Chemical Changes: Acids, Bases, Salts, Reactivity, and Electrolysis

The pH Scale and Acid-Base Fundamentals

The pH scale is a continuous numerical scale ranging from $0$ to $14$ that quantifies the acidity or alkalinity of an aqueous solution. Solutions with a pH value lower than $7$ are acidic, where lower numbers correspond to greater acidity and higher concentrations of hydrogen ions ($H^+$). A pH value of exactly $7$ represents a neutral solution, such as pure water, which is neither acidic nor alkaline. Solutions with a pH value greater than $7$ are alkaline, where higher numbers correspond to greater alkalinity and higher concentrations of hydroxide ions ($OH^-$).

Various household and biological substances occupy distinct positions across the pH spectrum. Concentrated acids, such as car battery acid and human stomach acid (which functions physiologically to kill ingested bacteria), exhibit pH values between approximately $1$ and $2$. Edible acidic substances, including lemon juice and vinegar, typically have pH values around $2$ to $3$. Acid precipitation (acid rain) has an average pH of approximately $4$. Human skin maintains a slightly acidic surface environment with an average pH of $5.5$, whereas unpolluted natural rain displays a pH of approximately $6$. In neutral and alkaline ranges, pure water is pH $7$, pancreatic secretions register at approximately pH $8$, household cleansers such as washing-up liquid and soap powder register at pH $9$ to $10$, domestic bathroom bleach displays a pH of approximately $12$, and concentrated caustic soda (sodium hydroxide, commonly used as a drain cleaner) exhibits a highly alkaline pH of $13$ to $14$.

The pH of a solution can be assessed experimentally using chemical indicators or electronic instruments. Chemical indicators are synthetic or organic dyes that change colour at specific pH thresholds. Wide-range indicators consist of mixtures of multiple dyes that exhibit gradual, progressive colour transitions across the entire pH scale from $0$ to $14$. Universal indicator is the archetypal wide-range indicator, displaying deep red in strongly acidic conditions, progressing through orange and yellow in weakly acidic environments, turning green at neutrality (pH $7$), and shifting through blue to deep purple in increasingly alkaline solutions. Indicators may also be infused into paper strips for field testing. For objective, high-precision measurement, a pH probe connected to an electronic pH meter is utilised. The probe is submerged into the test solution, directly measuring electrical potential to produce a real-time digital numerical pH reading. Electronic measurement eliminates subjective human errors associated with visually estimating colour nuances against comparison charts.

From a chemical perspective, an acid is defined as any substance that dissociates in aqueous solution to release hydrogen ions ($H^+$), yielding a solution with a pH of less than $7$. Because a standard neutral hydrogen atom consists of a single proton and a single electron, the loss of an electron to form an $H^+$ ion means that a hydrogen ion is fundamentally a bare proton. Common laboratory acids include hydrochloric acid ($HCl$), sulfuric acid ($H_2SO_4$), and nitric acid ($HNO_3$). A base is defined as any chemical compound that reacts chemically with an acid to neutralise it, yielding a salt and water. Bases are predominantly insoluble metal oxides, metal hydroxides, and metal carbonates, such as copper(II) oxide ($CuO$) and calcium carbonate ($CaCO_3$). An alkali represents a specific soluble subset of bases: an alkali dissolves in water to produce hydroxide ions ($OH^-$) in solution, resulting in an aqueous pH greater than $7$. Common alkalis include sodium hydroxide ($NaOH$) and potassium hydroxide ($KOH$).

When an acid and an alkali react together, they undergo a neutralisation reaction. In this process, the hydrogen ions originating from the acid combine with the hydroxide ions originating from the alkali to synthesize neutral water molecules. The corresponding ionic equation representing all acid-alkali neutralisation reactions is:

$H^+_{(aq)} + OH^-_{(aq)} \rightarrow H_2O_{(l)}$

Because the acidic and basic species neutralise one another, the resulting aqueous salt and water mixture reaches an end-point neutrality with an ideal pH of $7$.

Strong and Weak Acids versus Concentration

Acids in aqueous solutions are classified as either strong or weak depending upon their degree of ionisation and molecular dissociation. When an acid dissolves in water, its covalent bonds split to release aqueous hydrogen ions ($H^+$).

