**We can think about delta G as the instantaneous difference in free energy between reactants and products. And from the equation we see that****delta G is equal to delta naught plus RT ln of Q**. Delta G naught is the standard change in free energy between reactants and products.**4. Equilibrium is established when Delta G is at zero while lots of reactant at equilibrium when Delta G naught is positive: **If Delta G naught is positive at equilibrium, then we will have lots of reactants at equilibrium, meaning Q needs to be smaller (less than 1) to approach K. As Q gets smaller (i.e., as we get more reactants), the term ‘RT ln Q’ gets increasingly negative and eventually adding that term to a positive Delta G naught, will make Delta G = 0, equilibrium will be established and no further change occurs.

**3. At equilibrium Delta G is equal to zero while Delta G naught is negative: **If Delta G naught is negative at equilibrium, then we will have lots of products at equilibrium, which means that Q must be greater and also greater than 1 to approximate K. When Q gets bigger, it means more product is accumulated. The term ‘RT in Q’ gets increasingly positive, and eventually adding that term to a negative Delta G naught, will make Delta G = 0, equilibrium will be established and no further change occurs.

**5. When Delta G is zero at equilibrium, it will define which way the reaction proceeds while Delta G naught does not: **Note that it is Delta G and not Delta G naught that will be zero at equilibrium and the sign of it generated by the combination of Delta G naught and RT in Q, will define which way the reaction proceeds.

**Delta G Vs. Delta G Naught:** Delta g is used to find Gibbs free energy in nonstandard conditions while Delta G naught is used to determine Gibbs free chemical reaction energy under normal conditions. The standard condition means the pressure 1 bar and Temp 298K, Delta G naught is the measure of Gibbs free energy. The energy associated with a chemical reaction that can be used to do work change at 1 bar and 298 K, Delta G naught is not necessarily a non-zero value.

In other words, if product and reactant are equally favored at equilibrium, it’s because there is no difference. Delta G naught is always the same for a given reaction. But Delta G does depend on your conditions, but still relates to Delta H naught and Delta S naught. There are many differences between Delta G and Delta G naught, which are listed below:

If it so happens that products and reactants are equally favored at equilibrium, then ∆G° is zero, **BUT ∆G° is not *necessarily* ZERO at equilibrium.**

If ∆G° is positive at equilibrium, then we will have lots of reactants at equilibrium, meaning Q needs to be smaller (less than 1) to approach K. As Q gets smaller (i.e., as we get more reactants), the term ‘RT ln Q’ gets increasingly negative, and eventually adding that term to a positive ∆G°**, **will make ∆G = 0, equilibrium will be established and no further change occurs.

**Since this post was originally written in January 2012, the AP exam has changed. One of the changes was to remove equation #2 below from the equations & constants sheet. As such, I think that knowledge of it, and the consequences associated with it, are unlikely to be tested quantitatively on the exam in the future, but nevertheless, I still feel that understanding that conditions other than standard ones will cause ∆G to take on new values is a useful reference point.

Since Q is NOT the K, and we are NOT necessarily at the equilibrium position, the sign of ∆G can be thought of as a predictor about which way the reaction (that has reactants and products defined by Q), will go.

Since K is the equilibrium constant, we are at equilibrium, the amounts of products and reactants in the mixture are fixed, and the sign of ∆G**°** can be thought of as a guide to the ratio of the amount of products to the amount of reactants at equilibrium and therefore the thermodynamic favorability of the reaction.

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The prime usually denotes a standard free energy that corresponds to an apparent equilibrium constant where the concentration (or activity) of one or more constituents is held constant.

For example, for $ce{HA <=>A- + H+}$ the equilibrium constant is $K=frac{[ce{A-}][ce{H+}]}{[ce{HA}]}$ and the corresponding standard free energy change is $Delta G^circ=-RT ln(K)$. If you know the value of $K$ and the start concentration of $ce{HA}$ then you can compute the equilibrium concentrations of $ce{HA}$, $ce{A-}$, and $ce{H+}$.

However, if the pH is held constant then $[ce{H+}]$ is no longer a free variable and the apparent equilibrium constant is $K=frac{[ce{A-}]}{[ce{HA}]}$ and the corresponding standard free energy change is $Delta G^circ=-RT ln(K)$. So $Delta G^circ=-RT ln(K/[ce{H+}])=Delta G^circ+RTln[ce{H+}]$

In biochemistry there is often several important constituents in addition to $ce{H+}$ that are held constant such as $ce{Mg++}$, phosphate, etc.

