Units 3 & 4 Chemistry: Equilibrium, Acids, Redox, and Organic Analysis

Comprehensive vocabulary flashcards covering key terms and concepts across Unit 3 and Unit 4 Chemistry, including chemical equilibrium, acid-base systems, volumetric analysis, redox and electrochemistry, organic pathways, biomolecules, spectroscopy, and industrial chemical synthesis.

Open system

A reaction vessel lacking a boundary or lid, allowing matter and thermal energy to be freely exchanged with the surrounding atmosphere.

Closed system

A reaction vessel that allows the transfer of energy but prevents the exchange or loss of matter with the surroundings.

Dynamic equilibrium

A state achieved in a closed system where the rate of the forward reaction equals the rate of the reverse reaction, resulting in constant macroscopic concentrations.

Le Châtelier's principle

The rule stating that when an external change in concentration, temperature, or pressure is imposed on a system at equilibrium, the system will adjust to partially counteract the disturbance.

Equilibrium constant ($K_c$)

The numerical ratio formed by dividing the equilibrium concentrations of products by reactants, each raised to the power of their stoichiometric coefficients.

Reaction quotient ($Q$)

A calculated ratio of non-equilibrium or instantaneous species concentrations that determines the direction a mixture must shift to attain balance.

Solubility product constant ($K_{sp}$)

The equilibrium constant expression governing the dissolution of a sparingly soluble ionic solid into its constituent aqueous ions in a saturated solution.

RICE table

A tabular framework organizing the stoichiometric ratio, initial concentrations, changes occurring during the reaction, and final values at equilibrium.

Monoprotic acid

A substance capable of donating exactly one hydrogen ion ($\text{H}^+$) per chemical formula unit to a solution.

Polyprotic acid

A substance that can donate two or more protons per formula unit through stepwise ionisation stages.

Self-ionisation of water

An amphiprotic process wherein one water molecule donates a proton to another, generating equal amounts of hydronium ($\text{H}_3\text{O}^+_{(aq)}$) and hydroxide ($\text{OH}^-_{(aq)}$).

Ionisation constant of water ($K_w$)

The equilibrium value defined by the product $[\text{H}_3\text{O}^+][\text{OH}^-]$, which equals $1.00 \times 10^{-14}$ under standard conditions of $25\,^\circ\text{C}$.

Temperature effect on water autoionisation

Because the forward dissociation is endothermic, increasing thermal energy shifts the position forward, raising the equilibrium constant above $1.00 \times 10^{-14}$ and decreasing the neutral pH below $7.00$.

Brønsted-Lowry acid

Any chemical entity capable of donating a proton ($\text{H}^+$) to another reacting particle.

Brønsted-Lowry base

Any chemical entity capable of accepting a proton ($\text{H}^+$) from a reacting partner.

Conjugate acid-base pair

Two species linked in an acid-base equilibrium whose molecular formulas differ strictly by the presence or absence of a single proton.

Amphiprotic species

A molecule or ion that can donate or accept a proton depending on the surrounding medium, such as $\text{H}_2\text{O}$, $\text{HCO}_3^-$, or amino acids.

Difference between amphiprotic and amphoteric

The first term specifically describes species transferring protons, whereas the broader second term refers to any substance that can react with both acids and bases.

Acid dissociation constant ($K_a$)

An equilibrium parameter quantifying the extent to which a weak molecular acid dissociates into ions in aqueous solution.

Base dissociation constant ($K_b$)

An equilibrium parameter measuring the extent to which a weak base accepts a proton from water to produce hydroxide ions.

Mathematical relationship between $K_a$ and $K_b$

For any conjugate acid-base pair in aqueous solution at $25\,^\circ\text{C}$, their dissociation values satisfy the relationship $K_a \times K_b = K_w = 1.00 \times 10^{-14}$.

$$pK_a$$

The negative base-$10$ logarithm of an acid dissociation constant ($-\log_{10}(K_a)$), where smaller numerical values represent greater acid strength.

Acid-base indicator mechanism

Weakly acidic or basic dye molecules that display distinct colours between their un-ionised and ionised conjugate forms as hydrogen ion concentration varies.

