Food Science & Technology · Foundations · also General Chemistry, Organic Chemistry, Physical Chemistry, GATE XL-P Chemistry
The chemistry a food scientist uses every day: bonding, acids and bases, buffers, organic functional groups and reaction mechanisms, chemical kinetics and thermodynamics, and the physical chemistry of solutions, surfaces and colloids. Students who take GATE XL (Life Sciences) sit a compulsory chemistry section alongside the Food Technology section.
In the exams: GATE XL-P (compulsory chemistry section for XL candidates who choose XL-U Food Technology); background for ICAR AIEEA-PG Code 18 Unit I
Every food is a chemical system. Bread rises because yeast converts sugars into carbon dioxide and ethanol; jam sets because acid and sugar let pectin chains link up; milk curdles when acid drives casein to its isoelectric point; an apple browns because an enzyme oxidises phenols; oil turns rancid by a free-radical chain reaction. To understand, control and improve these changes a food scientist needs the general, organic and physical chemistry that a chemistry degree teaches in its first two years, applied to food.
This subject covers that foundation: atomic structure and bonding, stoichiometry and the preparation of laboratory solutions, acids, bases and buffers, chemical equilibrium and solubility, thermodynamics and kinetics, colligative properties, organic functional groups and nomenclature, the organic reactions most important in food (oxidation, hydrolysis, esterification and condensation), stereochemistry and optical activity, surface chemistry and colloids, electrochemistry and redox potential, and the coordination chemistry of metal ions in food. Food chemistry (the major and minor constituents) and food analysis build directly on it.
Chemistry is examined worldwide at every stage of a food science career. Life-science graduate entrance examinations usually carry a compulsory chemistry section for candidates who choose food technology; university entrance tests for food science degrees examine school chemistry; postgraduate admission tests in food science assume it throughout; and food analyst examinations test solution preparation, titrations, pH and the calculations of a quality-control laboratory. The questions with this manual follow those patterns, with full working for every calculation.
An atom has a nucleus of protons and neutrons surrounded by electrons. The atomic number Z is the number of protons and defines the element; the mass number is protons plus neutrons. Isotopes have the same Z but different numbers of neutrons: carbon has the stable isotopes ¹²C and ¹³C and the radioactive ¹⁴C. Because plants that fix carbon by the C3 and C4 pathways take up ¹³C in slightly different proportions, isotope-ratio mass spectrometry can detect, for example, cane or corn sugar added to fruit juice or honey; this is a standard authenticity test.
Electrons occupy orbitals described by four quantum numbers: the principal number n (shell), the azimuthal number l (subshell s, p, d, f), the magnetic number and the spin. Orbitals fill in order of increasing energy (the Aufbau principle), no two electrons in an atom share all four quantum numbers (the Pauli principle), and electrons occupy degenerate orbitals singly before pairing (Hund's rule). Sodium is 1s² 2s² 2p⁶ 3s¹; iron is [Ar] 3d⁶ 4s², and the Fe²⁺ and Fe³⁺ ions of haem pigments lose the 4s electrons first.
The periodic table arranges elements by Z so that elements in a group share outer-shell configurations and chemistry. Across a period, atomic radius decreases and ionisation energy and electronegativity increase; down a group, radius increases and ionisation energy decreases. On the Pauling scale fluorine is the most electronegative element (4.0), followed by oxygen (3.5), nitrogen (3.0), and carbon (2.5), with hydrogen at 2.1. These differences explain why O–H and N–H bonds are polar and why water, sugars and proteins form hydrogen bonds.
Ionic bonds form when electrons transfer from a metal to a non-metal, giving ions held by electrostatic attraction, as in sodium chloride. Covalent bonds share electron pairs; when the atoms differ in electronegativity the sharing is unequal and the bond is polar. A coordinate (dative) bond is a covalent bond in which both electrons come from one atom, as in metal-ligand complexes (chapter 12). Carbon forms four bonds using hybrid orbitals: sp³ carbon is tetrahedral with bond angles of about 109.5° (the saturated chains of fatty acids), sp² carbon is trigonal planar at 120° (C=C double bonds and carbonyl groups), and sp carbon is linear at 180°. Valence-shell electron-pair repulsion (VSEPR) predicts shapes from the electron pairs around the central atom: water is bent, with an H–O–H angle of about 104.5°, and ammonia is pyramidal.
Between molecules, weaker forces decide the physical behaviour of foods. Hydrogen bonds form between an H atom bonded to O or N and a lone pair on another O or N; they give water its high boiling point, high specific heat (about 4.18 kJ/(kg·K)) and high latent heats, and they hold together starch double helices, cellulose fibrils and protein secondary structure. Dipole-dipole forces act between polar molecules; London dispersion forces act between all molecules and dominate in non-polar substances, so the melting point of saturated fatty acids rises with chain length. Ionic interactions form salt bridges in proteins. The hydrophobic effect, the tendency of non-polar groups to cluster in water, is driven mainly by the entropy of water and drives protein folding, micelle formation and the adsorption of emulsifiers at oil-water interfaces. Knowing which force dominates explains why salt dissolves in water but not in oil, why fats crystallise, and why heating disrupts gels held by hydrogen bonds.
