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A catalytic surface site is a location on a solid surface capable of participating in a catalytic route. The schematic notation $\ast$ denotes a vacant site; $\mathrm{A}\ast$ denotes A bound there. A minimal illustrative catalytic cycle isThe first step represents adsorption; the second combines conversion and release. Cancelling $\ast$ and $\mathrm{A}\ast$ gives $\mathrm{A}\rightleftharpoons \mathrm{B}$. The site is regenerated even though its occupancy changes during the cycle. This is an original bookkeeping model, not a measured reaction mechanism for manganese dioxide.

If the only site states are vacant and occupied by A, the site inventory obeys $n_{\ast }+n_{\mathrm{A}\ast }=n_{\mathrm{sites}}$. A blocked or inaccessible part of a particle does not contribute available catalytic surface sites merely because it contributes mass. Different surface structures can have different abilities to bind and convert reactants; equal geometric surface area need not mean equal catalytic activity.

In heterogeneous catalysis, the catalyst and reactants occupy different thermodynamic phases; a solid surface can furnish the catalytic route.

For example, specify dissolved hydrogen peroxide as the reactant, water as the solvent, and solid manganese dioxide as the catalyst. The solid and liquid chemical solution are separate thermodynamic phases. A dispersion that looks thoroughly mixed still contains those separate phases; mixing distributes contact, rather than making the solid molecularly dissolved. The manganese dioxide peroxide contact demonstration makes the arrival of the liquid at the solid tangible.

A useful surface-route picture is approach from the liquid, adsorption at a catalytic surface site, conversion through surface-bound reaction intermediates, and release of reaction products. ^[ref] This sequence organizes possible roles; it is not a verified sequence of elementary reactions for every catalyst, or for the particular manganese dioxide demonstration. The net chemical equation supplies neither the binding geometry nor the identity of the working catalytic surface site.

This contact requirement makes surface area and diffusion practically important. More exposed solid can provide more catalytic surface sites, whereas diffusion through the surrounding liquid can limit their supply of reactants. Changing stirring or particle size can therefore change the observed reaction rate without identifying the microscopic rate constant. A recoverable solid is attractive for reuse, but separation alone does not establish retained activity: compare with catalyst deactivation.

In homogeneous catalysis, the catalyst and reactants occupy the same thermodynamic phase.

For a liquid example, specify hydrogen peroxide as the reactant, dissolved iodide as the catalyst, and water as the solvent. If both reacting chemical species remain dissolved in that one liquid thermodynamic phase, this is homogeneous catalysis. The University of Wisconsin–Madison's The Rate of a Reaction identifies iodide-catalysed hydrogen peroxide decomposition in its final example; the aqueous setting here makes the phase roles explicit.^[ref] Its net chemical equation isThe molecular oxygen is a reaction product: its bubbles do not make the dissolved catalytic route heterogeneous catalysis. The classification compares the participating catalyst and reactants, rather than demanding that everything in the vessel share a phase.^[ref]

Being in the same thermodynamic phase describes where the chemical species are; it does not specify their reaction mechanism. A uniform-looking liquid can contain unresolved solid particles, so appearance alone is insufficient. Conversely, suspended solid particles do not establish that their surfaces supply the observed catalysis: dissolved material released from them could contribute. Phase identification and catalytic participation need separate evidence.

The bookkeeping of a catalyst in homogeneous catalysis is particularly transparent in the schematic catalytic cycle A plus C forming AC, followed by AC releasing B plus C. This is the generic model in catalyst, not an asserted iodide reaction mechanism. Suppose a fixed volume initially contains $1.00\mathrm{mmol}$ of A and $0.20\mathrm{mmol}$ of C. At a chosen snapshot, let $0.60\mathrm{mmol}$ of B and $0.12\mathrm{mmol}$ of AC be present, with no other chemical reactions. Each AC contains one A unit and one C unit, soThus free A is $0.28\mathrm{mmol}$ and free C is $0.08\mathrm{mmol}$. B formation already exceeds the initial amount of substance of C by a factor of three, while the total C inventory remains $0.20\mathrm{mmol}$. Adding the two steps cancels C and AC, giving the net A-to-B chemical reaction; counting their constituents proves these balances. Conserved total catalyst does not mean constant free catalyst concentration.

