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How much Chemistry/Biochemistry is there on the MCAT? Content Category 5A: Unique nature of water and its solutions Content Category 5B: Nature of molecules and intermolecular interactions Content Category 5C: Separation and purification methods Content Category 5D: Structure, function, and reactivity of biologically-relevant molecules Content Category 5E: Principles of chemical thermodynamics and kinetics FAQs

How much Chemistry/Biochemistry is there on the MCAT?

The Chemical and Physical Foundations of Biological Systems section of the MCAT is comprised of two main components: MCAT Chemistry and MCAT Physics equations.

The chemistry part of this section is designed to test the knowledge of General Chemistry (30 %), Organic Chemistry (15 %), Introductory Biology (5 %), and first semester Biochemistry (25 %), as well as 25% MCAT Physics. The percentages are approximated to nearest 5% and can vary from one test to another test.

MCAT Chemistry/Biochemistry

What does this section test? All the processes happening in a biological body or our surroundings are basically chemical reactions which include bond breaking, bond formation, electron transfer, etc. Understanding of elements, compounds, structural formulas, properties, chemical and physical changes in terms of reactions, controlling the speed of a reaction etc. is vital in understanding almost all chemical and biochemical processes. This section tests your understanding of these fundamentals and subsequently their application to the functioning of biosystems. Now let’s get into the MCAT chemistry equations you need to know to ace this MCAT section!

Looking for MCAT physics equations as well? Check out our video!

5A Unique nature of water and its solution

Water is considered as a universal solvent. All the biochemical reactions are carried out in water as solvent. Water is also very important in maintaining buffering conditions in living system. Unique properties of water allow it to strongly interact with ions, biomolecules.

The content in this category covers solutions, solubility, acids, bases, buffers.

Acid and bases

Both Acids and Bases can be classified either as Arrhenius, Bronsted Lowry

Arrhenius acid and Bases

This definition exists only for aqueous solutions. An Arrhenius acid is a substance which increases hydronium ion concentration in water and an Arrhenius base is a substance that increases hydroxide ion concentration in water.

ACID: HCl(aq)  H+(aq) + Cl-(aq)

BASE: NaOH(aq)   OH-(aq)  + Na+(aq) 

Bronsted Lowry acid and base

Any substance that donates a proton is a Bronsted acid and any substance that is a proton acceptor is a Bronsted base. It is not specific to aqueous solutions only.

NH3(g) + H2O(l) ⇌ NH4+(aq) + OH-(aq)

Ammonia act both as Arrhenius and Bronsted Lowry Base

NH3(g) + HCl(aq) ⇌ NH4+(aq)+Cl(aq)

Ammonia act as only Bronsted Lowry Base

Autoionization of water

Water is amphoteric in nature acting as both Bronsted-Lowry acid and base. Since acid and bases can react with each other, that implies that water molecules can react with each other. Water molecules exchange proton with each other.


One molecule of water is donating a proton acting as Bronsted-Lowry acid and the other is accepting a proton thus acting as Bronsted-Lowry Base. Because water is a weak acid and a weak base, the hydronium and hydroxide ions exist in very small amount relative to that of non-ionized water and the reaction lies mostly towards the left-hand side. The equilibrium expression which is a ratio of products to reactants is therefore given as below. Remember that the concentrations of pure solids and liquids is not included in the expression.

Kw = [H3O]+[OH]-= 10-14

The value of Kw calculated at 25OC is 10-14. Taking negative logarithm on both sides of the equation we get:

−logKw​​=−log([H3​O+][OH])= -log 10-14

pKw=pH+ pOH​ = 14

pH = pOH = 7

pH is a measure of hydrogen ion concentration, a measure of the acidity or alkalinity of a solution. The pH scale usually ranges from 0 to 14. Aqueous solutions at 25°C with a pH less than 7 are acidic, while those with a pH greater than 7 are basic or alkaline

pH = -log[H3O+]

[H3O+] stands for the hydrogen ion concentration in units of moles per liter solution

Conjugate Acid-Base pair

NH3(g) + HCl(aq)NH4+(aq)+Cl(aq)

With Brønsted–Lowry acids and bases, in each acid base reaction, a conjugate acid–base pair consists of two substances that differ only by the presence of a proton (H⁺). A conjugate acid is formed when a proton is added to a base, and a conjugate base is formed when a proton is removed from an acid. In the above reaction, NH3(base) is accepting a proton to become NH4+ (conjugate acid of NH3). Similarly, HCl(acid) is donating a proton to become Cl-(conjugate base of HCl).

 Strong Acids and Bases

Strong Acids are the one which when dissolved in water dissociate completely. There are 7 strong acids known: HCl - hydrochloric acid, HNO3 - nitric acid, H2SO4 - sulfuric acid (HSO4- is a weak acid), HBr - hydrobromic acid, HI - hydroiodic acid, HClO4 - perchloric acid, HClO3 - chloric acid

HCl  H + + Cl- (complete dissociation)

Note: All the strong acids are monoprotic i.e they donate one proton(H+) except for H2SO4 which is a diprotic acid i.e donates two protons(H+)

Strong Bases are the one which dissociate completely when dissolved in water. Hydroxides of alkali and alkaline earth metals (NaOH, KOH, Ca(OH)2, LiOH etc.

