HYDROGEN OVERPOTENTIAL

5.5 HYDROGEN OVERPOTENTIAL

The polarization term that controls the corrosion rate of many metals in deaer- ated water and in nonoxidizing acids is hydrogen overpotential. In accord with the previously discussed defi nition of polarization, hydrogen overpotential is the difference of potential between a cathode at which hydrogen is being evolved,

φ measured , and a hydrogen electrode at equilibrium in the same solution; that is,

2 H overpotential = φ measured −− ( . 0 059 pH) Hydrogen overpotential, therefore, is measured in the same way as polarization

is measured. Ideally, hydrogen overpotential includes only the activation polar- ization term corresponding to the reaction 2H +

+ 2e − → H 2 , but reported values often include iR drop and sometimes concentration polarization as well.

64 KINETICS: POL ARIZATION AND CORROSION R ATES

The absolute values of hydrogen overpotential for a given metal decrease with:

1. Increasing temperature ( i 0 increases). For metals that corrode by hydro- gen evolution, decreasing hydrogen overpotential is one factor accounting for increase of corrosion as the temperature is raised.

2. Roughening of the surface. A sandblasted surface has a lower hydrogen overpotential than a polished surface. Greater area and improved catalytic activity of a rough surface account for the observed effect.

3. Decreasing current density. The Tafel equation,

i ηβ = log

is obeyed within the region of applied current density, i , below the limiting current density for concentration polarization, and above the exchange current density, i 0 . The term β is equal to 2.3 RT / α F , where α is a constant and the other terms have their usual signifi cance. The term α is approxi- mately 2 for platinum and palladium and is approximately 0.4 – 0.6 for Fe, Ni, Cu, Hg, and several other metals. Although hydrogen overpotential values usually differ in acid compared with alkaline media, values are not greatly sensitive to pH within either range.

Stern [4] showed that for a corroding metal we have

ii ++ r i corr ηβ = log

+ 2e − r 2 , which varies with potential and which, at equilibrium, is equal to i 0 . This equation describes the small observed slope of d η / d log i for small values of impressed current density,

where i is the reverse current for the reaction H → 2H +

i , with the slope increasing as i approaches i corr r + i and reaching the true Tafel slope, β , at i > > i r corr + i . Similarly, the overpotential for a noncorroding metal can

be represented by a modifi ed Tafel equation,

ii +

ηβ = log

This equation holds from zero to fi nite values of i (Fig. 5.5 ). Values of i r deter- mined from measured values of η also follow the Tafel equation, but with a slope of opposite sign.

The slow step in the discharge of hydrogen ions on platinum, or palladium, as described earlier, appears to be recombination of adsorbed hydrogen atoms. This assumption is consistent, from kinetic considerations, with an observed value

HYDROGEN OVERPOTENTIAL

of α = 2. For iron, the value of α is approximately 0.5 and, correspondingly, β =

0.1. To account for these values, the proposal has been made that the slow step in the hydrogen evolution reaction on iron is probably

H + H + − ads → H 2 − e

The same slow step may apply to various metals having intermediate values of hydrogen overpotential (e.g., iron, nickel, copper).

For metals of high hydrogen overpotential, such as mercury and lead, the slow discharge of the hydrated hydrogen ion is apparently the slow step:

H + → H ads − −e

For many metals at high current densities, this slow discharge step is also the controlling reaction. The slow discharge step may, instead, be the reduction

The reduction of water as the controlling reaction applies generally to metals in alkaline solutions at high and low current densities.

The rapidity with which H ads combines to form H 2 , either by combination with itself or with H + , is affected by the catalytic properties of the electrode surface. A good catalyst, such as platinum or iron, leads to a low value of hydro- gen overpotential, whereas a poor catalyst, such as mercury or lead, accounts for

a high value of overpotential. If a catalyst poison, like hydrogen sulfi de or certain arsenic or phosphorus compounds, is added to the electrolyte, it retards the rate of formation of molecular H 2 and increases the accumulation of adsorbed hydro- gen atoms on the electrode surface. * The increased concentration of surface hydrogen favors entrance of hydrogen atoms into the metal lattice, causing

hydrogen embrittlement (loss of ductility), and, in some stressed high - strength ferrous alloys, may induce spontaneous cracking, called hydrogen cracking (see Section 8.4 ). Catalyst poisons increase absorption of hydrogen whether the metal

* Increase of hydrogen overpotential normally decreases the corrosion rate of steel in acids, but pres- ence of sulfur or phosphorus in steels is observed instead to increase the rate. This increase probably results from the low hydrogen overpotential of ferrous sulfi de or phosphide, either existing in the

steel as separate phases or formed as a surface compound by reaction of iron with H 2 S or phosphorus compounds in solution. It is also possible [4] that the latter compounds, in addition, stimulate the anodic dissolution reaction, Fe → Fe 2+

− + 2e (reduce activation polarization), or alter the anode – cathode area ratio.

Similarly, arsenious oxide in small amount accelerates corrosion of steel in acids (e.g., H 2 SO 4 ), perhaps forming arsenides, but when present in larger amount (e.g., 0.05% As 2 O 3 in 72% H 2 SO 4 ), it is an effective corrosion inhibitor, probably because elementary arsenic, having a high hydrogen overpotential, deposits out at cathodic areas. Tin salts have the same inhibiting effect and have been used to protect steels from attack by pickling acids during descaling operations.

66 KINETICS: POL ARIZATION AND CORROSION R ATES

is polarized by externally applied current or by a corrosion reaction with accom- panying hydrogen evolution. For this reason, some oil - well brines containing H 2 S are diffi cult to handle in low - alloy steel tubing subject to the usual high stresses of a structure extending several thousand feet underground. Slight general cor- rosion of the tubing produces hydrogen, a portion of which enters the stressed steel to cause hydrogen cracking. In the absence of hydrogen sulfi de, general corrosion occurs, but without hydrogen cracking. High - strength steels, because of their limited ductility, are more susceptible to hydrogen cracking than are low -

strength steels, but hydrogen enters the lattice in either case, tending to blister low - strength steels instead. Values of β, i 0 , and of η at 1 mA/cm 2 for H + discharge are given in Table 5.1 . As can be seen from the table, values of hydrogen overpotential vary greatly with the metal. Values also change with concentration of electrolyte, but the effect, comparatively, is not large.