A strong acid ionises completely in aqueous solution. In a strong acid, all individual acid particles dissociate fully to release $H^+$ ions and their associated conjugate anions, meaning the reactants are completely converted into product ions. Exemplary strong acids and their dissociation equations include:

$HCl_{(aq)} \rightarrow H^+_{(aq)} + Cl^-_{(aq)}$

HNO_3_{(aq)} \rightarrow H^+_{(aq)} + NO_3^-_{(aq)}

H_2SO_4_{(aq)} \rightarrow 2H^+_{(aq)} + SO_4^{2-}_{(aq)}

A weak acid ionises only partially in aqueous solution, with only a small fraction of the dissolved molecules dissociating to release $H^+$ ions. The ionisation of a weak acid is a reversible chemical process that achieves a dynamic equilibrium between the undissociated molecular acid and the dissociated aqueous ions. Because only a minute proportion of the molecules dissociate, the position of chemical equilibrium lies far to the left, indicating that the concentration of intact, undissociated acid molecules vastly exceeds the concentration of free $H^+$ and conjugate anions at equilibrium. Common examples of weak acids include ethanoic acid ($CH_3COOH$, found in table vinegar), citric acid (found in citrus fruits), and carbonic acid ($H_2CO_3$). The reversible dissociation of ethanoic acid is represented by the equilibrium equation:

$CH_3COOH_{(aq)} \rightleftharpoons H^+_{(aq)} + CH_3COO^-_{(aq)}$

A strict scientific distinction exists between acid strength and acid concentration. Acid strength describes the intrinsic chemical tendency and proportion of acid particles to ionise and release protons in aqueous solution. Acid concentration represents a purely quantitative measurement of the total number of dissolved acid particles present per unit volume of solution (expressed in units such as $g\,dm^{-3}$ or $mol\,dm^{-3}$). Concentration measures the overall particle density regardless of whether those particles have dissociated into ions or remain as intact neutral molecules. Consequently, it is entirely possible to prepare a concentrated solution of a weak acid (containing a large quantity of dissolved weak acid molecules per $dm^3$, of which very few dissociate) or a dilute solution of a strong acid (containing very few dissolved strong acid particles per $dm^3$, all of which dissociate completely). For any given acid, increasing the concentration of dissolved solute introduces more free $H^+$ per unit volume, which systematically lowers the pH.

The pH scale is an inverse logarithmic representation of the molar concentration of hydrogen ions ($[H^+]$). Because of this inverse logarithmic relationship, an increase in hydrogen ion concentration causes a numerical decrease in pH. Specifically, every single integer decrease of $1$ on the pH scale corresponds directly to a ten-fold ($10\times$) increase in $H^+$ ion concentration. A solution of pH $1$ possesses $10$ times the $H^+$ ion concentration of a solution of pH $2$, and $100$ times ($10^2$) the $H^+$ ion concentration of a solution of pH $3$. The mathematical relationship governing this change is defined by the formula:

$\text{Factor by which } [H^+] \text{ changes} = 10^{-X}$

where $X$ represents the net change in pH, calculated as:

$X = \text{final pH} - \text{initial pH}$

For example, if the pH of an aqueous solution decreases from $7$ to $5$, the value of $X$ is $5 - 7 = -2$. Applying the formula yields:

$\text{Factor} = 10^{-(-2)} = 10^2 = 100$

This confirms that the hydrogen ion concentration has increased by a factor of $100$. Conversely, if the pH increases from $8$ to $11$, the net change is $11 - 8 = +3$, yielding a factor of $10^{-3} = \frac{1}{1000}$, meaning the $H^+$ concentration has decreased by a factor of $1000$.

Because the chemical reactions of acids are governed by the presence of free $H^+$ ions reacting with target substrates, the instantaneous rate of an acid reaction increases directly with increasing $H^+$ concentration. Therefore, if a strong acid and a weak acid are compared at identical molar concentrations, the strong acid will exhibit a significantly lower pH and a substantially faster rate of reaction, because its complete dissociation supplies a far greater concentration of active $H^+$ ions.

Quantitative Acid-Base Titrations and Indicator Chemistry

A titration is a high-precision volumetric analytical technique employed to determine the unknown concentration of an acidic or alkaline solution by finding the exact stoichiometric volume of a neutralising solution of known concentration required to reach the reaction end-point.

The rigorous laboratory procedure for executing an acid-alkali titration entails specific steps. A fixed volume of alkali is transferred into a clean conical flask using a volumetric pipette fitted with a safety pipette filler. The pipette is initially filled slightly above the graduation line, and the meniscus is carefully lowered until the bottom of the liquid curved surface aligns precisely with the calibration mark before draining the measured volume into the flask. Two to three drops of a suitable single indicator are added directly to the conical flask. A clean burette is filled with a standard solution of acid of precisely known concentration using a small funnel. To prevent facial chemical splash injuries, filling the burette must be performed below eye level, and laboratory safety glasses must be worn throughout. The funnel must be removed immediately after filling to prevent stray droplets of acid from draining into the burette during the titration. The initial liquid volume on the graduated burette scale is recorded to the nearest $0.05\,cm^3$.

A rough titration is performed first to establish the approximate end-point. The acid is rapidly dispensed from the burette into the conical flask while continuously swirling the mixture by hand. The flow of acid is stopped the instant the indicator undergoes an irreversible colour change, signifying that all alkali has been neutralised. The final volume reading is recorded, and the difference between the initial and final readings is calculated to determine the rough titre. Although the rough titre lacks high accuracy due to potential overshoot, it provides a benchmark volume for subsequent runs.