The mole is the amount of substance containing 6.022 × 10²³ entities (Avogadro's number); the molar mass in g/mol equals the relative formula mass. A balanced equation gives the mole ratios in which substances react, and the reagent that runs out first, the limiting reagent, sets the theoretical yield. In alcoholic fermentation, C₆H₁₂O₆ → 2 C₂H₅OH + 2 CO₂, so 180 g of glucose gives at most 92 g of ethanol and 88 g of carbon dioxide, a theoretical yield of 0.511 g ethanol per gram of glucose; real yields are lower because yeast uses some sugar for growth and by-products.
Food laboratories use many concentration units, and confusing them causes most calculation errors. Percentages may be w/w (grams per 100 g), w/v (grams per 100 mL) or v/v (millilitres per 100 mL, used for alcohol). Molarity (M) is moles of solute per litre of solution; molality (m) is moles per kilogram of solvent and does not change with temperature, which is why colligative calculations use it. Normality (N) is equivalents per litre, where an equivalent is the amount that supplies or reacts with one mole of H⁺ or electrons, so 1 M sulphuric acid is 2 N. Mole fraction is the moles of one component divided by total moles. Trace levels are expressed in parts per million (1 ppm = 1 mg/kg or, for dilute aqueous solutions, 1 mg/L) and parts per billion (1 µg/kg). Degrees Brix approximate the percentage of soluble solids by mass, measured as if they were sucrose, by refractometer or hydrometer.
To prepare a solution of known molarity, dissolve the calculated mass, molarity × volume × molar mass, and make up to the mark in a volumetric flask: 0.1 M sodium hydroxide needs 4.0 g per litre, and 0.1 M sodium chloride 5.844 g per litre. Dilutions follow C₁V₁ = C₂V₂, so 10 mL of a 1 M stock made up to 100 mL gives 0.1 M. A primary standard is a substance of high purity, stable in air, non-hygroscopic and of high molar mass, whose solution can be made by direct weighing; potassium hydrogen phthalate (for bases), anhydrous sodium carbonate (for acids) and potassium dichromate (for redox) are examples. Sodium hydroxide is not a primary standard, because it absorbs water and carbon dioxide, so its solution is standardised by titration against potassium hydrogen phthalate.
Titratable acidity is one of the most common food tests. A sample is titrated with standard alkali to a phenolphthalein end point (or to pH 8.1 to 8.2 with a meter for coloured samples) and the acidity is expressed as the predominant acid: acidity (%) = V × N × meq. wt × 100 / sample mass, where V is the titre in mL, N the normality of the alkali and the milliequivalent weight is the acid's equivalent weight divided by 1000. Standard milliequivalent factors are 0.064 for anhydrous citric acid (fruit juices), 0.090 for lactic acid (milk and fermented products), 0.067 for malic acid (apples), 0.075 for tartaric acid (grapes) and 0.060 for acetic acid (vinegar). For example, 10 g of juice needing 12.0 mL of 0.1 N sodium hydroxide has 12.0 × 0.1 × 0.064 × 100/10 = 0.768 per cent acidity as citric acid. The same pattern of calculation, titre × normality × equivalent weight, underlies the determination of salt, vitamin C, peroxide value and many other analyses.
In the Arrhenius definition an acid releases H⁺ in water and a base releases OH⁻. The Brønsted-Lowry definition is broader: an acid donates a proton and a base accepts one, so every acid has a conjugate base. The Lewis definition broadens it again: a Lewis acid accepts an electron pair (metal ions, for example) and a Lewis base donates one. Water ionises slightly, and at 25 °C its ion product Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴, so pH + pOH = 14 and neutral water has pH 7; Kw increases with temperature, so neutral pH is below 7 in hot water. pH is defined as −log₁₀[H⁺] (strictly the hydrogen-ion activity): a solution with [H⁺] = 10⁻³ M has pH 3, and each pH unit is a tenfold change in acidity.
Strong acids such as hydrochloric acid ionise completely, so 0.01 M HCl has pH 2. Weak acids ionise partially, described by the acid dissociation constant Ka and pKa = −log Ka; the smaller the pKa, the stronger the acid. For a weak acid alone in water, [H⁺] ≈ √(Ka·C) when the ionisation is small, so 0.1 M acetic acid (pKa 4.76) has pH about 2.9. Polyprotic acids lose protons in steps with successive pKa values: citric acid about 3.1, 4.8 and 6.4; phosphoric acid about 2.1, 7.2 and 12.3; carbonic acid about 6.4 and 10.3. Amino acids are amphoteric, carrying both an acidic carboxyl and a basic amino group; at the isoelectric point a molecule carries no net charge.
pH matters in food safety and quality. Foods with a pH of 4.6 or below are classed as acid foods for thermal processing, because spores of Clostridium botulinum do not germinate and produce toxin at or below that pH, so such foods can be processed at lower temperatures; low-acid foods above 4.6 need a full botulinum process. Weak-acid preservatives such as benzoic acid (pKa about 4.2) and sorbic acid (pKa about 4.8) act mainly in the undissociated form, which crosses microbial membranes, so they are far more effective at low pH.
