Fe(110)

nding structures (right). Reprinted with permission from Somorjai, G. A.;

.0

.0 N2,ad

TS

2Nad

.0

.0

f N2 on a clean Ru(0001) surface and the same surface covered withparing N2 dissociation on a terrace and at a step on Ru(0001). The upperthe terrace, whereas the lower curve shows the same energies at the step.printed figure with permission from Dahl, S.; Logadottir, A.; Egeberg, R. C.;99, 83, 1814. Copyright 1999 by the American Physical Society.

464 Ammonia Synthesis: State of the Bellwether Reaction

Van Hardeveld and Van Montfoort,71 these active sites are

called B5 sites. As can be seen in the top part of Figure 6(b),

on the terrace, one Ru surface atom is bonded to both nitrogen

atoms, which leads to the high barrier. The higher reactivity of

step atoms compared to terrace atoms is due to the lower

coordination number and higher lying d-states, which, alone,

however, cannot explain the large difference in the activation

barrier. The size of this electronic effect can be estimated from

the difference in adsorption energies of nitrogen atoms at the

right of the energy diagrams in Figure 6(b). It is possible to

obtain a lower activation barrier for N2 dissociation at the step

sites without affecting the adsorption strength of the nitrogen

atoms, which is the reason why the steps totally dominate the

ammonia synthesis reaction. This is seen in Figure 7(a), where

the calculated potential energy diagrams for ammonia synthe-

sis over flat and stepped Ru surfaces are shown. The compari-

son of rate constants presented in Figure 7(b) shows that N2

dissociation over the step edge is a rate-limiting step in ammo-

nia synthesis, despite the fact that NH2 hydrogenation has

higher activation energy. Had the nitrogen atoms been bound

much stronger to the step sites, the hydrogenation steps would

have been more uphill in energy, making them rate limiting.

Thus, the full benefit of the low activation barrier for N2

dissociation on the step sites would not have been so large.

This enormous difference in reactivity between steps and

terraces has two consequences. First, it means that the rate of

N2 dissociationonRu(0001) is completely dominated by as little

as a fraction of a percent of steps on the surface. The steps, thus,

dominate all thermal surface-science experiments of N2 adsorp-

tionon single-crystal surfaces. Since the number of step sites does

not varymuch between different surface facets, this explains why

the reaction was initially found to be structure insensitive.

Second, it means that the ammonia synthesis reaction should

be extremely structure sensitive onRu. For nanoparticles, the step

density depends strongly on the particle size, suggesting that the

catalytic activity of Ru catalysts for ammonia synthesis should be

dependent on the catalyst particle size.

200

3H2

N2

NH3

NH3TSterrace

TSstep

100

0

-1.00

Eto

t (eV

/mol

ecul

e)

-2.00

-3.00

-4.00

-5.00

*+ N 2

+3H 2

N 2 *+

3H 2

2N *+3H 2

2N *+6H *

N *+NH *+

5H *

N *+NH 2

*+4H *

N *+NH 3

*+3H *

N *+NH 3

+3H *

NH *+NH 3

+2H *

NH 2 *+

NH 3 +

H *

NH 3 *+

NH 3

*+2N

Figure 7 (left) The comparison of potential energy landscape for ammoniathat the step effect is less pronounced for the hydrogenation step, which resultThe comparison of the rate constants for NH2 hydrogenation and N2 dissociathan the dissociation barrier, N2 dissociation is by far the slowest step due totransition-state theory. Reprinted with permission from Logadottir, A.; Nørsk

Jacobsen et al. 73 were first to count the relative number of

B5-type sites present on Ru nanoparticles. By simple arguments,

it was estimated that the maximum probability for B5 sites was

obtained for particles in the range of 1.8–2.5 nm, and for larger

particles, the probability for B5 sites decreased monotonically.

For smaller particle sizes, the B5-type sites disappear, simply

because the crystals have to be of a certain size before these

sites are formed. Honkala et al. 31 calculated the surface energies

of the most relevant facets of Ru and used these to create

an atomistic Wulff construction for several Ru particles with

different diameters. After including edge effects, the B5 sites

appear along the (0001)/(10–11) intersection, and an estimate

for the number of active sites per particle size can be conducted

(Figure 8). It was concluded that there exists a lower limit to

the optimal ruthenium crystal size, which could explain the

measurements74,75 for promoted and unpromoted Ru/MgO cat-

alysts, and for promoted Ru/MgAl2O4 catalysts.73

An even more advanced estimate was presented by

Gavnholt and Schiotz.70 They took into account all cluster

sizes between 1.5 and 5 nm in diameter, and also included

the temperature dependence on the structure. The obtained

values for the number of sites are quite low, with a peak

value between 1.5 and 2.5 mol g�1 at particle diameter 3.5

nm and 30 mol g�1 at particle diameter 2 nm. These values

are approximately an order of magnitude lower than the esti-

mates made by Honkala et al.31 However, other sites included

in the study by Gavnholt and Schiotz70 gave significant contri-

butions to the overall rate. The graphs at 300 and 700 K in

Figure 9 indicate a maximum in the catalytic activity per vol-

ume catalyst at a cluster diameter of 3 nm, even though the

curve is not so smooth. At 300 K, it is clearly seen that all the

catalytic activity comes from step sites A and B, whereas step

site D starts contributing at 700 K. At 1200 K, the picture is a bit

clearer and smoother. There still seems to be an optimal cluster

diameter at 3 nm. The existence of such a maximum fits well

with the experimental observation that the catalytic activity

of a sample can increase after sintering of the smallest Ru

0.001 0.0015

500 ºC700 K

350 ºC

k1

k4

0

0

0

0

0

0

0.002 0.0025

1/T [K-1]

0.003 0.0035H 3

formation over terrace and step sites on Ru(0001). The plot showss from the fact that transitions states are similar on both surfaces. (Right)tion over stepped Ru. Despite that the hydrogenation barrier is higherthe entropic effects. The rate constants were estimated with harmonic

ov, J. K. J. Catal. 2003, 220, 273–279.

1

0.1

0.01

0 2 4 6

Active sites

Edge atoms

Smallcrystal

Ru(0001)surface

Step atom 1 nm

8Crystal size (nm)

Frac

tion

of t

otal

ato

ms

10

Figure 8 AWulff-shaped Ru particle, with an average diameter of 2.9 nm.Atoms that belong to active B5 sites are shown in red. Also depictured is aTEM image of a supported Ru particle with a step. The fraction of edgeatoms and active sites on small Ru crystals are shown relative to the totalnumber of atoms as a function of crystal size. Reprinted with permissionfrom Honkala, K.; Hellman, A.; Remediakis, I. N.; Logadottir, A.; Carlsson,A.; Dahl, S.; Christensen, C. H.; Nørskov, J. K. Science 2005, 307, 555–558and Jacobsen, C. J. H.; Dahl, S.; Hansen, P. L.; Tornqvist, E.; Jensen, L.;Topsoe, H.; Prip, D. V.; Moenshaug, P. B.; Chorkendorff, I., J. Mol. Catal. AChem. 2000, 163, 19–26.

Diameter/nm

1

2.5Step A

Step B

Step C

Step D

1.5 2 2.5Bond length / angstrom

Ene

rgy/

eV

step Astep Bstep Cstep D

3 3.5 4.54 5

Diameter/nm Diameter/nm

2

1.5

1

0.5

0

-0.5

-1

300 K

2.0 3.0 4.0 5.0 5.0 5.02.0 2.03.0 3.04.0 4.0

700 K 1200 K

Act

ivity

per

vol

ume

(arb

itrar

y un

it)

Act

ivity

per

vol

ume

(arb

itrar

y un

it)

A A A

B

B

B

D

D

Figure 9 Gavnholt and Schiotz70 define four different step sites.The calculated activation barriers range between 0.5 and 2 eV.Depending on the operational temperature, the contribution to theammonia production varies between the different step sites. Reprintedfigures with permission from Gavnholt, J.; Schiotz, J. Phys. Rev. B2008, 77, 035404, Copyright 2008 by the American Physical Society.