A series of careful, accurate titrations are executed next. The apparatus is reassembled with fresh alkali in the conical flask. Acid is added smoothly from the burette until the delivered volume approaches within approximately $1\,cm^3$ to $2\,cm^3$ of the rough titre benchmark. The stopcock is then adjusted to introduce the acid drop by drop, swirling continuously between drops. The titration is halted the instant a single drop produces an irreversible colour transition. The final burette volume is recorded, and the volume of acid consumed is calculated. The titration is repeated multiple times to ensure accuracy, identify anomalous results, and obtain concordant results, defined as titres agreeing within $0.1\,cm^3$ of one another. Concordant titres are averaged to calculate a reliable mean volume, which is subsequently used in quantitative stoichiometry calculations.

The selection of indicator is critical in volumetric titrations. Wide-range indicators, such as universal indicator, are unsuitable for titrations because they change colour gradually across a broad pH spectrum, rendering identification of the precise stoichiometric equivalence point impossible. In contrast, titrations require single indicators that undergo a sharp, sudden colour change at the exact neutralization end-point. The three standard single indicators utilised in laboratory titrations are:

Litmus: red in acidic conditions, blue in alkaline conditions.

Methyl orange: red in acidic conditions, yellow in alkaline conditions.

Phenolphthalein: colourless in acidic conditions, bright pink in alkaline conditions.

Neutralisation Reactions and the Preparation of Soluble Salts

Acids react with diverse bases, including metal oxides, metal hydroxides, and metal carbonates, through neutralisation pathways to form chemical salts. The identity of the salt produced is dictated by the specific metal cation supplied by the base and the characteristic conjugate anion supplied by the acid: hydrochloric acid ($HCl$) consistently forms chloride salts ($Cl^-$), sulfuric acid ($H_2SO_4$) forms sulfate salts ($SO_4^{2-}$), and nitric acid ($HNO_3$) forms nitrate salts ($NO_3^-$).

When acids react with insoluble metal oxides or soluble metal hydroxides, the products are exclusively a dissolved salt and neutral water:

$\text{Acid} + \text{Metal Oxide} \rightarrow \text{Salt} + \text{Water}$

$\text{Acid} + \text{Metal Hydroxide} \rightarrow \text{Salt} + \text{Water}$

Representative balanced chemical equations for metal oxide and metal hydroxide neutralisations include:

$2HCl_{(aq)} + Na_2O_{(s)} \rightarrow 2NaCl_{(aq)} + H_2O_{(l)}$

2HCl_{(aq)} + CuO_{(s)} \rightarrow CuCl_2_{(aq)} + H_2O_{(l)}

H_2SO_4_{(aq)} + 2KOH_{(aq)} \rightarrow K_2SO_4_{(aq)} + 2H_2O_{(l)}

HNO_3_{(aq)} + NaOH_{(aq)} \rightarrow NaNO_3_{(aq)} + H_2O_{(l)}

When acids react with metal carbonates, the reaction yields an aqueous salt, water, and gaseous carbon dioxide ($CO_2$):

$\text{Acid} + \text{Metal Carbonate} \rightarrow \text{Salt} + \text{Water} + \text{Carbon Dioxide}$

Representative balanced equations for carbonate neutralisations include:

2HCl_{(aq)} + Na_2CO_3_{(aq)} \rightarrow 2NaCl_{(aq)} + H_2O_{(l)} + CO_2_{(g)}

H_2SO_4_{(aq)} + CaCO_3_{(s)} \rightarrow CaSO_4_{(s/aq)} + H_2O_{(l)} + CO_2_{(g)}

2HNO_3_{(aq)} + CaCO_3_{(s)} \rightarrow Ca(NO_3)_2_{(aq)} + H_2O_{(l)} + CO_2_{(g)}

2HCl_{(aq)} + CaCO_3_{(s)} \rightarrow CaCl_2_{(aq)} + H_2O_{(l)} + CO_2_{(g)}

To prepare a pure, dry sample of a soluble salt from an insoluble base (such as preparing copper(II) chloride crystals from hydrochloric acid and copper(II) oxide), a standardised multi-step practical procedure is employed. A fixed volume of dilute acid is measured into a beaker and gently heated over a blue Bunsen burner flame to increase the initial rate of reaction; the burner is extinguished prior to adding reagents. The insoluble solid base is added to the warm acid incrementally using a spatula while stirring thoroughly with a glass rod. The solid dissolves as it neutralises the acid. Continued spatulas of the base are added until the substance is in excess, confirmed when solid powder remains visible and settles at the bottom of the beaker even after vigorous stirring, ensuring complete neutralisation of all acid molecules. The excess unreacted base is separated from the reaction mixture by filtration through filter paper seated within a filter funnel, collecting the clear aqueous salt filtrate in an evaporating basin.

