Oil and Gas Corrosion Prevention. From Surface Facilities to Refineries

199 Pages·2014·2.068 MB·English

by  James G. Speight (Auth.)

Oil and Gas Corrosion Prevention
Oil and Gas Corrosion
Prevention
From Surface Facilities to Refineries
James G. Speight PhD, DSc
CD&W Inc.,
Laramie, Wyoming, USA
AMSTERDAM(cid:129)BOSTON(cid:129)HEIDELBERG(cid:129)LONDON
NEWYORK(cid:129)OXFORD(cid:129)PARIS(cid:129)SANDIEGO
SANFRANCISCO(cid:129)SINGAPORE(cid:129)SYDNEY(cid:129)TOKYO
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PREFACE
Corrosion can occur at any point or at any time during petroleum and
natural gas recovery and processing. Thus is particularly true in ageing
refineries and gas processing plants which can present a serious safety
hazard.
Key findings into corrosion processes in refinery equipment and gas
processing plants as well as corrosion in offshore structures are
presented.
This book summarizes the key corrosion processes in refinery and
gas processing equipment (cid:1) such as storage tanks, reactors, and sour
water strippers (cid:1) and how it can be measured and controlled.
Methods of testing for corrosion are presented as well as preventative
measures. The book also contains a helpful glossary that will assist the
reader in understanding the terminology of corrosion.
Dr. James G. Speight
Laramie, Wyoming
December 2013
11
CHAPTER
Corrosion
Note: This chapter is available on the companion website: http://store.
elsevier.com/product.jsp?isbn=9780128003466&_requestid=1050998.
1.1 ABSTRACT
Corrosion is the deterioration of a material as a result of its interaction
with its surroundings and can occur at any point or at any time during
petroleum and natural gas processing. Although this definition is appli-
cable to any type of material, it is typically reserved for metallic alloys.
Furthermore, corrosion processes not only influence the chemical
properties of a metal or metal alloys, but they also generate changes in
their physical properties and mechanical behaviors.
It is the purpose of this chapter to present the terminology of the
different types of corrosion, as well as the key chemical aspects of cor-
rosion processes and the types of corrosion that can occur in a refinery
or gas processing plant.
1.2 CONTENT
1.1 Introduction
1.2 Corrosion Chemistry
1.3 Types of Corrosion
1.4 Effect of Temperature
ee11
CHAPTER
Corrosion
1.1 INTRODUCTION
Corrosion is a natural phenomenon and is the deterioration of a mate-
rial as a result of its interaction with its surroundings (Fontana, 1986;
Garverick, 1994; Shreir et al., 1994; Jones, 1996; Shalaby et al., 1996;
Peabody, 2001; Bushman, 2002; Landolt, 2007). Although this defini-
tion is applicable to any type of material, it is typically reserved for
metallic alloys. Of the known chemical elements, approximately 80 are
metals (Figure e1.1) and approximately 50% of these metals can be
alloyedwithothermetals,althoughthephysical,chemical,andmechan-
ical properties, and propensity for corrosion are all dependent upon the
composition of the alloys. Furthermore, corrosion processes not only
influence the chemical properties of a metal or metal alloys, but they
alsogeneratechangesinphysicalpropertiesandmechanicalbehavior.
Furthermore, identification and understanding of the chemistry of
corrosion can lead to the prevention of corrosion by preventing the pre-
dominantchemicalreactions(Wranglen,1985;UhligandRevie,1985).
It is the purpose of this chapter to present the chemistry, terminol-
ogy, and key chemical aspects of corrosion processes, along with the
types of corrosion that can occur in a refinery or gas processing plant.