Ammonia Synthesis: State of the Bellwether Reaction 465

particles.73 Furthermore, it is seen that the step site D also

contributes to the total activity at this higher temperature,

due to fact that it has a greater presence on the clusters than

the other two and the fact that the difference in barrier height

becomes less important as the temperature increases. The

ammonia synthesis typically runs at a temperature of around

700 K in industrial plant.

7.17.4 Micro-Kinetic Models of Catalytic AmmoniaSynthesis

As described earlier, N2 dissociation is the rate-limiting step in

catalytic ammonia synthesis. The full reaction follows a Lang-

muir–Hinshelwood (LH) mechanism, where H2 is also disso-

ciatively adsorbed on a surface before adsorbed nitrogen is

stepwise hydrogenated to form ammonia (see Figure 7). Each

reaction step obeys microscopic reversibility, and the full set

elementary reaction steps is:

H2 þ 2*P2H*N2 þ 2*P2N*N*þH*PNH*þ *NH*þH*PNH2*þ *NH2*þH*PNH3*þ *NH3*PNH3 þ *

where * and X* correspond to an empty site and an adsorbed

species X, respectively (see also Chapter 7.02).

Based on the reactionmechanism, it is possible to construct a

micro-kineticmodel.76,77 Suchmodels are extensively applied to

develop and verify an atomistic understanding of catalysis. The

success stories include industrially relevant catalytic reactions,

such as methanol synthesis78 and decomposition,79,80 water–

gas shift,81 ethylene oxidation,82,83 and ammonia synthesis.84,85

These examples show that the micro-kinetic models can bridge

the temperature, pressure, and material gaps that exist between

fundamental surface-science studies and industrially relevant

catalysis (see also Chapter 7.03).

A key step in constructing a micro-kinetic model is to for-

mulate an expression for a catalytic rate per active site based on

describing the elementary reaction steps.76,77,86 Usually, it is

necessary to make some approximations in order to obtain a

solvable model. Mean-field micro-kinetic models, where inter-

actions between adsorbates are neglected, are, in many cases,

adequate for a quantitative description of the reaction rate. To

analyze trends for the activity of different catalysts, the mean-

field models have distinct advantages, since the additional

assumptions, such as the inclusion of a rate-determining reac-

tion step and the steady-state approximation, make the model

entirely analytical.87 However, during typical reaction condi-

tions, the coverage of reactant, intermediate, and product adsor-

bates can be substantial, which can challenge the mean-field

assumption. To overcome this limitation, one can employ

extensions, such as mean-field expressions for lateral interaction

or even quasi-chemical approximations that include some local

structure information of the adsorbates.88 The other possibility

is to combine the micro-kinetic formulation with Monte Carlo

(MC) simulations of the surface coverages.31 Ammonia

466 Ammonia Synthesis: State of the Bellwether Reaction

synthesis, owing to N2 dissociation being the rate-determining

step and at equilibriumwith the products, is well suited for such

a combination.31

By combining activation barriers and adsorption energies

from density functional theory (DFT) calculations72 and a

micro-kinetic model, it is possible to predict the relative reac-

tion rates over different catalyst materials.67 This methodology

has been particularly successful for ammonia synthesis.19,31,72

The obtained results have provided the foundation for the

volcano curve16 and have, furthermore, given the basis for a

rational design approach to new catalysts. For more informa-

tion, see Section 7.17.5.

Micro-kinetic models of ammonia synthesis over iron-

based catalysts show a remarkable agreement with laboratory

measurements, both over single crystals and industrial

catalysts,56,58,76,86,89,90 see Figure 10(a) for an early example.

Different micro-kinetic models can give good overall agree-

ment with experimental results; therefore, the success of a

micro-kinetic model cannot be used as proof for a particular

reaction mechanism.77,91 Because of this, it is very important

to take as many input parameters as possible from indepen-

dent measurements, instead of obtaining them by fitting to the

overall rate data.

The predictive power of micro-kinetic models combined

with surface-science techniques are perhaps best exemplified

by the quantitative micro-kinetic model for a multipromoted

iron catalyst developed by Sehested et al.89 The parameters are

close to those derived from surface-science experiments and

based on Langmuiran adsorption. Here, the activity of a multi-

promoted iron catalyst is measured at total pressures from 1 to

100 bar and at temperatures from 320 to 440 �C, with the

hydrogen to nitrogen ratio varied by a factor of 10. The pre-

dictions of the model were compared to activity measurements

conducted during the study, and also to independent measure-

ments of Nielsen.26

Experimental exit NH3 mole fraction

1

10-1

10-2

1

1 atm

150 atm

300 atm10-3

10-3 10-2 10-1

Cal

cula

ted

exi

t N

H3

mol

e fr

actio

n

Figure 10 Ammonia concentration predicted from two micro-kinetic modelReprinted figure with permission from Stoltze, P.; Nørskov, J. K. Phys. Rev. LSehested, J.; Jacobsen, C. J. H.; Tornqvist, E.; Rokni, S.; Stoltze, P. J. Catal.

The micro-kinetic modeling of ammonia synthesis is easier

on Ru than on Fe,84,85 partly due the lack ofmolecular precursor

states of N2. Dahl and coworkers developed a micro-kinetic

model for nonpromoted ruthenium based predominantly on

surface-science observations.84,85 The main assumptions are

that N2 dissociation is a rate-determining step, and only the

step sites of an Ru(0001) surface are active in the reaction. The

model describes very well the rates of ammonia formation over a

ruthenium single crystal and over a supported Ru-based catalyst,

as shown in Figure 11. The input parameters for adsorption and

activation energies were, to a large extent, based on surface-

science results for the relevant species on ruthenium, although

some estimates were needed for parameters that are not directly

available from measurements. The model is also in excellent

agreement with an N2 isotope scrambling experiment over an

Ru/MgO catalyst (Figure 11). The fact that the micro-kinetic

results show good agreement with many widely different exper-

imental results strongly suggests that the model gives a funda-

mentally correct picture of ammonia synthesis over ruthenium.

Given the relative simplicity of the reaction mechanism on

Ru, it is natural that ammonia synthesis was the very first hetero-

geneous catalytic reaction that was fully characterized from first

principles. Its productivity was predicted for a supported techni-

cal catalyst with a surprisingly high level of accuracy.31,91,93 All

energy-related input parameters for the micro-kinetic model

were calculated with DFT. These include adsorption and activa-

tion energies, interaction energies, and entropy contributions.

The only input from the catalyst is the average particle size

distribution determined by high-resolution transmission elec-

tron microscopy. An MC technique was employed to determine

adsorbate coverages employing pair interactions from DFT

calculations, which takes the model beyond the ordinary

mean-field approximation. The absolute productivities are in

semi-quantitative agreement with the experimental data mea-

sured over an industrial, high-surface catalyst at various

10.0

1.0

1.0 10.00.1

0.1Measured ammonia concentration (%)

Cal

cula

ted

am

mon

ia c

once

ntra

tion

(%)

s,56,89 plotted as a function of the measured ammonia concentration.ett. 1985, 55, 2502. Copyright 1985 by the American Physical Society;1999, 188, 83–89.

Ammonia Synthesis: State of the Bellwether Reaction 467

temperatures, flows, and ratios; see Figure 12. For the depen-

dence of flow, temperature, and H2:N2 ratio, the first-principles

model even provides quantitative agreement with the experi-

mental data; see Figure 13.

70

60

50

40

30

20

10

0.010.01

Experimental output (NH3-%)

Mod

el o

utp

ut (N

H3-

%)

x (29

N2)

(pp

m)

0.1

0.1

1

1

10

10

Figure 11 Comparison between measured and calculated NH3 production ftemperature, 320–440 �C. The solid line is the calculated outlet concentrationpassed over 138 mg of catalyst containing 20 mmol g�1 of active sites. The fThe conditions are equal to the ones used by Hinrichsen et al.92 over the Ru/permission from Dahl, S.; Sehested, J.; Jacobsen, C. J. H.; Tornqvist, E.; Cho

10

10

1

1

0.1

0.1

0.01

0.010.001

0.001

10

10

1

1

Experimental ammonia output (NH3-%)

Experimental ammonia output (NH3-%)

0.1

0.1

0.01

0.010.001

0.001

(a)

(c)

Mod

el a

mm

onia

out

put

(NH

3-%

)M

odel

am

mon

ia o

utp

ut (N

H3-

%)

Figure 12 The calculated ammonia synthesis/decomposition output as calcuan industrially supported Ru-based catalyst. Reprinted with permission fromDahl, S.; Christensen, C. H.; Nørskov, J. K. Surf. Sci. 2009, 603, 1731–1739.

In addition, an explanation of why productivity can be

predicted with such a good accuracy, although the activation

energies of individual elementary steps are less accurate, was

provided by Honkala et al.31 It turns out that the network of

00

00

00

00

00

00

00

0450 500 550 600 650

T (K)700 750 800

or the Ru/MgA1204 catalyst: pressure, 1–100 bar; H2:N2 ratio, 6:1–1:4;of 29N2 when an inert gas containing 0.67% 29N2 and 0.60% 30N2 is

low is 50 Nml min�1 and the reactor is treated as a plug flow reactor.MgO catalyst, and the open circles are their results. Reprinted withrkendorff, I. J. Catal. 2000, 192, 391–399.

10

10

1

1

0.1

0.1

0.01

0.010.001

0.001

Experimental ammonia output (NH3-%)

Experimental nitrogen output (N2-%)

(b)

(d)

10

10

100

100

1

1

0.1

0.1

0.01

0.01

0.001

0.0010.0001M

odel

nitr

ogen

out

put

(N2-

%)

Mod

el a

mm

onia

out

put

(NH

3-%

)

0.0001

lated from the first-principles model and compared to measurements onHellman, A.; Honkala, K.; Remediakis, I. N.; Logadottir, A.; Carlsson, A.;

468 Ammonia Synthesis: State of the Bellwether Reaction

consecutive reactions is self-regulating in the sense that the

errors in the different parts of the reaction cancel each other.

This is related to the so-called compensation effect,87 which

gives hope that one may more generally be able to calculate

catalytic rates directly from first principles.

In the present discussion about reaction mechanism and

active sites in ammonia synthesis, we have neglected the fact

that all ammonia synthesis catalysts are promoted by alkali

metals, which is central for preparing the most active catalysts.

Alkali-metal promoters like K and Cs added to ruthenium

catalysts increase the activity for ammonia synthesis. In addi-

tion, the reaction order of ammonia changes from negative to

about zero, whereas the reaction order for hydrogen decreases

and can end up at �1.85 It is well known that alkali metals

adsorbed on a metal surface donate electrons to the surface.

Theoretical calculations have shown that the presence of alkali

metal on metal surfaces lowers the N2 dissociation barrier

through a direct electrostatic interaction.64 The rate will

increase due to the higher dissociation rate and smaller inhi-

bition by NH*, and the reaction order for ammonia will

increase as an effect of the latter. The model suggested by

Dahl et al.85 can qualitatively explain all the changes observed

when ruthenium catalysts are promoted with alkali metal due

to the rather high coverage of NHx, x ¼ 1–3, during synthesis

over the nonpromoted catalysts. So, as was the case for step

sites compared to terrace sites, the success of alkali-metal

Temperature (K) Flow (m

40

1

60 80 100

Pro

duc

tivity

(NH

3-%

)

580

(b)(a)

0.01

0.1

1

600 620 640 660

MeasuredCalculated

680 700 720

Figure 13 Ammonia productivity as a function of (a) temperature, (b) flow,(STP); (b) 713 K, N2:H2 (1:3); (c) 713 K, 30 N ml min�1 (STP), all at a total preRemediakis, I. N.; Logadottir, A.; Carlsson, A.; Dahl, S.; Christensen, C. H.; N

160

Rat

e (S

TP m

l NH

3/h

)

−ΔH: (kcal / metal atom)12080

100

10

1.0

0.1

0.01

MoRe

Ru

Os

Fe

Co

Rh

(a)

Pt

Ir

Ni

40

Figure 14 The rate of (a) ammonia synthesis and (b) ammonia decompositReproduced from Aika, K.; Yamazaki, K.; Ozaki, A. Chem. Lett. (Jap.) 1973, 2

promoters is based on their ability to lower the dissociation

barrier of the rate-limiting N2 dissociation step without

increasing the binding of N-containing species to the surface.

7.17.4.1 Activity Trends

In this section, it is shown how knowledge of the reaction

mechanism and structure of the active site has led to a funda-

mental understanding of the trends in catalytic ammonia syn-

thesis activity of transition metals.

In 1911, Sabatier formulated a principle of what constitutes

a good catalyst.94 He states that a catalyst should neither inter-

act too strongly or too weakly with the reactant. The reason is

that desorption from a reactive metal catalyst is slow and will

increase on less reactive metals. However, on very noble

metals, the large energy barrier for dissociation decreases the

dissociation rate. Thus, the best catalyst must be a compromise

between the two extremes.

Sabatier’s principle provides the rationale behind the vol-

cano curves95 often encountered in heterogeneous catalysis.

The volcano curves display the turnover frequency (or some

other activity-related property) against a property, such as the

heat of adsorption of the reactant by the catalyst. The typical

appearance looks like a triangle or an inverted parabola, and as

seen in Figure 14, this also holds for ammonia synthesis over

different transition metals.

l/min [STP]) Ratio (N2:H2)

120 140 160 180 4:1

0.1

1

1:1 1:3 1:5200

(c)

Measured MeasuredCalculated Calculated

and (c) ratio. The synthesis conditions are (a) N2:H2 (1:3), 30 Nml min�1

ssure of 50 bar. Reprinted with permission from Hellman, A.; Honkala, K.;ørskov, J. K. Surf. Sci. 2009, 603, 1731–1739.

16012080

16

14

12

10

8

0

Re

Ru

Fe

Co

NiRh

Pt

Log

k (m

olec

ules

/ c

m2 /

s)

40

−ΔH: (kcal / metal atom)(b)

ion as a function of the heat of chemisorption over various metals.96

, 161.

Ammonia Synthesis: State of the Bellwether Reaction 469

The fundamental reason for the existence of volcano curves

in heterogeneous catalysis is the general scaling that exists

between the interaction strength of surfaces with different

reactants, products, reaction intermediates, and transition

states.97,98 In chemistry, linear correlations between activation

(free) energies and reaction (free) energies are widespread,

dating back to Brønsted in 192899 and to Evans and Polanyi

a decade later.100 In heterogeneous catalysis, such relations

have been assumed to hold.28,101 However, establishing accu-

rate linear relationships over a sufficient range had to wait for

DFT calculations to become accurate enough. Figure 15(a)

shows the first published so-called Brønsted–Evans–Polanyi

(BEP) relationship based on DFT calculations between the reac-

tion energy for dissociative adsorption for N2 (the nitrogen

adsorption energy) and the N2 dissociation barrier on different

transition metals.19 The slope of the BEP lines in Figure 15 is

0.9. The reason for why this value is close to 1 is that the

transition state for N2 dissociation is very final-state-like. There-

fore, the transition-state energy essentially follows the nitrogen

adsorption energy from one metal to the next. Figure 15(b)

shows the BEP relation over a much wider range, that is, almost

12 eV in the chemisorption energy.102 This clearly indicates that

the occurrence of a BEP relation is a general manifestation of

the interaction of an adsorbate with the electronic structure of a

transition-metal surface (see also Chapter 7.15).