NaOH  Na+ + OH – (complete dissociation)

Weak Acid and Bases

All other acids are classified as weak acids. Few examples: HO2C2O2H - oxalic acid, H2SO3 - sulfurous acid, HSO4 - hydrogen sulfate ion, H3PO- phosphoric acid, HNO- nitrous acid, HF - hydrofluoric acid, C6H5COOH - benzoic acid, CH3COOH - acetic acid, HCOOH - formic acid.

               CH3COOH + H2O ⇌ H3O+ + CH3COO- (incomplete dissociation as shown by double sided arrow)

Weak Bases include hydroxide of group III metals, transition metal hydroxides and electron rich nitrogen containing compounds (amines). (NH3), Al(OH)3, Pb(OH)2, Fe(OH)3, Cu(OH)2, Zn(OH)2, N(CH3)3, CH3NH2

Al(OH)3 ⇌ Al+3 + 3OH- (incomplete dissociation as shown by double sided arrow)

Dissociation and Weak Acid /Bases

Incomplete dissociation of weak acids and bases means that the reaction not only proceeds in the forward direction but also happens simultaneously in the backward direction. Consider a weak acid HA dissociating partially

HA + H2O ⇌ H3O+ + A-

In the above generic weak acid/base reaction, all the species are present in a dynamic equilibrium with each other. The equilibrium expression is a ratio of products to reactants. The equilibrium constant for a weak acid dissociation reaction is specifically known as acid dissociation constant (Ka)

Ka = [H3O]+[A] -̸ [HA]

The acid dissociation constant Ka quantifies the extent of dissociation of a weak acid. The larger the value of Ka means more amount of acid dissociate into H+ and its conjugate base A- and hence the stronger the acid, and vice versa.

Similarly for a weak base, consider a weak base dissociating partially

B + H2O ⇌ BH+ + OH-

The equilibrium constant for a weak base dissociation reaction is specifically known as base dissociation constant (Kb)

Kb = [BH]+[OH]-̸ [B]

The base dissociation constant Kb quantifies the extent of dissociation of a weak base. The larger the value of Kb means the more amount of base dissociate into OH- and is conjugate acid BH+ and hence the stronger the base, and vice versa

For conjugate acid-base pair as shown below at 25oC

HA + H2O ⇌ H3O+ + A-

Ka.Kb = Kw = [H3O]+[OH]- = 10-14

Taking negative logarithm on both sides

pKa​+ pKb​=14

pKa= -log Ka

pKb = -logKb

We can use these equations to determine Kb or( pKb) of a weak base given Ka of conjugate acid. We can also calculate the Ka or(pKa) of a weak acid given Kb of its conjugate base.

Dissociation and Weak Acid /Bases in presence of added salts

Since weak acids and bases are only partially ionized and in equilibrium, addition of salts which increases the amount of product in an equilibrium reaction can suppress the dissociation of weak acids or bases and vice versa. This is known as Common ion effect

CH3COOH + H2O ⇌ H3O+ + CH3COO-

Adding CH3COONa (CH3COO- + Na+) shifts the equilibrium to left thus suppressing acid dissociation

Hydrolysis of salts of weak acids or bases

Salts are ionic compound. Hydrolysis of salt means reaction of salt with water. We can determine whether a salt solution will be acidic, basic, or neutral by considering the reactivity of both the cation and the anion with water. This in turn depends on the type cations and anions in a salt and its source.

Anions and Cations of the salt do NOT react with water = Neutral solution

Example includes salt coming from reaction of strong acid and strong base

NaOH + HCl NaCl + H2

Neither Na+(Group IA and group II A cations) and Cl- ( conjugate base of strong acid HCl) react with water. Since pH of water is 7 so the resulting solution is neutral

Anion reacts with water, but Cation does NOT react with water = Basic solution

Examples include salt coming from reaction of strong base and weak acid

Ba(OH)2 + CH3COOH  Ba(CH3COO)2 + H2O

Ba+2 (group II A cation) does not react with water whereas CH3COO anion (conjugate base of weak acid ) reacts with water making acetic acid and OH- therefore increasing concentration of OH- making the solution basic.

Cation reacts with water but Anion does NOT react with water = Acidic solution

Examples include salt coming from reaction of strong acid and weak base


NH4+ (not from group IA or group IIA) therefore reacts with water making NH4OH and H+. Thus, proton concentration increases making the solution acidic. NO3- (conjugate base of strong acid) does not react with water


A buffer is a solution which resists change in pH upon addition of acidic or basic component. Buffer solutions have a working pH range and buffer capacity which indicates how much acid/base can be neutralized before pH changes, and the amount by which it will change.