1.2 CORROSION CHEMISTRY
Metals occur either in the pure metallic state (the zero oxidation state:
for example, iron would be represented as Fe0 in this state) or in the
form of compounds with other elements (they acquire positive states of
oxidation). The tendency of metals to corrode is related to the stability
of the metallic state. Chemically, corrosion is an electrochemical reac-
tion involving the movement of electrons: when steel is first exposed to
the air, the original smooth surface of the zero oxidation state will be
covered with rust in a relatively short time—perhaps even within a
matter of hours (Figure e1.2). In actual fact, rust is not always com-
posed of a single iron oxide, but is generally a mixture of ferrous oxide
e2 OilandGasCorrosionPrevention
I II III IV V VI VII 0
1 2
H He
3 4 Transition Metals 5 6 7 8 9 10
Li Be B C N O F Ne
11 12 VIIIB 13 14 15 16 17 18
Na Mg IIIB IVB VB VIB VIIB IB IIB AI Si P S Cl Ar
19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
55 56 57–71 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86
Cs Ba I Hf Ta W Re Os Ir Pt Au Hg TI Pb Bi Po At Rn
87 8889–103 104 105 106 107 108 109
Fr Ra Rf Ha
Lanthides La57 Ce58 Pr59 Nd60 Pm61 Sm62 Eu63 Gd64 Tb65 Dy66 Ho67 Er68 Tm69 Yb70 Lu71
Actinides 89 90 91 92 93 94 95 96 97 98 99 100 101 102 103
Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
Metal Metalloid Nonmetal
Figuree1.1Periodictableoftheelements.
O
2
Water
Rust
Fe+2 and Fe+3 Fe2O3 – XH2O
“Pitting” Cathode:
Anode: O2 + 2H2O + 4e–
Fe ----> Fe+2 + 2e– ------> 4OH–
Fe+2 ---> Fe+3 + e–
Iron
Figuree1.2Representationofthechemistryofrustformation.
(FeO) and ferric oxide (Fe O ), as well as complex oxides such as mag-
2 3
netite (Fe O ).
3 4
1.2.1 General Chemistry
Generally, when iron is used in an oxidizing atmosphere, corrosion is
represented as the formation of rust (iron oxide), like what might occur
in a pipeline (Beavers and Thompson, 2006) and which is often repre-
sented by the formation of iron oxide and then hydration in humid
environments:
2Fe013O -2Fe O
2 2 3
2Fe O 1xH O-2Fe O :xH O
2 3 2 2 3 2
Corrosion e3
In reality, rust formation is a much more complex process and
occurs at the point of, or at some distance away from, the actual pit-
ting or erosion of iron (Figure e1.2). The involvement of water
accounts for the fact that rust formation occurs much more rapidly in
moist conditions than in a dry environment (such as a desert), where
water in the atmosphere or water in the ground is limited.
Similarly, corrosion in a pipe (a pipeline or an intrarefinery transfer
pipe) occurs when a negative area of metal (the anode) is connected to
a positive area (the cathode) by the pipe wall itself. As a result, elec-
trons can flow from the anode to the cathode (Chapter 3). In addition
to the anode, the cathode, and the connecting conductive material, the
electrochemical reaction requires one more element: the electrolyte.
The electrolyte is a conducting solution and—in the case of a pipe—is
the water within the pipe with its dissolved salts.
Many other factors affect the chemistry and the rate of corrosion.
For example, the presence of salt (sodium chloride, NaCl), like that
which occurs in marine environments (Chapter 5), greatly enhances the
rusting of metals because the dissolved salt increases the conductivity
of the aqueous solution formed at the surface of the metal and
enhances the rate of electrochemical corrosion. The sea and the salty
atmosphere are saline media that are highly aggressive to metals:
marine structures, such as drilling rigs, platforms, and marine pipe-
lines, usually show signs of severe corrosion unless they have been
properly protected (Chapter 5) (Dugstad et al., 1994).
Another way of explaining corrosion—using the same chemical
equations and resulting in the same effect—is through the agency of
anodic and cathodic reactions. Any electrochemical reaction requires
four elements, all of which must be in contact: (1) the anode, (2) the
cathode, (3) the conductive material, and (4) the electrolyte. In the
case of the corrosion of a pipe (Chapter 3), the anode, cathode, and
conductive material are all found in the pipe wall, while the electrolyte
is the water within the pipe. If any of these four elements, which make
up the corrosion cell, are absent or are not touching each other, then
corrosion cannot occur.
