Aluminum corrosion and the effect of chloride

Oxide layer becomes thicker

As can be seen from the decreasing current densities in Figure 1, the oxide layer becomes slightly thicker with storage time. This dependency is shown again in Figure 2 .

Fig. 2: Current density under condensation conditions with a voltage of 1.75 V applied as a function of the storage time in days

If one assumes that the voltage and the conductivity of the layer do not change, Ohm's law for the layer thickness follows:

Fig. 3: Simple potential-pH diagram of aluminumInthe course of twenty days, the layer thickness approximately doubles.

Unfortunately, it is also the case that passive layers on metals are always attacked - at least slightly. The passive layer of aluminum is only relatively stable in the pH range of around 4 to 9 (Fig. 3). At lower and higher pH values, it is severely attacked and can no longer guarantee protection. Stains for aluminum must therefore work in these ranges. However, as most applications are in the medium pH range, technical use is relatively unproblematic (Fig. 3).

Dense protective layers form on aluminum in the pH range of 4-10, which prevent corrosion. As expected, lowering the pH in the acidic range leads to an increase in the corrosion rate (Fig. 4).

The increase in corrosion rate with pH in the alkaline range is due to complexation. ^Al3+ ions do not dissolve, but [Al(OH₎₄] ions do.

Another form of attack is by chloride.
Unfortunately, the oceans contain around three percent chloride. As a result, chloride is practically omnipresent and thus a major component of aluminum corrosion. This is why (filiform) corrosion is particularly pronounced near the sea

 

Pitting corrosion is the most common type of corrosion of aluminum and is initiated, for example, when an oxygen atom in the oxide layer is replaced by a chlorine atom (Fig. 5).

Self-healing through oxygen

This makes the layer, which is only a few atomic layers thick, more permeable at this point. The layer is impermeable because it was able to grow just enough to become dense.

With the further replacement of oxygen ions by chloride ions, white corrosion efflorescence (hydrolyzed aluminium chloride) develops locally. In the best case scenario, the oxygen displaces the chloride again and self-healing takes place, i.e. the pores "suffocate" and the corrosion (almost) comes to a standstill (Fig. 6 and 7).

 

Pitting is also caused by the environment (Fig. 7). In industrial and sea air, corrosion is more severe than in land air. These environments contain chloride and other aggressive substances.

The composition and structure of the aluminum alloy is the second main factor influencing the corrosion of aluminum. For example, an alloy that is resistant to seawater (e.g. containing Mg; top layer: MgO) can fail completely in an acidic industrial atmosphere.

Less noble alloy partners dissolve

Pure aluminum naturally has the highest corrosion resistance. It decreases as the concentration of other metals in the alloy increases. Copper, with its high positive potential, has the strongest effect. In contrast, the more electronegative metals magnesium, zinc and chromium have a much smaller influence (Fig. 8).

Fig. 8: Influence of the alloy components on corrosion behavior

However, the highly electronegative lithium also has a significant effect. For this reason, its concentrations in today's alloys have been reduced compared to first-generation lithium alloys, which exhibit mechanically better values.

If precipitation by alloying elements or intermetallic phases occurs in an alloy, short-circuited local elements are formed. If a suitable electrolyte is also present, which is almost always the case due to humidity, the less noble partner dissolves. If pure aluminum is the less noble partner, the internal resistance of the element may become so high due to the dense oxide layer that forms (specific resistance ¹⁰¹² Ωm) that the current flow comes to a standstill. This is referred to as "self-healing". In the last column of Table 1, some normal potentials are also given. It can be seen that copper as an alloying element leads to a significant increase in the potential. This causes the neighboring copper-depleted aluminum in the local element to dissolve (Fig. 9).

Fig. 9: Local element on a precipitate

Contacts with foreign metals naturally also lead to potential formation. Some of these are listed in Table 1.

Tab. 1: Potentials of metals and alloys against pure aluminum, sorted by level

Metal	Potential against the normal hydrogen electrode mV	Potential against Pure aluminum 99.5 % mV
magnesium	-2372	-850
Zinc	-769	-300
AlZn 4 Mg 1	-719	-250
AlZn 1	-619	-150
AlMg 5	-520	-50 - -40
AlMg3	-520	-30 - -20
AlZnMgCu	-490	-20 - -10
AlMgMn	-480	-10 - ±0
Pure aluminum 99,5 %	-470	±0
AlMgMn	-470	±0 - +10
AlMn	-460	+10 - +20
Cadmium	-465	±0 - +20
EN-WC AlSi 12	-420 - -450	+30 - +60
EN-WC AlCu 4	-370	+100
iron	-370	+100
AlCuMg, cold hardened	-300 - -320	+150 - +180
Brass (50 % Zn)	-70	+400
nickel	+11	+480
copper	+80	+550
Chrome-nickel steel 18/8	-100	+850

As a valve metal, aluminum protects itself with its dense oxide layer. However, it is only 4-5 nm thick, i.e. only a few atomic layers. Its resistance is somewhat lower at surface inhomogeneities, so that a low current flow is possible at these points. This leads to attack at these points. There are inhomogeneities at roughnesses, inclusions and roll-ins. However, they can also occur when individual atoms are detached from the bond (see above). The latter naturally occurs from time to time due to the attack of the surrounding medium, especially at high or low pH values of the corrosion medium. Such defects have been detected electrochemically [3, 4] and by electron microscopy [5].