Equivalence point

The stage in a titration where the quantity of added titrant is stoichiometrically identical to the moles of analyte according to the balanced equation.

End point

The experimental stage during volumetric analysis where an added chemical indicator displays a distinct, permanent colour shift.

Buffer region

The flat portion of a weak acid or weak base titration curve where added titrant causes minimal pH change due to the coexistence of a conjugate pair.

Half-equivalence point

The midpoint of the buffer region where exactly half of the analyte has been neutralized, meaning $[\text{HA}] = [\text{A}^-]$ and the solution pH equals the $pK_a$.

Indicator selection for strong acid and weak base titration

Substances with acidic transition ranges such as methyl orange, bromocresol green, or methyl red, because the resulting salt hydrolysis produces an equivalence point below pH $7$.

Indicator selection for weak acid and strong base titration

Substances with alkaline transition ranges such as phenolphthalein, because the conjugate base hydrolysis creates an equivalence point above pH $7$.

Conductometric titration

An analytical method that tracks changes in electrical conductance as ions are consumed, neutralised, or replaced, identifying the equivalence point at the curve intersection.

Mohr method

A precipitation titration where chloride ions are quantified using silver nitrate titrant and potassium chromate indicator, forming a brick-red silver chromate precipitate at the endpoint.

Oxidation

A chemical process characterized by the loss of electrons, an elevation in oxidation number, or the addition of oxygen and removal of hydrogen in organic synthesis.

Reduction

A chemical process characterized by the gain of electrons, a decrease in oxidation number, or the removal of oxygen and gain of hydrogen in organic synthesis.

Oxidising agent

A chemical entity that extracts electrons from another reactant, inducing oxidation while undergoing reduction itself.

Reducing agent

A chemical entity that supplies electrons to another substance, bringing about reduction while undergoing oxidation itself.

Oxidation number

A formal bookkeeping charge assigned to an atom in a compound according to established electronegativity and bonding rules.

Oxidation number of free uncombined elements

A value of $0$, applicable to all elemental forms such as $\text{O}_2$, $\text{Fe}$, and $\text{Cl}_2$.

Oxidation number of hydrogen in metal hydrides

A value of $-1$, occurring when bonded to strongly electropositive metals such as in $\text{NaH}$ or $\text{CaH}_2$.

Oxidation number of oxygen in peroxides

A value of $-1$, found in compounds possessing an $\text{O}-\text{O}$ single bond like $\text{H}_2\text{O}_2$.

Galvanic cell

An electrochemical setup that converts chemical energy from spontaneous oxidation-reduction reactions into continuous electrical current.

Anode in an electrochemical cell

The electrode at which oxidation occurs, carrying a negative polarity in galvanic cells and a positive polarity in electrolytic cells.

Cathode in an electrochemical cell

The electrode at which reduction occurs, carrying a positive polarity in galvanic cells and a negative polarity in electrolytic cells.

Salt bridge

A conducting medium containing an inert electrolyte that completes an electrochemical circuit by allowing cations to migrate toward the cathode and anions toward the anode.

Standard cell potential ($E^\circ_{\text{cell}}$)

The maximum potential difference generated by a cell under standard states ($25\,^\circ\text{C}$, $1.0\,\text{mol}\,\text{dm}^{-3}$, and $100\,\text{kPa}$), calculated using $E^\circ_{\text{reduction}} - E^\circ_{\text{oxidation}}$

Condition for spontaneous redox reaction

A net cell electromotive force where $E^\circ_{\text{cell}} > 0\,\text{V}$.

Electrolytic cell

An electrochemical system that utilizes an external direct current power source to drive thermodynamically non-spontaneous chemical transformations.

Electrolysis products of molten sodium chloride

Liquid metallic sodium produced by reduction at the cathode and diatomic chlorine gas produced by oxidation at the anode.

Electrolysis products of concentrated aqueous sodium chloride

Hydrogen gas at the cathode due to preferential reduction of hydronium ions, chlorine gas at the anode due to high halide concentration, and aqueous sodium hydroxide in solution.

