Gibbs Free Energy and Equilibrium - Online Tutor, Practice Problems & Exam Prep

Gibbs free energy (ΔG) indicates the spontaneity of chemical reactions, with negative values signifying exergonic reactions that release energy, while positive values indicate endergonic reactions that absorb energy. The equilibrium constant (K_eq) measures the ratio of products to reactants, guiding the direction of reactions. Steady state differs from equilibrium, as it maintains stable concentrations without equal flow. Understanding these concepts is crucial for grasping cellular processes and metabolic pathways, where reactions strive for equilibrium but do not achieve it.

Gibbs Free Energy

Hi. In this video, we're going to talk about Gibbs free energy and equilibrium. I'm not going to lie, this is probably one of my least favorite cell biology topics, just because it's so chemistry-oriented, and it's not really my favorite thing. But don't fret because I'm leaving a lot of the heavy chemistry calculations to the chemistry courses, and we're just going to spend some time talking about how this relates to cell biology and understanding the concepts, so we know what we're talking about when we discuss topics like equilibrium.

The first concept we want to talk about is Gibbs free energy, which is a measure to determine if a chemical reaction will occur spontaneously. So, what does it mean if a reaction occurs spontaneously? Well, spontaneous reactions are those which are thermodynamically favorable. It just means that a reaction can occur without outside help. It doesn't need anything else to happen. A reaction will just occur. The reason it does that is because it's increasing the entropy, which is the disorder of the universe, and anything that really increases disorder doesn't necessarily need help because it's thermodynamically favorable.

Gibbs free energy can be calculated at a particular time point, which we really refer to as G, or it can be calculated regarding a change occurring during a reaction, which we measure as ΔG. This ΔG value changes a lot, just in a single reaction because the reaction is always moving towards equilibrium. So, these values aren't very static but instead are very dynamic depending on what point in a relation or what point in a reaction you're talking about.

Equilibrium is a state where the chemical reaction occurs equally in forward and reverse. Typically, in chemistry, we see reactants turning into products, and the reaction can occur this way or that way. Equilibrium is when this reaction is occurring equally in both directions. Why do we need to know this? How does this relate to biology? Cells do not exist at equilibrium, and instead, life depends on having reactions that are trying to reach equilibrium but not quite getting there. I think that phrase, trying to reach equilibrium but not quite getting there, expertly describes Cell Biology.

Now, let's talk about ΔG, the free energy calculation. The second calculation is the standard free energy change, which also measures the spontaneity of chemical reactions. What's the difference between the two? Well, ΔG calculates for a single direction with a known concentration of products and reactants. It measures real-life reactions and is sensitive to those changes because reactant and product concentrations are constantly changing in real-life reactions.

The standard free energy change, on the other hand, calculates spontaneity for a reaction occurring under standard conditions, which include the temperature being 25 degrees Celsius and the pressure being 1 atm. This calculation allows us to compare the thermodynamics of many reactions, so we can determine which reaction is more spontaneous under standardized conditions, rather than using ΔG which is highly dependent on changing concentrations.

Before we turn the page, let's talk about the concepts of exergonic and endergonic reactions. You might be familiar with these terms from chemistry. Exergonic reactions release energy whereas endergonic reactions absorb energy. Because energy is being released in exergonic reactions, they are more thermodynamically favorable. In such reactions, the free energy of the products will be less than the free energy of reactants, and ΔG will be negative. For endergonic reactions, ΔG will be positive, meaning they are less thermodynamically favorable, and the free energy of the products is more than the free energy of the reactants. At equilibrium, ΔG is zero, and no reactions occur.

Now, in addition to that, let's turn the page.

Equilibrium

Okay. So in this video, we're going to talk more about equilibrium, and specifically we're going to focus on this equilibrium constant. And what this is a measure of is the ratio of products to reactants. So if we go back to our chemistry drawing of a chemical reaction, equilibrium is when the forward and reverse reactions are occurring equally. And so how we can know, you know, whether or not that's going to happen is looking at the equilibrium constant, because that actually measures, you know, what's the proportion of products to reactants. So, this is a really important measure of directionality. And the reason that it is, is because if we have these reactants and we have these products, if there are more products than reactants, then it's going to move this way in order to equal out the amount of reactants and products. Whereas if there's more reactants than products, it's going to move this way because then it wants to equal out the amount of products. So, the number or the value of the k_eq can actually help determine the directionality. So, we can see that when k_eq is greater than 1, then the reaction favors products. If it's between 0 and 1, the reaction favors reactants. And then if it's at equilibrium, the ratio reflects a state where no net reaction occurs.

Let's look at a reaction here. So we have 3 reactions, so let's figure out, you know, what the k_eq is for them. So at equilibrium, the reactions are occurring both forward and reverse equally. And so the k_eq is going to be 0.5. Oh, let me move out of the way. Sorry for that. It equals point 5. And that's kind of easy to understand. But let's look at these other reactions. Now, when the reaction is moving forward more and is making more products, then what is the k_eq equal? So the k_eq is going to be less than 0.5. And then that's going to be opposite, of course, with the reverse reaction. So when the products are more and the reaction is moving in reverse, the k_eq is going to be greater than 0.5.