A buffer resists changes in pH when acid or base is added; it contains a weak acid and its conjugate base (or a weak base and its conjugate acid) in comparable amounts. The Henderson-Hasselbalch equation, pH = pKa + log([A⁻]/[HA]), gives the pH of a buffer from the ratio of base to acid forms. When the two are equal, pH = pKa; a tenfold excess of the base form raises the pH by one unit. Buffer capacity is greatest at pH = pKa and the useful range is about pKa ± 1. For example, an acetate buffer made with 0.2 mol of acetate and 0.1 mol of acetic acid has pH = 4.76 + log 2 = 5.06. The equation also gives the fraction ionised at any pH: at pH 3.2, two units below its pKa of 5.2, about 99 per cent of a weak acid is undissociated.
Foods are buffered by their own constituents. Milk is buffered by phosphates, citrates, casein and dissolved carbon dioxide, which is why lactic acid bacteria must produce considerable acid before the pH falls far enough to coagulate casein near its isoelectric point of about 4.6. Meat is buffered by proteins, phosphates and dipeptides; fruit juices by organic acids and their salts. Food formulators use buffer salts such as sodium citrate and phosphates to hold pH in processed cheese, beverages and baking powders, and acidulants (citric, malic, lactic, phosphoric acids) to set it. Acid-base indicators change colour over a pH range of about two units around their own pKa: phenolphthalein from colourless to pink between about 8.2 and 10, methyl orange from red to yellow between about 3.1 and 4.4. An indicator is chosen so that its range falls on the steep part of the titration curve: phenolphthalein for weak acid against strong base, methyl orange for strong acid against weak base.
A reversible reaction reaches equilibrium when forward and reverse rates are equal; the concentrations then stop changing, although both reactions continue. For aA + bB ⇌ cC + dD, the equilibrium constant is K = [C]ᶜ[D]ᵈ/([A]ᵃ[B]ᵇ), with pure solids and the solvent left out. A large K means products are favoured. The reaction quotient Q has the same form at any moment; if Q < K the reaction proceeds forward, if Q > K it runs backward. K is related to the standard free-energy change by ΔG° = −RT ln K, so a reaction with ΔG° = 0 has K = 1. K changes only with temperature: for an exothermic reaction K falls as temperature rises.
Le Chatelier's principle states that a system at equilibrium responds to a disturbance in the direction that partly offsets it. Adding a reactant or removing a product drives the reaction forward; this is why esterification yields more ester when water is removed, and why carbon dioxide escapes from an opened carbonated drink. Raising pressure favours the side with fewer gas molecules. Raising temperature favours the endothermic direction. In foods, the equilibria of calcium between soluble and colloidal phosphate in milk shift with temperature and pH: heating moves calcium phosphate into the casein micelles, cooling and acidification move it out.
For a sparingly soluble salt, the solubility product Ksp is the product of ion concentrations at saturation, each raised to its stoichiometric coefficient. For silver chloride, Ksp = [Ag⁺][Cl⁻] ≈ 1.8 × 10⁻¹⁰ at 25 °C, so its solubility is √Ksp ≈ 1.3 × 10⁻⁵ M; for a salt of type AB₂, Ksp = 4s³. A precipitate forms when the ion product exceeds Ksp. Adding a common ion reduces solubility (the common-ion effect). The Mohr titration for salt in foods uses these principles: chloride is titrated with silver nitrate, silver chloride precipitates first, and the first excess of silver forms red-brown silver chromate with the chromate indicator; the Volhard method back-titrates excess silver with thiocyanate in acid solution.
Solubility equilibria matter in many products. Potassium hydrogen tartrate crystallises in wine on cooling, so wines are cold-stabilised before bottling; calcium oxalate limits the availability of calcium from some leafy vegetables; struvite crystals (magnesium ammonium phosphate) can form in canned seafood; lactose, the least soluble common sugar, crystallises in sweetened condensed milk and ice cream, causing sandiness when crystals grow too large. Sugar solubility rises with temperature, which is used in crystallisation and confectionery.
Gas solubility follows Henry's law: the concentration of dissolved gas is proportional to its partial pressure above the liquid, c = k_H·p. Carbonated beverages are filled cold and under pressure because carbon dioxide is more soluble at low temperature and high pressure. Oxygen solubility also falls as temperature rises, which affects fermentation aeration and oxidation in stored liquids. The partition coefficient, the ratio of a solute's concentrations in two immiscible phases at equilibrium (often expressed as log P between octanol and water), decides how flavours and preservatives distribute between the fat and water phases of a food.
Thermodynamics deals with energy changes and the direction of change. The first law states that energy is conserved: the change in internal energy of a system is ΔU = q + w, the heat added to the system plus the work done on it (some engineering texts use the opposite sign for work). At constant pressure, the heat exchanged equals the enthalpy change, ΔH = ΔU + PΔV. Reactions that release heat are exothermic (ΔH < 0) and those that absorb it endothermic (ΔH > 0). Because enthalpy is a state function, Hess's law allows the enthalpy of a reaction to be obtained by adding the enthalpies of steps, or from standard enthalpies of formation: ΔH° = ΣΔHf°(products) − ΣΔHf°(reactants).
The energy in food is measured by calorimetry. A bomb calorimeter burns a sample in oxygen at constant volume and gives its gross energy; the body does not extract all of this energy, because some of it is lost in faeces and, for protein, in urea. The Atwater general factors allow for these losses: about 17 kJ (4 kcal) per gram of protein and of available carbohydrate, 37 kJ (9 kcal) per gram of fat and 29 kJ (7 kcal) per gram of ethanol, with lower factors for dietary fibre and polyols in many labelling rules. Sensible heat follows q = m·c·ΔT; latent heats are the large enthalpy changes of phase transitions: about 334 kJ/kg to melt ice and about 2257 kJ/kg to evaporate water at 100 °C. These two values dominate the energy bills of freezing and of drying and evaporation.
