Actual Free-Energy Change vs Standard Conditions

A cell keeps glucose-6-phosphate at roughly 200 μM. It keeps glucose at 5 mM. Neither matches the 1 M standard state chemists use on paper. That mismatch is not a footnote. It is the entire reason metabolism runs forward. The reaction catalyzed by phosphoglucose isomerase, glucose-6-phosphate ⇌ fructose-6-phosphate, has a standard free-energy change ΔG°′ close to +1.7 kJ/mol. That value looks mildly unfavorable. Yet inside a hepatocyte this isomerization proceeds essentially to completion in the glycolytic direction. This happens because actual substrate and product concentrations are nowhere near equal. They are also nowhere near 1 M. This article is about the gap between ΔG°′ and ΔG. ΔG°′ is a fixed reference number. It is defined for the Standard Free-Energy Change and the Biochemical Standard State. ΔG is the free-energy change that actually governs a reaction's direction. It determines whether a reaction proceeds forward, backward, or not at all. This happens inside a living cell at a given instant. The reaction quotient Q connects the two. Q is built from real, fluctuating intracellular concentrations. Cells exploit Q ruthlessly. By holding certain reactions far from equilibrium (Q ≪ Keq or Q ≫ Keq), they make ΔG large and negative. This happens even when ΔG°′ is unfavorable. They can also make an intrinsically favorable reaction run near-equilibrium. Such a reaction has large negative ΔG°′. It then contributes almost no net driving force. The big picture is a single equation. It is the van 't Hoff isotherm as applied to biochemistry: ΔG = ΔG°′ + RT ln Q. It looks almost too simple to carry so much biology. But it is the load-bearing beam under the entire logic. That logic governs Coupled Reactions and the Additivity of Free Energy. It also explains why certain steps in a pathway are singled out. Those steps are subject to Regulation at Nonequilibrium Steps. Every glycolytic enzyme is governed by this same relation. So is every step of oxidative phosphorylation. So is every transport process across a membrane. Only the definitions of Q and the standard state change. By the end of this article you will be able to compute ΔG. You will do this for any reaction given intracellular concentrations. You will be able to explain quantitatively why a reaction with unfavorable ΔG°′ can run forward in vivo. You will distinguish the equilibrium constant from the steady-state ratio actually observed in cells. You will also see why some enzymes sit at near-equilibrium. Others are held perpetually off-equilibrium as control points.

Key takeaways

ΔG°′ is a fixed constant defined at 1 M reactants/products, pH 7, 298 K; ΔG is the value that determines direction right now, in the cell's actual concentrations.
The bridge is ΔG = ΔG°′ + RT ln Q, where Q is the mass-action ratio of actual concentrations raised to their stoichiometric powers.
A reaction with ΔG°′ > 0 can still have ΔG ≪ 0 if the cell keeps [substrate] high and [product] low, i.e., Q ≪ Keq.
At equilibrium Q = Keq and ΔG = 0 by definition; cells generally avoid this state at key control steps.
Most metabolic reactions run near-equilibrium (ΔG ≈ 0, freely reversible); a few "committed step" reactions are held far from equilibrium and are the sites of allosteric and hormonal control.
Key idea: direction and driving force in vivo depend on the ratio of actual concentrations to equilibrium concentrations, not on the sign of ΔG°′ alone.