Crystallisation of the pure salt is carried out by heating the filtered salt solution over a boiling water bath or with an electric hotplate to evaporate a portion of the solvent water until the crystallisation point is reached (indicated by small crystals forming on a cool glass rod dipped into the liquor). Direct heating with a naked Bunsen flame is avoided to prevent thermal decomposition of the hydrated salt crystals. Heating is halted, and the saturated solution is allowed to cool slowly at room temperature, inducing the formation of large, pure salt crystals as solubility decreases. The resulting solid crystals are isolated from any residual liquor by secondary filtration, rinsed with a minimal volume of cold distilled water to remove superficial impurities, and dried thoroughly by pressing between absorbent filter papers or warming in a low-temperature drying oven.

The Reactivity Series and Reactions of Metals

The reactivity series arranges metallic elements in descending order of their chemical reactivity. When metals participate in chemical reactions, their atoms undergo oxidation by losing their valence shell electrons to achieve stable electron configurations, transforming into positively charged cations. The relative chemical reactivity of a metal is fundamentally defined by the ease with which its atoms lose valence electrons to form positive ions: metals that forfeit valence electrons with greater ease occupy higher positions in the series and react with greater vigour.

The definitive order of the reactivity series, including non-metal reference elements, comprises:

Potassium ($K$) — Very reactive

Sodium ($Na$)

Lithium ($Li$)

Calcium ($Ca$)

Magnesium ($Mg$) — Fairly reactive

Carbon ($C$) — [Non-metal reference standard]

Zinc ($Zn$)

Iron ($Fe$)

Hydrogen ($H$) — [Non-metal reference standard]

Copper ($Cu$) — Not very reactive

Gold ($Au$) — Extremely unreactive (native metal)

The position of a metal relative to hydrogen dictates its behaviour when placed in dilute acid solutions. The general word equation for an acid-metal reaction is:

$\text{Acid} + \text{Metal} \rightarrow \text{Salt} + \text{Hydrogen}$

The rate of reaction corresponds to the rate of effervescence of hydrogen gas bubbles. Highly reactive metals situated at the apex of the series (potassium, sodium, lithium, and calcium) react violently and explosively with dilute acids; potassium burns spontaneously and moves rapidly on the liquid surface. Magnesium reacts vigorously with cold dilute hydrochloric or sulfuric acid, producing effervescence as the metal dissolves:

Mg_{(s)} + 2HCl_{(aq)} \rightarrow MgCl_2_{(aq)} + H_2_{(g)}

Mg_{(s)} + H_2SO_4_{(aq)} \rightarrow MgSO_4_{(aq)} + H_2_{(g)}

Zinc and iron react steadily and slowly with cold dilute acids, producing a gentle stream of bubbles, though their rates increase upon heating:

Zn_{(s)} + 2HCl_{(aq)} \rightarrow ZnCl_2_{(aq)} + H_2_{(g)}

Zn_{(s)} + H_2SO_4_{(aq)} \rightarrow ZnSO_4_{(aq)} + H_2_{(g)}

Fe_{(s)} + 2HCl_{(aq)} \rightarrow FeCl_2_{(aq)} + H_2_{(g)}

Fe_{(s)} + H_2SO_4_{(aq)} \rightarrow FeSO_4_{(aq)} + H_2_{(g)}

Copper sits below hydrogen in the reactivity series and does not react with cold dilute acids.

Metals situated near the top of the reactivity series also react with cold liquid water at room temperature, forming a metal hydroxide and hydrogen gas:

$\text{Metal} + \text{Water} \rightarrow \text{Metal Hydroxide} + \text{Hydrogen}$

Potassium, sodium, lithium, and calcium react vigorously with water. For sodium and calcium, the reactions proceed according to:

2Na_{(s)} + 2H_2O_{(l)} \rightarrow 2NaOH_{(aq)} + H_2_{(g)}

Ca_{(s)} + 2H_2O_{(l)} \rightarrow Ca(OH)_2_{(aq)} + H_2_{(g)}

Magnesium reacts extremely slowly with cold water, whereas zinc, iron, and copper do not react with water at room temperature.

The relative reactivity of metals can also be investigated quantitatively through calorimetry. Acid-metal displacement reactions are exothermic, meaning that the displacement of hydrogen releases thermal energy. When conducting a comparative thermal experiment, strict experimental controls must be maintained: the mass of metal, particle surface area, volume of acid, initial liquid temperature, and acid molar concentration must be kept constant. Under these controlled conditions, the magnitude of the measured temperature rise ($\Delta T$) serves as a direct quantitative indicator of relative reactivity, with the most reactive metals producing the greatest thermal increases.