The BEP curves in Figure 15 are key to understanding the

trends in the ammonia synthesis activity of different transition-

metal catalysts, since a good catalyst for ammonia synthesis is

characterized by a low activation energy for N2 dissociation and

weak bonding of the most abundant surface intermediates,

which all bond to the surface via a nitrogen atom.19 For this

reason, the BEP relation does not only provide a method for

estimating activation barriers from adsorption energies, but it

governs the ammonia synthesis activity completely, as it deter-

mines the relation between the rate of the slowest step and the

amount of free sites on the surface, which is limiting when the

ammonia product gives rise to a high surface coverage on more

reactive surface sites.

From Figure 15, it is clear that the BEP relation is not

restricted to a single active site. The two lines represent terrace

-3.0 -2.0

--2.0

-1.0

-1.0

0.0

0.0

DE [eV/N2]

Ea

[eV

/N2]

ETS

[eV

]

1.0

1.0

2.0

2.0 fcc-Fe(111)Ru(0001)

Ru(0001)step

Fe/Ru(0001)

Fe(110)step

Fe(110)

Mo(110)step

Mo(110)Fe(111)

Pd(111)

Cu(111)

Pd(211)step

3.0

3.0

4.0

4.0

5.0

6.0

Figure 15 Correlation between the energy in the transition state and the disfrom Logadottir, A.; Rod, T. H.; Nørskov, J. K.; Hammer, B.; Dahl, S.; JacobsChristensen, C. H.; Norskov, J. K. Phys. Chem. Chem. Phys. 2008, 10, 5202.

and step sites, respectively. The linearity of the plots is a result

of a fixed geometry of the adsorption site for N and the tran-

sition state when the electronic structure of the transition

metals give rise to variation in both energies. That the two

lines have the same slope shows the transition state is final-

state-like to the same extent. The parallel shift of the two lines

is a manifestation of the geometrical effect, in the present case,

the requirement for dissociation sites with at least five metal

atoms at the step site, as discussed previously.

By combining a micro-kinetic model for the ammonia syn-

thesis with some assumptions on how the adsorption energies

of NHx and H scale with the N adsorption energy and, finally,

the calculated BEP relation, the catalytic activity as a function

of nitrogen-binding strength can be calculated.19 Figure 16

displays the results; they clearly form a volcano curve with Fe

and Ru near the maximum, closely resembling what is found

experimentally, as seen when comparing to Figure 14. When

using the micro-kinetic model to analyze the surface coverage

and N2 dissociation rate as a function of the nitrogen adsorp-

tion energy, it is clear how the maximum in activity is a

compromise between a high N2 dissociation rate (low barrier

for N2 dissociation) and a large number of free sites (lowN and

NH coverage) where the reaction can occur. On the right side

of the volcano curve, the rate is low due to a too high barrier for

N2 dissociation and, on the left side, the rate is low due to a too

high coverage of N and NH, leaving too few free sites for N2

dissociation, which is still the rate-limiting step here.

The effect of catalyst promotion with alkali metals can

easily be incorporated into the model based on the ability to

reduce the N2 dissociation barrier and lower the binding

strength of NHx species to the surface.18 The result of the

analysis is that promotion is most effective for the best non-

promoted catalysts and that promotion will always be essential

for obtaining an optimal ammonia synthesis catalyst.

It has now been established that the ability to describe

catalytic trends for different surfaces as a function of a single

key adsorption energy parameter (the descriptor) not only

holds for ammonia synthesis.103 The method is very general.

For instance, the BEP relation has been established for many

reactions and, for some reactions sharing some classes of

-6 -4 -2

2

0

0

2

2

4

4

6

6 Fit to pure metalsPure metals fcc(211)Decorated stepsBulk alloys

Ediss (N2) [eV]

sociative adsorption energy of nitrogen. Reprinted with permissionen, C. J. H. J. Catal. 2001, 197, 229–231 and Munter, T. R.; Bligaard, T.;

-75.0 -50.0 -25.0 0.0

Mo

Fe

Ru

Os

Co

Ni

“CoMo”

TOF(

s-1 )

10-5

10-4

10-3

10-2

10-1

100

101

[DE-DE(Ru)](kJ/mol N2)

25.0 50.0 75.0-100.0

Figure 16 Turnover frequencies for ammonia synthesis as a function ofthe adsorption energy of nitrogen. Reprinted from Jacobsen, C. J. H.;Dahl, S.; Clausen, B. S.; Bahn, S.; Logadottir, A.; Nørskov, J. K. J. Am.Chem. Soc. 2001, 123, 8404–8405, with permission Copyright 2001American Chemical Society.

NH3 concentration (%)

TOF

[s−1

]

0

2.5

5.0

7.5

10.0

12.5

15.0

17.5

20.0

1 2 3 4 5 6

Fe

Ru

Co3Mo3N

7 8

Figure 17 Measured turnover frequencies for promoted Ru, Co3Mo3N,and Fe catalysts. Reprinted from Jacobsen, C. J. H.; Dahl, S.; Clausen,B. S.; Bahn, S.; Logadottir, A.; Nørskov, J. K. J. Am. Chem. Soc. 2001, 123,8404–8405, with permission Copyright 2001 American Chemical Society.

470 Ammonia Synthesis: State of the Bellwether Reaction

reaction, a universal behavior has been established, for exam-

ple, the dissociative adsorption of the diatomic molecules O2,

N2, NO, and CO98 and (de)hydrogenation reactions.97 Fur-

thermore, linear scaling relations between the adsorption ener-

gies of adsorbates that bind in a similar fashion to a surface has

been established,104 for example, the adsorption energy of NH,

NH2 that binds via the N atom to the surface scales with the

adsorption energy of N, and the slope of these linear scaling

relations can be approximately derived from simple bond

order conservation arguments.104

By combining BEP relations, scaling relations with micro-

kinetic models, it has now been possible to construct volcano

curves for a number of reactions.105 For some reactions, it is

necessary to use not only one but two descriptors, for example,

reactions where CO is hydrogenated, since the reaction inter-

mediates bind via both C and O, and the adsorption energy of

these does not correlate linearly. However, this still results in a

large reduction of parameters that need to be known in order

to predict the activity of a catalyst in a given reaction.

7.17.5 Discovery of New Catalysts for AmmoniaSynthesis

The ability to use computational methods to screen new cata-

lyst materials is very valuable, since the preparation and testing

of new catalysts is a time-consuming task. The full kinetic

description of a given catalyst from theoretical calculation is a

quite demanding task, and screening a large number of systems

using a procedure that requires such an approach for each

system is, at the moment, too time-consuming, even with the

best computers. Therefore, a great achievement is that a direct

link can be created between measured catalytic activity and

easily-calculable parameter(s) describing the surface site, for

example, the adsorption energy of N on a step site. In this

section, it will be shown how this idea can be used to search

for more optimal catalysts and, furthermore, how it leads to a

method to find the most optimal loading of an ammonia

synthesis reactor with the best combination of catalysts.

One of the first examples of where the results from

electronic-structure calculations were exploited to find a new

catalyst is ammonia synthesis. The volcano curve in Figure 16

presents a starting point, as it shows that none of the pure

transition metals are located at the top of the volcano curve.

Ruthenium is nearest to the maximum, but it is a scarce and

expensive metal. The volcano curve also reveals what is needed

in order to obtain a more optimal catalyst: a material with an

active site that has optimal nitrogen adsorption energy must be

found. It is natural to assume that this can be achieved by

creating an active site where two different metals from either

side of the volcano are present.

Following this interpolation principle, it was realized that

Mo-Co-based materials are suitable candidates to be tested.