Where this lives today

Metabolic flux analysis and systems biology depend on exactly this calculation at genome scale: modern flux-balance and thermodynamics-based metabolic models compute ΔG for every reaction in a reconstructed metabolic network using measured or estimated intracellular metabolite concentrations, flagging which reactions are thermodynamically feasible in a given direction and which are effectively irreversible in vivo — directly informing which enzymes are viable drug targets, since a reaction pinned far from equilibrium is far more sensitive to inhibition than one sitting near equilibrium (where the reverse reaction can simply pick up slack). Diabetes and cancer metabolism research routinely uses ¹³C and ³¹P NMR or LC-MS metabolomics to measure actual intracellular concentrations of glycolytic and TCA-cycle intermediates, feeding them into ΔG = ΔG°′ + RT ln Q to identify which steps are held nonequilibrium (and are candidate control/regulatory nodes, often the same enzymes — hexokinase, phosphofructokinase-1, pyruvate kinase — long known from classical biochemistry to be allosterically regulated). In bioengineering, synthetic-biology pathway design tools (such as thermodynamic-based flux analysis, TFA/MDF—max-min driving force—algorithms) explicitly search for metabolite concentration profiles that keep every engineered reaction step's actual ΔG negative, since a pathway with even one reaction stuck at ΔG ≥ 0 under the chosen concentrations will not run, no matter how favorable its ΔG°′ looks on paper. Pharmacological interventions that alter cellular Pᵢ, NAD⁺/NADH, or ADP/ATP ratios (for example in ischemia or in mitochondrial disease) change actual ΔG values throughout central metabolism even though every ΔG°′ in the cell remains unchanged, which is why clinical and cellular bioenergetics measurements (phosphorylation potential, NADH/NAD⁺ ratio) are tracked as functional readouts of metabolic state rather than the fixed standard free energies.

Ask the tutor: How does this same ΔG = ΔG°′ + RT ln Q logic combine two reactions sharing a common intermediate to make an otherwise uphill biosynthetic step proceed?

The physical picture — thorough prose

Think of ΔG°′ as the free-energy difference between products and reactants when everyone is present at a fixed reference concentration (1 M for solutes, pH 7, 55.5 M for water, 1 atm for gases). It is a property of the reaction itself — bond energies, solvation, entropy of the standard state — and does not change no matter what a particular cell is doing at a particular moment. ΔG, in contrast, is a property of the reaction and the current mixture. It asks: given what is actually dissolved in this cytosol right now, which way does the reaction want to go, and how strongly? The two are connected because free energy is a logarithmic function of concentration. Raising the concentration of a reactant lowers its chemical potential's contribution to the "downhill" side less than you might think linearly — it is a logarithmic penalty/bonus, μ = μ° + RT ln[X]. Summing these chemical potentials over products minus reactants, weighted by stoichiometric coefficients, produces the reaction quotient Q and the correction term RT ln Q. When Q is small (products scarce, reactants abundant relative to their equilibrium ratio), ln Q is a large negative number, and this term can overwhelm even a positive ΔG°′, making ΔG very negative. This is exactly how the cell "spends" concentration gradients as a form of stored free energy, alongside the more familiar currency of ATP hydrolysis. A second key intuition: Keq is the special value of Q at which ΔG = 0. If the cell's actual Q sits below Keq, the reaction has a thermodynamic driving force to proceed forward, converting reactants to products (and thereby increasing Q toward Keq). If Q sits above Keq, the reverse happens. The forward rate at a given ΔG°′ tells you nothing about spontaneity; only the relation of actual Q to Keq — equivalently, the sign and magnitude of RT ln(Q/Keq) — does that job.

Figure

ΔG vs ΔG°′: the reaction-quotient correction

The picture is not just qualitative. Because ln is a strictly monotonic function, the correction term RT ln Q at 310 K (37 °C, body temperature) is about 2.58 kJ/mol per unit of ln Q, and per decade (factor of 10) in Q it contributes 2.58 × ln(10) ≈ 5.94 kJ/mol. A cell that maintains a product/substrate ratio even a factor of 100 below equilibrium is buying itself roughly 11.9 kJ/mol of extra forward driving force — often enough to flip the sign of ΔG entirely relative to ΔG°′.

Figure

Free-energy profile: standard vs actual state

The core ideas — the governing equations/relationships

For a general reaction

aA + bB \rightleftharpoons cC + dD

the actual free-energy change is:

\Delta G = \Delta G^{\circ\prime} + RT\ln Q

Q = \dfrac{[C]^{c}[D]^{d}}{[A]^{a}[B]^{b}}

where:

$\Delta G$ — actual free-energy change under the real, current concentrations (kJ/mol); its sign tells you the direction of spontaneous change right now.
$\Delta G^{\circ\prime}$ — biochemical standard free-energy change, fixed for the reaction (1 M solutes, pH 7, 298 or 310 K, 1 atm).
$R$ — gas constant, 8.314 J mol⁻¹ K⁻¹.
$T$ — absolute temperature (K); 310 K for mammalian cytosol is common in worked problems.
$Q$ — reaction quotient, the mass-action ratio evaluated at the actual instantaneous concentrations, not at equilibrium.