The second law states that the entropy (S) of an isolated system never decreases; entropy measures how many ways the energy and matter of a system can be arranged. A process is spontaneous at constant temperature and pressure when the Gibbs free energy change, ΔG = ΔH − TΔS, is negative. A reaction with ΔH < 0 and ΔS > 0 is spontaneous at all temperatures; one with ΔH > 0 and ΔS > 0 becomes spontaneous above T = ΔH/ΔS. Melting of ice is the example: ΔH and ΔS are both positive, and ΔG changes sign at 0 °C. A negative ΔG says nothing about speed: sucrose and oxygen are thermodynamically unstable together, yet sugar is stable in air because the reaction is very slow without a catalyst or a flame.
Food processes are full of entropy-driven changes. Protein denaturation on heating has large positive ΔH and ΔS: the unfolded chain has many more conformations, so above the denaturation temperature the TΔS term wins. The hydrophobic effect is entropy-driven, because water molecules ordered around non-polar groups are released when those groups cluster. Starch gelatinisation, fat melting and the dissolution of sugar are endothermic processes that proceed because entropy increases.
Thermodynamics also defines water activity. The chemical potential of water in a food is μ = μ° + RT ln a_w, where a_w, the water activity, equals the ratio of the vapour pressure of water over the food to that of pure water at the same temperature (strictly, the ratio of fugacities). Water moves from regions of high to low chemical potential, so a moist filling loses water to a dry crust, and microbes cannot draw water from a food whose a_w is below their limit. The energy needed for drying rises sharply at low moisture because the remaining water is held with a binding energy above the latent heat of free water.
Chemical kinetics studies how fast reactions go. The rate law expresses rate as a function of concentrations, rate = k[A]ᵐ[B]ⁿ, where k is the rate constant and the exponents m and n, the orders with respect to each reactant, are found by experiment, not from the balanced equation; the overall order is m + n. Molecularity, the number of molecules in an elementary step, applies only to single steps of a mechanism. The units of k reveal the order: concentration per time for zero order, per time for first order, and per concentration per time for second order.
The integrated rate laws give concentration against time. Zero order: C = C₀ − kt, with half-life C₀/(2k), which depends on the starting amount. First order: ln C = ln C₀ − kt, with half-life 0.693/k, independent of the starting amount. Second order (one reactant): 1/C = 1/C₀ + kt, with half-life 1/(kC₀). The order is identified by which plot is linear. When one reactant is in large excess, its concentration barely changes and the reaction behaves as pseudo-first order: the acid-catalysed hydrolysis of sucrose in water to glucose and fructose, the inversion used to make invert syrup, is the classic example. Many quality changes in food are modelled with zero-order kinetics (some non-enzymatic browning, some lipid-oxidation indicators) or first-order kinetics (loss of vitamins, pigments and microbial populations).
Collision theory explains that molecules react only when they collide with enough energy, the activation energy Ea, and in a suitable orientation. Transition-state theory describes the high-energy arrangement at the top of the energy barrier. Because the fraction of molecules with energy above Ea grows exponentially with temperature, rate constants follow the Arrhenius equation, k = A·exp(−Ea/RT), where A is the pre-exponential (frequency) factor, R = 8.314 J/(mol·K) and T is the absolute temperature in kelvin. A plot of ln k against 1/T is a straight line of slope −Ea/R. Between two temperatures, ln(k₂/k₁) = (Ea/R)(1/T₁ − 1/T₂). A reaction with Ea of about 50 kJ/mol roughly doubles its rate between 25 and 35 °C; reactions with higher Ea are more sensitive to temperature. The temperature quotient Q₁₀, the ratio of rates 10 °C apart, is used in shelf-life work; values of 2 to 3 are common for chemical deterioration.
The difference in activation energies explains why high-temperature short-time processing works. Microbial spore inactivation has a very high apparent activation energy (often well above 200 kJ/mol), whereas most vitamin and colour losses have lower ones, so raising the temperature speeds microbial death more than quality loss; a shorter time at a higher temperature therefore achieves the same sterility with better quality. A catalyst speeds a reaction by providing a path with lower activation energy, without changing the equilibrium position or ΔG. Enzymes are biological catalysts of great specificity; metal ions such as copper and iron catalyse lipid oxidation; acids catalyse sucrose inversion and the hydrolysis of glycosidic bonds. Inhibitors, such as chelating agents that bind pro-oxidant metals, slow reactions by removing the catalyst.
Colligative properties depend on the number of solute particles in a solution, not on their identity. Raoult's law states that the vapour pressure of the solvent over an ideal solution is p = x·p₀, where x is the mole fraction of solvent and p₀ its vapour pressure when pure. For water in a food, x is an ideal estimate of water activity: a solution of 1 mol of sucrose in 1 kg of water (55.5 mol) has x = 55.5/56.5 = 0.982. Real food solutions deviate from ideal behaviour, especially when concentrated, so equations with correction terms (such as the Norrish equation) or measured values are used in practice.