Buffer can be made either by combination of

Weak acid and salt of its conjugate base : CH3COOH/CH3COONa

The Henderson-Hasselbalch approximation allows us one method to approximate the pH of a buffer solution from a weak acid. The basic equation is as follows:

pH= pKa+ log [A] ̸ [HA]

[HA] = Concentration of weak acid

[A-] = Concentration of Conjugate base

pKa = acid dissociation constant of weak acid

Weak base and salt of its conjugate base: NH3/NH4Cl

pOH = pKb + log[HB+] ̸ [B]

[HB+] = Concentration of conjugate acid

[B] = Concentration of weak base

pKb = base dissociation constant of weak base


Solubility tells a limit to how much solute can be dissolved in given amount of solvent. Concentration of a solution is a quantitative measure of amount of solute dissolved in solvent/solution

Common way to express concentration:

Molarity(M):    moles of solute /one liter of solution

Molality (m):    moles of solute / one kilogram of solvent

Mass% (m/m%): (mass of solute/mass of solution) x100

Volume % (v/v %): (Volume of solute/volume of solution) x100

Parts per million(ppm): (mass of solute/mass of solution) x100

Mole fraction (χ) : χA = moles of A/ moles of all the substances

Solubility Product

The solubility product constant, Ksp, is the equilibrium constant for a solid substance dissolving in an aqueous solution. It represents the maximum level at which a solute can dissolve in solution. The more soluble a substance is, the higher the Ksp value. For a general dissolution reaction

aA (s) ⇌ cC(aq) + dD (aq)

KSp = [C]c/[D]d

Note that pure solids are not included in an equilibrium expression.

5B Nature of molecules and intermolecular interactions

  • Concept of covalent bonding. Types of covalent bonds (polar, non-polar covalent bonds). Lewis dot symbol of neutral atoms and ions Lewis electron dot diagrams use dots to represent valence electrons around an atomic symbol.
  • Lewis electron dot diagrams for ions have fewer (for cations) or more (for anions) dots than the corresponding atom. Writing the Lewis structure of a molecule.
  • Valence shell electron pair repulsion theory. The valence shell electron pair repulsion (VSEPR) theory is a model used to predict 3-D molecular geometry based on the number of valence shell electron bond pairs among the atoms in a molecule or ion. According to this model electron pairs will arrange themselves to minimize repulsion effects from one another. The molecule or polyatomic ion is given an AXmEn designation, where A is the central atomX is a bonded atomE is a nonbonding valence electron group (usually a lone pair of electrons), and m and n are integers. The number of groups is equal to the sum of m and n. Using this information, we can describe the molecular geometry around a central atom.
  • Resonating structures: Set of 2 or more Lewis structures which have same number of electrons and atoms but only differ in the type of connectivity between atoms.
  • Formal charge calculation on Lewis structure. Formal charge indicates whether a molecule is overall neutral or a charged species(ions)

FC = V-N-B/2

FC= formal charge, V= number of valence electrons, N= Number of nonbonding electrons, B= number of bonding electrons

  • Polar Vs non-polar bonds in a covalent compound. When bonds are formed between two similar atoms, the pull towards the shared bonding electron is same. However, the bond between two unlike atoms, which differ in their affinities for electrons (electronegativity) is said to be a polar covalent bond. When a covalent bond is formed between two atoms of different elements, the bonding pair of electrons will lie more towards the atom, which has more affinity for electrons (i.e. higher electronegativity value).
  • A polar covalent bond behaves as if it were partially ionic. The greater the difference in electronegativity (EN) the greater the ionic character of the bond. Remember ionic bonds are forms by complete transfer of electrons
  • Dipole moments occur when there is a separation of charge. This can happen between two ions in an ionic bond or between atoms in a covalent bond(polar covalent bonds); Since dipole moments are a result of electronegativity difference, larger the difference in electronegativity, the larger the dipole moment. The distance between the charge separation is also a deciding factor into the size of the dipole moment. The dipole moment is a measure of the polarity of the molecule.

Dipole Moment (µ) = Charge (Q) * distance of separation (r)

It is measured in Debye units denoted by ‘D’. 1 D = 3.33564 × 10-30 C.m, where C is Coulomb and m denotes a meter.

  • When a bond is formed, there are two different types of overlaps that occur: Sigma (σ) and Pi (π). Sigma (σ) Bonds are formed between two atoms by head-to-head overlapping or bond. Pi (π) Bonds form when two un-hybridized p-orbitals overlap. This is called "side-by-side" bond. To overlap effectively, the orbitals must match each other in energy. The process by which all of the bonding orbitals become the same in energy and bond length is called hybridization (sp, sp2, sp3, sp3d, sp3d2 etc.). We can determine the hybridization by only counting the number of bonds or lone pairs of electrons around a central atom to determine its hybridization.
  • When double and triple bonds are present between two atoms, there is additional bonding holding the atoms together. While a sigma(σ) bond is always the first covalent bond between two atoms, a pi (π) bond is always the second bond between two atoms.

Single bond = one sigma bond

Double bond = one sigma and one pi bond

Triple bond = one sigma and two pi bonds

  • Multiple bonding its effect on bond length and bond energies. Multiple bonding decreases bond length. Multiple bonding increases bond energy. Rigidity in molecular structure. Multiple bonding increases rigidity in molecular structure. Single bonds can rotate, but double and triple bonds cannot

Stereochemistry of Covalently bonded molecules.

Isomers: Compounds with the same formula but different arrangements of atoms

Conformational isomers: differ by rotation around a single sigma bond They don't require bond breaking

Configurational isomers: can only be interchanged by breaking and reforming bonds

Stereoisomers isomers: differ in spatial arrangement of atoms, rather than order of atomic connectivity. They can be further classified cis-trans isomers, enantiomers, diastereomer

Enantiomers isomers: non-superimposable mirror images of each other rotate equal but opposite directions in plane polarized light react differently in chiral environments

Diastereomers: Isomers which are not mirror images of each other

R and S nomenclature: Around the chiral carbon atom, groups can be prioritized by their atomic mass.