Reading: ΔG°′ is a constant offset; RT ln Q is a variable correction that depends entirely on what the cell is actually doing with its concentrations. When Q < 1 and small, ln Q is very negative and pulls ΔG down; when Q > 1 and large, ln Q is positive and pushes ΔG up. The link to the equilibrium constant is obtained by setting ΔG = 0 (the defining condition of equilibrium) and Q = Keq:

0 = \Delta G^{\circ\prime} + RT \ln K_{eq}

\Delta G^{\circ\prime} = -RT \ln K_{eq}

where:

$K_{eq}$ — the true equilibrium constant, the only value of Q at which the net reaction rate is zero.

Substituting this back gives the most physically transparent form of the whole relationship:

\Delta G = RT \ln\!\left(\dfrac{Q}{K_{eq}}\right)

Reading: everything ΔG°′ contributes is already baked into Keq. The actual driving force depends only on how far the current ratio Q has drifted from the equilibrium ratio Keq — nothing about the 1 M reference state appears explicitly anymore. This form makes the central lesson unmissable: a reaction can have any ΔG°′ you like and still run forward, backward, or sit still, entirely depending on where the cell keeps Q relative to Keq.

Derivation AFrom chemical potentials to ΔG = ΔG°′ + RT ln Q

Goal: derive the reaction-quotient correction from the concentration-dependence of chemical potential, for a general reaction aA + bB ⇌ cC + dD.

Step 1 — chemical potential of a solute. For a dilute solute X at concentration [X], the chemical potential (partial molar free energy) is

\mu_X = \mu_X^{\circ} + RT \ln[X]

Here $\mu_X^{\circ}$ is the chemical potential at the reference concentration (1 M, biochemical standard state), and the $RT\ln[X]$ term is the entropic cost/benefit of being at some other concentration — more dilute species have lower chemical potential because they are more disordered per mole, more concentrated species have higher chemical potential.

Step 2 — total free energy of a reacting mixture. The free-energy change for converting a A + b B into c C + d D, at whatever concentrations currently exist, is the stoichiometric sum of chemical potentials of products minus reactants:

\Delta G = \big(c\,\mu_C + d\,\mu_D\big) - \big(a\,\mu_A + b\,\mu_B\big)

Step 3 — substitute the concentration dependence. Insert $\mu_X = \mu_X^\circ + RT\ln[X]$ for each species:

\Delta G = \big[c\mu_C^\circ + d\mu_D^\circ - a\mu_A^\circ - b\mu_B^\circ\big] + RT\big[c\ln[C] + d\ln[D] - a\ln[A] - b\ln[B]\big]

Step 4 — recognize the two pieces. The bracketed term of standard chemical potentials is, by definition, ΔG°′ (the free-energy change when everything is at the reference state). The second bracket, using $\ln x^n = n\ln x$ and combining logs, is:

c\ln[C] + d\ln[D] - a\ln[A] - b\ln[B] = \ln\!\left(\frac{[C]^c[D]^d}{[A]^a[B]^b}\right) = \ln Q

Step 5 — assemble. Substituting back:

\Delta G = \Delta G^{\circ\prime} + RT\ln Q

This confirms the master equation from first principles: it is nothing more than bookkeeping of the logarithmic concentration-dependence of chemical potential, summed with the correct stoichiometric signs. No new physics enters; the entire content is that free energy per mole of solute grows logarithmically, not linearly, with concentration.

Go deeper · Derivation B · 1 minWhy an "unfavorable" ΔG°′ reaction can run forward in the cell

Derivation BWhy an "unfavorable" ΔG°′ reaction can run forward in the cell

Goal: show explicitly, using the isomerization glucose-6-phosphate (G6P) ⇌ fructose-6-phosphate (F6P), how maintaining [F6P]/[G6P] below Keq converts a positive ΔG°′ into a negative ΔG.