Because the solute lowers the vapour pressure, the solution boils at a higher temperature and freezes at a lower temperature than the pure solvent. For dilute solutions, ΔT_b = i·K_b·m and ΔT_f = i·K_f·m, where m is the molality, K_b and K_f are constants of the solvent (for water 0.512 and 1.86 K·kg/mol) and i, the van 't Hoff factor, is the number of particles each formula unit gives: 1 for sucrose and glucose, close to 2 for sodium chloride and close to 3 for calcium chloride (somewhat less in practice because of ion pairing). So 1 molal sucrose freezes at about −1.86 °C and 0.5 molal sodium chloride also at about −1.86 °C. Molar mass can be found from the freezing-point depression of a known mass of solute: M = K_f·w·1000/(ΔT_f·W), with w grams of solute in W grams of solvent.
Osmotic pressure is the pressure that must be applied to a solution to stop the flow of pure solvent into it through a semipermeable membrane. For dilute solutions, the van 't Hoff equation gives π = i·C·R·T, with C in mol/m³ (or mol/L with R = 0.08206 L·atm/(mol·K)). A 0.1 M sucrose solution at 25 °C has an osmotic pressure of about 0.1 × 0.08206 × 298 = 2.4 atm, roughly 0.25 MPa. Sea water and concentrated juices have osmotic pressures of tens of atmospheres, which reverse-osmosis membranes must exceed to push water out.
Colligative effects are used and measured across food technology. Ice cream does not freeze solid at freezer temperatures because the sugars and milk salts depress the freezing point, and the formulation of sweeteners controls how much water is frozen at serving temperature, and so the scoopability. Freezing-point measurement (cryoscopy) of milk detects added water, because the freezing point of genuine milk falls within a narrow range a little above −0.5 °C (the exact reference value is set by each authority and varies slightly by region and season), and added water brings it closer to 0 °C. Boiling-point elevation reduces the temperature difference available in evaporators, so a concentrated juice boils several degrees above pure water at the same pressure. Osmotic dehydration of fruit in sugar syrup, the preservation of jams and syrups, salting and brining, and the plasmolysis of microbial cells in high-sugar or high-salt foods all depend on the osmotic pressure difference across cell membranes. Electrolytes such as salt are particularly effective per gram because they dissociate into more particles and have low molar mass.
Organic chemistry is the chemistry of carbon compounds, and the reactivity of an organic molecule is set by its functional groups. Hydrocarbons contain only carbon and hydrogen: alkanes are saturated, alkenes contain C=C double bonds, alkynes C≡C triple bonds, and aromatic compounds contain a benzene ring with delocalised electrons. Terpenes, built from five-carbon isoprene units, are the hydrocarbons of citrus and herb aromas; limonene is a monoterpene, and carotenoids are tetraterpenes.
Oxygen-containing groups dominate food chemistry. Alcohols carry –OH on a saturated carbon (ethanol, glycerol, the sugar alcohols); phenols carry –OH on an aromatic ring (the polyphenols of tea, wine and fruit, and the antioxidants tocopherol and butylated hydroxyanisole). Ethers contain C–O–C. Aldehydes contain –CHO and ketones C=O between two carbons; reducing sugars contain one or the other in their open-chain form, and aldehydes such as hexanal are markers of lipid oxidation while diacetyl (butane-2,3-dione) gives the aroma of butter. Carboxylic acids contain –COOH (acetic, lactic, citric acids and the fatty acids), and esters –COO– formed from an acid and an alcohol (the fruity flavours, such as ethyl acetate and isoamyl acetate, and triacylglycerols). Nitrogen groups include amines (–NH₂, as in amino acids and the biogenic amines histamine and tyramine) and amides (–CONH–, the peptide bond). Sulphur groups include thiols (–SH, as in cysteine) and sulphides and disulphides, which give the aromas of onion, garlic and cooked meat and the disulphide cross-links of gluten.
The International Union of Pure and Applied Chemistry (IUPAC) system names a compound from its longest carbon chain containing the main functional group, numbered to give that group the lowest locant, with substituents as prefixes and the main group as a suffix. Priority for the suffix runs approximately: carboxylic acid > ester > amide > nitrile > aldehyde > ketone > alcohol > amine, with lower-priority groups named as prefixes (hydroxy-, oxo-, amino-). Chain lengths are named meth-, eth-, prop-, but-, pent-, hex- and so on; -ane, -ene and -yne show single, double and triple bonds, with the double bond's position given by a locant.
Many food compounds have trivial names that every food scientist uses alongside the systematic ones: acetic acid is ethanoic acid; lactic acid is 2-hydroxypropanoic acid; citric acid is 2-hydroxypropane-1,2,3-tricarboxylic acid; acetaldehyde is ethanal; acetone is propanone; vanillin is 4-hydroxy-3-methoxybenzaldehyde. Fatty acids are named by chain length and double bonds: stearic acid is octadecanoic acid (C18:0); oleic acid is cis-9-octadecenoic acid (C18:1), with the double bond nine carbons from the carboxyl end; linoleic acid (C18:2) and α-linolenic acid (C18:3) are the essential fatty acids. The omega (n−) system counts from the methyl end instead, so linoleic acid is an n−6 and α-linolenic acid an n−3 fatty acid. The functional groups present decide a molecule's polarity, and therefore its solubility, volatility and flavour release: short-chain acids and alcohols mix with water, while long hydrocarbon chains make fats insoluble in it.