Faraday's first law of electrolysis

The principle stating that the mass of substance deposited or liberated at an electrode is directly proportional to the total electrical charge ($Q = I \times t$) delivered to the cell.

Faraday's second law of electrolysis

The principle stating that producing one mole of substance requires consuming a whole number of moles of electrons determined by the ionic charge stoichiometry.

Faraday's constant ($F$)

The absolute electrical charge carried by one mole of electrons, approximately equal to $96\,485\,\text{C}\,\text{mol}^{-1}$.

Homologous series

A family of organic molecules sharing the same functional group and general formula, where each sequential member differs by a $-\text{CH}_2-$ unit.

Boiling point trend in straight-chain alkanes

Values increase progressively with carbon chain length because larger electron clouds enhance temporary dipoles and overall London dispersion forces.

Effect of chain branching on hydrocarbon boiling points

Branched isomers exhibit lower boiling points than straight-chain equivalents because compact molecular shapes reduce contact surface area and weaken dispersion interactions.

Alcohol classification (primary, secondary, tertiary)

Categorisation based on whether the hydroxyl-bearing carbon is directly attached to one, two, or three other carbon atoms.

Hydrogen bonding in organic molecules

A strong dipole-dipole attraction occurring when hydrogen is covalently linked to strongly electronegative atoms ($\text{N}$, $\text{O}$, or $\text{F}$), markedly elevating boiling points in alcohols, carboxylic acids, and amides.

Catalytic hydrogenation of alkenes

An addition reaction wherein molecular hydrogen adds across a carbon-carbon double bond in the presence of finely divided metal catalysts like $\text{Ni}$, $\text{Pt}$, or $\text{Pd}$ to form an alkane.

Bromine test for unsaturation

A diagnostic procedure where rapid decolourisation of reddish-brown aqueous bromine without catalyst confirms the presence of carbon-carbon multiple bonds.

Hydration of ethene

An industrial electrophilic addition reaction combining ethene with steam at $300\,^\circ\text{C}$ and $70\,\text{atm}$ over solid phosphoric acid catalyst to produce ethanol.

Markovnikov's rule

The empirical rule dictating that when a hydrogen halide adds to an asymmetric alkene, the hydrogen atom attaches to the alkene carbon holding more hydrogen substituents, while the halogen binds to the more substituted carbon.

Elimination reaction of haloalkanes

The removal of a halogen and an adjacent hydrogen atom under reflux with concentrated ethanolic base ($\text{NaOH}$ or $\text{KOH}$) to yield an alkene and a salt.

Oxidation of primary alcohols

A stepwise reaction using acidified potassium dichromate that converts the compound first into an aldehyde under gentle distillation, and subsequently into a carboxylic acid under reflux.

Oxidation of secondary alcohols

A conversion employing acidified dichromate or permanganate under reflux to produce a ketone, which resists further oxidative cleavage.

Oxidation resistance of tertiary alcohols

The inability to undergo typical oxidation reactions with acidified dichromate or permanganate because the hydroxyl-bearing carbon atom lacks a hydrogen atom.

Benedict's test

A diagnostic redox assay where an aldehyde reduces copper(II) complexes to yield an insoluble red-orange copper(I) oxide precipitate, while ketones show no change.

Condensation reaction

A transformation linking two organic molecules together via a new covalent bond, accompanied by the concurrent elimination of a small stable molecule such as water.

Esterification

A reversible equilibrium condensation between a carboxylic acid and an alcohol heated under reflux in the presence of concentrated sulfuric acid catalyst.

Amide formation

A condensation process where a carboxylic acid reacts with an amine or ammonia at elevated temperatures ($> 100\,^\circ\text{C}$) to yield an amide and water.

Nucleophilic substitution with potassium cyanide

A synthetic pathway where a haloalkane is heated under reflux with ethanolic cyanide, displacing the halide to form a nitrile while extending the carbon skeleton by one atom.

Reduction of nitriles

A catalytic or hydride-mediated transformation using $\text{LiAlH}_4$ or hydrogen gas with metal catalysts to convert cyano groups into primary amines.