Dissolved contact is useful because reactants can encounter the catalytic chemical species throughout the mixed liquid, without first reaching a macroscopic solid surface. Recovering a dissolved catalyst from the same chemical solution can instead require a separation. Neither ease of contact nor phase classification determines an activation energy or an equilibrium constant; those depend on the relevant chemical kinetics and endpoint thermodynamics.

The University of Washington's Decomposition of H₂O₂ with MnO₂ describes a supervised demonstration using solid manganese dioxide and hydrogen peroxide listed as 30%. Its cylinder version uses a spatula to introduce the solid; the stated readout is immediate molecular oxygen bubbles and heat production. An alternate bottle version replaces the contents of a tea bag with the solid and suspends it above the liquid. Releasing the bag brings it into contact with the liquid. The bag-above-liquid stage supplies a before-contact comparison; the page gives no timed gas collection, replicated blank trials or quantitative reaction rates.^[ref]

The balanced net chemical equation is $2{\mathrm{H}}_2{\mathrm{O}}_2⟶2{\mathrm{H}}_2\mathrm{O}+{\mathrm{O}}_2$. Appearance of gas after contact supports a contact-dependent acceleration of decomposition under these conditions. It does not distinguish surface conversion from a contribution of dissolved material, establish regeneration, or measure an activation energy. The simultaneous heat release also means that temperature can change after contact. A matched no-solid comparison and an empty-bag comparison would test background decomposition and the bag's contribution; these are proposed controls, not reported trials. The catalysed peroxide comparison supplies a complementary inspected teaching experiment with separate single-addition and paired-addition comparisons.

Oxygen can be obtained by decomposing an oxygen-rich chemical compound, or by separating it from air. For the historical laboratory, heating potassium chlorate, heating certain metal oxides, and electrolysis of water are useful examples. The substance collected is molecular oxygen, ${\mathrm{O}}_2$.

A common preparation heats potassium chlorate, often with manganese dioxide as a catalyst:A heat-resistant vessel and delivery tube conduct the evolved gas to an inverted collecting vessel over water, after the initial displaced air has been excluded. Oxygen is only slightly soluble in water, making this collection principle useful. The manganese dioxide changes the reaction rate and required heating conditions; it is not the net source of the oxygen in this catalysed balance. Without the catalyst, potassium perchlorate can occur as an intermediate.

Two older preparations use heated metal oxides:The first explains the historical decomposition of mercury(II) oxide, with metallic Mercury deposited on cooler parts of the apparatus. It is historical chemistry, not a preferred present-day teaching preparation. The second is one stage in the reduction of manganese dioxide; at still higher temperature further oxygen loss can give ${\mathrm{Mn}}_3{\mathrm{O}}_4$, so the oxide residue depends on the conditions. Strong heating of potassium nitrate provides another net balance:In electrolysis, suitable separated electrodes collect oxygen at the anode and hydrogen at the cathode. With a suitable electrolyte and electrodes, the net chemical equation isThe gases must be kept in separate collecting spaces if the experiment is to demonstrate their separate identities and the approximately $2:1$ volume ratio. Choice of electrolyte matters: a chloride-containing electrolyte can give chlorine instead of clean oxygen evolution.

For a modern small-scale comparison, catalytic decomposition of hydrogen peroxide givesThis is a different feedstock from potassium chlorate. Large-scale separation of oxygen from air by cryogenic fractionation belongs to later industrial practice and should not be substituted for an 1866 laboratory answer. A collected sample can be tested by its vigorous support of combustion, conventionally the relighting of a glowing splint; this observation identifies a property, while the preparation and exclusion of interfering gases establish the sample's provenance.

Iodine is a dark, lustrous crystalline halogen which gives a violet vapour; historically, an important source was the ash of iodine-rich seaweeds. Elemental iodine commonly occurs as diatomic ${\mathrm{I}}_2$. It melts at about $114{}^{\circ }\mathrm{C}$ and boils at about $184{}^{\circ }\mathrm{C}$ at ordinary atmospheric pressure. It also undergoes sublimation below its melting point, depositing crystals again on a cooler surface.^[ref]

Iodine is sparingly soluble in water, giving a brownish solution, and dissolves more freely in ethanol and in an aqueous iodide solution. In the latter case the chemical equilibriumforms triiodide. Iodine with starch, under suitable conditions, gives a sensitive dark-blue colour; this is a useful recognition test, but oxidants that liberate iodine from added iodide must be distinguished from iodine initially present. Reaction with reactive metals gives iodides; in such compounds the iodine usually has oxidation state $-1$.