Hydration energy determines ion stability

Hydrated ^Al3+ is always [Al(_H2O)6^]3+. This means that the inner water shell is chemically hydrated (stoichiometric composition). The outer shell is then predominantly physically bound (Coulombic forces). These hexaaquaaluminum ions act as a weak acid (pKs = 4.97) [8]. The level of hydration energy ultimately determines whether an ion is stable or not.

Fig. 10: Size ratios of the aluminum ions

Tab. 2: Shows the thermodynamic data of the hydration of aluminum ions

ion	z	r0	n	-∆Gh0	-∆Hh0	-∆Sh0	ri	z2/r
-	e	nm	-	kJ/mol	kJ/mol	J/(mol-K)	nm	nm-1
Al+	1		6	464	485	70,7	0,129	7,752
Al2+	2			1941	2021	270,7	0,107	37,383
Al3+	3	0,057	6	4577	4707	441,0	0,085	105,882

Fig. 11: Energy diagram of the formation of aluminum ionsThermodynamicdata of the hydration of aluminum ions at infinite dilution are summarized in Table 2. The determining parameter for the free energy is ^z2/r (Fig. 10). The energy ratios of the formation of aluminum ions are shown in Figure 11 (but without the extremely short-lived ^Al2+). As you can undoubtedly see, energy is only released in the trivalent ions. It is therefore the most stable of the aluminum ions.

The unstable monovalent ions are stabilized by complexation. Chloride is an important stabilizer. The hydrated aluminium ions also hydrolyse. The coordination number 6 results in the following equilibrium for [Al(_H2O)6]³⁺:

As a result, the corrosion electrolyte becomes more acidic. The cathodic reaction is:

 

in relation to the dissolved oxygen.

If corrosion products precipitate, the electrolyte is so highly concentrated that the water activity must also be taken into account.

It is difficult to explain the mechanism of aluminum dissolution because several reactions take place in parallel. The formation of ^Al3+ ions not only takes place via permeation reactions, but also via various chemical reactions (Fig. 12).

1. the disproportionation (DPR) of the ^Al+ ions formed

2. the reaction with hydrogen ions and

3. the reaction with water.

GT10-20_alu-gl-3.jpg GT10-20_alu-gr-12.jpg Fig. 12: Reaction diagram of the anodic reactions of aluminum

Effect of the chloride

Chloride is pure poison for aluminum, even in small doses. Why?

A through reaction such as Al ↔ ^Al3+ + 3 e- will never take place in this way because 4 particles will never collide or fly apart at the same time. Such a summation reaction must therefore take place in stages. The first stage is then:

In the case of aluminum, this is the detachment of the ^3p1electron, the two 3s electrons are bound much more tightly. The Al+ ion is now relatively unstable and short-lived. It continues to react immediately (Fig. 12).

In acidic solution with the hydrogen ions^Al++ 2 H+ ^{↔ Al3++} H2_and in alkaline solution with the hydrogen ions of water ^Al++ 2 _H2O↔ ^Al3++_H2+ 2 OH- .

If chloride is present, it stabilizes the Al+ ion. But that's not all, it also catalyzes the permeation reaction of the aluminium. Protective, stable reaction products are formed during the reaction with water. As a result, the corrosion site heals itself. This is not the case with the permeation reaction involving chloride.

If chloride is present, this reaction is preferred, as it takes place at a more negative potential than without.

The aluminum chloride continues to react as long as free water is present.

Dissociation of the basic salt:

The resulting hydrogen partially escapes to the outside as a gas. However, as it is initially formed atomically, the other part diffuses into the aluminum and leads to local embrittlement, which can have an accelerating effect on stress corrosion cracking.

With dissociation, the chloride is available again for the next passage reaction. Chloride therefore acts as a catalyst for aluminum dissolution. Corrosion can progress. The consequences are, for example, pitting corrosion and filiform corrosion.

The potential of the pitting reaction depends on the chloride concentration with 60 mV per order of magnitude (1 Cl- and 1 e-) [6, 7]

Conclusion for the electroplater from the mode of action of the chloride: If corrosion problems occur, city water (chlorinated!) should never be used as the final rinsing water.