Optical isomers: A molecule which has a chiral center or asymmetric carbon atom and differ in the placement of substituted group around one central atom.

Optically active: Capability to rotate the plane of vibration of polarized light to right or left. Chiral molecules are optically active molecules. It would be useful to have a common standard for optical rotation that allows to compare samples collected under slightly different concentrations and path lengths The specific rotation of a molecule is the rotation in degrees observed upon passing polarized light through a path length of 1 decimetre (dm) at a concentration of 1 g/mL.

[α] = αobserved /c x l

[α] = specific rotation in degrees

αobserved = observed rotation

l = pathlength (dm)

c = concentration (g/mL)

Intermolecular forces:

Intermolecular forces are the forces that exist between molecules. These forces determine the physical properties of substances. For example: Density, viscosity, boiling point, melting point, thermal conductivity, thermal expansion, heat of vaporization.

Types of intermolecular forces:

Ion-dipole forces exist between ions and polar (dipole) molecules

Ion-induced dipole forces exist between ions and non-polar molecules

Dipole-dipole forces exist between two polar (dipole) molecules.

Dipole-induced dipole forces exist between a polar molecule and a non-polar molecule.

Induced dipole forces exist between two non-polar molecules.

Hydrogen bonds are a type of dipole-dipole force that occurs when a hydrogen atom is attached to a highly electronegative atom (oxygen, fluorine, nitrogen). A hydrogen atom on one molecule is attracted to the electronegative atom on a second molecule.

5C Separations and Purifications 

Complex mixtures of substances especially biomolecules (peptides, proteins), small organic molecules(chiral, achiral, enantiomers) typically requires separation of the components to be analyzed. Obtaining pure compounds is of great practical importance in chemistry. Most synthetic reactions in chemistry yield mixtures of products and it is important to have a reasonably clear idea of how mixtures of compounds can be separated. Almost all biochemical compounds occur naturally as components of very complex mixtures from which they should be separated and isolated. Many methods have been developed to accomplish this task, and the method used is dependent on the types of substances and type of intermolecular forces that exists between different components of mixtures. The topic in this category covers separation and purification methods involving extraction, chromatography, and electrophoresis. The topics and subtopics are below:

  • Differential extraction is the method of separation of any organic component present in an aqueous solution
  • Column Chromatography: This is a separation method where the analyte is combined within a liquid or gaseous mobile phase pumped through a stationary phase. Utilized in the separation of small organic molecules, amino acids, carbohydrates, and fatty acids.
  • There are two main types of column chromatographic techniques:
  • Gas-Liquid chromatography: chromatography employing a gas as the moving carrier medium and a liquid as the stationary medium. Utilized in the separation of alcohol, ester, lipid, amino groups, and observation of enzymatic interactions
  • High-pressure liquid chromatography (HPLC): In HPLC, mobile phase passes through columns under 10–400 atmospheric pressure, and with a high (0.1–5 cm//sec) flow rate.
  • Ion-exchange chromatography: It is based on electrostatic interactions between charged protein groups, and solid support material.
  • Size-exclusion chromatography: It separates macromolecules(proteins, polysaccharides, nucleic acids) based on their differences in molecular sizes. This procedure is basically used to determine molecular weights of proteins, and to decrease salt concentrations of protein solutions.
  • Affinity chromatography: It is a liquid chromatographic method in which a biological agent or biomimetic ligand is used for the selective retention of complementary compounds. Example antibody/antigen, enzyme/substrate or enzyme/inhibitor 

  • Thin layer Chromatography: It is a “solid-liquid adsorption” chromatography. stationary phase is an adsorbent coated on silica plate on which analyte is loaded and mobile phase travels upward.
  • Paper Chromatography: It is a “liquid-liquid” chromatography. A thick filter paper with water drops settled in its pores made up the stationary “liquid phase.” Mobile phase consists of an appropriate fluid placed in a developing tank. Utilized in the separation of proteins, and in studies related to protein synthesis.
  • Electrophoresis: It is a method which is used to separate mixtures of DNA, RNA, or proteins according to their molecular size pushed by an electrical field through a gel that contains small pores. Electrophoresis is based on the phenomenon that most biomolecules exist as electrically-charged particles, possessing ionizable functional groups.
  • Quantitative Analysis of proteins and peptides: Use of Mass spectroscopy in the separation and quantification of proteins, peptides.
  • Racemic Mixtures: These are 50:50 mixtures of two enantiomers of chiral small molecules. Since enantiomers have similar physical properties, it is difficult to separate the. Methods to separate racemic mixtures are:
  • Reaction with pure chiral compound to create diastereomers. Diastereomers have different physical properties and thus can be separated easily.
  • Chiral Chromatography with a chiral stationary phase react with one of the two enantiomers in a racemic mixture.

5D Structure, function, and reactivity of biologically relevant molecules

Biological macromolecules (carbohydrates, lipids, proteins, and nucleic acids) are an important component of the cell and performs a wide array of functions. The structure and hence function of these macromolecules is governed by the type of functional groups and foundational principles of chemistry such as: covalent bonds and polarity, bond rotations and vibrations, non-covalent interactions, the hydrophobic effect and dynamic aspects of molecular structure. The content in this category covers the primary, secondary, tertiary structures of different biomolecules, their function and reactivity in various biochemical pathways.