Oxidation in organic chemistry usually means gaining oxygen or losing hydrogen. Primary alcohols oxidise to aldehydes and then to carboxylic acids (ethanol to acetaldehyde to acetic acid, the basis of vinegar making and of wine spoilage), and secondary alcohols oxidise to ketones. The most important oxidation in food is the autoxidation of unsaturated lipids, a free-radical chain reaction. In initiation, a hydrogen atom is removed from a carbon next to a double bond, most easily from the methylene group between two double bonds in linoleic and linolenic acids, forming a lipid radical; light, heat and metal ions promote this step. In propagation, the radical adds oxygen to form a peroxyl radical, which takes a hydrogen from another lipid molecule, forming a hydroperoxide and a new radical, so one initiation can oxidise many molecules. In termination, radicals combine to form non-radical products. Hydroperoxides break down into aldehydes, ketones and alcohols with rancid odours. Phenolic antioxidants interrupt propagation by donating a hydrogen atom and forming a stable radical; singlet-oxygen oxidation, sensitised by chlorophyll or riboflavin in light, is a separate, faster mechanism.
Hydrolysis is the cleavage of a bond by water. Esters hydrolyse to acids and alcohols; in acid the reaction is reversible, but with alkali it goes to completion as saponification, forming carboxylate salts (soaps), the basis of the saponification value of fats. Enzymatic lipolysis releases free fatty acids, causing hydrolytic rancidity in butter and milled rice bran. Glycosidic bonds hydrolyse to give sugars (starch to glucose, sucrose to invert sugar), and peptide bonds to give peptides and amino acids, as in protein hydrolysates and soy sauce.
Esterification combines a carboxylic acid and an alcohol into an ester and water. Fischer esterification is acid-catalysed and reversible, so yields rise when one reactant is in excess or water is removed. Esters form during fermentation and ageing of wines, spirits and fruit products and contribute much of their fruity aroma. In the fat industry, interesterification rearranges fatty acids on the glycerol backbone, with a chemical catalyst or a lipase, to change melting behaviour without creating trans fatty acids. Catalytic hydrogenation, by contrast, is an addition reaction: hydrogen adds across C=C bonds on a nickel catalyst, raising the melting point, and partial hydrogenation also isomerises some cis bonds to trans, which is why trans fats are now restricted in many countries.
A condensation reaction joins two molecules with the loss of a small molecule, usually water. Peptide bonds, glycosidic bonds and ester bonds all form by condensation, and their hydrolysis is the reverse. The aldol condensation joins two carbonyl compounds and occurs in the later stages of browning. The Maillard reaction begins with a condensation: the carbonyl group of a reducing sugar reacts with an amino group of an amino acid, peptide or protein to give a glycosylamine (via a Schiff base), which rearranges to an Amadori compound (from aldoses) or a Heyns compound (from ketoses). These intermediates then fragment and react further, including through Strecker degradation of amino acids by dicarbonyls, which yields aldehydes with malty, roasted aromas, and finally polymerise into brown melanoidins. The reaction is fastest at intermediate water activity and alkaline pH, and it also consumes lysine, reducing protein quality, and can form acrylamide from asparagine and reducing sugars in fried and baked starchy foods. The underlying mechanism of all these carbonyl reactions is nucleophilic addition: a nucleophile such as an amine or an alcohol attacks the electron-poor carbon of a C=O group.
Isomers have the same molecular formula but different arrangements. Structural isomers differ in connectivity (glucose and fructose are both C₆H₁₂O₆, one an aldose, the other a ketose). Stereoisomers have the same connectivity but differ in spatial arrangement. Geometric (cis-trans) isomers arise around a double bond or a ring: in cis-fatty acids the chain bends at the double bond, which lowers the melting point, whereas trans isomers are straighter and pack like saturated chains. The E/Z system names geometric isomers using priority rules when cis and trans are ambiguous.
A carbon atom bonded to four different groups is a stereocentre (a chiral centre). A molecule with one stereocentre exists as two enantiomers, non-superimposable mirror images, with identical physical properties except the direction in which they rotate plane-polarised light and their interactions with other chiral molecules, such as enzymes and taste and odour receptors. Stereoisomers that are not mirror images are diastereomers and differ in physical properties. A molecule with n stereocentres has at most 2ⁿ stereoisomers; fewer exist when a meso form, with internal symmetry, is possible. Glucose in its open chain has four stereocentres, so there are 16 aldohexoses. Two naming systems are used. The D/L system, which relates the configuration of the highest-numbered stereocentre to D- or L-glyceraldehyde, is standard for sugars and amino acids; the R/S system of Cahn, Ingold and Prelog applies to any stereocentre by ranking substituents. Neither label predicts the sign of rotation, which is written separately as (+) for dextrorotatory and (−) for laevorotatory: D-fructose is laevorotatory.
Optically active compounds rotate the plane of polarised light. A polarimeter measures the observed rotation α, and the specific rotation is [α] = α/(l·c), with l in decimetres and c in grams per millilitre, at a stated temperature and wavelength (usually the sodium D line). Sucrose has a specific rotation of about +66.5°. When sucrose is hydrolysed to equal amounts of glucose (+52.7° at equilibrium) and fructose (about −92°), the mixture becomes laevorotatory, about −20°: the sign of rotation inverts, which is why the product is called invert sugar. Polarimetry is a standard method for sugar in the sugar industry, measuring sucrose as degrees on an international sugar scale.
