University of technology

Chemical engineering department

Effect of pH on corrosion rate

Yassir abdulkareem Shallal

2021
Introduction

In chemistry, pH historically denoting "potential of hydrogen" (or "power of hydrogen")] is a

scale used to specify the acidity or basicity of an aqueous solution. Acidic solutions (solutions
with higher concentrations of H+ ions) are measured to have lower pH values than basic
or alkaline solutions.

The pH scale is logarithmic and inversely indicates the concentration of hydrogen ions in the
solution. This is because the formula used to calculate pH approximates the negative of the base
10 logarithm (i.e. the decimal logarithm) of the molar concentration of hydrogen ions in the
solution. More precisely, pH is the negative of the base 10 logarithm of the activity of the H+ ion.

At 25 °C, solutions with a pH less than 7 are acidic, and solutions with a pH greater than 7 are
basic. Solutions with a pH of 7 at this temperature are neutral (e.g. pure water). The neutral value
of the pH depends on the temperature – being lower than 7 if the temperature increases. The pH
value can be less than 0 for very strong acids, or greater than 14 for very strong bases.

1-pH
pH is defined as the decimal logarithm of the reciprocal of the hydrogen ion activity, aH+, in a
solution

For example, for a solution with a hydrogen ion activity of 5×10−6 (at that level, this is essentially
the number of moles of hydrogen ions per litre of solution) there is 1/(5×10−6) = 2×105, thus such
a solution has a pH of log10(2×105) = 5.3. Consider the following example: a quantity of 107 moles
of pure (pH 7) water, or 180 metric tonnes (18×107 g), contains close to 18 g
of dissociated hydrogen ions.
Note that pH depends on temperature. For instance at 0 °C the pH of pure water is about 7.47. At
25 °C it is 7.00, and at 100 °C it is 6.14.

This definition was adopted because ion-selective electrodes, which are used to measure pH,
respond to activity. Ideally, electrode potential, E, follows the Nernst equation, which, for the
hydrogen ion can be written as

where E is a measured potential, E0 is the standard electrode potential, R is the gas constant, T is
the temperature in kelvins, F is the Faraday constant. For H+ number of electrons transferred is
one. It follows that electrode potential is proportional to pH when pH is defined in terms of
activity. Precise measurement of pH is presented in International Standard ISO 31-8 as follows:
A galvanic cell is set up to measure the electromotive force (e.m.f.) between a reference electrode
and an electrode sensitive to the hydrogen ion activity when they are both immersed in the same
aqueous solution. The reference electrode may be a silver chloride electrode or a calomel
electrode. The hydrogen-ion selective electrode is a standard hydrogen electrode.

Reference electrode | concentrated solution of KCl || test solution | H2 | Pt

Firstly, the cell is filled with a solution of known hydrogen ion activity and the emf, ES, is
measured. Then the emf, EX, of the same cell containing the solution of unknown pH is measured.

2-Electrochemical Aspects of Corrosion

Corrosion in aqueous environment and atmospheric environment is an electrochemical process
because corrosion involves the transfer of electrons between a metal surface and an aqueous
electrolyte solution [Zvandasara, 2009]. The current flows from a higher potential to a lower
one. Hence, there are two reactions taking place simultaneously in the system. One reaction
occurs as the electrons are discharged from the surface, called the anode. The released electrons
are consumed in the other circuit of the corrosion cell shown schematically in Fig. 1. The
corrosion cell consists of the following four components: anode, cathode, electrolyte, and
Electronic connector. Several cathodic reactions are possible depending on what reducible
species are present in the solution. Typical reactions are the reduction of dissolved oxygen gas
or the reduction of hydrogen ions 2H+ + 2e- ↔ H2 (in an acidic solution) (1) However, if
oxygen is present, two other reactions may occur: O2 + 4H+ + 4e- ↔ 2H2O (acid solutions)
(2.2) O2 + 2H2O + 4e- ↔ 4OH- (neutral and alkaline solutions) (2)

Figure 1 Electrochemical corrosion of metal [Lyon, 1996]

*The relationship between the corrosion rate and the ph value cannot be
established by a general law for all minerals and acids, but it is a special case for
each metal.

Therefore, in this report, we will adopt the corrosion of iron as an example,

because iron is the most widely used and corrosive metal in the industry.

*It should be noted that the Pourbaix diagram is a thermodynamic diagram

It is possible to predict the occurrence of corrosion or not

However, corrosion rate cannot be calculated using pourbaix diagram .