Extraction of Metals and Reduction by Carbon

Most metals exist naturally within the Earth's crust in chemical combination with oxygen or other non-metals, forming minerals within rock formations known as metal ores. An ore is a geological deposit containing a sufficient concentration of a metal compound to justify economic extraction. The natural spontaneous reaction of exposed metal surfaces with atmospheric oxygen produces metal oxides in an oxidation process; for example, the corrosion of iron yields iron(III) oxide (rust):

4Fe_{(s)} + 3O_2_{(g)} \rightarrow 2Fe_2O_3_{(s)}

2Mg_{(s)} + O_2_{(g)} \rightarrow 2MgO_{(s)}

Only the most unreactive metals, such as gold ($Au$), are found uncombined in their native, elemental state because they do not spontaneously react with atmospheric oxygen.

In chemical terminology involving oxygen transfer, oxidation is defined as the chemical gain of oxygen, whereas reduction is defined as the chemical loss of oxygen. To isolate a pure elemental metal from a mined metal oxide ore, the compound must undergo chemical reduction to remove the bonded oxygen atoms.

The method chosen to extract a metal is governed by the position of the metal relative to carbon in the reactivity series. Carbon is a cheap, readily available reducing agent (derived from coal as industrial coke). A metal oxide can be reduced by heating with carbon if the parent metal is situated below carbon in the reactivity series. In this process, carbon displaces the less reactive metal by taking its oxygen to form carbon dioxide gas, leaving behind the pure metal:

$\text{Metal Oxide} + \text{Carbon} \rightarrow \text{Metal} + \text{Carbon Dioxide}$

During this reaction, the metal oxide loses oxygen and is reduced, while carbon gains oxygen and is oxidized. Industrial examples include the extraction of copper, lead, and iron:

2CuO_{(s)} + C_{(s)} \rightarrow 2Cu_{(s)} + CO_2_{(g)}

2PbO_{(s)} + C_{(s)} \rightarrow 2Pb_{(s)} + CO_2_{(g)}

2Fe_2O_3_{(s)} + 3C_{(s)} \rightarrow 4Fe_{(l)} + 3CO_2_{(g)}

Because carbon is more reactive than zinc, iron, and copper, carbon reduction is the standard, cost-effective extraction method for these elements. In contrast, metals situated above carbon in the reactivity series (potassium, sodium, lithium, calcium, magnesium, and aluminium) cannot be extracted using carbon because carbon cannot displace metals that have a greater affinity for oxygen. These electropositive metals must instead be extracted through molten electrolysis, an energy-intensive and expensive electrical process.

Oxidation, Reduction, and Ionic Half Equations

In comprehensive chemical theory, oxidation and reduction are defined by electron transfer using the mnemonic OIL RIG: Oxidation Is Loss of electrons, Reduction Is Gain of electrons. Because free electrons cannot exist independently in bulk solution, oxidation and reduction occur concurrently in coupled redox reactions: the electrons lost by the oxidised species are accepted by the reduced species.

The reaction between a solid metal and an acid is an example of a redox process. Solid magnesium atoms are oxidised, losing two electrons to form magnesium cations ($Mg^{2+}$), while aqueous hydrogen ions ($H^+$) are reduced, gaining those electrons to form diatomic hydrogen gas ($H_2$):

Mg_{(s)} + 2H^+_{(aq)} \rightarrow Mg^{2+}_{(aq)} + H_2_{(g)}

Displacement reactions represent another category of redox reactions. In a single displacement reaction, an uncombined, more reactive metal displaces a less reactive metal from an aqueous solution containing its salt. For example, solid iron placed into a blue solution of copper(II) sulfate displaces copper, producing an aqueous solution of iron(II) sulfate and solid copper precipitate:

Fe_{(s)} + CuSO_4_{(aq)} \rightarrow FeSO_4_{(aq)} + Cu_{(s)}

To identify the participating chemical species, the complete equation is expanded into an all-ion format:

Fe_{(s)} + Cu^{2+}_{(aq)} + SO_4^{2-}_{(aq)} \rightarrow Fe^{2+}_{(aq)} + SO_4^{2-}_{(aq)} + Cu_{(s)}

Sulfate ions ($SO_4^{2-}$) remain unchanged in aqueous solution on both sides of the reaction. Because they do not participate in electron transfer or alter their oxidation state, they are classified as spectator ions and are eliminated to generate the net ionic equation:

$Fe_{(s)} + Cu^{2+}_{(aq)} \rightarrow Fe^{2+}_{(aq)} + Cu_{(s)}$

In every metal displacement reaction, the elemental metal atom loses electrons (oxidised), while the dissolved metal cation gains electrons (reduced). The overall redox process can be broken down into individual balanced half equations displaying electron loss and gain:

Oxidation half equation: $Fe_{(s)} \rightarrow Fe^{2+}_{(aq)} + 2e^-$

Reduction half equation: $Cu^{2+}_{(aq)} + 2e^- \rightarrow Cu_{(s)}$

Similarly, when calcium metal is introduced to iron(II) sulfate solution, calcium displaces iron because calcium is more reactive:

$Ca_{(s)} + Fe^{2+}_{(aq)} \rightarrow Ca^{2+}_{(aq)} + Fe_{(s)}$

Oxidation half equation: $Ca_{(s)} \rightarrow Ca^{2+}_{(aq)} + 2e^-$

Reduction half equation: $Fe^{2+}_{(aq)} + 2e^- \rightarrow Fe_{(s)}$

When writing balanced half equations, charges are balanced by placing electrons ($e^-$) on the side with the higher positive charge. To combine individual half equations into an overall ionic redox equation, both half equations must be multiplied by integers until their electron counts are equal, allowing the electrons on opposite sides to cancel out completely.

