Equilibrium is maintained because the container is sealed.
We need a way to express how the amount of reactants and products affect the equilibrium of the system with a symbol.
The brackets are used to show the concentrations of reactants and products.
The reaction quotient is equal to the concentrations of the products of the chemical equation over the reactants, with each concentration raised to the power of the coefficients of that substance in the balanced chemical equation.
The value of the reaction quotient doesn't change when the reaction reaches equilibrium.
The reaction quotient always has the same value when a mixture of reactants and products reaches equilibrium.
The equilibrium constant is equal to the reaction quotient.
1.6 x 102 is the equilibrium constant.
The yield of a reaction when it reaches equilibrium is measured by the magnitude of an equilibrium constant.
equilibrium is attained when a small proportion of reactants are converted into products.
The direction that establishes equilibrium will be followed by a system that is not at equilibrium.
The equilibrium constant for the reaction is 0.640
It is important to know that an equilibrium can be established from either reactants or products.
One technique used to determine whether a reaction is truly at equilibrium is to start with reactants in one experiment and products in another.
We may be certain that the system has reached equilibrium if the same value of the reaction quotient is observed when the concentrations stop changing.
Determine which direction the reaction proceeds in each of the three experiments shown.
The reaction will move to the right.
The reaction will move to the left.
The reaction will move to the right.
Determine the direction in which each of the following reactions will reach equilibrium by calculating the reaction quotient.
It should be pointed out that using concentrations in these computations is a convenient but simplified approach that sometimes leads to results that conflict with the law of mass action.
While a detailed discussion of this important quantity is beyond the scope of an introductory text, it is necessary to be aware of a few important aspects.
The activity and concentration of a substance are 888-609- 888-609- 888-609- 888-609-
Pure Condensed phases are equal to 1.
This section will show examples of equilibria yielding such expressions.
In this chapter, we will look at the two most common types of equilibria, those occurring in liquid-phase solutions and those involving exclusively gaseous species.
One type of equilibria is the reaction between solutes in liquid solutions.
A mixture of both can be the chemical species involved.
There are several examples here.
The second class of equilibria consists of reactions in which all reactants and products are gases.
We will see soon that partial pressures of the gases may be used as well.
When gases are involved in a reaction, the partial pressure of each gas can be used instead of its concentration in the equation for the reaction quotient because the partial pressure of a gas is directly proportional to its temperature.
The pressure of a gas is proportional to its concentration.
Solid, liquid, or gas phases are possible.
When dealing with equilibria, remember that the activities of pure liquids are not in equilibrium constant expressions.
We can use partial pressures instead of concentrations in equations for reaction quotients of heterogeneous equilibria.
Reactions go in both directions, as we saw in the previous section.
Le Chatelier's principle can be used to predict changes in equilibrium concentrations when a system is under stress.
The changes needed to reach equilibrium may not be obvious if we have a mixture of reactants and products.
Adding or removing reactants or products can temporarily shift a chemical system out of equilibrium.
Changes to the concentrations of reactants and products return the system to equilibrium.
Le Chatelier's principle leads us to predict that the concentration of Fe(SCN)2+ should decrease, increasing the concentration of SCN- part way back to its original concentration, and increasing the concentration of Fe3+ above its initial equilibrium concentration.
The first equilibrium in the solution is shifted to the left due to the decrease in the Fe concentration.
The values for this example have been determined.
The system has been stressed by introducing additional H2.
The stress is relieved when the reaction shifts to the right, using up some of the excess H2, reducing the amount of uncombined I2, and forming additional HI.
Changing the pressure of a system can change the position of equilibrium.
Changes in pressure only have a measurable effect in systems in which gases are involved, and then only when the chemical reaction produces a change in the total number of gas molecule in the system.
An easy way to see a system like this is to look for different moles of gas on the reactant and product sides.
Adding an inert gas that is not part of the equilibrium will change the total pressure but not the partial pressures of the gases in the equilibrium constant expression.
The equilibrium won't be perturbed by the addition of a gas not involved in it.
There is a dramatic demonstration of equilibrium changing with pressure.
When we increase the pressure of a gaseous system at equilibrium, either by decreasing the volume of the system or by adding more of one of the components, we introduce a stress by increasing the partial pressures of one or more of the components.
The shift in the equilibrium that reduces the total number of molecules per unit of volume will be favored because it relieves the stress.
A decrease in pressure would favor the reverse reaction.
This doesn't completely relieve the stress of the increased pressure, but it does reduce it.