Sugars in solution show mutarotation. D-Glucose cyclises to two anomers that differ in configuration at C1, the new stereocentre: α-D-glucopyranose (+112°) and β-D-glucopyranose (+18.7°). Either pure anomer dissolved in water changes its rotation over time to the equilibrium value of +52.7°, a mixture of roughly one-third α and two-thirds β. Epimers are diastereomers that differ at a single stereocentre: D-mannose is the C2 epimer and D-galactose the C4 epimer of D-glucose.
Stereochemistry shapes taste, odour, nutrition and authenticity. Proteins are built from L-amino acids (glycine alone is achiral), and enzymes act on one enantiomer of their substrate. R-(−)-carvone smells of spearmint, while S-(+)-carvone smells of caraway. Lactic acid bacteria produce L(+), D(−) or mixed lactic acid depending on the species. Because natural flavours are often a single enantiomer while synthetic ones are racemic, chiral gas chromatography helps detect undeclared synthetic flavour additions in products sold as natural. Processing at high pH and temperature can racemise amino acids, slightly lowering protein quality.
Molecules at a surface or an interface lack neighbours on one side, so the surface has excess free energy; per unit area this is the surface (or interfacial) tension, about 72 mN/m for water at 25 °C against air. Surface-active agents (surfactants), molecules with a hydrophilic head and a hydrophobic tail, adsorb at interfaces and lower the tension, which makes it easier to create new interface when forming emulsions and foams. The contact angle of a liquid on a solid describes wetting: an angle below 90° means the liquid wets the surface, which matters in the cleaning of equipment and the reconstitution of powders. Food emulsifiers are characterised by their hydrophilic-lipophilic balance (HLB): low values (about 3 to 6) favour water-in-oil emulsions such as margarine, higher values (about 8 to 18) favour oil-in-water emulsions such as mayonnaise and beverages.
Adsorption is the accumulation of a substance on a surface. Physisorption involves weak van der Waals forces, is reversible and can build multilayers; chemisorption involves chemical bonds, is specific and limited to a monolayer. The Langmuir isotherm assumes a monolayer of equivalent sites: θ = KP/(1 + KP), where θ is the fraction of sites covered. The Freundlich isotherm, x/m = kC^(1/n), is empirical and fits heterogeneous surfaces. The Brunauer-Emmett-Teller (BET) isotherm extends Langmuir's model to multilayers and is used both to measure the surface area of powders and to estimate the monolayer moisture content from food sorption isotherms, often the moisture of best storage stability for dry foods. Adsorption is used in food processing: bleaching earths and activated carbon remove pigments and oxidation products during oil refining, and activated carbon decolourises sugar syrups.
A colloid is a dispersion of one phase in another in which the dispersed particles are roughly 1 nm to 1 µm across, too small to settle quickly and large enough to scatter light. The Tyndall effect, the visible scattering of a light beam, distinguishes colloids from true solutions, and the Brownian motion of the particles slows their settling. Colloids are classified by the phases involved: sols (solid in liquid, such as cocoa particles in a drink), emulsions (liquid in liquid, such as milk, cream and mayonnaise), foams (gas in liquid, such as whipped cream and beer head), gels (a liquid held in a continuous solid network, such as gelatin dessert and set yoghurt), and aerosols. Lyophilic (solvent-loving) colloids such as gelatin and gums dissolve to form stable dispersions; lyophobic colloids such as metal sols are stable only when charged or protected.
Colloidal stability depends on the balance between attraction (van der Waals forces) and repulsion (electrostatic and steric). Charged particles carry an electrical double layer, and the zeta potential, the potential at the slipping plane, indicates stability: larger magnitudes, roughly above 30 mV, usually mean stable dispersions. DLVO theory combines van der Waals attraction and double-layer repulsion. Adding salt compresses the double layer and can cause coagulation, more strongly the higher the charge of the counter-ion (the Schulze-Hardy rule). Adsorbed proteins and polysaccharides add steric stabilisation. Particles also cream or settle under gravity at a rate given by Stokes' law, v = 2r²(ρ_p − ρ_f)g/(9η), so halving the droplet radius cuts the creaming rate fourfold, which is why milk is homogenised, and increasing the viscosity of the continuous phase slows separation, which is why beverages and dressings contain thickeners. The food-structure-encapsulation subject treats emulsions, foams and gels in depth.
Oxidation is loss of electrons and reduction is gain; the oxidising agent is itself reduced. Oxidation numbers track the electrons: in KMnO₄ manganese is +7, and it is reduced to +2 in acid solution, gaining five electrons. A galvanic cell separates the oxidation and reduction half-reactions so that the electrons flow through an external circuit. Each half-reaction has a standard electrode potential E° measured against the standard hydrogen electrode (0 V); the more positive E°, the stronger the oxidising agent. The cell potential E°cell = E°cathode − E°anode is positive for a spontaneous reaction, and ΔG° = −nFE°, where F = 96 485 C/mol is the Faraday constant.