The topic and subtopics are as follow:

Nucleic Acids and Nucleotides

  • Difference between Nucleoside and Nucleotide.
  • Nucleoside: Nitrogenous base linked covalently to a sugar(ribose or deoxyribose)
  • Nucleotide: Nucleoside attached to an additional phosphate group
  • RNA and DNA (structure, function and chemistry)
  • RNA: Ribonucleic acid: Single stranded polynucleotide chains with alternating ribose sugar and phosphate group backbone. Attached to each sugar is one of four bases: adenine (A), uracil (U), cytosine (C) or guanine (G). Involved in synthesis of proteins via translation. mRNA, rRNA, and tRNA are the three main types of RNA involved in protein synthesis.
  • DNA: Deoxyribonucleic acid: Double stranded polynucleotide chains with deoxyribose sugar phosphate group backbone. Attached to each sugar is one of four bases: adenine (A), thymine(T), cytosine (C) or guanine (G). It is a genetic information molecule that stores information to synthesis proteins. Functions of DNA include replication via transcription, encoding, cell management and the ability to mutate.

Amino Acids, Peptides, Proteins

  • Amino acids: Amino acids are the building blocks of proteins and hence functioning of cells.
  • Structure and configuration of all twenty amino acids( alpha and beta amino acids, R or S absolute configuration)
  • Classification as either lewis acid or bases
  • Classification as hydrophilic or hydrophobic
  • Synthesis of alpha amino acid( amino and carboxylic group attached to the alpha carbon)
  • Gabriel Synthesis: Alkylation of phthalimidomalonic ester followed by hydrolysis and decarboxylation to give amino acid
  • Strecker synthesis. Just three starting components are needed. Ammonia, potassium cyanide and aldehyde or ketone to give the amino acid. Much simpler and more efficient
  • Peptide and proteins: Most abundant biomolecules.
  • Reactions
  • Peptide linkage: Bond connecting two amino acids is called peptide bond. Condensation reaction of amino and carboxylic group.
  • Sulfur linkage: Disulfide bond between sulfur containing side group of two cysteine amino acid molecules. It affects the folding and stabilization of protein structures.
  • Hydrolysis: Reverse of peptide bond formation where peptide bonds are cleaved with addition of water.
  • General principles of protein structure
  • Primary, secondary, tertiary and quaternary structures of proteins
  • Isoelectronic point: pH at which the protein is electrically neutral. Determined by isoelectric focusing.
  • Three-dimensional structure of proteins
  • Denaturing and folding: Most stable structure of a protein depends on its environment, intramolecular and intermolecular forces. A protein folds to minimize its intrinsic energy and to maximize the entropy of the system. Changing elements of its environment causes denaturing of the protein and hence its function.
  • Hydrophobic interaction and solvation layer: Since proteins largely function in aqueous environment, folding of a proteins happens in a way to sequester hydrophobic amino acids towards the inside core of protein. As a result, water molecules surrounding the folded protein ( solvation layer) have favourable interactions with the exposed hydrophilic groups.
  • Non-enzymatic protein functions
  • Immunity
  • Transportation
  • Nourishment, movement and regulation
  • Structural

Small organic biomolecules


  • Lipids tend to be hydrophobic, nonpolar, and made up mostly of hydrocarbon chains. These are categorized into three main types
  • Storage Lipids: Responsible for providing energy, transporting fat soluble vitamins and insulate body organs. Consist of two components: Glycerol and fatty acid
  • Free fatty acids: Saturated and unsaturated fatty acids, trans fat, omega fatty acids
  • Triglycerides: Used in the formation of soap by the process called saponification
  • Structural Lipids: Major component of cell membranes
  • Phospholipids: Amphipathic molecule with one hydrophilic phosphate group head and two hydrophobic uncharged fatty acid tails.
  • Phosphatidate: Precursor in the biosynthesis of phospholipids.
  • Sphingolipids: Sphingolipids are a class of lipids containing a backbone of sphingoid bases, a set of aliphatic amino alcohols that includes sphingosine rather than a glycerol molecule.
  • Waxes: Contain long fatty acid chains connected to alcohols by ester linkages. However plant waxes contain long chain hydrocarbon containing functional groups.
  • Signal/Cofactor lipids: Lipid molecules serve as intracellular signals or secondary messengers, which can bind to intracellular enzymes to activate intracellular processes.
  • Fat soluble vitamins: Vitamin A, D, E, K
  • Steroids: Four fused carbon rings with a short side chain e.g cholesterol. Precursors to vitamin D, hormones, bile salts.
  • Prostaglandins: Prostaglandins are compounds in the body made of fats that have hormone-like effects. These are twenty carbon fatty acids.


Macromolecules containing carbon, oxygen and hydrogen.