For the historical seaweed route, collect and burn iodine-accumulating seaweeds to produce kelp ash, then extract the soluble salts with water. Crystallization removes much of the common salt content, leaving an iodide-enriched mother liquor. Oxidation releases iodine, which can be distilled or sublimed and collected on a cool surface. One representative net chemical equation for liberation from iodide isThe manganese dioxide is the oxidizing agent; the seaweed does not contain a supply of ready-made violet iodine crystals. This describes the chemical principle of a historical extraction, rather than prescribing one unique factory recipe.

Iodine also occurs in natural brines and in iodate-bearing deposits associated with nitrate ores. Modern industrial supply relies heavily on brines and iodate ores; it should not be equated with the seaweed source emphasized in nineteenth-century teaching.^[ref]

A proposed modern teaching measurement puts a pipette-measured portion of an iron ${\mathrm{Fe}}^{2+}$ chemical solution in a conical flask, acidifies with excess dilute sulfuric acid, and adds standardized potassium permanganate from a clamped burette while swirling. The documented endpoint principle is that added purple permanganate is consumed until the first small excess leaves a pale purple colour; no separate colour indicator is required. Insufficient acidity can produce brown manganese dioxide, invalidating the assumed five-electron half-reaction.^[ref]

Record initial and final burette readings; their difference is the delivered volume $V$. The following controls are proposed additions to test the inference. Run a reagent blank with the same water, sulfuric acid and endpoint criterion but no sample, obtaining $V_b$. Use a separately established potassium permanganate concentration $c$, clear the burette tip of air, and repeat sample portions. A known added amount of substance of ${\mathrm{Fe}}^{2+}$ should increase consumption by one fifth of that amount. A blank estimates reagent demand and the small endpoint excess; it cannot remove an interfering sample reactant.

Assuming complete, selective reaction according to the acidic permanganate-iron balance, the corrected inference isWith $c$ in $\mathrm{mol}{\mathrm{L}}^{-1}$, use $V$ and $V_b$ in litres. As an original stipulated calculation, not an experimental result, take $c=0.02000\mathrm{mol}{\mathrm{L}}^{-1}$, $V=12.60\mathrm{mL}$ and $V_b=0.10\mathrm{mL}$. Then the permanganate consumption is $0.2500\mathrm{mmol}$ and the inferred ${\mathrm{Fe}}^{2+}$ amount is $1.250\mathrm{mmol}$. In a $25.00\mathrm{mL}$ sample portion this corresponds to $0.05000\mathrm{mol}{\mathrm{L}}^{-1}$.

In this example the corrected volume is $12.50\mathrm{mL}=0.01250\mathrm{L}$. The acidic permanganate-iron balance gives $z_A=1$ for ${\mathrm{Fe}}^{2+}$ and $z_T=5$ for permanganate. Thus $0.2500\mathrm{mmol}$ of permanganate accepts $1.250\mathrm{mmol}$ of formal electrons, exactly what $1.250\mathrm{mmol}$ of ${\mathrm{Fe}}^{2+}$ supplies. The amount of substance ratio is five to one, despite equality of the exchanged electron amounts.

The denominator is the original $25.00\mathrm{mL}$ aliquot, not its final volume after adding sulfuric acid, water and titrant. Such dilution changes concentrations in the flask without changing the initial sampled amount of substance of ${\mathrm{Fe}}^{2+}$. Omitting the $0.10\mathrm{mL}$ titration blank would instead give $1.260\mathrm{mmol}$ and $0.05040\mathrm{mol}{\mathrm{L}}^{-1}$, an overestimate of $0.80\%$ relative to the corrected constructed value. The displayed digits express the stipulated calculation, rather than establish its measurement uncertainty.

The endpoint detects residual permanganate, not the iron identity directly. Other sample reactants that consume permanganate raise the apparent ${\mathrm{Fe}}^{2+}$ amount; prior oxidation of sample ${\mathrm{Fe}}^{2+}$ to ${\mathrm{Fe}}^{3+}$ lowers it. A recovery check probes response to the added iron but does not prove absence of every interference. The readout determines the reacting ${\mathrm{Fe}}^{2+}$ inventory under the stated conditions, rather than total iron in an arbitrary mixture. No measured titres are asserted here.