Principles of Electrolysis and Molten Ionic Compounds

Electrolysis is the process of chemically decomposing an ionic compound into its constituent elements by passing a direct electric current (DC) through an electrolyte. An electrolyte is a molten liquid or aqueous solution containing mobile ions capable of conducting electricity. Solid ionic compounds cannot undergo electrolysis because their constituent ions are locked into fixed positions within a rigid three-dimensional crystal lattice and cannot migrate under the influence of an electric field. When melted into a liquid state or dissolved in water, the lattice breaks down, freeing the ions to move towards oppositely charged electrodes.

An electrolysis cell consists of two conducting rods (electrodes) submerged in the electrolyte, connected externally to a DC electrical power supply (such as a battery). Electrodes are constructed from inert conducting materials (typically graphite carbon or platinum) to prevent them from participating chemically in the reaction. The negative electrode is designated the cathode, and the positive electrode is designated the anode. During electrolysis, the electrical potential causes migration of ions:

Positive ions (cations) migrate toward the negative cathode, where they gain electrons (reduction) and are discharged as neutral elements.

Negative ions (anions) migrate toward the positive anode, where they lose electrons (oxidation) and are discharged as neutral elements.

The term discharge refers to the neutralisation of an ion by electron transfer at an electrode surface. The continuous movement of ions through the bulk electrolyte combined with the movement of electrons through the external circuit wiring creates a sustained flow of electrical charge.

In the electrolysis of a molten binary ionic compound containing only two elements, the products are predictable: the metal cation is reduced to elemental metal at the cathode, and the non-metal anion is oxidised to elemental gas or liquid at the anode. For example, during the electrolysis of molten lead(II) bromide (PbBr_2_{(l)}), the molten liquid contains mobile $Pb^{2+}$ cations and $Br^-$ anions:

At the cathode ($-\text{ve}$), lead ions each gain two electrons, undergoing reduction to form molten lead metal, which pools at the bottom of the vessel:

$Pb^{2+}_{(l)} + 2e^- \rightarrow Pb_{(l)}$

At the anode ($+\text{ve}$), bromide ions each lose one electron, undergoing oxidation to form neutral bromine atoms that pair covalently into diatomic molecules, releasing brown bromine vapour:

2Br^-_{(l)} \rightarrow Br_2_{(g)} + 2e^-

Overall, the battery functions as an electron pump, extracting electrons from the bromide ions at the anode and supplying them to the lead ions at the cathode, thereby decomposing the compound into its constituent elements:

PbBr_2_{(l)} \rightarrow Pb_{(l)} + Br_2_{(g)}

Industrial Extraction of Aluminium

Aluminium is situated above carbon in the reactivity series and must be extracted from its purified ore by molten electrolysis. The primary geological ore of aluminium is bauxite, which is chemically refined to yield pure aluminium oxide ($Al_2O_3$, alumina). Solid aluminium oxide has an exceptionally high melting point exceeding $2000\,^\circ\text{C}$. Melting large tonnages of pure $Al_2O_3$ directly would require vast amounts of thermal energy, making industrial production economically unfeasible.

To overcome this economic and thermal barrier, purified aluminium oxide is dissolved in molten cryolite ($Na_3AlF_6$, an aluminium-bearing mineral). The resulting cryolite-alumina mixture melts at approximately $900\,^\circ\text{C}$ to $1000\,^\circ\text{C}$, significantly lowering the operating temperature. This substantial reduction in working temperature conserves vast quantities of thermal energy and decreases electrical heating costs. The molten mixture functions as the electrolyte, containing mobile aluminium cations ($Al^{3+}$) and oxide anions ($O^{2-}$).

The industrial electrolysis cell utilizes massive blocks of carbon (graphite) for both the cathode lining and the suspended anode rods. Direct current supplied to the cell drives the following electrode processes:

At the negative cathode lining ($-\text{ve}$), aluminium cations undergo reduction by gaining three electrons each, producing liquid aluminium metal:

$Al^{3+}_{(l)} + 3e^- \rightarrow Al_{(l)}$

Liquid aluminium is denser than the molten cryolite electrolyte, so it sinks to the base of the cell, where it is periodically tapped off as molten metal.

