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Chemical alteration of silicate minerals by water and dissolved carbon dioxide. Coupled to transport and subsequent carbonate burial, this provides a long-term geological sink for added carbon. Its response is much slower than atmospheric mixing or upper-ocean exchange; a rapid pulse can persist while this adjustment proceeds.

Release of water-rich fluid by reactions that destabilize hydrated minerals in a descending slab. It can promote flux melting of the mantle wedge and contribute to dehydration embrittlement. The released water is not evidence that the whole slab has melted.

Two equal circular rings of radius $R$ in parallel planes at $z=\pm L$ provide a boundary for an open film. In the ideal axisymmetric model, a stationary connected film has the catenoid profileThe catenoid existence threshold for equal rings explains why this equation can have two solutions, one solution, or none. The cancellation of principal curvatures holds for both solutions whenever they exist; the mean curvature equation alone does not select a stable minimal surface.

This ideal geometry uses $R=1$ and $L/R=0.60$. The two profiles have different waist radii $a$, although both have $H=0$. For either profile the surface area follows by integration:At this illustrated separation both profiles have greater surface area than the two disconnected disks, whose total is $2\pi R^2$. This comparison disproves a claim of global area minimization in a class allowing disconnected competitors; it does not determine the sign of the second variation for small variations of a connected film.

A reported experiment used coaxial wire rings of radius $4\mathrm{cm}$, drawn from a surfactant solution and slowly separated until collapse. The solution contained tetradecyltrimethylammonium bromide in deionized water, with glycerol added to vary viscosity. Fluorescein enabled fluorescence imaging under cyan LED arrays; a Phantom V641 camera recorded 4,000 or 5,600 frames per second. Extracted contours gave neck radius, attachment angle, and inferred surface area. They showed neck narrowing and eventual breakup, rather than indefinite persistence of a stationary catenoid. Comparing solutions across viscosities helped distinguish liquid dissipation from the surrounding air's inertia, but the ring displacement beyond the last stable state was not precisely controlled. These are collapse observations, not direct measurements of equilibrium $H$. ^[ref]

A proposed classroom measurement would instead hold two equal rings coaxial on a translation stage, leave both sides open to the same gas, and photograph the settled silhouette beside a scale. Fit $r(z)=a\cosh (z/a)$ and inspect fit residuals before inferring cancellation of the principal curvatures. Repeat at several separations and with the camera rotated about the axis to test rotational symmetry. Keep solution composition, ring wetting, and elapsed drainage time comparable, and shield the film from drafts. Film weight, drainage, evaporation, surfactant gradients, and rupture can invalidate the ideal soap film pressure balance. A silhouette fit alone neither measures the pressure jump nor tests all possible perturbations of a stable minimal surface.

A fluid colloidal system. A solid-in-liquid example is fine silicon dioxide dispersed in water. The term can also include other fluid colloidal systems, so the dispersed and continuous constituents should be specified.

Solubility specifies the amount that can remain dissolved at chemical equilibrium under stated conditions; it is different from how quickly a sample dissolves. Separating ions from an ionic crystal costs energy, but solvation establishes new interactions with the solvent. In water, hydration of ions couples charged ions to polar water molecules. The net enthalpy of dissolution can have either sign; entropy and concentration also enter the Gibbs free energy balance. Consequently, being a salt does not guarantee high solubility in water or any other solvent.^[ref]

Low solubility must also be distinguished from weak electrolyte behavior. A small amount of an ionic compound can dissolve essentially completely as ions, giving few charge carriers because little material dissolved. Calling that a weak electrolyte solely because its chemical solution conducts poorly would conflate two different questions.^[ref]

A solubility product constrains a product of thermodynamic activities; the available inventories determine which combinations of concentrations are possible. Their intersection determines the molar solubility. Consider an original, abstract AB solid, with the preceding one-to-one dissolution chemical equation. Stipulate $c^{\circ }=1\mathrm{M}$, ${\gamma }_{\mathrm{A}}={\gamma }_{\mathrm{B}}=1$ and $K_{\mathrm{sp},c}=10^{-4}{\mathrm{M}}^2$ at fixed temperature and pressure. Thus $K_{\mathrm{sp}}=10^{-4}$ is dimensionless. These are chosen model parameters, not data for an identified material.

Initially, a completely dissociated soluble XB salt supplies B anions at molar concentration $c_0\ge 0$ and X cations at the same molar concentration; A is absent. All three ions have unit electric charge magnitude. X remains a spectator ion, and only AB can form a solid. Assume constant volume, no additional chemical reactions, and negligible contributions from self-ionization of water. Dissolving AB raises both reacting-ion concentrations by the same $s$:Charge neutrality is satisfied: $c_{\mathrm{A}}+c_{\mathrm{X}}-c_{\mathrm{B}}=s+c_0-(s+c_0)=0$. Adding B alone without its accompanying cation would omit the initial electric charge balance.

In the plane of the two reacting-ion concentrations, this inventory constraint is the line $c_{\mathrm{B}}=c_{\mathrm{A}}+c_0$. Saturation supplies the hyperbola $c_{\mathrm{B}}=K_{\mathrm{sp},c}/c_{\mathrm{A}}$ for $c_{\mathrm{A}}>0$. Their positive intersection solves a quadratic equation:The quadratic formula supplies two roots of a polynomial; only the positive one is a physical concentration. Uniqueness also follows because $s(s+c_0)$ strictly increases for $s\ge 0$, starting at zero and growing without bound. The second expression, obtained by multiplying by the conjugate, avoids subtraction of nearly equal numbers at large $c_0$.

For pure water, $c_0=0$ gives $s_{\ast }=\sqrt{K_{\mathrm{sp},c}}=0.010\mathrm{M}$. For $c_0=0.030\mathrm{M}$,Their product is $10^{-4}{\mathrm{M}}^2$, although neither concentration equals the pure-water value. This is the useful geometrical distinction: an inventory line and an equilibrium constant specify different constraints.

Check the solid inventory before imposing saturation. If the initially added AB has amount of substance $n_0$ in volume $V$, the saturated branch requires $n_0>V{s}_{\ast }$ for excess solid to remain. For example, choose $n_0=0.0020\mathrm{mol}$ and $V=0.100\mathrm{L}$. In pure water, $0.0010\mathrm{mol}$ dissolves and the same amount of substance remains as solid. At $c_0=0.030\mathrm{M}$, approximately $0.0003027756\mathrm{mol}$ dissolves, leaving $0.0016972244\mathrm{mol}$ as solid. In this model, an insufficient supply instead givesIf $n_0/V<s_{\ast }$, all AB dissolves and the ion product is below the solubility product. At $n_0/V=s_{\ast }$, the last solid disappears exactly at saturation.

How can an invisible dissolved inventory be measured? The unsigned laboratory handout Solubility Product Constant, Ksp, for Calcium Hydroxide by Acid-Base Titration, Procedure, starts with saturated calcium hydroxide, supplied by the teacher or prepared from calcium and water. Its apparatus includes filter paper, a funnel, stoppered flasks, a volumetric pipette, a volumetric flask and a burette. Two successive rounds of filtration remove particles. The filtrate is stoppered to reduce carbon dioxide exposure. A $10.00\mathrm{mL}$ aliquot is diluted with deionized water and titrated with standardized hydrochloric acid; bromothymol blue gives a pale-green titration endpoint. The readout is the difference between final and initial burette readings; four consistent trials are required.^[ref]

Unlike AB, calcium hydroxide has one-to-two dissolution stoichiometry:For molar solubility $s$ in the ideal model with no background calcium or hydroxide, no additional chemical reactions, and negligible self-ionization of water, $[{\mathrm{Ca}}^{2+}]=s$ and $[{\mathrm{OH}}^{-}]=2s$. Hence $K_{\mathrm{sp},c}=4s^3$, with units ${\mathrm{M}}^3$; it is not the AB constant used in the figures. The neutralization chemical equation is ${\mathrm{H}}^{+}+{\mathrm{OH}}^{-}⟶{\mathrm{H}}_2\mathrm{O}$. Let $C_H$ be the known titrant molar concentration, $v$ its delivered volume at the equivalence point, and $V_a$ the original aliquot volume. Then the inventory calculation isDilution of the aliquot changes its concentration during the acid-base titration, but preserves its initial amount of substance of hydroxide. Therefore $V_a$, rather than the diluted flask volume, belongs in the denominator.

The handout identifies two biases: carried-over calcium hydroxide solid dissolves during acid-base titration, consuming extra titrant; carbon dioxide consumes hydroxide and forms calcium carbonate. Filtration and stoppering address these effects, while repeat trials test consistency. Its review questions also ask about dissolved ion pairs, which filtration cannot remove.^[ref]

Further controls would compare a water blank, independently verify titrant standardization, and record the saturation temperature; these are proposed checks, not reported measurements. A titration endpoint estimates the equivalence point and may introduce a systematic offset. Even perfect inventory measurement does not determine activity coefficients. Recovering the dimensionless thermodynamic solubility product additionally requiresThis explains both the power and the limit of the experiment: acid-base titration measures a dissolved inventory, while the link from that inventory to a thermodynamic equilibrium constant still depends on a chemical model.

Solute potential is the contribution to water potential from dissolved material, with pure water assigned zero. In the nearly incompressible description, ${\Psi }_s=(RT/{\bar{V}}_w)\ln a_w$. In a dilute ideal chemical solution, ${\Psi }_s\simeq -RT{c}_{\mathrm{osm}}$, where $c_{\mathrm{osm}}$ is the osmotic concentration of dissolved particles. Thus solute potential is negative while the corresponding osmotic pressure is positive: ${\Psi }_s\simeq -\Pi$. This sign difference distinguishes a contribution to water potential from the compensating mechanical pressure.^[ref]

A chemical solution is a homogeneous mixture in which designated solutes are distributed in a solvent. Representative equal-volume samples have the same composition after mixing. Individual molecules and ions remain distinguishable. The solute/solvent labels describe their roles in this mixture.^[ref]

A dissolved solute is dispersed as molecules, chemical atoms, or ions, rather than as the larger solid particles of a particle suspension. For example, sucrose retains its molecular formula during dissolution in water; a chemical reaction is not required. Dissolved sodium chloride supplies sodium cations and chloride anions, surrounded by water molecules through solvation. Dissolved does not mean chemically destroyed or converted into water.^[ref]

This makes a useful connection between microscopic identity and laboratory sampling. If a chemical solution is well mixed, the amount of substance in a measured portion follows from its concentration and volume. A scoop from a poorly mixed particle suspension need not represent the whole vessel. Visual clarity alone cannot establish composition or exclude a fine dispersion; the quantitative claim needs representative sampling.

The denominator of molar concentration is the final volume of the chemical solution, including the dissolved solute. It is not the volume of solvent initially poured in. A volumetric flask fixes the final total volume; weighing the solute supplies its mass, and a specified molar mass converts that mass to amount of substance.^[ref]

Consider an original illustrative account, not reported measurements. Suppose $2.922\mathrm{g}$ of pure dry sodium chloride, with stipulated molar mass $58.44\mathrm{g}\mathrm{mo}{\mathrm{l}}^{-1}$, is completely dissolved and quantitatively transferred into a chemical solution whose final volume is $250.0\mathrm{mL}$. Assume no losses, no additional chemical reactions, and complete dissociation of an ionic compound. ThenHere analytical concentration counts the sodium chloride inventory, not intact dissolved $\mathrm{NaCl}$ molecules. Under these assumptions, the amount of substance of sodium cations and of chloride anions is separately $0.05000\mathrm{mol}$; each has molar concentration $0.2000\mathrm{mol}{\mathrm{L}}^{-1}$. Their combined ionic amount of substance is $0.1000\mathrm{mol}$, but this does not double either the salt inventory or the mass. Each sodium cation still contains a sodium atomic nucleus, and each chloride anion a chlorine atomic nucleus. Counting chemical elements survives dissociation of an ionic compound.

The mass ledger is different from the volume ledger. If the final chemical solution had an illustrative net mass of $252.922\mathrm{g}$, the water contribution would be $252.922-2.922=250.000\mathrm{g}$, assuming nothing entered or escaped except those two components. The ratio $2.922/252.922\approx 0.011553$ expresses the solute share of total mass; the separate mass density would be $252.922/250.0=1.011688\mathrm{g}\mathrm{m}{\mathrm{L}}^{-1}$ before rounding. These stipulated quantities are mutually consistent bookkeeping inputs, not a measured mass density for saltwater. Neither establishes that $250.0\mathrm{mL}$ of water was added.

More generally, with full transfer and conserved solute, $n=m/M$ and $c=m/(MV)$. No additive volume or ideal chemical solution assumption is needed when $V$ is the measured final volume. If only a fraction $f$ of the weighed solute reaches the flask, the actual result is $c=fm/(MV)$. This short calculation explains why an undissolved residue or an incomplete rinse creates a systematic deficit even when the final calibration mark is read correctly.

A calibration mark can be tested against an independent mass measurement. The Buffalo State-hosted Laboratory Experiment 1. Calibration of Volumetric Glassware uses a $25\mathrm{mL}$ pipette, filler, conical flask, weighing bottle, balance, and thermometer. It specifies clean, freely draining glassware, bubble-free filling, an eye-level liquid meniscus reading, gravity drainage, and retaining the residual liquid allowed for by the pipette's calibration. Degassed water is delivered to a tared bottle; ten weighings and initial/final temperature readings give repeated delivered volumes using the water mass density. Their mean and standard deviation assess delivery and repeatability. The balance and thermometer must themselves be checked before claiming accuracy. The same handout describes weighing an empty volumetric flask and then weighing it filled to its mark.^[ref]

The independent inference is $V=m_{\mathrm{water}}/{\rho }_{\mathrm{water}}(T)$. Agreement between repeats constrains scatter; agreement with the nominal delivery additionally needs trustworthy mass density, temperature, and balance readings. This experiment tests delivered or contained volume under its conditions. It does not prove that a subsequent solute transfer is complete or a prepared chemical solution is homogeneous.

The University of Wisconsin–Madison teaching module supplies the complementary preparation demonstration. Dissolved solute and vessel rinsings reach the volumetric flask; the contents reach room temperature before a dropper adjusts the eye-level liquid meniscus to the calibration mark. The module excludes a stirring bar during filling because its displaced volume is absent from the calibration, and requires mixing the neck liquid with the bulk chemical solution. Its self-check identifies incomplete dissolution and transfer as a cause of lower-than-intended molar concentration.^{[ref][ref][ref]}

The observable endpoints are a liquid boundary at a mark and a delivered mass, not a direct count of solute particles. An overfilled flask has excessive final volume; removing some well-mixed chemical solution also removes solute, so it cannot recover the intended inventory merely by restoring the mark. Measurement uncertainty, contamination, incomplete transfer, and sampling before full mixing limit the inferred molar concentration. The apparatus demonstrates why the words “dilute to” refer to final chemical solution volume.

Solvation is stabilization through interactions between a solute and its solvent. It involves both the solute's surroundings and reorganization of the solvent: neighbors move, orientations change, and existing interactions compete with new ones. Electric forces, Van der Waals forces, and specific interactions such as hydrogen bonds can contribute. Solvation can also occur at groups on an insoluble material, so it does not require complete dissolution.^[ref]

Solvation by water is called hydration. Hydration of ions is the charged special case, but neutral molecules such as sucrose can also interact with water without becoming ions. A hydrated cation retains its positive net electric charge; orienting neutral water molecules around it does not supply an extra electron. Dissolution, desolvation, and ionization describe different changes, as the examples below show.^[ref][ref]

The useful question is therefore not just whether a solute attracts a solvent, but how much the whole system gains after paying for the rearrangements. Solvation free energy describes a specified transfer into a solvent; solvation thermodynamic balance explains why even cooling during dissolution can accompany favorable mixing. The existing solubility discussion concerns how much remains dissolved at chemical equilibrium.

At fixed temperature and pressure, compare fully specified initial and final states usingThe entropy change is for the whole selected system, including both solute and solvent. Dispersing a solute can increase accessible arrangements while orienting nearby solvent molecules restricts others. One picture of an ordered hydration shell cannot determine the sign of the total entropy change. A negative Gibbs free energy change favors the specified process; it does not determine its reaction rate.^[ref]

An illustrative enthalpy accounting path separates solute constituents, reorganizes the solvent, and establishes their interactions:This is a chosen decomposition between compatible states, not a measured sequence. If a gaseous-ion-to-chemical solution solvation term already includes solvent reorganization, do not add that reorganization a second time. Lattice enthalpy from solution calorimetry gives the alternative Hess's law route that combines lattice separation with the complete hydration of ions term.^[ref]

The invented inputs are $+100$ for separation and $+20$ for solvent reorganization, in $\mathrm{kJ}\mathrm{mo}{\mathrm{l}}^{-1}$. Interaction terms of $-150$ in solvent A and $-90$ in solvent B give $\Delta H_A=-30$ and $\Delta H_B=+30$. Stipulated total entropy changes of $-50$ and $+150\mathrm{J}\mathrm{mo}{\mathrm{l}}^{-1}{\mathrm{K}}^{-1}$ give $-T\Delta S_A=+15$ and $-T\Delta S_B=-45\mathrm{kJ}\mathrm{mo}{\mathrm{l}}^{-1}$ at $300\mathrm{K}$. Both selected processes therefore have $\Delta G=-15\mathrm{kJ}\mathrm{mo}{\mathrm{l}}^{-1}$. This deliberately constructed comparison shows why the signs of the enthalpy and Gibbs free energy changes need not agree; it does not predict either solvent's solubility.

For incremental dissolution at a particular composition, the relevant chemical potential difference also depends on thermodynamic activities. At chemical equilibrium that driving difference vanishes. A quoted standard Gibbs free energy change and the finite change between arbitrary mixtures answer different questions; use compatible standard states and the reaction quotient when discussing solubility.

The motivating calorimetry experiment is especially instructive. A weighed potassium chloride sample dissolves in a known amount of water in a platinum-lined, shielded calorimeter; a quartz-oscillator thermometer and measured electrical heating provide the readouts. Electrical calibration before and after the reaction, comparison with a resistance thermometer, and checks on residual water and particle size constrain errors. The inspected authored article reports positive enthalpy of solution at $298.15\mathrm{K}$ for the specified dilution, despite complete dissolution. This is a combined enthalpy balance, not an isolated solvation free energy or direct single-ion hydration of ions measurement. Shield tracking and final concentration matter; the existing lattice enthalpy from solution calorimetry discussion gives the numerical result and the conversion limits.^[ref]

A solvent is the component designated as the medium of a chemical solution, commonly the component present in greatest proportion. Water is a common solvent, not a necessary one. A mixture of mostly ethanol with a little water can designate ethanol as solvent and water as solute. Chemical solutions can also be gas or solid mixtures; molecular nitrogen is the principal medium of ordinary well-mixed air. When proportions are comparable, state the convention explicitly.^[ref]

The solvent is part of the final sample. Choosing it influences solubility and solvation; measuring its added volume alone does not determine the final chemical solution volume. Solvent and solute masses add in a closed account, whereas the final volume should be measured rather than inferred by adding pre-mixing volumes.

Solvent-specific stabilization means that the same species need not have the same solvation free energy in different solvents. Chemical polarity, hydrogen bonds, and the cost of reorganizing each solvent all matter. Water can stabilize ions by ion-dipole attractions; a favorable isolated attraction alone still cannot decide the total dissolution balance.^[ref]

With a common gaseous reference and compatible dissolved standard states, subtraction removes that reference:A negative result favors solvent B under those reference conventions. This comparison assumes the same species $X$: if protonation, ionization, or formation of ion pairs changes its identity, include a chemical equilibrium for each corresponding chemical reaction.

The specific heat capacity at fixed pressure and composition is $c_P=(\partial h/\partial T{)}_P$, where $h$ is specific enthalpy. Equivalently, it is heat capacity at constant pressure divided by mass. Under the process assumptions in heat capacity, it includes expansion work as well as the internal energy change. The sensible heat example uses this response separately for ice and water.

What separates a hydration shell from the surrounding water experimentally? Electrical conductivity addresses a different question. The documented lamp demonstration immerses metal electrodes in a chemical solution, rinses them between samples, and records brightness; adding sodium chloride to distilled water makes it conductive. A water blank and dissolved sucrose comparison test whether dissolved material alone explains the signal. Fixed geometry, temperature, and stated concentrations are additional proposed controls. This reveals mobile electric charge, with limits from contamination, electrode interfaces and detection threshold; it supplies no shell positions. Conductivity experiment on salt solutions explains the apparatus and controls in detail.^[ref]

In the authored study Influence of ions on the hydrogen-bond structure in liquid water by Anne Willem Omta and others, pump-probe spectroscopy uses tunable mid-infrared pulses traversing a sample. The probe's wave polarization is set at $45^{\circ }$ to the pump; after the sample, splitting and polarizers select parallel and perpendicular transmission. Measuring each transmission with and without the pump gives $D_{\parallel }$ and $D_{\perp }$, the respective logarithms of the transmission ratios. Their dimensionless anisotropy isIts decay probes reorientation rather than merely loss of vibrational excitation. Dilute semiheavy water in heavy water separates the O–H and O–D bands. Varying the semiheavy water concentration checked that transfer of excitation between nearby O–H groups did not determine the decay. These apparatus and control details are from the inspected author-uploaded article, not a new experiment.^[ref]