Step 1 — the standard-state number. Literature value: $\Delta G^{\circ\prime} = +1.7\ \text{kJ/mol}$ at 298 K. This describes the (hypothetical) situation where [G6P] = [F6P] = 1 M.

Step 2 — convert to Keq. Using $\Delta G^{\circ\prime} = -RT\ln K_{eq}$ at T = 298 K, R = 8.314 J mol⁻¹K⁻¹ = 0.008314 kJ mol⁻¹K⁻¹:

\ln K_{eq} = -\frac{\Delta G^{\circ\prime}}{RT} = -\frac{1.7}{(0.008314)(298)} = -\frac{1.7}{2.478} = -0.686

K_{eq} = e^{-0.686} = 0.504

So at equilibrium the cell would need [F6P]/[G6P] ≈ 0.50 — products a bit less abundant than reactant, consistent with a mildly "uphill" forward reaction under standard conditions.

Step 3 — insert real cellular concentrations. Measured intracellular values in liver cytosol are approximately [G6P] ≈ 200 μM and [F6P] ≈ 60 μM (i.e., glycolytic flux keeps F6P being consumed by phosphofructokinase almost as fast as it forms, holding the ratio below equilibrium). Then:

Q = \frac{[\text{F6P}]}{[\text{G6P}]} = \frac{60\times10^{-6}}{200\times10^{-6}} = 0.30

Step 4 — compute ΔG.

\Delta G = \Delta G^{\circ\prime} + RT\ln Q = 1.7 + (2.478)\ln(0.30)

\ln(0.30) = -1.204

\Delta G = 1.7 + (2.478)(-1.204) = 1.7 - 2.983 = -1.28\ \text{kJ/mol}

Step 5 — interpret. ΔG°′ said "mildly unfavorable" (+1.7 kJ/mol); the actual free-energy change is −1.28 kJ/mol, clearly favorable, because Q = 0.30 sits below Keq = 0.50. The isomerase is not fighting thermodynamics — it is riding a driving force created entirely by the cell's concentration management, chiefly the fact that phosphofructokinase downstream continuously drains F6P. Notice the reaction is close to equilibrium (Q and Keq differ only by a factor of ~1.7), consistent with phosphoglucose isomerase being a near-equilibrium, freely reversible step rather than a committed regulatory point.

Worked examples

Example 1 (foundational): sign flip from ΔG°′ to ΔG for a simple isomerization. Take a generic isomerization A ⇌ B with ΔG°′ = +2.5 kJ/mol at T = 310 K (37 °C, physiological temperature). Suppose the cell maintains [A] = 4.0 mM and [B] = 0.80 mM. Step 1: compute RT.

RT = (0.008314\ \text{kJ mol}^{-1}\text{K}^{-1})(310\ \text{K}) = 2.577\ \text{kJ/mol}

Step 2: compute Q.

Q = \frac{[B]}{[A]} = \frac{0.80}{4.0} = 0.20

Step 3: compute ln Q.

\ln(0.20) = -1.609

Step 4: apply the master equation.

\Delta G = 2.5 + (2.577)(-1.609) = 2.5 - 4.147 = -1.65\ \text{kJ/mol}

The reaction, "uphill" by ΔG°′, is downhill by 1.65 kJ/mol under these actual concentrations. Limiting-case check: if the cell instead kept [A] = [B] (equal concentrations, Q = 1), then ln Q = 0 and ΔG = ΔG°′ exactly — the two quantities coincide only when all species sit at equal, unit-consistent concentrations relative to each other in the ratio dictated by stoichiometry (here, simply equal concentrations since stoichiometry is 1:1). This is the correct sanity check: the concentration correction vanishes exactly when Q = 1, regardless of what those equal concentrations actually are (as long as both are in molar units matching the standard-state definition), confirming the formula collapses to ΔG = ΔG°′ at Q = 1 as it must.

Example 2 (applied): the ATP hydrolysis reaction inside a real cell. ATP + H₂O → ADP + Pᵢ has ΔG°′ = −30.5 kJ/mol at 298 K (this already folds in the biochemical convention that water activity is fixed at 1, not 55.5 M, in Q). A typical hepatocyte maintains approximately [ATP] = 3.0 mM, [ADP] = 0.010 mM (achieved partly via adenylate kinase and creatine phosphate buffering), and [Pᵢ] = 5.0 mM. Step 1: write Q. Since water is omitted (standard-state convention absorbs its near-constant activity into ΔG°′):

Q = \frac{[\text{ADP}][\text{P}_i]}{[\text{ATP}]}

Step 2: insert numbers (convert all to consistent units, here molar):

Q = \frac{(0.010\times10^{-3})(5.0\times10^{-3})}{3.0\times10^{-3}} = \frac{5.0\times10^{-8}}{3.0\times10^{-3}} = 1.67\times10^{-5}

Step 3: compute RT at 310 K and ln Q.

RT = 2.577\ \text{kJ/mol}, \qquad \ln(1.67\times10^{-5}) = -11.00

Step 4: apply the master equation.

\Delta G = -30.5 + (2.577)(-11.00) = -30.5 - 28.35 = -58.8\ \text{kJ/mol}

This is the number usually quoted for the actual free energy of ATP hydrolysis in vivo, roughly double the standard value in magnitude — and it is this larger, actual ΔG that is available to drive endergonic processes coupled to ATP hydrolysis, motor proteins, active transport, and biosynthesis, not the textbook −30.5 kJ/mol. Limiting-case check: if instead the cell allowed ADP and Pᵢ to accumulate until Q rose to equal $K_{eq} = e^{-\Delta G^{\circ\prime}/RT} = e^{30.5/2.478}\approx e^{12.3}\approx 2.2\times10^{5}$ , ΔG would fall to exactly zero — the cell would have "spent" all its ATP free energy and hydrolysis would stop being useful, which is precisely why maintaining a low ADP/Pᵢ background (far from Keq) is metabolically essential.

Example 3 (exam/competition level): back-calculating a required concentration ratio for a target ΔG. A biosynthetic enzyme catalyzes X ⇌ Y with ΔG°′ = +8.0 kJ/mol at 310 K. The cell needs this reaction to proceed with an actual driving force of ΔG = −4.0 kJ/mol to keep pace with downstream demand. What ratio [Y]/[X] must be maintained (assuming 1:1 stoichiometry, Q = [Y]/[X])? Step 1: rearrange the master equation for ln Q.

\ln Q = \frac{\Delta G - \Delta G^{\circ\prime}}{RT}

Step 2: insert numbers.

\ln Q = \frac{(-4.0) - (8.0)}{2.577} = \frac{-12.0}{2.577} = -4.657

Step 3: exponentiate.

Q = e^{-4.657} = 9.5\times10^{-3}

So the cell must keep [Y]/[X] ≈ 0.0095, i.e. Y roughly 100-fold less abundant than X. Step 4 — sanity-check against Keq: compute Keq for comparison.

K_{eq} = e^{-\Delta G^{\circ\prime}/RT} = e^{-8.0/2.577} = e^{-3.104} = 0.045

Q (0.0095) is indeed well below Keq (0.045), consistent with a forward-favorable ΔG, and the ratio Q/Keq = 0.0095/0.045 ≈ 0.21 reproduces the target ΔG independently:

\Delta G = RT\ln(Q/K_{eq}) = (2.577)\ln(0.21) = (2.577)(-1.560) = -4.02\ \text{kJ/mol}

matching the −4.0 kJ/mol target within rounding, confirming internal consistency of the two equivalent formulas ΔG = ΔG°′ + RT ln Q and ΔG = RT ln(Q/Keq).

Limiting cases, scope & common traps

Situation	What happens to ΔG	Interpretation
Q = 1 (all species at unit reference activity, matching stoichiometric convention)	ΔG = ΔG°′ exactly	The correction term vanishes; standard and actual coincide only here
Q = Keq	ΔG = 0	True chemical equilibrium; no net forward or reverse flux
Q ≪ Keq	ΔG large and negative	Strong forward driving force regardless of the sign of ΔG°′
Q ≫ Keq	ΔG large and positive	Reaction driven in reverse regardless of ΔG°′
T → 0 K	RT ln Q → 0, ΔG → ΔG°′	At absolute zero the concentration correction vanishes (unphysical for solutions, but the correct algebraic limit)
Reaction held continuously far from equilibrium by flux (e.g., committed steps)	ΔG stays large and negative in steady state	Hallmark of a regulated, "irreversible" step — a candidate control point
Reaction near-equilibrium in vivo (many isomerases, mutases)	ΔG ≈ 0	Freely reversible, flux direction set by mass action elsewhere in the pathway, not by this step's own thermodynamics

The chemist behind it

The modern language of free energy and chemical potential traces to Josiah Willard Gibbs, the American mathematical physicist who, working largely in isolation at Yale in the 1870s, built the entire framework of chemical thermodynamics — the free energy that bears his name, the phase rule, and the concept of chemical potential that makes $\mu = \mu^\circ + RT\ln[X]$ possible in the first place. Gibbs published his masterwork On the Equilibrium of Heterogeneous Substances in the Transactions of the Connecticut Academy of Sciences, a journal so obscure that Maxwell and Boltzmann in Europe took years to notice it, yet it quietly contained essentially the complete apparatus later chemists and biochemists would use to describe reacting mixtures at fixed temperature and pressure. Gibbs never applied his theory to biological systems himself — the biochemical standard state (pH 7, defined ΔG°′) was a twentieth-century adaptation by biochemists needing a convention that made sense for reactions in aqueous, near-neutral, dilute cellular environments — but every calculation in this article, from the ATP hydrolysis number to the phosphoglucose isomerase ratio, is a direct descendant of his 1876 chemical potential.

Origins · 1 minJosiah Willard Gibbs and how this was discovered

Josiah Willard Gibbs

1839–1903 · Yale University, mathematical physics and chemical thermodynamics

Gibbs formalized the free-energy function G = H − TS and the concept of chemical potential, showing that a substance's contribution to a system's free energy depends logarithmically on its concentration or partial pressure. This single insight — that dilution or concentration shifts free energy by RT ln[X] — is the entire mechanistic content connecting a reaction's fixed standard free-energy change to its variable, condition-dependent actual free-energy change, the relationship at the heart of cellular bioenergetics.

History

- 1873–1878 — Gibbs publishes On the Equilibrium of Heterogeneous Substances, introducing free energy and chemical potential.

- 1884 — van 't Hoff formalizes the relationship between equilibrium constants, temperature, and reaction free energy (the van 't Hoff isotherm/equation).

- 1930s — Fritz Lipmann and colleagues establish the central bioenergetic role of ATP and phosphoryl-transfer free energies in cellular metabolism.

- 1950s–1960s — biochemists standardize the pH-7 biochemical convention (ΔG°′) to make tabulated free energies directly usable for aqueous, near-neutral cellular reactions.

- 1970s–present — direct measurement of intracellular metabolite concentrations by ³¹P-NMR and mass spectrometry allows routine calculation of actual in vivo ΔG values for pathways such as glycolysis and oxidative phosphorylation.

Is ΔG°′ ever directly useful for predicting what happens in a real cell?

Yes, but only combined with information about actual concentrations. ΔG°′ tells you the intrinsic thermodynamic tendency of the reaction under a fixed, standardized reference state, which lets you compute Keq via ΔG°′ = −RT ln Keq. Keq in turn tells you the equilibrium ratio of products to reactants — the target that Q would reach if the reaction were left alone. Comparing the cell's actual Q to that Keq, via ΔG = RT ln(Q/Keq), is what actually predicts direction and magnitude of driving force in vivo. Used alone, without concentration data, ΔG°′ only tells you the reaction's behavior in an artificial 1 M reference mixture — informative for tabulating and comparing reactions, but not a direct predictor of cellular behavior.

Why doesn't water's concentration (55.5 M) appear in Q for reactions like ATP hydrolysis?

By convention, in the biochemical standard state, water is treated as a "special" species whose activity is fixed near 1 because its concentration (55.5 M) barely changes during any biochemical reaction taking place in dilute aqueous solution. Its near-constant contribution is folded into the definition of ΔG°′ itself rather than left as a variable in Q. The same convention applies to H⁺ at pH 7 (that is the entire point of the prime, ′, in ΔG°′ — it means "at pH 7," fixing [H⁺] rather than leaving it at the chemical standard state of 1 M). This is why ATP + H₂O → ADP + Pi uses Q = [ADP][Pi]/[ATP] with no explicit [H₂O] term, even though water is technically a reactant.

Can ΔG be positive for a reaction that is actually happening in a cell right now?

Individually, no — a reaction with ΔG > 0 as written cannot proceed spontaneously in that direction; if it appears to be "running" it is either running in reverse (making ΔG for the reverse direction negative) or it is being driven by coupling to a second, more exergonic reaction sharing a common intermediate (the subject of Coupled Reactions and the Additivity of Free Energy). Many biosynthetic steps that look uphill in isolation (e.g., some steps of gluconeogenesis or amino acid biosynthesis) proceed only because they are thermodynamically linked to ATP or GTP hydrolysis, whose large negative ΔG more than compensates when the two half-reactions' ΔG values are summed.

How is ΔG measured or estimated for a reaction inside a real cell?

Researchers first obtain ΔG°′ from calorimetry, redox potentials, or compiled thermodynamic databases (such as eQuilibrator), then measure actual intracellular metabolite concentrations using techniques like ³¹P-NMR (for phosphorylated intermediates and pH), LC-MS/MS metabolomics, or isotope-labeling flux studies. These concentrations are substituted into Q, and ΔG = ΔG°′ + RT ln Q is evaluated numerically. Because intracellular volumes are tiny and metabolite pools are compartmentalized (cytosol vs. mitochondrial matrix have very different ratios), this calculation is always an approximation using bulk-averaged concentrations, but it remains the standard way to estimate in vivo thermodynamic feasibility for genome-scale metabolic models.

Does temperature change ΔG°′ itself, or only the RT ln Q term?

Both, in general — ΔG°′ = ΔH°′ − TΔS°′, so ΔG°′ is itself temperature-dependent unless ΔH°′ and ΔS°′ happen to be temperature-independent over the range considered (a common simplifying assumption over the narrow biological range 25–40 °C). In most introductory calculations, tabulated ΔG°′ values are given at a specific reference temperature (298 K or 310 K) and simply looked up rather than recalculated, while the RT ln Q term is explicitly recomputed at the temperature of interest since R and T appear directly in it. For rigorous work spanning a wide temperature range, both the ΔG°′ temperature-dependence and the RT ln Q term must be tracked.

Why do cells bother keeping some reactions near equilibrium and others far from it?

Near-equilibrium reactions (ΔG ≈ 0, freely reversible, like triose phosphate isomerase or phosphoglucose isomerase) are metabolically "cheap" to run in either direction and respond passively to mass-action pressure from neighboring reactions — they are not good control points because a small change in enzyme activity barely shifts flux when the reaction is already near its thermodynamic limit. Reactions held far from equilibrium (large negative ΔG maintained continuously, like hexokinase, phosphofructokinase-1, and pyruvate kinase in glycolysis) are functionally irreversible under cellular conditions and are the sites where allosteric effectors, covalent modification, and hormonal signals actually control net pathway flux — the subject taken up directly under Regulation at Nonequilibrium Steps.

Is the relationship ΔG = ΔG°′ + RT ln Q exact, or an approximation for dilute solutions?

It is exact for ideal solutions, where activity coefficients equal 1 and chemical potential is exactly $\mu^\circ + RT\ln[X]$ . Real cellular solutions are crowded (macromolecular crowding, ionic strength effects, and specific ion binding all shift activity coefficients away from 1), so a fully rigorous treatment would use activities rather than concentrations: $\mu = \mu^\circ + RT\ln a_X$ where $a_X = \gamma_X[X]$ . For most biochemical purposes, concentrations are used directly as a good approximation, since activity coefficients in cytosol are reasonably close to those in the dilute reference conditions used to tabulate ΔG°′ — but for high-precision thermodynamic modeling of crowded environments, this is a recognized source of systematic error.