1 The Chemistry of Peels

A Hypothesis of Action Mechanisms and a Proposal of a New Classification of Chemical Peelings

The following definition of chemical peels found in the literature has been chosen and adapted by the authors for the purposes of this chapter. A chemical peel is a treatment technique used to improve and smooth the facial and/or body skin’s texture using a chemical solution that causes the dead skin to slough off and eventually peel off. The regenerated skin is usually smoother and less wrinkled than the old skin.

It is advised to seek training with a specialist such as a dermatologist, plastic surgeon, otorhinolaryngologist (facial plastic surgeon) or maxillofacial plastic surgeon who is experienced in the specific type of peel you wish to perform.

Introduction

This chapter proposes a new classification of chemical peels based on the mechanism of action of chemical peel solutions. The traditionally accepted mechanism has been based on the concept that the effect of a peeling solution on the skin is based purely on its acidity. By using elementary concepts in chemistry three separate mechanisms of action for chemical peeling solutions will be explained:

1 Acidity

2 Toxicity

3 Metabolic interactions.

The literature devoted to chemical peels is full of information about the methodology, indications, contraindications, side effects, as well as the results obtained. Without any proof, acidity has always been assumed to be the sole mechanism of action of peeling agents. All peeling agents were assumed to induce the three stages of tissue replacement: destruction, elimination, and regeneration, all accompanied by a controlled stage of inflammation.

A brief study of the chemistry of the molecules and solutions used in chemical peels immediately questions the hypothesis that acidity is the only basis for the action of peeling solutions. In fact, with the exception of trichloroacetic acid (TCA) and non-neutralized glycolic acid solutions, the most commonly used peeling solutions are only weakly acidic, and phenol and resorcinol mixtures may not be acidic at all, having a pH greater than 7 in some formulations.

You will find detailed below descriptions of some elementary chemistry concepts that, along with a review of the chemistry of the skin, should help to explain the possible interactions between different peelings solutions and the skin. Finally, two new classifications of solutions for peelings will be proposed, one according to their mechanisms of action (classification of L. Dewandre), and the other according to chemical parameters (structure of the molecula, pK_a, etc; or classification of A. Tenenbaum).

Useful Elements of Basic Chemistry

Understanding some of the basic concepts of chemistry is necessary to truly understand chemical peels. Mineral and organic chemistry are taught as biochemistry to medical students, but most practicing physicians do not remember these fundamental sciences.

Also chemistry has been unfortunately evicted in cosmetic dermatology from aesthetic medicine courses, masters, workshops and congresses. A brief review of useful information should help to update most practitioners.

• Acids

An acid (from the Latin acidus meaning sour) is traditionally considered any chemical compound that, when dissolved in water, gives a solution with a hydrogen ion activity greater than in pure water, i.e., a pH less than 7.0. That approximates the modern definition of Johannes Nicolaus Brønsted and Martin Lowry, who independently defined an acid as a compound which donates a hydrogen ion (H⁺) to another compound (called a base). Acid/base systems are different from redox reactions in that there is no change in oxidation state. Acids can occur in solid, liquid or gaseous form, depending on the temperature. They can exist as pure substances or in solution.

Chemicals or substances having the property of an acid are said to be acidic (adjective).

• Polarity and the inductive effect

The polarity of the H—A bond is the first factor contributing to the acid strength.

As the electron density on hydrogen decreases, it is more acidic. Moving from left to right across a row on the periodic table elements become more electronegative (excluding the noble gases).

In several compound classes, collectively called carbon acids, the C—H bond can be sufficiently acidic for proton removal. Unactivated C—H bonds are found in alkanes and are not adjacent to a heteroatom (O, N, Si, etc). Such bonds usually only participate in radical substitution.

Polarity refers to the distribution of electrons in a bond, the region of space between two atomic nuclei where a pair of electrons is shared. When two atoms have roughly the same electronegativity (ability to attract electrons) the electrons are shared evenly and spend equal time on either end of the bond. When there is a significant difference in electronegativities of two bonded atoms, the electrons spend more time near the nucleus of the more electronegative element and an electrical dipole, or separation of charges, occurs, such that there is a partial negative charge localized on the electronegative element and a partial positive charge on the electropositive element. Hydrogen is an electropositive element and accumulates a slightly positive charge when it is bonded to an electronegative element such as oxygen or chlorine.

The electronegative element need not be directly bonded to the acidic hydrogen to increase its acidity. An electronegative atom can pull electron density out of an acidic bond through the inductive effect. The electron-withdrawing ability diminishes quickly as the electronegative atom moves away from the acidic bond.

Carboxylic acids are organic acids that contain an acidic hydroxyl group and a carbonyl (C—O bond). Carboxylic acids can be reduced to the corresponding alcohol; the replacement of an electronegative oxygen atom with two electropositive hydrogens yields a product which is essentially non-acidic. The reduction of acetic acid to ethanol using LiAlH₄ (lithium aluminum hydride or LAH) and ether is an example of such a reaction.

The pK_a for ethanol is 16, compared to 4.76 for acetic acid.

• Atomic radius and bond strength

The size of the atom A or atomic radius is the second factor contributing to the acid strength.

Moving down a column on the periodic table, atoms become less electronegative but also significantly larger, and the size of the atom tends to dominate its acidity when sharing a bond to hydrogen.

Hydrogen sulfide, H₂S, is a stronger acid than water, even though oxygen is more electronegative than sulfur. Just, this is because sulfur is larger than oxygen and the H—S bond is more easily broken than the H—O bond.

Another factor that contributes to the ability of an acid to lose a proton is the strength of the bond between the acidic hydrogen and the atom that bears it. This, in turn, is dependent on the size of the atoms sharing the bond. For an acid HA, as the size of atom A increases, the strength of the bond decreases, meaning that it is more easily broken, and the strength of the acid increases. Bond strength is a measure of how much energy it takes to break a bond. In other words, it takes less energy to break the bond as atom A grows larger, and the proton is more easily removed by a base.

• Chemical characteristics

It is important to keep in mind the difference between monoprotic acids (having one unique pK_a) and polyprotic acids (having two or more pK_a).

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 (CH₃COOH), 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 H₂A) 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 (H₂CO₃) can lose one proton to form bicarbonate anion (HCO₃⁻) and lose a second to form carbonate anion (CO₃²⁻). Both Ka values are small, but K_a1 > K_a2.

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

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

A triprotic acid (H₃A) 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 since it is energetically less favorable to lose a proton if the conjugate base is more negatively charged.

Weak Acid/Weak Base Equilibria

In order to lose a proton, it is necessary that the pH of the system rise above the pK_a of the protonated acid. The decreased concentration of H⁺ in that basic solution shifts the equilibrium towards 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.

• Buffer solution

A buffer solution is an aqueous solution consisting of a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid. It has the property that the pH of the solution changes very little when a small amount of acid or base is added to it. Buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of chemical applications. Many life forms thrive only in a relatively small pH range; an example of a buffer solution is blood.

Buffer Capacity

Figure 1.1 Buffer capacity for pK_a = 7 as percentage of maximum

Buffer capacity is a quantitative measure of the resistance of a buffer solution to pH change on 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 C_A 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 = pK_a.

2 At pH = pK_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 = pK_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.