The ammonia synthesis turnover frequency turned out to be

much better over a mixed-metal Co3Mo3N catalyst than over

its constituents.20 It is even better than the optimal Fe catalyst

at all reaction conditions, and also better than Ru at low NH3

concentrations; see Figure 17.

Central to the interpolation principle working experimen-

tally is that the mixed site must be stable under reaction con-

ditions. This might not work due to the preferential segregation

of one metal to the surface or phase separation of the alloy

under the reaction conditions, which can be induced by strong

binding of one of the components to the reaction intermedi-

ates. This could have been expected for the Co-Mo system,

since N binds considerably stronger to the Mo than to Co;

however, importantly, the ordering energy of the ternary

nitride prevents this.20

This rational approach for the computational screening of

materials for better catalysts is of general validity and has now

been used for a number of catalytic reactions, as reviewed

recently.105 The approach can be used at different levels. In

the simplest form, the interpolation principle is used to suggest

new catalysts candidates, while in more time-consuming

schemes, the volcano curve descriptors are calculated for a

large number of possible active sites, exposing different

Ammonia Synthesis: State of the Bellwether Reaction 471

elements.106 In this way, it is possible to take into account

discrepancies from BEP and scaling relations.

The position of the maximum of the volcano curve is

not the same at different reaction conditions. As an example,

Figure 18 shows how the position of the volcano curve changes

as the approach to equilibrium changes. At low ammonia con-

centrations, Fe is the best catalyst among the pure metals, while

Ru is the best choice near equilibrium concentrations of ammo-

nia. This is consistentwith experimental observations and results

from the fact that Fe catalysts and other catalysts on the left side

of the volcano are strongly inhibited by ammonia, which gives

rise to a high coverage of N and NHx species and, therefore, the

number of free sites available for N2 dissociation decrease with

increasing ammonia concentration. Since the coverage of N-

containing species in equilibrium ammonia is low for catalysts

on the right side of the volcano curve, their activities are not

affected significantly by the approach to equilibrium.

Based on the volcano curves and the simplified assumption

that the density of active sites in different catalysts is the same,

an optimal value of the nitrogen-binding energy can be calcu-

lated, giving rise to the so-called optimal catalyst curves, which

can be used to choose the catalysts to load into a given ammo-

nia synthesis reactor for the optimal ammonia output.33 In the

inlet of the reactor, a reactive catalyst like Fe is preferred, while

as the ammonia concentration increases, less reactive catalysts

are preferred. This implies that it is optimal to load an infinite

number of different catalysts with carefully chosen nitrogen-

binding energies. In practice, this is, of course, not possible.

Generally, the smallest possible total catalyst volume results

when the reactor is designed to follow the optimal operating

line (see Section 7.17.2), and simultaneously use catalysts

with nitrogen-binding energies as close to the optimum as

possible.33 Consequently, it means that, with our current

understanding of the ammonia synthesis reaction, it is possible

102

101

100

10-1

10-2

10-3

10-4

-100 -50 500EN*-EN* (Ru) (kJ/mol N2)

Mo

TOF

(s-1

)

Fe

Co

90%

20%

h = 5%

RuOs

10-5

Figure 18 (Left panel): calculated volcano curves at 420 �C, 80 bar, 2:1 H2

450 �C, 200 bar, 3:1 H2:N2; equilibrium 25.4% NH3. In both cases, the volcan90, 20, and 5% approach to equilibrium. Reprinted with permission from JacLogadottir, A.; Nørskov, J. K. J. Catal. 2002, 205, 382–387.

to link density functional calculations with industrial reactor

design and catalyst selection.33

It is, of course, important to balance the potential savings in

the reactor volume against the cost of available catalysts and, in

practice, iron catalysts are totally dominating the market due

to their low price. One attempt was made to take advantage of

the high activity of ruthenium catalysts close to equilibrium. In

the Kellogg advanced ammonia process,107,108 a promoted

carbon-supported ruthenium catalyst is used in the last bed

of the process, where the ammonia concentration is high. This

part of the reactor is operated at lower pressure than normal

and with a H2:N2 ratio below 3 to further promote the activity

of Ru. However, it is doubtful that the benefits of this process

design can compensate for the higher cost and shorter lifetime

of the Ru/C catalysts compared to the traditional Fe catalyst.

From a price perspective, the Co3Mo3N catalyst is more attrac-

tive than an Ru catalyst; however, so far, a process using the

catalyst has not been realized. One reason for this is that it is

the product of activity per active site and the density of these

that determines the technically relevant activity of a catalyst,

and, so far, it has been difficult to prepare the Co3Mo3N

catalyst with a high surface area, since high-temperature nitri-

fication is needed in order to produce the catalyst.

7.17.6 New Perspectives in Ammonia Synthesis

Depleting fossil fuel resources together with severe environmen-

tal consequences caused by the combustion of carbon-

containing molecules enforce us to look for new energy solu-

tions. There are several possible options for currently employed

hydrocarbons both for storage and for use as energy carriers,109

of which the most prominent ones are hydrogen, methanol,

ethanol, and methane. Also, synthetic hydrocarbons could

-100 -50 50EN*-EN* (Ru) (kJ/mol N2)

Mo

Fe

Co

Ru

Os

0

:N2, equilibrium 17.4%NH3. (Right panel): calculated volcano curves ato curves are calculated at ammonia concentrations corresponding to aobsen, C. J. H.; Dahl, S.; Boisen, A.; Clausen, B. S.; Topsøe, H.;

Synthesis Decomposition

Ammonia concentration (%)

Op

timal

DE

(eV

)

0.001

-0.4

-0.6

-0.8

-1.0

-1.2

-1.4

-1.60.01 0.1 1 10 100

Figure 19 Dissociative N2 adsorption energy of optimal catalysts forammonia synthesis/decomposition at 773 K, 1 bar, and 3:1 H2/N2.Equilibrium corresponds to ca. 0.13% ammonia. Reprinted withpermission from Boisen, A.; Dahl, S.; Nørskov, J. K.; Christensen, C. H.,J. Catal. 2005, 230, 309–312.

472 Ammonia Synthesis: State of the Bellwether Reaction

gradually replace fossil energy carriers, with the further benefit of

keeping the present infrastructure intact.

Ammonia has a high hydrogen content, making it a viable

candidate for serving as a hydrogen source for standardhydrogen

fuel cells. It can also be used in internal combustion engines and

direct ammonia fuel cells. All these together make ammonia a

promising option for new energy solutions.3,109 While nitrogen

is taken from the air, the Haber–Bosch process requires a hydro-

gen source. By far the most common source is currently natural

gas, but coal, petroleum coke, or heavy petroleum fractions can

be used to produce hydrogen as well. Electrocatalytic water split-

ting offers a sustainable way to obtain hydrogen, where any

source of electricity can be employed, including hydro, wind,

and solar cell power, or even nuclear power.

The ability to become a liquid at moderate pressure allows

ammonia to actually store more hydrogen per unit volume than

compressed hydrogen or cryogenic liquid hydrogen. However,

before ammonia can be used as a transportation fuel, design

standards for on-board ammonia fuel tanks must be established,

as well as procedures for ammonia transfer from storage to vehi-

cle tanks. Compared to gasoline, ammonia has higher ignition

energy, higher flashpoint, and a narrower explosive range when

mixedwith air, thus, an explosion or fire would be less likely with

a ruptured ammonia tank than with a gasoline tank. Both gaso-

line and diesel fuel can contain carcinogenic components in their

vapors, while ammonia does not. Furthermore, combusted

ammonia does not produce greenhouse gases nor soot, since it

does not contain carbon.However, the downside is the toxicity of

ammonia: at low concentration, vapors are highly irritating, with

sharp suffocating odor, and at high concentration, it is actually

life threatening. It has been shown that a convenient way to

alleviate this safety problem is to reversibly store the ammonia

with high density in an ammine complex such as Mg(NH3)6Cl2,

Ca(NH3)8Cl2, Mn(NH3)6Cl2, and Ni(NH3)6Cl2.110 The ammo-

nia is bound in these compounds with a strength that makes the

partial pressure of ammonia so low that the storage material can

be handled safely at room temperature and allows the desorption

of ammonia at reasonable temperatures. Hence, although alter-

natives exist, ammonia has several interesting properties that

make it a possible option for special applications in energy tech-

nology. These properties include high energy density, a well-

established synthesis, and high hydrogen content.