1.2.2 Anodic Reactions
The chemical reactions that take place in corrosion processes are redox
reactions (reduction(cid:1)oxidation reactions), which include all chemical
e4 OilandGasCorrosionPrevention
reactionsinwhichatomshavetheoxidationstatechangedandgenerally
involve the transfer of electrons between species. Such reactions require
a material that is oxidized (the metal) and another that is reduced (the
oxidizing agent). In the oxidation sequence (at the anode), the metal
loses electrons. In the reduction reaction (at the cathode), the oxidizing
agentgainstheelectronsthathavebeenreleasedfromthemetal.
Oxidation of elemental iron occurs at the anode. First, the elemen-
tal iron breaks down, reacts, and then leaves the pipe, so pits form in
the pipe’s surface at the anode:
elemental iron-ferrous iron 1 electrons
Fe0 Fe21 2e2
The following redox reaction is dictated by the availability of water
and oxygen which are crucial to the formation of rust:
4Fe211O -4Fe3112O22
2
In addition, the following multistep acid(cid:1)base reactions and dehy-
dration equilibria affect the course of rust formation:
Fe2112H O"FeðOHÞ 12H1
2 2
Fe3113H O"FeðOHÞ 13H1
2 3
FeðOHÞ "FeO1H O
2 2
FeðOHÞ "FeOðOHÞ1H O
3 2
2FeOðOHÞ"Fe O 1H O
2 3 2
With limited dissolved oxygen, iron (Fe21)-containing materials are
favored, including ferrous oxide (FeO) and magnetite (Fe O ). High-
3 4
oxygen concentrations favor the formation of ferric (Fe31) materials.
The nature of rust changes with time, reflecting the slow rates of the
reactions of solids.
On an electrochemical basis, the rust builds up a coating over the
anode’s surface. Ferrous hydroxide may then react with more water to
produce another form of rust: ferric hydroxide [Fe(OH) ]. These layers
3
of rust create tubercles (mounds on a metal surface that form a micro-
environment), which can become problematic because they decrease
the carrying capacity of the pipe and can be dislodged during high
water flows, resulting in red water (water containing rust as a suspen-
sion). It should be noted that this is not the same reaction that turned
Corrosion e5
the Nile red in Biblical times, which was due to an algal infestation of
the river. The tubercles actually slow the rate of corrosion by prevent-
ing the electrolyte from interacting with the anode. If the tubercles are
dislodged, the anode again comes in contact with water and the corro-
sion rate increases.
1.2.3 Cathodic Reactions
In cathodic reactions, the electrons from the breakdown of elemental
iron flow through the pipe wall to the cathode. There, they leave the
metal and enter the water by reacting with hydrogen ions and forming
hydrogen gas:
Hydrogenions 1 Electrons2Hydrogengas
2H1 2e2 H2
The hydrogen gas remains on the cathode and thereby isolates the
cathode from the water by polarization. Just as the buildup of a tuber-
cle breaks the connection between the anode and the electrolyte and
slows the corrosion process, polarization breaks the connection
between the cathode and the electrolyte and slows corrosion.
Dissolved oxygen in the water is able to react with the hydrogen
gas surrounding the cathode:
Hydrogengas 1 Oxygen-Water
2H2 O2 2H2O
Thisreaction(depolarization)removesthehydrogengassurrounding
the cathode and speeds up the corrosion process and, as a result, water
thathasahighcontentofdissolvedoxygenismorecorrosivethanwater
withalowcontentofdissolvedoxygen.Bycombining theelectrical and
chemical reactions, pipeline corrosion, in particular, involves a series of
electrochemicalevents.Hydrogengaswillcoatthecathodeandseparate
itfromthewaterbypolarization.
1.3 TYPES OF CORROSION
While corrosion (specifically rust formation) can be represented by rel-
atively simple chemical reactions (shown above), that is not the com-
plete story. There are various forms of corrosion, some of which
involve redox reactions and many of which do not involve redox