  • Nomenclature and classification and common names: Aldose or ketose, number of carbon atoms in the chain.
  • Monosaccharides: Straight chain forms ( Fischer projections) undergo cyclization to create cyclic sugar molecules( Haworth projections).
  • Absolute configuration (D, L)
  • Cyclic structure and conformations of hexoses (alpha, beta)
  • Epimers and anomers
  • Keto-enol tautomerism
  • Disaccharides: Two monosaccharides chained together through glycosidic linkage.
  • Hydrolysis of glycosidic linkage
  • Polysaccharides

Aldehydes and Ketones

Nomenclature and physical properties

  • Important reactions
  • Nucleophilic addition reactions at C=O bond. Products and stereochemistry of addition product
  • Acetal, hemiacetal: Addition of 2 or 1 equivalent of alcohol group.
  • Imine and Enamine: Addition of primary and secondary amines.
  • Hydride reagent: Addition of lithium aluminum hydride and sodium borohydride.
  • Cyanohydrin: Addition of hydrogen cyanide
  • Oxidation of aldehydes: Formation of carboxylic acids
  • Reactions at adjacent positions: enolate chemistry; Kinetic vs thermodynamically stable enolate
  • Keto-enol tautomerism (α-racemization)
  • Aldol condensation, retro-aldol
  • Wolff-Kishner reaction
  • Grignard reagents
  • Organometallic reagents
  • Effect of substituents on reactivity of C=O; steric hindrance: Bulky groups on either side of C=O blocks access to the electrophilic carbon, so reactivity goes down
  • Acidity of α-H; carbanions: Alpha proton is acidic because the resulting carbanion is stabilized by resonance.


Nomenclature and physical properties

  • Important reactions
  • Oxidation
  • Substitution reactions: SN1 or SN2
  • Protection of alcohols
  • Preparation of mesylates and tosylates

Carboxylic Acids

Nomenclature and physical properties

  • Important reactions
  • Carboxyl group reactions
  • Amides (and lactam), esters (and lactone), anhydride formation
  • Reduction
  • Decarboxylation
  • Reactions at 2-position, substitution

Acid Derivatives

Nomenclature and physical properties

  • Important reactions
  • Nucleophilic substitution
  • Transesterification
  • Hydrolysis of amides
  • Relative reactivity of acid derivatives
  • Steric effects
  • Electronic effects
  • Strain (e.g., β-lactams)


Nomenclature and physical properties

  • Hydroquinones
  • Ubiquinones
  • Biological two electron redox centers

Polycyclic and heterocyclic aromatic compounds

  • Biological aromatic heterocylces: Purines, Pyrimidines
  • Huckle’s rule and delocalization in defining aromaticity

5E Principles of chemical thermodynamics and kinetics

All the reactions happening in our body, around us are dynamic in nature following the principles of thermodynamics and reaction kinetic. Thermodynamics determines the energy changes happening in a reaction and its effect on the position of chemical equilibrium. Kinetics is the study to determine the speed of reaction. Enzymes are the biological catalyst known to carry out biochemical reactions to sustain life. Without enzymes, we cannot visualize basic reactions like metabolism, anabolism, and even maintaining homeostasis. The catalysts/biocatalysts facilitate both rapid and efficient progression of a reaction under set conditions.

The content in this category covers principles of chemical thermodynamics and kinetics, including enzyme catalysts. The topics are listed as below:

Principles of Bioenergetics


  • Free energy/Keq: When the reaction attains equilibrium, certain ratio of reactants and products is obtained. This ratio is called equilibrium constant(Keq). It is ratio of molar concentration of products and reactant to the exponent of their coefficients.

aA + bB ⇌ cC + dD 

Keq = [C]c[D]d/[A]a[B]b 

From Keq , we can calculate standard energy change( ∆Go ) called gibbs free energy associated with a chemical reaction at any temperature.( molar gas constant R= 8.314Jmol-1K-1)

T is temperature in kelvin

∆Go = -RT ln Keq

Gibs free energy is the energy available in a system to do work.

  • Concentration: Le Châtelier's Principle


  • ATP hydrolysis ΔG << 0: Adenosine triphosphate is an energy powerhouse. The hydrolysis of three different phosphate group generating free phosphate group is an energy releasing process

ATP + H2O → ADP +Pi + free energy

  • ATP group transfers: Biological reactions having large positive gibbs free energy, receive a portion of ATP molecule to make the reaction thermodynamically favourable and easy to proceed.

Biological oxidation–reduction

Reactions involving electron transfers.

  • Half-reactions: Oxidation half reaction involving losing electrons. Reduction half reaction involving gaining electrons
  • Soluble electron carriers: Small organic molecules acting as electron shuttles between oxidized and reduced species. NAD+( nicotinamide adenine dinucleotide) and FAD( flavin adenine dinucleotide).
  • Flavoproteins: These are the proteins with FAD or FMN(flavin mononucleotide) occurring in eukaryotic and prokaryotic electron transfer systems

Energy Changes in Chemical reactions

Thermodynamic system

Understanding of system, surroundings. Definition of state functions and concept (∆H, ∆G, ∆S)

Laws of Thermodynamics

  • Zeroth Law: In case of thermal equilibrium, a system has unchanging state of temperature. If a =b and b=c then b=c
  • First law of thermodynamics: States the conservation of energy. Internal energy of the system is constant. Energy cannot be created or destroyed or added or taken away from a system, it only can be transferred or converted in the form of heat and work. Units of energy are joules or kilojoules. Internal energy(U) is related to heat(q) and work(w) as

∆U = q + w

  • PV diagram and work done: These are used to describe pressure volume relationship in a dynamic system. The enclosed area under the curve gives the work done in the system. Under constant pressure( atmospheric pressure) work done by system is P∆V. ∆v is the volume change. This gives:

U = q – w = q - P∆V

  • Second law of thermodynamics: Entropy: The entropy (state of disorder) of a system is constantly increasing. Unit of entropy is joules per kelvin. Entropy (S) of a system depends on the heat transferred in the system (q) and its temperature(T) in kelvin. For reversible processes ΔS = q / T.