Ammonia decomposition into nitrogen and hydrogen, a

reverse process to ammonia formation, will be a key reaction

in any future ammonia-based energy infrastructure. Today,

ammonia decomposition plays a role in two different areas

related to energy and environmental science, namely, to pro-

vide carbon-free hydrogen for fuel cells and to remove ammo-

nia from the reformate of internal gasification combined cycle

power plants.

Much of our current understanding of ammonia synthesis

can be directly employed to build an understanding of ammonia

decomposition.21 Owing to deviations in operating conditions,

for example, pressure, temperature, and gas composition,

ammonia formation and decomposition processes present

some clear and important differences. For instance, the volcano

curve for ammonia decomposition shifts towards catalyst mate-

rials that bind to nitrogen weaker, that is, towards more noble

metals. Similar to the ammonia synthesis reaction, the position

of the maximum of the volcano in ammonia decomposition

sensitively depends also on the reaction conditions. Interest-

ingly, it is observed that the optimal ammonia synthesis catalyst

is never the optimal ammonia decomposition catalyst; see

Figure 19.21 This does not indicate that the principle of micro-

scopic reversibility does not hold, but, rather, highlights that

widely different reaction conditions in ammonia synthesis and

decomposition result in very different optimal binding energies

for the two reactions, except, of course, at equilibrium. Also, for

ammonia decomposition, a combination of first-principle cal-

culations, micro-kinetics, and surface science can lead to the

discovery of new catalysts formulations.111

One can imagine several scenarios where it would be advan-

tageous to produce ammonia locally on a smaller scale, even if the

production costs would be higher than in the Haber–Bosch pro-

cess. This point of view is becomingmore relevant as ammonia is

used in the SCR process for removing nitrogen oxides from diesel

engine exhaust, with the recent proposal of using ammonia salts

as an energy carrier,3 or with the increased political focus on

transporting safety for hazardous chemicals. Hinnemann and

Nørskov have performed a theoretical study on the possibility to

produce ammonia under ambient temperatures and pressures. As

mentioned in the introduction, microorganisms exist in nature,

which use the enzyme nitrogenase to form ammonia from pro-

tons, electrons, and atmospheric nitrogen. In the enzyme, the

active site is an MoFe7S9 cluster, which catalyzes the reaction:

N2þ8Hþþ8e� P2NH3 þH2;

where the nitrogen molecule is hydrogenated stepwise. The

source of energy for this reaction consists of at least 16 ATP

molecules.6 These presumably increase the chemical potential

of the electrons. It has been hypothesized that part of the

enzyme functions just like a battery.6

Since protonation reactions in particular appear to be so

well-described within the DFTmethodology, it offers hope that

it is adequate also for describing electrochemical ammonia

synthesis.112 In an electrochemical cell, ammonia synthesis

on a ruthenium electrode in aqueous solution cannot proceed

via the same reaction mechanism as the Haber–Bosch process,

because all available ruthenium step sites for N2 dissociation

will be poisoned by strongly adsorbing oxygen atoms and OH

groups.112 The calculations show that the barrier for N2 disso-

ciation is too high on the close-packed terraces at low electro-

chemical temperature. This leaves the biomimetic mechanism

Reaction coordinate

U = -1.08 V

U = 0 V

*NHNH*NHNH2

8

6

4

2

0

N 2 +

6(H+ + e

- )

*N 2 +

6(H+ + e

- )

*N 2H +

5(H+ + e

- )

*NNH 2 +

4(H+ + e

- )

*N + N

H 3 +

3(H+ + e

- )

*NH + N

H 3 +

2(H+ + e

- )

*NH 2 +

NH 3

+ (H

+ + e- )

*NH 3 +

NH 3

2NH 3

Free

ene

rgy

[eV

]

Figure 20 Free energy for the associative mechanism of electrochemicalammonia synthesis on a flat Ru(0001) surface obtained from DFTcalculations of the binding energy and vibrational frequencies, as well asentropy of the gas molecules. For an electrolyte with pH¼0 at 300 K, thisalso gives the free energy for the electrochemical reaction when there is noapplied bias, U ¼ 0 V. The most difficult step is the addition of the firstproton to the adsorbed N2 molecule. With an applied potential of U ¼�1.08 V, all the elementary steps involve either no change or a decrease inthe free energy. Reprinted with permission from Sørensen, R. Z.;Hummelshøj, J. S.; Klerke, A.; Reves, J. B.; Vegge, T.; Nørskov, J. K.;Christensen, C. H. J. Am. Chem. Soc. 2008, 130, 8660–8668.

Ammonia Synthesis: State of the Bellwether Reaction 473

strongly favored, especially because the tendency for direct

protonation of the nitrogen molecule can be tuned directly

by varying the voltage over the electrochemical cell. Figure 20

shows the full free-energy reaction diagram for the biomimetic

ammonia synthesis route at zero potential with respect to the

standard hydrogen electrode. The plot demonstrates that the

first hydrogenation process has, by far, the highest barrier. An

energy diagram is also shown at a potential of�1.08 V. At such

a low potential, all protonation reactions are thermodynami-

cally downhill. If such a negative potential were applied,

ammonia could be produced electrochemically.112

7.17.7 Conclusion

We have, throughout this chapter, shown many of the reasons

why ammonia synthesis has been, and still is, the bellwether

reaction in heterogeneous catalysis. In particular, we have

described how, in recent years, it has been possible to provide

a full understanding of the ammonia synthesis reaction at the

atomic level through the combined use of experiments and

quantum mechanical electronic structure calculations. This

level of understanding led to the understanding of catalytic

trends, which was used for the rational design of a new catalyst

for ammonia synthesis. It is now clear that the approach is of

general validity and it has been used successfully to design new

catalysts for a number of reactions.

Finally, we have reviewed how ammonia synthesis over

ruthenium in industrial conditions can be understood only

from input from first-principle calculations. The fact that we

can treat one reaction on one surface in such enormous

theoretical detail reveals promise. There is a linear correspon-

dence between the computational power and the number of

catalyst materials that we can screen for a given reaction at the

current high level of accuracy. With the historical increase in

available computer power, it is difficult to remain skeptical of

the importance of theory in the field of catalysis, already in the

near future. It can be envisioned that, soon, theoretical model-

ing will not only be used for reproducing the experimentally

known facts about a given reaction on a given surface, but

could become the standard choice as the first starting point

when a new catalyst for a known reaction is desired, or even

when an unknown reaction to obtain a given product is

needed. This would be the beginning of the era of ‘Catalysis

Informatics’,113 and there are strong indications that ammonia

synthesis will continue to be the bellwether reaction in this

development.