3-Effect of PH on corrosion

The effect of the pH of solution to which iron or steel is exposed is influenced by temperature.
The potential of hydrogen or symbol pH is defined as the negative logarithm of the hydrogen
concentration, represented as [H+ ] in moles/liter.

pH = - log [H+ ] (1.3)

The pH value is used to represent the acidity of a solution. First, consider the exposure of iron to
aerated water at room temperature (aerated water will contain dissolved oxygen). The corrosion
rate for iron as a function of pH is illustrated in Fig. 2. In the range of pH = 4 to pH =10, the
corrosion rate of iron is relatively independent of the pH of the solution. In this pH range, the
corrosion rate is governed largely by the rate at which oxygen reacts with absorbed atomic
hydrogen, thereby depolarizing the surface and allowing the reduction reaction to continue. For
pH values below 4.0, ferrous oxide (FeO) is soluble. Thus, the oxide dissolves as it is formed
rather than depositing on the metal surface to form a film. In the absence of the protective oxide
film, the metal surface is in direct contact with the acid solution, and the corrosion reaction
proceeds at a greater rate than it does at higher pH values. It is also observed that hydrogen is
produced in acid solutions below a pH of 4, indicating that the corrosion rate no longer depends
entirely on depolarization by oxygen, but on a combination of the two factors [hydrogen evolution
and oxygen reduction reaction (depolarization)]. For pH valuesabove about pH 10, the corrosion
rate is observed to fall as pH is increased. This is believed to be due to an increase in the rate of
the reaction of oxygen with Fe (OH) 2 n (Hydrated FeO) in the oxide layer to form the more
protective Fe2O3 [Gedeon, 2000]. Iron is weakly amphoteric, at very high temperatures such as
those encountered in boilers, the corrosion rate increases with increasing basicity. PH has no
effect on noble metals such as gold and platinum, but amphoteric metals dissolve rapidly in either
acidic or basic solutions such as aluminum and zinc [Perry and Green, 1997].
4-Corrosion of Iron in Acid

In strong acids, such as hydrochloric and sulfuric acids, the diffusion - barrier oxide film on the
surface of iron is dissolved below pH 4. In weaker acids, such as acetic or carbonic acids,
dissolution of the oxide occurs at a higher pH; hence, the corrosion rate of iron increases
accompanied by hydrogen evolution at pH 5 or 6. This difference is explained by the higher total
acidity or neutralizing capacity of a partially dissociated acid compared with a totally dissociated
acid at a given pH. In other words, at a given pH, there is more available H + to react with and
dissolve the barrier oxide film using a weak acid compared to a strong acid. The increased
corrosion rate of iron as pH decreases is not caused by increased hydrogen evolution alone; in
fact, greater accessibility of oxygen to the metal surface on dissolution of the surface oxide favors
oxygen depolarization, which is often the more important reason. In more concentrated acids, the
rate of hydrogen evolution is so pronounced that oxygen has difficulty in reaching the metal
surface. Hence, depolarization in more concentrated acids contributes less to the overall corrosion
rate than in dilute acids, in which diffusion of oxygen is impeded to a lesser extent. Potential and
polarization measurements indicate that oxygen in small concentrations at the metal surface
increases cathodic polarization, thereby decreasing corrosion; in higher concentrations, oxygen
acts mainly as a depolarizer, increasing the rate. The important depolarizing action of dissolved
oxygen suggests that the velocity of an acid should markedly affect the corrosion rate. This effect
is observed, particularly with dilute acids, In addition, the inhibiting effect of dissolved oxygen
is observed within a critical velocity range, with the critical velocity becoming higher the more
rapid the inherent reaction rate of steel with the acid. Relative motion of acid with respect to metal
sweeps away hydrogen bubbles and reduces the thickness of the stagnant liquid layer at the metal
surface, allowing more oxygen to reach the metal surface. In the absence of dissolved oxygen,
only hydrogen evolution occurs at cathodic areas, and an effect of velocity is no longer observed.
This result is expected because hydrogen overpotential (activation polarization) is insensitive to
velocity of the electrolyte For aerated acid, the minimum rate occurs at higher velocities the more
concentrated the acid because the rate of hydrogen evolution is more pronounced, thereby
impeding oxygen diffusion to the metal surface [Revie and Uhlig, 2008].
 References

1-Gedeon G, “Corrosion Overview”, corrosion, vol. 4 pp. 37- 43, 2000.

2-Lyon W., “Standard Handbook of Petroleum and Natural Gas Engineering”, vol.1, Gulf
Company, Houston, Texas, 1996.

3-Perry R.H, and D.W. Green, “Perry Chemical Engineers Handbook”, 7thed, M.C. Graw-Hill,
United States, 1997.

4-Revie R.W. and H. H. Uhlig, “Corrosion and Corrosion Control an Introduction to Corrosion
Science and Engineering” fourth edition, John Wiley & Sons, Inc., Hoboken New Jersey, 2008.

5-Zvandasara T., “Influence of Hydrodynamics on Carbon Steel Erosion Corrosion and

Inhibitor Efficiency in Simulated Oil Field Brines”, PhD thesis 2009

6-Wikipedia