Why equilibrium is important

The results highlight the importance of the dynamics of equilibrium fluctuations between most stable conformers in the control of the reaction mechanism, i promoting the nucleophilic attack in 1AC2NA by allowing the most stable conformers to equilibrate only via rotation in a direction that intercepts the reaction coordinate and ii favoring a general base-catalyzed water attack in 3AC2NA by favoring equilibration via rotation that allows inclusion of a water molecule in a proper position for reaction.

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The energy quanta can distribute themselves in a greater variety of ways as well. Like a solid, the molecules can vibrate; but because they have more freedom of movement, molecules have rotational and some translational energy, too.

Gases have the highest entropy values because they have the greatest freedom of movement. Gas particles are separate and distribute themselves throughout their container see Section 2: Microstates, Macrostates, and Matter. And, gases also possess all three types of energy: translational, rotational, and vibrational. Finally, a substance that is dissolved in a liquid also has a high level of entropy for reasons similar to gases. Dissolved particles are free to not only move throughout the volume of the liquid, but also to move in all three ways.

Generally speaking, the entropy value of a dissolved substance is higher than pure liquids but less than gases. To summarize, the entropy of the phases of matter are:. Using this information, we can make an educated guess about whether entropy is increasing or decreasing when a chemical reaction occurs.

Consider the following reaction in which solid table salt dissolves in water:. In this case, a solid low entropy is turning into two aqueous ions higher entropy. We can reasonably assume that entropy increases. As we know, this reaction is spontaneous; salt dissolves in water. Consider another spontaneous process, the sublimation of dry ice solid CO 2 :.

Both of the above reactions are spontaneous, and both produce products with higher entropy than the reactants. It may be tempting, then, to assume that all such reactions are spontaneous. But consider the burning of hydrogen gas:. Yet, the reaction is spontaneous; hydrogen is highly flammable. As the previous sections have shown, chemical reactions that increase entropy tend to be spontaneous. Section 3 demonstrated that exothermic reactions increase entropy by allowing a wider distribution of energy quanta.

Section 4 illustrated how reactions that create liquids and gases also tend to increase entropy. What now remains is to combine these two factors to create a complete picture—a way to definitively determine if a reaction will be spontaneous. When they both oppose spontaneity, the reaction is never spontaneous. In the s, American mathematician Josiah Willard Gibbs — developed the concept of Gibbs free energy given the variable G and an equation that determines whether or not a reaction will happen spontaneously.

The Gibbs free energy of a system depends on the enthalpy, entropy, and absolute temperature of the system the derivation of this equation is beyond the scope of this text :.

Although the entropy of the system decreases in this reaction, it is more than offset by the heat released to the surroundings. Again, this conforms to experience; below its melting point, a solid will not melt.

Indeed, the reverse reaction is spontaneous at these lower temperatures; liquid water turns to ice. Neither the forward nor the reverse reaction is spontaneous. At this particular temperature, the reaction is held in a kind of limbo between the forward and reverse. Osmosis and Entropy The natural progression toward greater entropy sometimes produces surprising results. The diagram below shows a U-shaped tube divided into two halves by a semipermeable membrane.

When the tube is filled with water, the levels on the right and the left are the same, just as one would expect. However, if salt is added to the left side, a strange thing occurs.

Water will pass from the right side to the left side across the membrane, and the two levels will become uneven. Greater mixing occurs when water moves toward the side with salt, and therefore the entropy increases. It may seem that this kind of diffusion due to entropy is a passive process and not important outside of the laboratory.

In fact, semipermeable membranes are ubiquitous in living things. The membrane that surrounds all cells is semipermeable; water passes freely but larger molecules do not. The consequences are enormous.

Without knowing how it works, humankind has been taking advantage of osmosis for centuries by salting food for preservation. Putting salt on the surface of food pulls moisture out, making it a less hospitable environment for bacteria. A high-salt environment also kills bacteria outright, as osmosis pulls their water out causing the bacteria to shrivel and die. Sugar need not be refrigerated for the same reason; although it is a rich source of energy for microorganisms, such a high concentration of sugar would kill any bacteria attempting to live on it.

Osmosis also plays an important role in some human diseases. The bacterium that causes cholera, Vibrio cholerae , secretes a toxin that binds to the surface of cells in the intestine. The toxin stimulates cells to secrete large amounts of Cl — ions into the intestinal cavity.

While too much osmosis is a bad thing, too little can be bad too. Patients suffering from cystic fibrosis have a defect in a chloride channel in the cell membrane. Various mutations in this channel cause it to malfunction, and it does not release enough Cl — ions onto the membranes lining the pancreas, lungs, sweat glands, salivary glands, and other organs.

Because Cl — ions remain trapped inside the cells, water also remains inside the cell. Thus, the mucus coating these membranes lacks water and becomes too viscous. The symptoms of cystic fibrosis are the result of thick, sticky mucus clogging the channels of the affected organs. As mentioned in Section 1 of this unit, the rusting of iron is a spontaneous chemical reaction.

Iron mines extract various types of iron ore, which consist largely of compounds made of iron and oxygen. To get pure iron metal for use in manufacturing, there must be some way to reverse the spontaneous reaction of iron with oxygen. In other words, the following reaction must occur:.

Iron Ore Smelting The smelting of iron in a blast furnace requires the coupling of chemical reactions. The process of smelting iron ore—processing the ore in a blast furnace—reverses the reaction.

Figure In the intense heat of the blast furnace, the reaction is coupled to another reaction, the spontaneous reaction of carbon with oxygen:. These reactions are coupled because the O 2 produced by the first reaction is consumed in the second reaction. Combining the two reactions above, we get:. Coupled reactions are not just important in heavy industry; all living things rely on coupled reactions inside their cells.

For example, the following reaction shows the synthesis of glutamine, an amino acid, from the glutamate ion and ammonia in the form of ammonium:. Many reactions like this are coupled with the hydrolysis of adenosine triphosphate ATP :. From Diamond to Graphite The transformation of diamond into graphite is spontaneous at room temperature and pressure, but the reaction proceeds extremely slowly. Author: Mario Sarto, 4 February Pencil: Wikimedia Commons, Public Domain.

So, overall, the process is spontaneous. In this way, ATP drives thousands of biological reactions. Instead of making more ATP from scratch, organisms simply reverse the breakdown of ATP in the nonspontaneous reaction:.

What about spontaneous reactions: Must they occur? Figure The reaction is so slow, however, that it is not noticeable over human scales of time. However, if we accidentally lost a diamond ring in the batter while baking a cake, the temperatures in the oven would speed up the process and turn the diamond into worthless graphite. If the products are more stable than the reactants, this does not mean that the reaction will continue until all of the reactants have changed into products. Some of the time, the most stable situation is when there is a mixture of reactants and products.

The reaction will continue until it reaches this state of maximum thermodynamic stability lowest Gibbs free energy. Equilibrium Chemical reactions spontaneously progress toward equilibrium where free energy is at a minimum, just as gravity pulls a ball to the lowest possible level.

The situation is analogous to a ball rolling into a valley, as shown in Figure Note that position B is more stable than A because it is lower. The most stable position is in the valley between A and B. At this point, the system has reached minimum free energy—maximum thermodynamic stability.

This position is called equilibrium. At equilibrium, the reaction appears to have stopped, because the amounts of reactants and products are constant. They are constant not because the reactions have stopped, but because the forward and reverse reactions are happening at the same rate. At equilibrium, products are formed via the forward reaction and destroyed by the reverse reaction.

Because the rates of formation and destruction are equal, the amount of products remains constant. The same is true for the reactants; the forward reaction destroys the reactants, and the reverse reaction creates them, but the amount remains constant. Energy Diagrams Energy diagrams for two reactions. To quantify exactly where the equilibrium lies, we can calculate the ratio of products to reactants in a very specific way.

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List of Partners vendors. Equilibrium is the state in which market supply and demand balance each other, and as a result prices become stable. Generally, an over-supply of goods or services causes prices to go down, which results in higher demand—while an under-supply or shortage causes prices to go up resulting in less demand. The balancing effect of supply and demand results in a state of equilibrium. The equilibrium price is where the supply of goods matches demand.

When a major index experiences a period of consolidation or sideways momentum, it can be said that the forces of supply and demand are relatively equal and the market is in a state of equilibrium. For example, a dearth of any one good would create a higher price generally, which would reduce demand, leading to an increase in supply provided the right incentive.

The same would occur in reverse order provided there was excess in any one market. Modern economists point out that cartels or monopolistic companies can artificially hold prices higher and keep them there in order to reap higher profits.

The diamond industry is a classic example of a market where demand is high, but supply is made artificially scarce by companies selling fewer diamonds in order to keep prices high. As noted by Paul Samuelson in his work Foundations of Economic Analysis, the term equilibrium with respect to a market is not necessarily a good thing from a normative perspective and making that value judgment could be a misstep.

Markets can be in equilibrium, but it may not mean that all is well. For example, the food markets in Ireland were at equilibrium during the great potato famine in the mids. Higher profits from selling to the British made it so the Irish and British market was at an equilibrium price that was higher than what consumers could pay, and consequently many people starved. Suppose the price being charged for the good in question is below the market clearing price P Eq.

This is represented in the diagram above where the consumer is being charged the price P LOW. This results in a shortage of goods on the market. More has been supplied than demanded and as such, consumers are unable to buy all that they want at the current price.

With too many buyers chasing too few goods, sellers find that they cannot satisfy all their customers needs at this price so they respond to the shortage by raising their prices without losing sales. This rise in price has two separate effects. Firstly, rising prices reduce the amount of the good that consumers wish to buy as the price of the good itself is now more expensive. Therefore Quantity Demanded falls following the rise in price.

Secondly, as the price of the good itself is now more expensive, sellers receive a greater financial reward for selling the good now than they did before. As a result of this greater financial reward, sellers increase the amount they are willing to offer for sale. Therefore, quantity supplied increases following the rise in price.

As the price rises, quantity demanded falls, quantity supplied rises and the market reaches equilibrium. It is important to understand that when the market is not in equilibrium, as a result of the price being below the market price above , there are natural forces at work Price Changes to bring the market back into the desired situation of Market Equilibrium.

Thus the activities of many buyers and many sellers always push market price towards the equilibrium price. Once the market reaches its equilibrium, all buyers and sellers are satisfied and there is no upward or downward pressure on the price.

Why is Market Equilibrium a Desired Outcome? The above question has the potential to be one of the most important questions in economics. And, over the years, economists have written at length about it. It is very important for you to understand why Market Equilibrium matters if you wish to go on to study economics but I think a simple analogy will suffice in order to help you understand. For me at least, an economy is simply the way we organise how peoples material desires are met.