References

1. Smil, V. Nature 1999, 400, 415.2. Ullmann, F. Ullmann’s Encyclopedia of Industrial Chemistry; WILEY-VCH:

Weinheim, 2010.3. Schuth, F.; Palkovits, R.; Schlogl, R.; Su, D. S. Energ. Environ. Sci. 2012, 5,

6278–6289.4. Tamaru, K. In Catalytic Ammonia Synthesis Fundamentals and Practise; Jennings,

J. P., Ed.; Plenum Press: New York, 1991.5. Schrock, R. R. Proc. Natl. Acad. Sci. 2006, 103, 17087.6. Hinnemann, B.; Nørskov, J. K. Top. Catal. 2006, 37, 55–70.7. Ammonia: Catalysis and Manufacture; Nielesen, A., Ed.; Springer-Verlag:

New York, 1995.8. Hopper, C. W. In Catalytic Ammonia Synthesis Fundamentals and Practise;

Jennings, J. P., Ed.; Plenum Press: New York, 1991.9. Haber, F. Nobel Prize Lectures 1918; Elsevier: Amsterdam, The Netherlands,

1966.10. Bosch, C. Nobel Prize Lecture 1931; Elsevier: Amsterdam, The Netherlands, 1966.11. Haber, F.; van Oordt, G. Z. Fur Anorg. Chem. 1905, 44, 341–378.12. Mittasch, A. Adv. Catal. 1950, 2, 81.13. Mittasch, A. U.S. Patent 1,173,532, 1913.14. Aika, K.-I.; Hori, H.; Ozaki, A. J. Catal. 1972, 27, 424–431.15. Ozaki, A.; Aika, K.; Furuta, A.; Okagami, A. U.S. Patent 3,770,658, 1973.16. Ozaki, A.; Aika, K.-I. In Catalysis, Science and Technology; Boudart, M., Ed.;

Springer: Berlin, 1981.17. Aika, K.; Ozaki, A. In Ammonia: Catalysis and Manufacture; Nielesen, A., Ed.;

Springer-Verlag: New York, 1995.18. Dahl, S.; Logadottir, A.; Jacobsen, C. J. H.; Nørskov, J. K. Appl. Catal. Gen. 2001,

222, 19–29.19. Logadottir, A.; Rod, T. H.; Nørskov, J. K.; Hammer, B.; Dahl, S.; Jacobsen, C. J. H.

J. Catal. 2001, 197, 229–231.20. Jacobsen, C. J. H.; Dahl, S.; Clausen, B. S.; Bahn, S.; Logadottir, A.;

Nørskov, J. K. J. Am. Chem. Soc. 2001, 123, 8404–8405.21. Boisen, A.; Dahl, S.; Nørskov, J. K.; Christensen, C. H. J. Catal. 2005, 230,

309–312.22. Foster, A.; James, P. G.; McCarroll, J. J.; Tennison, S. R. U.S. Patent 4,163,775,

1979.23. Chementator, Chem. Eng. 1993, 100, 19.24. McCarroll, J. J.; Tennison, S. R. U.S. Patent 4,600,571, 1986.25. Boudart, M. Top. Catal. 1994, 1, 405–414.26. Nielsen, A. An Investigation on Promoted Iron Catalysts for the Synthesis of

Ammonia; Gjellerup: Copenhagen, 1968.27. Jennings, J. R. Catalytic Ammonia Synthesis: Fundamentals and Practice; Plenum

Press: New York, 1994.28. Lundqvist, B. I.; Bligaard, T.; Nørskov, J. K. In Handbook in Surface Science vol 3:

Dynamics; Hasselbrink, E., Lundqvist, B. I., Eds.; 2008; p 269; Series editorsRichardson N. V.; Holloway, S.

29. Grunze, M. In The Chemistry and Physics of Solid Surfaces and HeterogeneousCatalysis; King, D. A., Woodruff, D. P., Eds.; Elsevier: Amsterdam, TheNetherlands, 1982; Vol 4.

474 Ammonia Synthesis: State of the Bellwether Reaction

30. Tamaru, K. Catalytic Ammonia Synthesis, Fundamentals and Practice; PlenumPress: New York, 1991.

31. Honkala, K.; Hellman, A.; Remediakis, I. N.; Logadottir, A.; Carlsson, A.; Dahl, S.;Christensen, C. H.; Nørskov, J. K. Science 2005, 307, 555–558.

32. Dybkjaer, I. In Ammonia: Catalysis and Manufacture; Nielesen, A., Ed.; Springer-Verlag: New York, 1995; p 199.

33. Jacobsen, C. J. H.; Dahl, S.; Boisen, A.; Clausen, B. S.; Topsøe, H.; Logadottir, A.;Nørskov, J. K. J. Catal. 2002, 205, 382–387.

34. Gramatica, G.; Pernicone, N. In Catalytic Ammonia Synthesis; Jennings, J. R., Ed.;Plenum: New York, 1991; p 211.

35. Somorjai, G. A.; Li, Y. Proc. Natl. Acad. Sci. 2011, 108, 917–924.36. Somorjai, G. A. J. Phys. Chem. B 2000, 104, 2969.37. Somorjai, G. A.; Park, J. Y. Surf. Sci. 2009, 603.38. Hammer, B.; Nørskov, J. K. Adv. Catal. 2000, 45, 1253.39. Nørskov, J. K.; Scheffler, M.; Toulhoat, H. MRS Bull. 2006, 31, 669.40. Hammer, B.; Nørskov, J. K. Nature 1995, 376, 238.41. Lopez, N.; Nørskov, J. K. J. Am. Chem. Soc. 2002, 124, 11262.42. Nørskov, J. K.; et al. Chem. Soc. Rev. 2008, 37, 2163.43. Nakata, T.; Matsushita, S. J. Phys. Chem. 1968, 72, 458.44. Shvachko, V. I.; Fogel, Y. M.; Kolot, Y. Kinet. Katal. 1966, 7, 834.45. Ertl, G. J. Vac. Sci. Technol. A 1983, 1, 1247.46. Ertl, G.; Huber, M.; Lee, S. B.; Paal, Z.; Weiss, M. Appl. Surf. Sci. 1981, 8, 373.47. Ertl, G.; Lee, S. B.; Weiss, M. Surf. Sci. 1982, 114, 527–545.48. Ertl, G.; Weiss, M.; Lee, S. B. Chem. Phys. Lett. 1979, 60, 391–394.49. Emmett, P. H.; Brunauer, S. J. Am. Chem. Soc. 1933, 55, 1738–1739.50. Emmett, P. H.; Brunauer, S. J. Am. Chem. Soc. 1934, 56, 35–41.51. Boszo, F.; Ertl, G.; Grunze, M.; Weiss, M. J. Catal. 1977, 49, 18.52. Paal, Z.; Ertl, G.; Lee, S. B. Appl. Surf. Sci. 1981, 8, 231.53. Spencer, N. B.; Schoonmaker, R. C.; Somorjai, G. A. J. Catal. 1982, 74, 129.54. Strongin, D. R.; Carrazza, J.; Bare, S. R.; Somorjai, G. A. J. Catal. 1987, 103,

213.55. Strongin, D. R.; Somorjai, G. A. J. Catal. 1988, 109, 51.56. Stoltze, P.; Nørskov, J. K. Phys. Rev. Lett. 1985, 55, 2502.57. Dumesic, J. A.; Trivino, A. A. J. Catal. 1989, 116, 119.58. Dumesic, J. A.; Trevino, A. A. J. Catal. 1989, 116, 119–129.59. Stoltze, P.; Nørskov, J. K. J. Catal. 1988, 110, 1–10.60. Hinrichsen, O. Cat. Today 1999, 53, 177–188.61. Ertl, G. Catal. Rev. Sci. Eng. 1980, 21, 201.62. Strongin, D. R.; Bare, S. R.; Somorjai, G. A. J. Catal. 1988, 109, 51.63. Mortensen, J. J.; Ganduglia-Pirovano, M. V.; Hansen, L. B.; Hammer, B.;

Stoltze, P.; Nørskov, J. K. Surf. Sci. 1999, 422, 8–16.64. Mortensen, J. J.; Hansen, L. B.; Hammer, B.; Nørskov, J. K. J. Catal. 1999, 182,

479.65. Dietrich, H.; Geng, P.; Jacobi, K.; Ertl, G. Chem. Eng. Sci. 1996, 51, 1683.66. Mortensen, Y. M.; Hammer, B.; Norskov, J. K. J. Catal. 1997, 169, 85.67. Dahl, S.; Logadottir, A.; Egeberg, R. C.; Larsen, J. H.; Chorkendorff, I.;

Tornqvist, E.; Nørskov, J. K. Phys. Rev. Lett. 1999, 83, 1814.68. Diekhoner, L. M.; Mortensen, H.; Baurichter, A.; Luntz, A. C.; Hammer, B. Phys.