At the positive carbon anodes ($+\text{ve}$), oxide anions undergo oxidation by losing two electrons each, discharging to form diatomic oxygen gas molecules:

2O^{2-}_{(l)} \rightarrow O_2_{(g)} + 4e^-

The overall balanced equation for the cell reaction is:

2Al_2O_3_{(l)} \rightarrow 4Al_{(l)} + 3O_2_{(g)}

A significant operational consequence of this configuration is the degradation of the carbon anodes. Because the operating temperature exceeds $900\,^\circ\text{C}$, the oxygen gas discharged at the positive electrodes reacts with the hot carbon (graphite) blocks to produce carbon dioxide gas:

C_{(s)} + O_2_{(g)} \rightarrow CO_2_{(g)}

This continuous combustion gradually erodes the carbon anodes, requiring them to be replaced regularly, which adds to the operating costs of aluminium extraction.

Electrolysis of Aqueous Solutions

The electrolysis of aqueous ionic solutions is more complex than that of molten salts due to the presence of water molecules. Water undergoes a small degree of self-ionisation, establishing a dynamic equilibrium:

$H_2O_{(l)} \rightleftharpoons H^+_{(aq)} + OH^-_{(aq)}$

Consequently, an aqueous electrolyte always contains four distinct ionic species: the metal cation and non-metal anion from the dissolved salt, alongside hydrogen ions ($H^+$) and hydroxide ions ($OH^-$) from the water. Because multiple ions migrate toward each electrode, preferential discharge rules determine which species reacts.

At the negative cathode ($-\text{ve}$), positive metal cations ($M^{n+}$) and hydrogen ions ($H^+$) compete for reduction. The outcome depends on the relative position of the metal compared to hydrogen in the reactivity series:

If the metal is more reactive than hydrogen (such as potassium, sodium, lithium, calcium, magnesium, aluminium, zinc, or iron), the $H^+$ ions are preferentially discharged, producing bubbles of hydrogen gas:

2H^+_{(aq)} + 2e^- \rightarrow H_2_{(g)}

If the metal is less reactive than hydrogen (such as copper, silver, or gold), the metal cations are preferentially discharged, depositing a solid layer of pure metal over the cathode surface:

$Cu^{2+}_{(aq)} + 2e^- \rightarrow Cu_{(s)}$

At the positive anode ($+\text{ve}$), non-metal anions and hydroxide ions ($OH^-$) compete for oxidation. The outcome depends on whether a halide ion is present:

If halide ions ($Cl^-$, $Br^-$, $I^-$) are present in the solution, the halide ions are preferentially discharged, forming elemental halogen molecules ($Cl_2$, $Br_2$, or $I_2$):

2Cl^-_{(aq)} \rightarrow Cl_2_{(g)} + 2e^-

2Br^-_{(aq)} \rightarrow Br_2_{(aq/g)} + 2e^-

If no halide ions are present (such as when the salt contains polyatomic anions like sulfate, $SO_4^{2-}$, or nitrate, $NO_3^-$), the hydroxide ions ($OH^-$) from water are preferentially discharged, yielding oxygen gas, water, and releasing electrons:

4OH^-_{(aq)} \rightarrow O_2_{(g)} + 2H_2O_{(l)} + 4e^-

Consider the electrolysis of aqueous copper(II) sulfate (CuSO_4_{(aq)}). The solution contains four ions: $Cu^{2+}$, $SO_4^{2-}$, $H^+$, and $OH^-$. At the cathode, copper is less reactive than hydrogen, so $Cu^{2+}$ ions are discharged to form a salmon-pink solid copper coating ($Cu^{2+} + 2e^- \rightarrow Cu$). At the anode, sulfate is not a halide, so $OH^-$ ions are discharged, producing oxygen gas effervescence and water ($4OH^- \rightarrow O_2 + 2H_2O + 4e^-$).

Consider the electrolysis of aqueous sodium chloride ($NaCl_{(aq)}$, brine). The solution contains $Na^+$, $Cl^-$, $H^+$, and $OH^-$. At the cathode, sodium is more reactive than hydrogen, so $H^+$ ions are discharged to yield hydrogen gas bubbles ($2H^+ + 2e^- \rightarrow H_2$). At the anode, chloride is a halide, so $Cl^-$ ions are discharged to produce pale green chlorine gas bubbles ($2Cl^- \rightarrow Cl_2 + 2e^-$).

To capture and test the gases produced during aqueous electrolysis, inverted test tubes or gas collection burettes filled with electrolyte are mounted over the inert electrodes. Standard laboratory diagnostic tests identify the resulting gases:

Hydrogen gas ($H_2$): Holding a burning wooden splint near the mouth of the collection tube produces a distinctive squeaky pop.

Oxygen gas ($O_2$): Introducing a glowing wooden splint into the collected gas results in the splint relighting.

Chlorine gas ($Cl_2$): Placing moist blue litmus paper into the gas causes the paper to turn temporarily red before being bleached white.

Review Questions and Worked Solutions

Question: A student uses universal indicator to test the pH of some lemon juice. What colour would you expect the indicator to turn?