The decomposition of NO2 into NO and O2 tends to restore the pressure when there is a decrease in it.
The reaction quotient is shifted away from the equilibrium value when concentration or pressure is changed.
The value of the equilibrium constant can be changed by a change in temperature.
We can qualitatively predict the effect of the temperature change by applying Le Chatelier's principle and treating it as a stress on the system.
When hydrogen reacts with a radioactive substance, heat is created.
We can write this reaction with heat because it is exothermic.
The internal energy of the system can be increased by increasing the temperature of the reaction.
Increasing the temperature increases the amount of one of the products of the reaction.
There is an increase in the concentration of H2 and I2 and a reduction in the concentration of HI when the reaction shifts to the left.
Lowering the temperature of this system reduces the amount of energy present, favors the production of heat, and favors the formation of hydrogen iodide.
The equilibrium constant for the reaction changes when we change the temperature of the system.
The equilibrium constant is increased by lowering the temperature in the HI system.
The value of the equilibrium constant decreases when the temperature is raised.
The equilibrium shifts to the left if we put a stress on the system by cooling the mixture.
The brown color fades when the concentration of brown NO2 decreases and the concentration of N2O4 increases.
Predicting the effects of changes in concentration, pressure, and temperature on reactant and product concentrations is possible with this interactive.
A catalyst can speed up a reaction.
The increase in reaction rate may cause a system to reach equilibrium more quickly, but it has no effect on the equilibrium constant or equilibrium concentrations.
Ammonia is one of the top 10 chemicals in the world.
2 billion pounds are manufactured in the United States each year.
Ammonia is important in our global economy.
It is an importantfertilizer for the growth of corn, cotton, and other crops.
Ammonia is converted to nitric acid, which is used in the production of explosives, dyes, and fibers, as well as in the steel industry.
The process for converting diatomic nitrogen, which cannot be used by plants as a nitrogen source, was developed by a German chemist in the early 20th century.
Diatomic nitrogen is unavailable due to the stability of the nitrogen-nitrogen triple bond.
Nitrogen must be converted to a more bioavailable form for plants to use atmospheric nitrogen.
In December 1868, he was born in Breslau, Prussia.
He went on to study chemistry at the University of Karlsruhe and developed a process for making ammonia from hydrogen and nitrogen under high temperatures and pressures.
The 1918 Nobel Prize in Chemistry was awarded for the synthesis of ammonia from its elements.
The production of fertilizers was no longer dependent on mined feed stocks thanks to the Haber process.
Synthetic nitrogenfertilizer production has increased the number of humans that arable land can support from 1.9 persons perhectare in 1908 to 4.3 in 2008.
In the early 20th century, the work of the winner of the Nobel Prize changed agricultural practices.
Adding chemical weapons to the army's war strategies were affected by his work.
One of the fathers of chemical warfare was a man named Haber, who worked in ammonia production.
During World War I, he was involved in the development of poisonous gases.
"During peace time a scientist belongs to the World, but during war time he belongs to his country," he said.
The use of gas warfare was defended against accusations that it was inhumane.
He is an example of an ethical dilemma faced by scientists in times of war and the sword of science.
ammonia products can be more than one.
Nitrogen compounds can be used to achieve destructive ends.
Ammonium nitrate has been used in explosives.
The bomb used in the attack on the Alfred P. Murrah Federal Building in Oklahoma City was made of Ammonium nitrate.
Nitrogen and hydrogen form ammonia.
After the factors that influence its equilibrium were understood, it was possible to make ammonia in useful quantities by the reaction of nitrogen and hydrogen.
An industrial process that gives a large yield of product quickly is practical.
Increasing the pressure on the system in which N2, H2, and NH3 are at equilibrium is one way to increase the yield of ammonia.
The rate of formation of ammonia is slow when the pressure of a mixture of N2, H2, and NH3 is increased.
If we prepared a mixture of N2 and H2 at room temperature, no ammonia would form during our lifetime.
Increasing the temperature to increase the rate lowers the yield.
If we lower the temperature to shift the equilibrium to favor the formation of more ammonia, it will take more time to reach equilibrium.
A catalyst can be used to recover part of the rate of formation lost.
The catalyst's effect on the reaction is to cause equilibrium to be reached more quickly.
The best compromise among rate, yield, and the cost of equipment necessary to produce and contain high-pressure gases at high temperatures is achieved by using conditions of 500 degC, 150-900 atm, and the presence of a catalyst.