Monosaccharide structural isomers ($\alpha$- vs $\beta$-glucose)

Cyclic hexose forms differing at the anomeric carbon, where the first isomer features the hydroxyl group pointing downwards and the second features it pointing upwards.

Glycosidic bond

A covalent ether linkage connecting monosaccharide units through the condensation of their hydroxyl groups.

Peptide bond

A rigid covalent amide linkage formed by condensation between the carboxyl carbon of one amino acid and the amino nitrogen of another.

Four levels of protein structure

Primary (amino acid sequence), secondary (local hydrogen-bonded $\alpha$-helices and $\beta$-pleated sheets), tertiary (overall 3D conformation from side-chain interactions), and quaternary (assembly of multiple polypeptide subunits).

Tertiary protein stabilization interactions

Hydrophobic associations, hydrogen bonding, ionic salt bridges between charged residues, and covalent disulfide cross-links between cysteine units.

Induced fit model

An enzymatic mechanism where substrate binding causes dynamic conformational adjustments in the catalytic active site to optimize transition-state stabilization.

Base peak in mass spectrometry

The most intense spectral line assigned an arbitrary relative abundance of $100\%$, representing the most stable and prevalent charged fragment.

Molecular ion peak ($M^+$)

The spectral peak corresponding to the intact, unfragmented radical cation of the analyte, whose mass-to-charge ratio gives the nominal molecular weight.

Chlorine isotope peak pattern

A distinctive molecular ion cluster showing an $M$ to $M+2$ intensity ratio of approximately $3:1$ due to the natural abundance of chlorine-35 and chlorine-37.

Infrared active molecular vibrations

Periodic stretching and bending movements that cause a net change in the molecular dipole moment as infrared radiation is absorbed.

Carbonyl ($\text{C}=\text{O}$) infrared absorption band

A sharp, intense absorption peak appearing between $1650\,\text{cm}^{-1}$ and $1850\,\text{cm}^{-1}$ across aldehydes, ketones, esters, and carboxylic acids.

Hydroxyl ($\text{O}-\text{H}$) infrared absorption in carboxylic acids vs alcohols

The acid features an extremely broad absorption band spanning $2500\,\text{cm}^{-1}$ to $3300\,\text{cm}^{-1}$ centered near $3000\,\text{cm}^{-1}$, whereas the neutral alcohol produces a smooth, parabolic peak near $3300\,\text{cm}^{-1}$.

Retention factor ($R_f$)

The dimensionless ratio calculated by dividing the migration distance of an individual solute spot by the distance reached by the solvent front.

Isoelectric point ($pI$) in gel electrophoresis

The specific pH at which an amphiprotic molecule carries no net electrical charge, causing its electrokinetic migration to cease.

Haber process compromise conditions

A compromised operating temperature of $400\text{ to }500\,^\circ\text{C}$ balancing rate and exothermic yield, a high pressure of $200\,\text{atm}$, and a porous iron catalyst.

Contact process conversion reaction

The catalytic oxidation of sulfur dioxide according to $2\text{SO}_{2(g)} + \text{O}_{2(g)} \rightleftharpoons 2\text{SO}_{3(g)}$ at $450\,^\circ\text{C}$ and $1\text{ to }2\,\text{atm}$ over vanadium(V) oxide.

Electrode reactions in an acidic hydrogen fuel cell

Hydrogen is oxidised at the anode ($2\text{H}_{2(g)} \rightarrow 4\text{H}^+_{(aq)} + 4e^-$) and oxygen is reduced at the cathode ($\text{O}_{2(g)} + 4\text{H}^+_{(aq)} + 4e^- \rightarrow 2\text{H}_2\text{O}_{(l)}$).

Electrode reactions in an alkaline hydrogen fuel cell

Hydrogen is oxidised at the anode ($2\text{H}_{2(g)} + 4\text{OH}^-_{(aq)} \rightarrow 4\text{H}_2\text{O}_{(l)} + 4e^-$) and oxygen is reduced at the cathode ($\text{O}_{2(g)} + 2\text{H}_2\text{O}_{(l)} + 4e^- \rightarrow 4\text{OH}^-_{(aq)}$).