Rev. Lett. 2000, 84, 4906.69. Shetty, S.; Jansen, A. P. J.; van Santen, R. A. J. Phys. Chem. C 2008, 112,

17768–17771.70. Gavnholt, J.; Schiotz, J. Phys. Rev. B 2008, 77, 035404.71. Van Hardeveld, R.; Van Montfoort, A. Surf. Sci. 1966, 4, 396.72. Logadottir, A.; Nørskov, J. K. J. Catal. 2003, 220, 273–279.73. Jacobsen, C. J. H.; Dahl, S.; Hansen, P. L.; Tornqvist, E.; Jensen, L.; Topsoe, H.;

Prip, D. V.; Moenshaug, P. B.; Chorkendorff, I. J. Mol. Catal. A Chem. 2000, 163,19–26.

74. Aika, K.; Kumasaka, M.; Oma, T.; Matsuda, H.; Watanabe, N.; Yamazaki, K.;Ozaki, A.; Onishi, T. Appl. Catal. 1986, 28, 57.

75. Bossi, A.; Garbassi, F.; Petrini, G.; Zanderighi, L. J. Chem. Soc. Faraday Trans.1982, 78, 1029.

76. Stoltze, P. Prog. Surf. Sci. 2000, 65, 65.77. Dumesic, J. A.; Rudd, D. F.; Aparicio, L. M.; Rekoske, J. E.; Trivino, A. A. The

Microkinetics of Heterogeneous Catalysis; ACS: Washington, 1993.

78. Askgaard, T.; Nørskov, J. K.; Ovesen, C. V.; Stoltze, P. J. Catal. 1995, 156,229.

79. Greeley, J.; Mavrikakis, M. J. Am. Chem. Soc. 2004, 126, 3910.80. Kandoi, S.; Greeley, J.; Sanchez-Castillo, M. A.; Evans, S. T.; Gokhale, A. A.;

Dumesic, J. A.; Mavrikakis, M. Top. Catal. 2006, 37, 17.81. Ovesen, C. V. C.; Hammershøi, B. S.; Steffensen, G.; Askgaard, T.;

Chorkendorff, I.; Nørskov, J. K.; Rasmussen, J. P.; Stoltze, P.; Taylor, P. J. Catal.1996, 158, 170.

82. Linic, S.; Barteau, M. A. J. Am. Chem. Soc. 2003, 125, 4034.83. Stegelmann, C.; Schiødt, N. C.; Campbell, C. T.; Stoltze, P. J. Catal. 2004, 221,

630.84. Nørskov, J. K. Prog. Surf. Sci. 1991, 38, 103–144.85. Dahl, S.; Sehested, J.; Jacobsen, C. J. H.; Tornqvist, E.; Chorkendorff, I. J. Catal.

2000, 192, 391–399.86. Stoltze, P. Phys. Scripta 1987, 36, 824.87. Bligaard, T.; Nørskov, J. K.; Dahl, S.; Matthiesen, J.; Christensen, C. H.;

Sehested, J. J. Catal. 2004, 224, 206–217.88. Hellman, A.; Honkala, K. J. Chem. Phys. 2007, 127, 194704.89. Sehested, J.; Jacobsen, C. J. H.; Tornqvist, E.; Rokni, S.; Stoltze, P. J. Catal.

1999, 188, 83–89.90. Fastrup, B. J. Catal. 1997, 168, 235–244.91. Hellman, A.; Honkala, K.; Remediakis, I. N.; Logadottir, A.; Carlsson, A.; Dahl, S.;

Christensen, C. H.; Nørskov, J. K. Surf. Sci. 2009, 603, 1731–1739.92. Hinrichsen, O.; Rosowski, E.; Hornung, A.; Muhler, M.; Ertl, G. J. Catal. 1997,

165, 33.93. Hellman, A.; Honkala, K.; Remediakis, I. N.; Logadottir, A.; Carlsson, A.; Dahl, S.;

Christensen, C. H.; Nørskov, J. K. Surf. Sci. 2006, 600, 4264–4268.94. Sabatier, P. Ber. Deutsch. Chem. Ges. 1911, 44, 1984.95. Balandin, A. A. Adv. Catal. 1969, 19, 1.96. Aika, K.; Yamazaki, K.; Ozaki, A. Chem. Lett. (Jap.) 1973, 2, 161.97. Wang, S.; Petzold, V.; Tripkovic, V.; Kleis, J.; Howalt, J. G.; Skulason, E.;

Fernandez, E. M.; Hvolbæk, B.; Jones, G.; Toftelund, A.; Falsig, H.; Bjorketun, M.;Studt, F.; Abild-Pedersen, F.; Rossmeisl, J.; Nørskov, J. K.; Bligaard, T. Phys.Chem. Chem. Phys. 2011, 13, 20760.

98. Nørskov, J. K.; Bligaard, T.; Logadottir, A.; Bahn, S. R.; Hansen, L. B.;Bollinger, M. V.; Bengaard, H. S.; Hammer, B.; Sljivancanin, Z.; Mavrikakis, M.;Xu, Y.; Dahl, S.; Jacobson, C. J. H. J. Catal. 2002, 209, 275.

99. Brønsted, N. Chem. Rev. 1928, 5, 231.100. Evans, M. G.; Polanyi, M. Trans. Faraday Soc. 1937, 33, 448.101. Boudart, M. In Handbook of Heterogeneous Catalysis; Ertl, G., Knozinger, H.,

Weitkamp, J., Eds.; Wiley-VCH: Weinheim, 1997.102. Munter, T. R.; Bligaard, T.; Christensen, C. H.; Norskov, J. K. Phys. Chem. Chem.

Phys. 2008, 10, 5202.103. Nørskov, J. K.; Abild-Pedersen, F.; Studt, F.; Bligaard, T. Proc. Natl. Acad. Sci.

2011, 108, 937–943.104. Abild-Pedersen, F.; Greeley, J.; Studt, F.; Rossmeisl, J.; Munter, T. R.;

Moses, P. G.; Skullason, E.; Bligaard, T.; Nørskov, J. K. Phys. Rev. Lett. 2007, 99,016105.

105. Nørskov, J. K.; Bligaard, T.; Rossmeisl, J.; Christensen, C. H. Nat. Chem. 2009,1, 37.

106. Studt, F.; Abild-Pedersen, F.; Bligaard, T.; Sørensen, R. Z.; Christensen, C. H.;Nørskov, J. K. Science 2008, 320, 1320–1322.

107. Czuppon, T. A.; Knez, S. A.; Schneider, R. W.; Worobets, G. Ammonia Plant SafetyRelat. Facil. 1994, 34, 236.

108. Strait, R. Nitrog. Methan. 1999, 238, 37.109. Schuth, F. Chem. Eng. Technol. 2011, 83, 1984.110. Sørensen, R. Z.; Hummelshøj, J. S.; Klerke, A.; Reves, J. B.; Vegge, T.;

Nørskov, J. K.; Christensen, C. H. J. Am. Chem. Soc. 2008, 130, 8660–8668.111. Hansgen, D. A.; Vlachos, D. G.; Chen, J. G. Nat. Chem. 2010, 2, 484–489.112. Skulason, E.; Bligaard, T.; Gudmundsdottir, S.; Studt, F.; Jan Rossmeisl, F. A.-P.;

Vegge, T.; Jonsson, H.; Nørskov, J. K. Phys. Chem. Chem. Phys. 2012, 14,1235.

113. Hummelshøj, J. S.; Abild-Pedersen, F.; Studt, F.; Bligaard, T.; Nørskov, J. K.Angew. Chem. Int. Ed. 2012, 51, 272–274.