The Chemistry of Peels: A Hypothesis of Action Mechanisms and a Proposal of a New Classification of Chemical Peelings

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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.


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:

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, pKa, 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.


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).

Brønsted Acids

While 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 at the difference of 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 (CH3COOH), (used as 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: CH3COOH acts as an Arrhenius acid because it acts as a source of H3O+ when dissolved in water, and it acts as a Brønsted acid by donating a proton to water. In the second example CH3COOH undergoes the same transformation, donating a proton to ammonia (NH3), but 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 hydronium ion is formed. The next reaction does not involve the formation of ions but can still be viewed as 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 which 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 while 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 H3O+ 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 image 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 image 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 [H2O] means the concentration of H2O. The acid dissociation constant Ka is generally used in the context of acid-base reactions. The numerical value of Ka 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 Ka than the weaker acid; the ratio of hydrogen ions to acid will be higher for the stronger acid as the stronger acid has a greater tendency to lose its proton. Because the range of possible values for Ka spans many orders of magnitude, a more manageable constant, pKa is more frequently used, where pKa = −log10Ka. Stronger acids have a smaller pKa than weaker acids. Experimentally determined pKa at 25°C in aqueous solution are often quoted in textbooks and reference material.

Acid Strength

For peelers, this notion is very important because stronger acids have a higher Ka and a lower pKa than weaker acids.

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

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.

According to the classification of A. Tenenbaum, which is described later in this chapter, peelers should be careful with the dangerous distinction between so called ‘cosmetic’, peelings for acids with pKa > 3 and ‘medical’, peelings for acids with pKa < 3, because some acids like salicylic acid with a pKa near 3, as the phenol, toxic substance with a pKa > 3 need to be exclusively used by trained physicians.

Polarity and the inductive effect

The polarity of the HA 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 pKa for ethanol is 16, compared to 4.76 for acetic acid.

Chemical characteristics

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

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, Ka1 and Ka2.



The first dissociation constant is typically greater than the second; i.e., Ka1 > Ka2. 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 Ka1 > Ka2.

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 Ka1 > Ka2 > Ka3.




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 Ka values will differ since it is energetically less favorable to lose a proton if the conjugate base is more negatively charged.

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

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 Kw is the self-ionization constant of water and CA is the analytical concentration of the acid, equal to [HA]+[A]. The term Kw/[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: