Topic 6: Reactivity 3: What are the mechanisms of chemical change

Chemical change happens through a small number of recurring mechanisms: moving protons, moving electrons, or sharing electrons. This theme groups acid-base chemistry, redox and electrochemistry, and organic mechanisms under those three big ideas, giving you a unified way to predict and explain reactions.

Proton transfer: Bronsted-Lowry acids and bases

A Bronsted-Lowry acid is a proton (H+) donor and a Bronsted-Lowry base is a proton acceptor. Every acid-base reaction is therefore a transfer of one H+ from the acid to the base. Water is amphiprotic, meaning it can act as either an acid or a base depending on its partner: it accepts a proton from HCl but donates one to ammonia. When an acid donates its proton, what remains is a species capable of accepting a proton back, which leads directly to the idea of conjugate pairs.

Conjugate acid-base pairs

A conjugate pair differs by exactly one H+. When an acid HA loses a proton it becomes its conjugate base A-, and when a base B gains a proton it becomes its conjugate acid BH+. For example, in the reaction of ethanoic acid with water, CH3COOH and CH3COO- form one pair while H2O and H3O+ form the other. The stronger an acid, the weaker its conjugate base, because a species that readily gives up H+ has little tendency to take it back.

pH and the ionic product of water

Pure water self-ionizes slightly: 2 H2O gives H3O+ plus OH-. The ionic product Kw equals [H+] times [OH-] and is 1.0 x 10^-14 at 298 K, so in neutral water both ions are 1.0 x 10^-7 mol/dm3. The pH scale is defined as pH = -log[H+], so a tenfold change in [H+] shifts pH by one unit. Because Kw is fixed at a given temperature, pH + pOH = 14 at 298 K. Kw increases with temperature, so the pH of neutral water falls below 7 when heated even though it stays neutral.

Strong versus weak acids and bases

Strong acids and bases ionize essentially completely in water (HCl, HNO3, H2SO4; group 1 hydroxides), so their concentration of dissolved acid roughly equals the concentration of H+ released. Weak acids and bases ionize only partially and exist mostly as undissociated molecules in equilibrium (ethanoic acid, ammonia). At equal concentration a strong acid has a lower pH, higher electrical conductivity, and faster initial reaction rate than a weak acid, but both can fully neutralize the same amount of base because total available protons are what matter for neutralization.

HL: Ka, Kb, pKa and pKb

For a weak acid HA, the acid dissociation constant Ka = [H+][A-] / [HA] measures the position of its ionization equilibrium; a larger Ka means a stronger acid. The corresponding base constant Kb describes a weak base. For a conjugate pair, Ka times Kb equals Kw, so pKa + pKb = 14 at 298 K. Because the values span many orders of magnitude, we use pKa = -log Ka and pKb = -log Kb: a smaller pKa means a stronger acid. These constants let you calculate the pH of weak acid or weak base solutions from concentration alone.

HL: Buffer solutions

A buffer resists changes in pH when small amounts of acid or base are added. An acidic buffer is made from a weak acid and a salt of its conjugate base (for example ethanoic acid plus sodium ethanoate). Added H+ is mopped up by the conjugate base and added OH- is neutralized by the weak acid, so pH stays nearly constant. The pH is governed by the Henderson-Hasselbalch relationship, pH = pKa + log([A-]/[HA]); when acid and salt concentrations are equal, pH equals pKa, which is the point of maximum buffering capacity.

Electron transfer: redox and oxidation states

Redox reactions transfer electrons. Oxidation is loss of electrons and reduction is gain of electrons (remember OIL RIG). The oxidizing agent is itself reduced and the reducing agent is itself oxidized. Oxidation states are bookkeeping numbers assigned using rules: elements are 0, hydrogen is usually +1, oxygen usually -2, and the states in a neutral compound sum to zero. An increase in oxidation state signals oxidation and a decrease signals reduction, which is how you identify what is being oxidized and reduced even in complex equations.

Half-equations and balancing redox

A redox reaction can be split into a reduction half-equation and an oxidation half-equation, each balanced separately. In acidic solution you balance the atoms other than O and H first, then add H2O to balance oxygen, H+ to balance hydrogen, and finally electrons to balance charge. The two halves are multiplied so the electrons cancel when added, giving the balanced overall equation. This method makes it possible to balance reactions, such as those involving permanganate or dichromate, that are very hard to balance by inspection.

Voltaic and electrolytic cells

A voltaic (galvanic) cell converts chemical energy into electrical energy from a spontaneous redox reaction. Oxidation occurs at the negative anode and reduction at the positive cathode; electrons flow through the external wire from anode to cathode while a salt bridge maintains charge balance. An electrolytic cell does the opposite: an external power supply drives a non-spontaneous reaction, with oxidation at the now positive anode and reduction at the negative cathode. In both cells reduction always occurs at the cathode; only the relative signs and energy direction differ.

HL: Standard electrode potentials

Each half-cell has a standard electrode potential measured against the standard hydrogen electrode, which is defined as 0.00 V. More positive potentials correspond to species more easily reduced (stronger oxidizing agents). For a cell, E_cell = E(cathode) - E(anode), using the more positive electrode as the cathode. A positive E_cell indicates a spontaneous reaction (linked to a negative free energy change), while a negative E_cell means the reaction is non-spontaneous and would require electrolysis to proceed.

Electron sharing: radicals and free-radical substitution

Some bonds break homolytically, with each atom keeping one electron, producing radicals: species with an unpaired electron that are highly reactive. Alkanes react with halogens such as chlorine in ultraviolet light by free-radical substitution. The mechanism has three stages: initiation, where UV light splits Cl2 into two chlorine radicals; propagation, a chain of steps that regenerates radicals and forms the product; and termination, where two radicals combine to remove them from the chain. Because propagation can run many times and form multiple products, this mechanism is hard to control.

HL: Nucleophiles, electrophiles and SN1/SN2

Heterolytic bond breaking gives ions, and reactions are driven by electron-rich nucleophiles attacking electron-poor electrophiles. Halogenoalkanes undergo nucleophilic substitution by two pathways. SN2 is a single concerted step, second order overall, favored by primary substrates and proceeding through a transition state with inversion of configuration. SN1 goes through a carbocation intermediate, is first order in the substrate, and is favored by tertiary substrates because the intermediate is more stable. Electrophilic addition adds across the C=C double bond of alkenes, while condensation reactions, such as ester or amide formation, join molecules with the loss of a small molecule like water.