The Chemistry of Peels: A Hypothesis of Mechanism of Action and Classification of Peels

Arrhenius Acids

The Arrhenius concept is the easiest one retained by most peelers, because most peeling acids are ionic compounds, acting as a source of H ₃ O ⁺ when dissolved in water.

The Swedish chemist Svante Arrhenius attributed the properties of acidity to hydrogen in 1884. An Arrhenius acid is a substance that increases the concentration of the hydronium ion, H ₃ O ⁺ , when dissolved in water. This definition stems from the equilibrium dissociation of water into hydronium and hydroxide (OH ⁻ ) ions:

In pure water most molecules exist as H ₂ O, but a small number of molecules are constantly dissociating and reassociating. Pure water is neutral with respect to acidity or basicity, because the concentration of hydroxide ions is always equal to the concentration of hydronium ions. An Arrhenius base is a molecule that increases the concentration of the hydroxide ion when dissolved in water. Note that chemists often write H ⁺ (aq) and refer to the hydrogen ion when describing acid–base reactions, but the free hydrogen nucleus, a proton, does not exist alone in water; it exists as the hydronium ion, H ₃ O ⁺ .

Brønsted Acids

Although the Arrhenius concept is useful for describing many reactions, it is also quite limited in its scope. Brønsted acids act by donating a proton to water and, differently than Arrhenius acids, can also be used to describe molecular compounds, whereas Arrhenius acids must be ionic compounds.

In 1923 chemists Johannes Nicolaus Brønsted and Thomas Martin Lowry independently recognized that acid–base reactions involve the transfer of a proton. A Brønsted–Lowry acid (or simply Brønsted acid) is a species that donates a proton to a Brønsted–Lowry base. Brønsted–Lowry acid–base theory has several advantages over Arrhenius theory. Consider the following reactions of acetic acid (CH ₃ COOH) (used as a chemical peel for the décolleté by some great peelers like L. Wiest), the organic acid that gives vinegar its characteristic taste:

Both theories easily describe the first reaction: CH ₃ COOH acts as an Arrhenius acid because it acts as a source of H ₃ O ⁺ when dissolved in water, and it acts as a Brønsted acid by donating a proton to water. In the second example CH ₃ COOH undergoes the same transformation, donating a proton to ammonia (NH ₃ ), but it cannot be described using the Arrhenius definition of an acid because the reaction does not produce hydronium.

As with the acetic acid reactions, both definitions work for the first example, where water is the solvent and a hydronium ion is formed. The next reaction does not involve the formation of ions but can still be viewed as a proton transfer reaction.

Lewis Acids

The Brønsted–Lowry definition is the most widely used definition; unless otherwise specified, acid–base reactions are assumed to involve the transfer of a proton (H ⁺ ) from an acid to a base.

A third concept was proposed by Gilbert N. Lewis that includes reactions with acid–base characteristics that do not involve a proton transfer. A Lewis acid is a species that accepts a pair of electrons from another species; in other words, it is an electron pair acceptor. Brønsted acid–base reactions are proton transfer reactions, whereas Lewis acid–base reactions are electron pair transfers. All Brønsted acids are also Lewis acids, but not all Lewis acids are Brønsted acids. Contrast the following reactions, which could be described in terms of acid–base chemistry:

In the first reaction a fluoride ion, F ⁻ , gives up an electron pair to boron trifluoride to form the product tetrafluoroborate. Fluoride “loses” a pair of valence electrons because the electrons shared in the B–F bond are located in the region of space between the two atomic nuclei and are therefore more distant from the fluoride nucleus than they are in the lone fluoride ion. BF3 is a Lewis acid because it accepts the electron pair from fluoride. This reaction cannot be described in terms of Brønsted theory, because there is no proton transfer. The second reaction can be described using either theory. A proton is transferred from an unspecified Brønsted acid to ammonia, a Brønsted base; alternatively, ammonia acts as a Lewis base and transfers a lone pair of electrons to form a bond with a hydrogen ion. The species that gains the electron pair is the Lewis acid; for example, the oxygen atom in H ₃ O ⁺ gains a pair of electrons when one of the H–O bonds is broken and the electrons shared in the bond become localized on oxygen. Depending on the context, Lewis acids may also be described as a reducing agent or an electrophile.

Dissociation and Equilibrium

Reactions of acids are often generalized in the form HA ⇌ H ⁺ + A ⁻ , where HA represents the acid and A ⁻ is the conjugate base. Acid–base conjugate pairs differ by one proton and can be interconverted by the addition or removal of a proton (protonation and deprotonation, respectively). Note that the acid can be the charged species and the conjugate base can be neutral, in which case the generalized reaction scheme could be written as HA ⇌ H ⁺ + A. In solution there exists an equilibrium between the acid and its conjugate base. The equilibrium constant K is an expression of the equilibrium concentrations of the molecules or the ions in solution. Brackets indicate concentration, such that [H ₂ O] means the concentration of H ₂ O. The acid dissociation constant K _a is generally used in the context of acid–base reactions. The numerical value of K _a is equal to the concentration of the products divided by the concentration of the reactants, where the reactant is the acid (HA) and the products are the conjugate base and H ⁺ .