One form of the fundamental equation of thermodynamics is:

$$dU = TdS – P dV + sum_{i}mu_i dN_i$$

In this equation, the total internal energy has canonical variables $V$, $S$, and $dN_i$, where $S$ is the total entropy (in units of $frac{mathrm J}{mathrm K}$), $V$ is the total volume, and $N_i$ is the number of moles of each molecular species present. $T$ is temperature; $P$ is pressure, and $mu_i$ is the chemical potential of species $i$. The equation implies that if we were to know an equation that gave $U$ as a function of $S$, $V$, and all the $N_i$ we would know everything about the system. However, this is inconvenient for two reasons. First, $S$ and $V$ are extensive variables. Make the system bigger without changing its composition, and $S$ and $V$ increase. Second and more importantly, it is often difficult to hold $S$ constant when doing experiments. The same is true of $V$. (We live in a constant pressure atmosphere.)

Taking the Legendre transform of $U$ with respect to variables $S$ and $V$ gives a new fundamental equation:

$$dG = -S dT + V dP + sum_{i}mu_i dN_i$$

This equation means that if we knew a function that gave the Gibbs free energy as a function of $T$, $P$, and all the $N_i$, we could easily compute all thermodynamic properties of the system.

Say were interested in the thermodynamics of ATP hydrolysis:

$ce{ATP + H2O <=> ADP + Pi}$

This equation is really better written as

$ce{A-P3O10H3 + H2O -> A-P2O7H2 + H3PO4}$

But of course in a buffer at pH 7, the conditions where many biochemical reactions occur, there really wont be $ce{H3PO4}$ etc., there will be dissociation of protons $ce{H+}$ and formation of anions like $ce{H2PO4-}$ etc. So now all those reactions will have to be tracked too. The number of protons released by ATP is not the same as released by inorganic phosphate, and this is generally true. During a reaction, it is difficult to hold the number $N_{ce{H+}}$ of protons constant, but through judicious choice of buffers etc. it is possible to hold the chemical potential of protons constant (i.e. do experiments at constant pH). Under such conditions, it makes sense to continue the Legendre transformations one step further:

$$dG^prime = -S dT + V dP – N_{ce{H+}} dmu_{ce{H+}} + sum_{i eq ce{H+}}mu_i dN_i$$

$Delta G^{circ prime}$ is a Legendre transform of $Delta G^{circ}$ with respect to the number of protons in the system.

Robert Albertys paper from 1994 is a good place for further reading.

The prime has nothing to do with whether a concentration is held constant. The prime is used to indicate that some species, typically protons or ions such as Mg, are being set to a value other than the standard 1M concentration for use in a reference free energy. Although the prime has been adopted by some in the biochemical community and was endorsed in 1994 by IUPAC, the prime is not used in other communities for good reason – it is much better to explicitly state what concentrations are being used for the reference free energy. When using a reference free energy other than the standard free energy, concentrations can vary or be held fixed. In most biochemical situations, it is useful to keep the pH fixed since buffers are used. But this is a separate issue.

The biochemistry convention does not assume all solutions are 1 M. If you did this then you would have [H+] = 1 M. This simply never happens in biochemical reactions. Instead, we assume pH = 7. We also assume that water has an activity of 1 even though its concentration is 55 M.

Delta G naught means that the reaction is under standard conditions (25 celsius, 1 M concentraion of all reactants, and 1 atm pressure). Delta G naught prime means that the pH is 7 (physiologic conditions) everything else is the same. The concentration of [H+] now isnt 1 molar because 1 molar concentration would be an extremely low pH (0). Delta G naught prime is just like Delta G naught but for biology.

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## FAQ

**What happens when Delta G equals Delta G naught?**

^{0}‘ (pronounced “delta G naught prime”) as the free energy change of a reaction under “standard conditions” which are defined as: All reactants and products are at an initial concentration of 1.0M. Pressure of 1.0 atm. Temperature is 25°C.

**What does Delta G naught mean?**

**How do you find Delta G from Delta G naught?**

**What is Delta G naught at equilibrium?**