For irreversible processes ΔS > q / T.

ΔS ≥ q / T

  • Relative entropy of gas, liquid and crystal states

Measurement of heat changes(calorimetry), heat capacity and specific heat

  • Calorimetry is the study of transferring energy as heat resulting in temperature change. Relationship between internal energy and heat ( q= thermal energy) with no other form of energy transfer is

∆U = q

Given the temperature change of the system ∆T, measured by recording initial and final temperatures, energy transfer as heat is calculated as: C = heat capacity, m = mass of substance

Q = mC∆T

  • Heat capacity(C ): the amount of heat required to raise the temperature of something by 1oC
  • Molar heat capacity: heat capacity per mole
  • Specific heat capacity: heat capacity per mass
  • It takes 4.2 J of heat energy to raise the temperature of 1 gram of water by 1 °C.
  • Some useful conversion factors: 1 calorie = 4.2 J; 1 Calorie (with capital C) = 1000 calorie = 4200 J.

Endothermic and Exothermic reactions

Enthalpy, H, and standard heats of reaction and formation

  • enthalpy or H is the heat content of a reaction. H stands for heat.
  • ΔH is the change in the heat content of a reaction. + means heat is taken up, - means heat is released.
  • Enthalpy changes can be measured by calculating the Standard heat of reaction, ΔHrxn, is the change in heat content for any reaction. Standard enthalpy changes refer to reactions done under standard conditions, and with everything present in their standard states. The standard state is where things are in their natural, lowest energy, state. For example, oxygen is O2 (diatomic gas) and carbon is C (solid graphite).

ΔHorxn = ∑ ΔHo products - ∑ΔHoreactants

  • Standard heat of formation, ΔHf, is the change in heat content a formation reaction.
  • Hess’ Law of Heat Summation: If a reaction involves multiple intermediate steps, then the enthalpy of the final reaction can simply be calculated using the sum of standard enthalpies of the intermediate reactions.

Bond dissociation energy as related to heats of formation

  • Bond dissociation is the energy required to break bonds.
  • ΔHrxn = Bond dissociation energy of all the bonds in reactants - bond dissociation energy of all the bonds in products.
  • ΔHrxn = Enthalpy of formation of all the bonds in products - Enthalpy of formation of all the bonds in reactants.
  • Bond dissociation energy is positive because energy input is required to break bonds.
  • The enthalpy of formation of bonds is negative because energy is released when bonds form.

Free energy

  • Free energy also known as Gibbs free energy(G) is the energy available that can be converted to do work. H = enthalpy, S = entropy, T is temperature in Kelvin.


  • Standard free energy of a reaction is given as

ΔGorxn = ∑ ΔGo products - ∑ΔGoreactants

  • Spontaneous reactions and ΔG°
  • Coefficient of expansion
  • Heat of fusion(Hfus), heat of vaporization(Hvap)

ΔHfus = -ΔHsolid   ΔHvap = -ΔHcondensation

ΔHfus + ΔHvap = ΔHsubstance

  • Phase diagram: pressure and temperature

Rate Processes in Chemical Reactions - Kinetics and Equilibrium

  • Reaction Rates: It tells the speed of the reaction. It is measured either by decrease in the concentration of reactant or increase in concentration of products over time. For a reaction the rate is measured as:

A + 2B = 3C

-d[A]/dt = -1/2 d[B]/dt = +1/3 d[C]/dt

[A]. [B], [C] are the respective molar concentrations of reactants and products

  • Dependence of reaction rate on concentration of reactants: As the concentration of reactant molecules is increased, the number of molecules with minimum required energy ( Activation energy) increase.
  • Rate law( r): Equation that relates reaction rate with the concentration or partial pressure of the reactants.

aA + bB → cC 

r = k[A]x[B]y

k = rate constant

[A], [B] = molar concentration of reactants A and B

x, y = vary for each reaction and are determined experimentally

  • Reaction order: Tells the relation between concentration of reactants and reaction rate.
  • Order of reaction is the sum of exponent x and y and is determined experimentally.
  • Rate determining step: Some reactions happen in multiple steps with multiple intermediates. However, the rate of a multistep reaction depends on the slowest step also known as rate determining elementary step.
  • Dependence of reaction rate upon temperature:
  • Interpretation of energy profiles showing energies of reactants, products, activation energy, and ΔH for the reaction
  • Arrhenius Equation relates the rate of a chemical equation to its activation energy, which in turn can be obtained from the heat energy  from the surroundings. This equation can be used to understand how the rate of a chemical reaction depends on temperature

k = A e -Ea/RT

A = frequency or pre-exponential factor

e^(-Ea/RT) = fraction of collisions that have enough energy to overcome the activation barrier

Ea = activation energy

R = universal gas constant

T = temperature

Kinetic control versus thermodynamic control of a reaction


  • Homogeneous and heterogeneous catalysts
  • Type of catalysis
  • Covalent catalysis
  • Acid-base catalysis
  • Metal ion catalysis
  • Approximation