The stronger of two acids will have a higher K _a than the weaker acid; the ratio of hydrogen ions to acid will be higher for the stronger acid because the stronger acid has a greater tendency to lose its proton. Because the range of possible values for K _a spans many orders of magnitude, a more manageable constant, p K _a , is more frequently used, where p K _a = −log ₁₀ K _a . Stronger acids have a smaller p K _a than weaker acids. Experimentally determined p K _a at 25°C in aqueous solution are often quoted in textbooks and reference material.

Acid Strength

For peelers, the notion of acid strength is very important, because stronger acids have a higher K _a and a lower p K _a than weaker acids.

For our classification, two parameters have to be taken into consideration for peelers:

• 1.

  The p K _a, a synonym of the acid’s strength.

• 2.

  The pH, a synonym of the penetration for the selected acid.

For chemists, the strength of an acid refers to its ability or tendency to lose a proton. A strong acid is one that completely dissociates in water; in other words, one mole of a strong acid, HA, dissolves in water, yielding one mole of H ⁺ and one mole of the conjugate base, A ⁻ , and none of the protonated acid HA. In contrast a weak acid only partially dissociates, and at equilibrium both the acid and the conjugate base are in solution. In water each of these essentially ionizes 100%. The stronger an acid is, the more easily it loses a proton, H ⁺ . Two key factors that contribute to the ease of deprotonation are the polarity of the H–A bond and the size of atom A, which determines the strength of the H–A bond. Acid strengths are also often discussed in terms of the stability of the conjugate base.

Caution is advised against simply classifying “cosmetic peels” for acids with p K _a > 3 and “medical peels” for acids with p K _a < 3, because some acids like phenol can be toxic substances even with a p K _a > 3.

Monoprotic Acids

Monoprotic acids are those acids that are able to donate one proton per molecule during the process of dissociation (sometimes called ionization ), as shown below (symbolized by HA):

Common examples of monoprotic acids in organic acids indicate the presence of one carboxyl group, and mostly these acids are known as monocarboxylic acid. Examples in organic acids include acetic acid (CH ₃ COOH), glycolic acid, and lactic acid.

Polyprotic Acids

Polyprotic acids are able to donate more than one proton per acid molecule, in contrast to monoprotic acids that only donate one proton per molecule. Specific types of polyprotic acids have more specific names, such as diprotic acid (two potential protons to donate) and triprotic acid (three potential protons to donate).

A diprotic acid (here symbolized by H ₂ A) can undergo one or two dissociations depending on the pH. Each dissociation has its own dissociation constant, K _a1 and K _a2 .

The first dissociation constant is typically greater than the second; i.e., K _a1 > K _a2 . For example, the weak unstable carbonic acid (H ₂ CO ₃ ) can lose one proton to form bicarbonate anion (HCO ₃ ⁻ ) and lose a second to form carbonate anion (CO ₃ ²⁻ ). Both K _a values are small, but K _a1 > K _a2 .

Diprotic acids used for peelings are malic, tartaric, and azelaic acids.

Two dissociations mean that such acids can generate two peelings, depending on the pH, with the second one less acidic than the first one. In this case, we consider one peeling reaction per one dissociation.

A triprotic acid (H ₃ A) can undergo one, two, or three dissociations and has three dissociation constants, where K _a1 > K _a2 > K _a3 .

An organic example of a triprotic acid is citric acid, which can successively lose three protons to finally form the citrate ion. Even though the positions of the protons on the original molecule may be equivalent, the successive K _a values will differ, because it is energetically less favorable to lose a proton if the conjugate base is more negatively charged.

Weak Acid–Weak Base Equilibria

To lose a proton, it is necessary that the pH of the system rise above the p K _a of the protonated acid. The decreased concentration of H ⁺ in that basic solution shifts the equilibrium toward the conjugate base form (the deprotonated form of the acid). In lower-pH (more acidic) solutions, there is a high enough H ⁺ concentration in the solution to cause the acid to remain in its protonated form, or to protonate its conjugate base (the deprotonated form).

Solutions of weak acids and salts of their conjugate bases form buffer solutions.

Le Chatelier’s Principle

In a solution there is an equilibrium between a weak acid, HA, and its conjugate base, A ⁻ :

• When hydrogen ions (H ⁺ ) are added to the solution, equilibrium moves to the left, as there are hydrogen ions (H ⁺ or H ₃ O ⁺ ) on the right-hand side of the equilibrium expression.