The Nernst equation gives the potential away from standard conditions: E = E° − (RT/nF) ln Q, which at 25 °C is E = E° − (0.0592/n) log₁₀ Q. The glass pH electrode responds in this way to hydrogen-ion activity, so an ideal electrode changes by 59.2 mV per pH unit at 25 °C, and pH meters are calibrated with two or three buffers and compensated for temperature because the slope changes with temperature. Ion-selective electrodes for sodium, fluoride and nitrate work on the same principle.
The redox potential (Eh) of a food describes its oxidising or reducing tendency and affects which microbes can grow. Aerobic organisms need positive Eh; obligate anaerobes such as clostridia grow at negative Eh. Fresh plant and muscle tissues contain reducing substances such as ascorbic acid and sulphydryl groups; vacuum packing, cooking and the activity of microbes lower Eh. Redox titrations are workhorses of food analysis: iodometric titration with sodium thiosulphate for the peroxide value of fats; permanganate and dichromate titrations; and the 2,6-dichlorophenolindophenol titration for ascorbic acid, in which the blue dye is reduced to colourless by vitamin C.
A coordination compound has a central metal ion bonded to ligands, molecules or ions that donate electron pairs (Lewis bases). The coordination number is the number of donor atoms bound, commonly four or six. Ligands that bind through two or more atoms are chelating agents and form ring structures that are much more stable than complexes with comparable single-donor ligands (the chelate effect). Ethylenediaminetetraacetic acid (EDTA) is hexadentate and binds most divalent and trivalent metal ions in a 1 : 1 ratio, which is the basis of complexometric titration of water hardness with Eriochrome Black T indicator; stability constants measure how strongly a ligand binds a metal.
Metal ions shape the colour, texture and stability of foods. In myoglobin, iron bound in a porphyrin ring (haem) gives meat its colour: purple deoxymyoglobin and bright-red oxymyoglobin contain Fe²⁺, while brown metmyoglobin contains Fe³⁺, and cured meats owe their pink colour to nitric oxide bound to the iron. In chlorophyll, magnesium sits in a similar porphyrin-type ring; acid replaces it with hydrogen to give olive-brown pheophytin when green vegetables are overcooked, and zinc or copper can replace it to give a stable green colour (regreening), a practice that some jurisdictions restrict. Calcium ions bridge the carboxyl groups of low-methoxyl pectin and alginate in the egg-box structure of their gels, and calcium phosphate nanoclusters hold casein micelles together. Calcium salts also firm canned vegetables and fruit by cross-linking pectin.
Transition metals, especially copper and iron, are potent pro-oxidants: they decompose lipid hydroperoxides into radicals and, by the Fenton reaction of Fe²⁺ with hydrogen peroxide, form hydroxyl radicals. Even traces matter, which is why equipment in contact with oils avoids copper alloys. Chelating agents such as citric acid, phosphates and EDTA bind these metals and act as synergists with phenolic antioxidants. Metal-polyphenol complexes cause discolouration (iron with tannins turns tea and some canned foods dark), and tin and iron dissolving from cans can react with anthocyanins and sulphur compounds. Mineral nutrition also depends on coordination chemistry: phytate and oxalate chelate iron, zinc and calcium and reduce their absorption.
A student uses this subject in every laboratory practical: preparing and standardising solutions, calculating titratable acidity, salt and peroxide values, calibrating a pH meter, choosing an indicator, and interpreting why a buffer holds or fails. In food chemistry lectures it explains Maillard browning, lipid oxidation, inversion of sucrose and the colour of meat and vegetables. In examinations, the fastest gains come from the standard numbers (pKa values, Kf of water, Atwater factors, specific rotations) and from never confusing molarity, molality and normality.
Five years into a quality-control or product-development job, a technologist uses chemistry to set pH and acidity specifications, choose acidulants and buffer salts, select emulsifiers by HLB, predict freezing points of frozen desserts, troubleshoot rancidity (metal contamination, light exposure, missing chelator), and explain to a supplier why a crystal or a colour defect appeared. Laboratory supervisors check calculations, standard solution records and electrode calibrations as part of accreditation.
A senior technical leader with twenty-five years' experience applies chemistry to decisions with consequences: whether a reformulated low-acid product needs a full botulinum process, whether a cheaper oil blend will meet shelf life, how far a packaging change alters oxidation, and whether a supplier's analytical results are chemically plausible. They also judge the chemical basis of regulatory limits and claims and represent the company in technical discussions with regulators.
A researcher at this level studies reaction mechanisms and physical chemistry in complex food matrices: radical chemistry of lipid oxidation in emulsions, Maillard pathways and process contaminants, thermodynamics of biopolymer phase behaviour, metal-ligand chemistry of mineral bioavailability, and chirality in authenticity testing. The core chemistry here is the vocabulary for that work and for reviewing whether kinetic and equilibrium models are applied correctly.
This manual covers general, organic and physical chemistry at the level food science needs; it does not cover spectroscopy and instrumental methods (in food-analysis and advanced-analytical-omics), quantum chemistry, detailed reaction mechanisms beyond the four food-relevant types, or polymer chemistry of food biopolymers in depth (in food-chemistry-major-constituents). Constants are quoted at 25 °C where they vary with temperature; reference values that differ by authority, such as the freezing point of milk, are described without a single number. Tables of pKa, Ksp and electrode potentials are not provided.
Original practice questions written to the syllabus and pattern of the exams named in each question's examTags. They are not past papers and do not reproduce any board's questions.
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