Equilibrium in reversible chemical reactions

  • Law of mass action: Rate of any chemical reaction is proportional to the product of the masses of the reacting substances, with each mass raised to a power equal to the coefficient that occurs in the chemical equation. In a chemical equilibrium the rate of forward reaction is same as the rate of backward reaction

aA + bB ⇌ cC + dD 

rate of forward = rate of backward

k1[A]a[B]b = k2[C]c[D]d

k1/k2 = Keq = [A]a[B]b/ [C]c[D]d

k1 and k2 are rate constants

Keq is the equilibrium constant


  • Le Châtelier’s Principle: Concept in an equilibrium reaction and its application
  • Relationship of the equilibrium constant and ΔG°

ΔG° = -RTlnKeq

ΔG° = standard Gibbs free energy change

R = universal gas constant

T = temperature

Keq = equilibrium constant


  • Classification by the type of reaction they catalyze
  • Mechanism
  • Substrates and enzyme specificity
  • Active site model
  • Induced-fit model
  • Cofactors, coenzymes, and vitamins
  • Kinetics
  • General (catalysis)
  • Michaelis–Menten: Kinetic model for examining the reaction of substrates and enzymes. Michaelis-Menten plot is obtained under two specific conditions:

Enzyme concentration is held constant

Substrate concentration is increased

Reaction velocity is plotted Vs substrate concentration

Vmax = [E] * kcat

Where [E] = concentration of enzyme in the experiment

kcat = turnover number of a single enzyme

vmax as the maximum speed at which the reaction can proceed. The Michael Menten equation is given as :

V0 = (Vmax x [S]) / (KM + [S])

V0 = reaction velocity,how fast products are being formed.

Vmax describes the maximum reaction velocity, and

[S] is the substrate concentration.

Km is the Michaelis constant. It is equal to the substrate concentration when the reaction velocity (V) is equal to half the maximum reaction velocity, or half Vmax.

  • Cooperativity
  • Effects of local conditions on enzyme activity
  • Inhibition: Use of Lineweaver-Burk plot (1/[S] vs 1/Vo). This plot is used to identify the type of enzyme inhibition
  • Competitive inhibition
  • Uncompetitive inhibition
  • Mixed and non-competitive inhibition
  • Regulatory enzymes
  • Allosteric
  • Covalently modified


1. How long is the chemistry and physics section on MCAT?

The Chemical and Physical Foundations of Biological Systems is the first section on the exam. You are given 95 minutes to complete 59 questions combined. The chemical/biochemical foundational concepts of the section cover approximately 60 percent of this section. The questions are combination of both passage-based and discrete questions.

2. What are some techniques to study, memorize and apply the chemistry concepts and equations?

Chemistry and biochemistry are two branches of science which involve not only memorizing formulas, structures, and reactions but understanding the rationale behind them. With that said, how can you grasp the concepts strongly? Here are few tips you can follow:

  1. Reading alone is not sufficient. Write along. Information passed through hands is retained for longer time. Do not hesitate to write notes.
  2. Have you ever wondered why visuals stay in our memory than written text? How is this applicable here? Animations are the answers. Various concepts ( e.g. phase diagrams, gas laws, intermolecular forces, bond formation , bond breaking, chemical changes etc.) are very well depicted as animations. Try to visualize the process happening and it can be understood easily. There are many free resources which include scientific animations, and you can make use of them if you are finding it hard to understand by just reading.
  3. Apply your knowledge. Do practice problems, quizzes, and tests. Break down the problem into smaller problems and try to go step by step. For e.g. finding the molecular formula of an unknown powder given the percent mass composition (A,B,C). How can you approach it? You know that mass number must be converted to some integers x, y, z to look something like empirical formula (AxByCz). How can you achieve it? By converting mass to moles for each atom and then to molecular formula.

Remember it is not about how much you know, it is about how much can you apply. 

3. Will I be provided with a periodic table/calculator in the MCAT?

During the exam, you are provided with the periodic table. However, you are expected to do calculations without the help of calculator. So, it is advised that during your preparation you are doing all the calculations related to the concepts without the aid of a calculator. That way you give yourself enough practice to efficiently and correctly solve numerical problems.

4. How early can I start preparing for the section?

When to start studying for the MCAT will often depend on how much time you have before your MCAT test date. If you are very confident with your content knowledge, then 2-3 months are good enough to practice and test yourself. Only knowing the content is not sufficient. You should be able to apply it under stressful conditions in a restricted timeframe. What does that mean? This means you test yourself with actual practice exams in one sitting as if you are appearing in the actual test. However, if you are not very confident with your content knowledge then it is recommended to give yourself at least 4-6 months. This includes reviewing the content covered under the section along with practice exams, practice question banks etc. Remember to create a thorough MCAT study schedule to help you prepare.

5. What is the hardest section of the MCAT?

Often, students find the MCAT CARS section the most challenging because it’s more difficult to prepare for. Since one cannot really prepare by memorizing equations or concepts, CARS can be a challenging section to prepare for. But don’t worry, with the right MCAT CARS strategy you will be able to ace this section!

To your success,

Your friends at BeMo

BeMo Academic Consulting

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