• When hydroxide ions (OH ⁻ ) are added to the solution, equilibrium moves to the right, as hydrogen ions are removed in the reaction (H ⁺ + OH ⁻ → H ₂ O).

Thus, in both cases, some of the added reagent is consumed in shifting the equilibrium in accordance with Le Chatelier’s principle, and the pH changes by less than it would if the solution were not buffered.

Henderson–Hasselbach Equation

The acid dissociation constant for a weak acid, HA, is defined as:

Simple manipulation with logarithms gives the Henderson–Hasselbach equation, which describes pH in terms of p K _a :

In this equation [A ⁻ ] is the concentration of the conjugate base and [HA] is the concentration of the acid. It follows that when the concentrations of acid and conjugate base are equal, often described as half-neutralization, pH = p K _a . In general, the pH of a buffer solution may be easily calculated, knowing the composition of the mixture, by means of an ICE table. An ICE (initial, change, equilibrium) table is a simple matrix formalism that used to simplify the calculations in reversible equilibrium reactions (e.g., weak acids and weak bases or complex ion formation).

One should remember that the calculated pH may be different from measured pH.

Buffer Capacity

Buffer capacity ( Fig. 1.1 ) is a quantitative measure of the resistance of a buffer solution to pH change with the addition of hydroxide ions. It can be defined as follows:

where dn is an infinitesimal amount of added base and d(pH) is the resulting infinitesimal change in pH. With this definition the buffer capacity can be expressed as:

where K _w is the self-ionization constant of water and CA is the analytical concentration of the acid, equal to [HA] + [A ⁻ ]. The term K _w /[H ⁺ ] becomes significant at pH greater than about 11.5, and the second term becomes significant at pH less than about 2. Both these terms are properties of water and are independent of the weak acid. Considering the third term, it follows that:

• 1.

  Buffer capacity of a weak acid reaches its maximum value when pH = p K _a .

• 2.

  At pH = p K _a ± 1 the buffer capacity falls to 33% of the maximum value. This is the approximate range within which buffering by a weak acid is effective. Note: at pH = p K _a − 1, the Henderson–Hasselbach equation shows that the ratio [HA]:[A ⁻ ] is 10:1.

• 3.

  Buffer capacity is directly proportional to the analytical concentration of the acid.

Useful Buffer Mixtures

• Citric acid, sodium citrate, pH range 2.5 to 5.6.

• Acetic acid, sodium acetate, pH range 3.7 to 5.6.

Neutralization

There is among physicians a big confusion between a buffered peel (see above) and a neutralized peel. In chemistry, neutralization is a chemical reaction whereby an acid and a base react to form water and a salt.

In an aqueous solution, solvated hydrogen ions (hydronium ions, H ₃ O ⁺ ) react with hydroxide ions (OH ⁻ ) formed from the alkali to make two molecules of water. A salt is also formed. In nonaqueous reactions, water is not always formed; however, there is always a donation of protons (see Brønsted–Lowry acid–base theory).

Often, neutralization reactions are exothermic, giving out heat to the surroundings (the enthalpy of neutralization). On the other hand, an example of endothermic neutralization is the reaction between sodium bicarbonate (baking soda) and any weak acid—for example, acetic acid (vinegar).

Neutralization of the chemical peeling agent is an important step, the timing of which is determined by either the frost in the skin or how much contact time the peel has with the skin. Neutralization is achieved by a majority of peelers applying cold water or wet, cool towels to the face following the frost. According to physical chemistry, using water just after the frost provokes an exothermic reaction that can provoke a “cold” burn. Other neutralizing agents that can be used include bicarbonate spray or soapless cleansers. Peeling agents for which this neutralization step is less important include salicylic acid, Jessner’s solution, TCA, and phenol.

In partially neutralized alpha-hydroxy acid (AHA) solutions, the acid and a lesser amount of base are combined in a reversible chemical reaction that yields unneutralized acid and a salt.

The resulting solution has less free acid and a higher pH than a solution that has not been neutralized. In partially neutralized formulations, the salt functions as a reservoir of acid that is available for second-phase penetration. This means that partially neutralized formulas can deliver as much, if not more, AHA than free acid formulas, but in a safer, “time-released” manner. Therefore the use of partially neutralized glycolic acid solutions seems prudent, because they have a better safety profile than low-pH solutions containing only free glycolic acid.

Clinical studies have shown that a partially neutralized lactic acid preparation improves the skin, both in appearance and histologically. Other studies using skin tissue cultures showed that partially neutralized glycolic acid stimulates fibroblast proliferation—an index of tissue regeneration. Looking at electrical conductance of the skin (an indicator of water content or moisturization), higher pH products (those that have been partially neutralized) are better moisturizers than lower pH preparations.