The related thesis chapter documents two O–H absorption bands in a sodium perchlorate chemical solution: approximately $3400\mathrm{c}{\mathrm{m}}^{-1}$ for groups forming hydrogen bonds to water, and $3575\mathrm{c}{\mathrm{m}}^{-1}$ for groups forming hydrogen bonds to the anion. Selecting pump and probe bands, then a later delay window, separated their responses. The anion-associated anisotropy decayed more slowly; the water-bonded component showed no resolved change in decay time across the tested sodium perchlorate concentrations. The chapter also used O–D excitation to extend observation beyond the shorter O–H vibrational lifetime. These results support distinct local dynamics rather than a permanently fixed hydration shell. Spectral overlap, finite observation windows, and the selected salts limit the inference: this does not prove that every ion leaves all distant water properties unchanged.^[ref]

Consider an unnamed model substance with molar latent heat $L_m=40.0\mathrm{kJ}\mathrm{mo}{\mathrm{l}}^{-1}$ and saturation pressure $p_1=10.0\mathrm{kPa}$ at $T_1=300\mathrm{K}$. Stipulate the ideal-vapor Clausius-Clapeyron approximation throughout $280\le T\le 340\mathrm{K}$, constant $L_m$, and $R=8.314\mathrm{J}\mathrm{mo}{\mathrm{l}}^{-1}{\mathrm{K}}^{-1}$. These parameters define an illustrative model, not a measured dataset or a fit to water. At $T_2=330\mathrm{K}$,At the reference state the gas molar volume is $R{T}_1/p_1=0.24942{\mathrm{m}}^3\mathrm{mo}{\mathrm{l}}^{-1}$, and the local slope is $dp/dT\simeq 0.5346\mathrm{kPa}{\mathrm{K}}^{-1}$. The tangent estimate $p_2\approx 10.0+30(0.5346)=26.04\mathrm{kPa}$ is much lower than the integrated result: a local derivative should not be extrapolated over this whole interval. By contrast, over a small temperature increment, the fractional sensitivity is $d\ln (p/p_1)/dT=L_m/(R{T}_1^2)\simeq 0.05346{\mathrm{K}}^{-1}$.

A Stirling engine repeatedly heats and cools a confined gas, whose expansion and compression drive a moving mechanical load. This makes it a particularly accessible example of a heat engine: both the repeating motion and the two regions exchanging heat can be observed.

In the published Demonstration experiments with a Stirling engine by Christopher G Deacon and others, an electrical heater warms one end of a cylinder and flowing water cools the other at $12{}^{\circ }\mathrm{C}$. A moving displacer shuttles gas through copper wool that temporarily stores and returns internal energy. A laser spot traces a closed pressure-volume path on a screen: mechanical motion controls one coordinate, pressure the other. Calibrating both scales converts the enclosed area to work per thermodynamic cycle; multiplying by rotation frequency gives gas-cycle power. At $97\mathrm{W}$ heater input, the paper reports $4.4\mathrm{W}$ gas-cycle power, while a separate Prony brake measures $3.4\mathrm{W}$ shaft output. The cold-side temperature is controlled and two higher heater settings are compared. The gas has a nonuniform temperature, and heater energy also escapes to the room: neither output/input ratio measures $W/Q_h$ for an ideal two-thermal reservoir engine. These observations demonstrate work production and losses, not a direct measurement of the bound in Carnot's theorem. The inspected apparatus, readouts and results are on pp. 180–182, Figures 1–3 and Table 1.

Stoichiometry relates the quantities consumed and formed in a chemical reaction. Its coefficients are ratios of entity counts, or equivalently of amount of substance, rather than ratios of masses. For the water chemical equation, consuming two molecular hydrogen molecules per molecular oxygen molecule gives the same $2:1$ ratio as consuming two moles of molecular hydrogen per mole of molecular oxygen. Multiplying every entity count by the same factor preserves the ratio. See Flowers et al., Chemistry 2e, reaction stoichiometry.

Coefficients and subscripts do different jobs. The coefficient in $2{\mathrm{H}}_2\mathrm{O}$ counts two water molecules; the subscript says that each contains two hydrogen chemical atoms. Changing coefficients scales quantities. Changing a subscript changes the specified chemical substance.

This distinction catches a common counting error: one mole of water molecules contains two moles of hydrogen chemical atoms and one mole of oxygen chemical atoms. Always specify what is counted.

Stoichiometry records reactant and product multiplicities in a chemical reaction or a chemical reaction network. The change caused by one reaction is its stoichiometric vector; multiplying by the reaction rate gives its contribution to the concentration equations. For two identical reactants the rate-constant convention absorbs any combinatorial factor.