Now, let me come back and talk about another concept, and that is steady state. So, in steady state is not equilibrium. And I think that's really important to understand, because we're very familiar with steady state, or very familiar with equilibrium, but not so much with the term steady state. So what is steady state? Well, that's actually the stability of concentrations of reactants and products. So, for equilibrium, what we can see, I'm going to draw this again here, that these reactions are occurring the same. But at steady state, there can be a much higher concentration of products or reactants and a low concentration of products. But, as long as these concentrations stay relatively stable, this is occurring at steady state. So, in cell biology, what does this mean? Well, we can talk about this in terms of nutrients, for example, coming into a cell. Now, the reaction of nutrients actually coming into the cell is not in equilibrium, because they're not equally flowing in and out of the cell. They're only flowing in one direction. So, they're flowing into the cell. So, not at equilibrium, but they can be at steady state if the concentration of nutrients outside the cell and inside the cell remains relatively stable. So when it moves in, there could be a really high concentration of nutrients outside a cell and a really low concentration inside. But as long as it maintains those concentrations, it's fine. It's still at steady state, but it's definitely not happening in equilibrium because there's not an equal flow, and there's not this movement towards equilibrium. We're trying to get, you know, the products and reactants to have the same concentration on either side, because that's not going to happen in terms of nutrients crossing a cell membrane. So that's equilibrium and steady state. So now, let's move on.

The standard free energy change (ΔG°) measures which of the following?

• A. Spontaneity of a chemical reaction under any condition
• B. Spontaneity of a chemical reaction under standard conditions
• C. Ratio of products to reactants
• D. Disorder in the universe

The above energy diagram represents what type of reaction?

• A. Exergonic
• B. Endergonic

True or False:In an exergonic reaction, ΔG will be positive because energy is being released.

• A. True
• B. False

Which of the following ΔG values represents a reaction that is most likely to occur?

• A. 30
• B. 250
• C. -30
• D. -250

Here’s what students ask on this topic:

What is Gibbs free energy and how does it determine the spontaneity of a reaction?

Gibbs free energy (ΔG) is a thermodynamic quantity that indicates whether a chemical reaction will occur spontaneously. It combines enthalpy (ΔH) and entropy (ΔS) changes in the system, and is given by the equation:

$\mathrm{\Delta G}=\mathrm{\Delta H}-T\mathrm{\Delta S}$

where T is the temperature in Kelvin. A negative ΔG indicates a spontaneous, exergonic reaction that releases energy, while a positive ΔG indicates a non-spontaneous, endergonic reaction that requires energy input. Understanding ΔG helps predict the direction and feasibility of biochemical reactions in cellular processes.

What is the difference between standard free energy change (ΔG°) and Gibbs free energy (ΔG)?

Gibbs free energy (ΔG) measures the spontaneity of a reaction under actual conditions, considering the real concentrations of reactants and products. It is dynamic and changes as the reaction progresses. Standard free energy change (ΔG°), on the other hand, measures spontaneity under standard conditions: 25°C, 1 atm pressure, and 1 M concentrations of reactants and products. ΔG° allows for the comparison of different reactions' thermodynamics under consistent conditions. The relationship between ΔG and ΔG° is given by:

$\mathrm{\Delta G}=\mathrm{\Delta G°}+\mathrm{RT}\ln \left(Q\right)$

where R is the gas constant, T is the temperature, and Q is the reaction quotient.

How does the equilibrium constant (Keq) relate to Gibbs free energy?

The equilibrium constant (Keq) is a measure of the ratio of products to reactants at equilibrium. It is related to Gibbs free energy change (ΔG°) by the equation:

$\mathrm{\Delta G°}=-\mathrm{RT}\ln \left(K\mathrm{eq}\right)$

where R is the gas constant and T is the temperature in Kelvin. A large Keq (greater than 1) indicates a reaction that favors product formation and has a negative ΔG°, meaning it is spontaneous. Conversely, a small Keq (less than 1) indicates a reaction that favors reactants and has a positive ΔG°, meaning it is non-spontaneous.

What is the difference between equilibrium and steady state in cellular processes?

Equilibrium in cellular processes refers to a state where the forward and reverse reactions occur at equal rates, resulting in no net change in the concentrations of reactants and products. Steady state, however, refers to a condition where the concentrations of reactants and products remain constant over time, but not necessarily equal. In steady state, the input and output rates are balanced, but the system is not at equilibrium. This is crucial for cellular processes, as cells often maintain steady states to ensure continuous function and avoid reaching equilibrium, which would halt metabolic activities.

What are exergonic and endergonic reactions, and how do they relate to Gibbs free energy?

Exergonic reactions are chemical reactions that release energy, characterized by a negative Gibbs free energy change (ΔG). These reactions are spontaneous and thermodynamically favorable. Endergonic reactions, on the other hand, absorb energy and have a positive ΔG, making them non-spontaneous and less thermodynamically favorable. The relationship between these reactions and Gibbs free energy is crucial for understanding cellular metabolism, as cells often couple exergonic and endergonic reactions to drive necessary biological processes.