Solution: Lemon juice contains citric acid and has an acidic pH of approximately $2$ to $3$. Universal indicator displays a red or orange-red colour in solutions within this pH range.

Question: The pH of an unknown solution is found to be $8$. Is the solution acidic or alkaline?

Solution: The solution is alkaline, because any aqueous solution with a pH value strictly greater than $7$ is alkaline.

Question: How do you determine when the end-point of a titration has been reached?

Solution: The end-point is identified by an immediate, permanent colour change of the single indicator upon the addition of a single drop of titrant.

Question: Give two reasons why you should repeat a titration until you have consistent readings.

Solution: First, repeating titrations allows the identification and exclusion of anomalous results to ensure accuracy. Second, obtaining concordant results (titres within $0.1\,cm^3$ of one another) allows the calculation of an accurate, representative mean titre volume.

Question: Name a strong acid.

Solution: Hydrochloric acid ($HCl$), sulfuric acid ($H_2SO_4$), or nitric acid ($HNO_3$).

Question: A student added strong acid to a weakly acidic solution of pH $6$. The pH of the new solution was found to be pH $3$. By how many times did the concentration of $H^+$ increase?

Solution: The decrease in pH is $6 - 3 = 3$ pH units. Because each unit decrease corresponds to a ten-fold ($10\times$) increase in hydrogen ion concentration, the factor increase is calculated as $10^3 = 1000$. The concentration of $H^+$ ions increased by $1000$ times.

Question: Calcium carbonate is added to hydrochloric acid. Write the word equation and the balanced symbol equation for the reaction that occurs.

Solution: The word equation is:

$\text{calcium carbonate} + \text{hydrochloric acid} \rightarrow \text{calcium chloride} + \text{water} + \text{carbon dioxide}$

The balanced symbol equation is:

$CaCO_{3(s)} + 2HCl_{(aq)} \rightarrow CaCl_{2(aq)} + H_2O_{(l)} + CO_{2(g)}$

Question: Give the balanced equation, including state symbols, for the reaction of sodium and water.

Solution: Sodium reacts with water to produce aqueous sodium hydroxide and hydrogen gas:

$2Na_{(s)} + 2H_2O_{(l)} \rightarrow 2NaOH_{(aq)} + H_{2(g)}$

Question: Write a balanced equation for the reduction of lead oxide, $PbO$, by carbon, $C$.

Solution: Carbon reduces lead(II) oxide to elemental lead, oxidising to carbon dioxide:

$2PbO_{(s)} + C_{(s)} \rightarrow 2Pb_{(s)} + CO_{2(g)}$

Question: Explain why calcium cannot be extracted from calcium oxide by reduction using carbon.

Solution: Calcium is higher in the reactivity series than carbon. Because carbon is less reactive than calcium, it cannot displace calcium from its oxide bonds; calcium can only be extracted via electrolysis.

Question: Zinc reacts with acid in a redox reaction according to the ionic equation:

$Zn_{(s)} + 2H^+_{(aq)} \rightarrow Zn^{2+}_{(aq)} + H_{2(g)}$

State which species is being oxidised and which is being reduced, and write a balanced half equation for the oxidation reaction.

Solution: Zinc ($Zn$) is oxidised because each neutral atom loses two electrons to form a $Zn^{2+}$ cation. Hydrogen ions ($H^+$) are reduced because each ion gains an electron to form neutral hydrogen atoms in $H_2$. The balanced oxidation half equation is:

$Zn_{(s)} \rightarrow Zn^{2+}_{(aq)} + 2e^-$

Question: The equation for the reaction of zinc and iron sulfate is:

$Zn_{(s)} + FeSO_{4(aq)} \rightarrow ZnSO_{4(aq)} + Fe_{(s)}$

Write an ionic equation for the reaction, and state which species is being reduced and which is being oxidised.

Solution: Sulfate ($SO_4^{2-}$) acts as a spectator ion. Eliminating spectator ions gives the ionic equation:

$Zn_{(s)} + Fe^{2+}_{(aq)} \rightarrow Zn^{2+}_{(aq)} + Fe_{(s)}$

Zinc ($Zn$) is oxidised (loses electrons to become $Zn^{2+}$), and iron(II) ions ($Fe^{2+}$) are reduced (gain electrons to become $Fe$).

Question: A student carries out electrolysis on molten calcium chloride. What is produced at the anode and the cathode?

Solution: At the anode ($+\text{ve}$), chloride ions are oxidised to chlorine gas ($Cl_2$). At the cathode ($-\text{ve}$), calcium ions are reduced to molten calcium metal ($Ca$).

Question: An aqueous solution of copper bromide, $CuBr_2$, is electrolysed using inert electrodes. State what is produced at the anode and the cathode.

Solution: At the anode ($+\text{ve}$), bromide is a halide, so it is discharged to produce bromine ($Br_2$). At the cathode ($-\text{ve}$), copper is less reactive than hydrogen, so copper cations are discharged to produce solid copper metal ($Cu$).