Explain How 3 M Sulfuric Acid Caused the Equilibrium to Shift Back Again
Learning Objectives
Past the end of this module, you will be able to:
- Describe the means in which an equilibrium system can exist stressed
- Predict the response of a stressed equilibrium using Le Châtelier's principle
As we saw in the previous department, reactions proceed in both directions (reactants go to products and products go to reactants). We tin tell a reaction is at equilibrium if the reaction quotient (Q) is equal to the equilibrium constant (Thousand). We adjacent address what happens when a system at equilibrium is disturbed so that Q is no longer equal to Chiliad. If a system at equilibrium is subjected to a perturbance or stress (such as a alter in concentration) the position of equilibrium changes. Since this stress affects the concentrations of the reactants and the products, the value of Q volition no longer equal the value of Grand. To re-establish equilibrium, the system volition either shift toward the products (if Q < K) or the reactants (if Q > K) until Q returns to the same value equally 1000.
This procedure is described by Le Châtelier's principle: When a chemical organisation at equilibrium is disturbed, it returns to equilibrium by counteracting the disturbance. As described in the previous paragraph, the disturbance causes a modify in Q; the reaction will shift to re-establish Q = K.
Predicting the Direction of a Reversible Reaction
Le Châtelier'south principle can be used to predict changes in equilibrium concentrations when a organization that is at equilibrium is subjected to a stress. However, if we have a mixture of reactants and products that have non nevertheless reached equilibrium, the changes necessary to attain equilibrium may not be so obvious. In such a case, we tin compare the values of Q and Thousand for the system to predict the changes.
Effect of Change in Concentration on Equilibrium
A chemical system at equilibrium tin can be temporarily shifted out of equilibrium by adding or removing i or more of the reactants or products. The concentrations of both reactants and products then undergo additional changes to return the system to equilibrium.
The stress on the arrangement in Figure ane is the reduction of the equilibrium concentration of SCN– (lowering the concentration of 1 of the reactants would cause Q to be larger than K). Every bit a outcome, Le Châtelier's principle leads united states to predict that the concentration of Iron(SCN)2+ should decrease, increasing the concentration of SCN– office fashion back to its original concentration, and increasing the concentration of Fethree+ above its initial equilibrium concentration.
Figure 1. (a) The test tube contains 0.1 Thou Iron3+. (b) Thiocyanate ion has been added to solution in (a), forming the ruby Fe(SCN)2+ ion. Atomic number 263+(aq) + SCN−(aq) ⇌ Fe(SCN)ii+(aq). (c) Silver nitrate has been added to the solution in (b), precipitating some of the SCN− as the white solid AgSCN. Ag+(aq) + SCN−(aq) ⇌ AgSCN(south). The decrease in the SCN− concentration shifts the outset equilibrium in the solution to the left, decreasing the concentration (and lightening colour) of the Fe(SCN)ii+. (credit: modification of work by Marking Ott)
The event of a modify in concentration on a organization at equilibrium is illustrated further by the equilibrium of this chemical reaction:
[latex]{\text{H}}_{2}\left(g\right)+{\text{I}}_{ii}\left(chiliad\right)\rightleftharpoons2\text{HI}\left(g\correct){One thousand}_{c}=l.0\text{ at }400^\circ\text{C}[/latex]
The numeric values for this example have been determined experimentally. A mixture of gases at 400 °C with [H2] = [Itwo] = 0.221 M and [Hullo] = one.563 Chiliad is at equilibrium; for this mixture, Qc = Chiliadc = 50.0. If H2 is introduced into the system so rapidly that its concentration doubles earlier it begins to react (new [Hii] = 0.442 M), the reaction will shift so that a new equilibrium is reached, at which [H2] = 0.374 M, [I2] = 0.153 Grand, and [HI] = 1.692 One thousand. This gives:
[latex]{Q}_{c}=\frac{{\left[\text{HI}\right]}^{2}}{\left[{\text{H}}_{2}\right]\left[{\text{I}}_{2}\right]}=\frac{{\left(1.692\correct)}^{2}}{\left(0.374\right)\left(0.153\right)}=l.0={K}_{c}[/latex]
We have stressed this system by introducing additional H2. The stress is relieved when the reaction shifts to the right, using upward some (simply not all) of the excess Hii, reducing the corporeality of uncombined Itwo, and forming additional How-do-you-do.
Effect of Change in Pressure on Equilibrium
Sometimes we can modify the position of equilibrium past irresolute the force per unit area of a arrangement. Withal, changes in pressure take a measurable effect merely in systems in which gases are involved, and then only when the chemical reaction produces a change in the total number of gas molecules in the system. An easy mode to recognize such a system is to look for different numbers of moles of gas on the reactant and production sides of the equilibrium. While evaluating pressure (besides as related factors like volume), it is of import to remember that equilibrium constants are divers with regard to concentration (for Kc ) or fractional pressure level (for YardP ). Some changes to total pressure, similar calculation an inert gas that is not function of the equilibrium, will alter the total pressure simply non the partial pressures of the gases in the equilibrium abiding expression. Thus, add-on of a gas not involved in the equilibrium volition not perturb the equilibrium.
Check out this video to meet a dramatic visual demonstration of how equilibrium changes with pressure changes.
As nosotros increase the pressure of a gaseous organization at equilibrium, either past decreasing the book of the organisation or by adding more than of ane of the components of the equilibrium mixture, we introduce a stress past increasing the partial pressures of 1 or more of the components. In accordance with Le Châtelier'south principle, a shift in the equilibrium that reduces the total number of molecules per unit of volume will be favored considering this relieves the stress. The opposite reaction would be favored by a decrease in pressure.
Consider what happens when we increase the pressure on a system in which NO, O2, and NO2 are at equilibrium:
[latex]2\text{NO}\left(g\right)+{\text{O}}_{2}\left(m\correct)\rightleftharpoons2{\text{NO}}_{2}\left(yard\right)[/latex]
The germination of boosted amounts of NOtwo decreases the full number of molecules in the arrangement, because each time two molecules of NO2 form, a total of three molecules of NO and O2 are consumed. This reduces the total pressure exerted by the system and reduces, but does not completely relieve, the stress of the increased force per unit area. On the other mitt, a subtract in the pressure on the system favors decomposition of NO2 into NO and Oii, which tends to restore the pressure.
Now consider this reaction:
[latex]{\text{N}}_{2}\left(1000\correct)+{\text{O}}_{2}\left(g\right)\rightleftharpoons2\text{NO}\left(chiliad\correct)[/latex]
Because there is no modify in the total number of molecules in the system during reaction, a change in pressure does non favor either germination or decomposition of gaseous nitrogen monoxide.
Upshot of Change in Temperature on Equilibrium
Changing concentration or force per unit area perturbs an equilibrium because the reaction quotient is shifted away from the equilibrium value. Changing the temperature of a system at equilibrium has a different effect: A alter in temperature really changes the value of the equilibrium abiding. Withal, we can qualitatively predict the effect of the temperature change by treating it as a stress on the system and applying Le Châtelier's principle.
When hydrogen reacts with gaseous iodine, oestrus is evolved.
[latex]{\text{H}}_{2}\left(g\correct)+{\text{I}}_{two}\left(g\right)\rightleftharpoons2\text{Hullo}\left(chiliad\correct)\Delta H=-nine.4\text{kJ}\left(\text{exothermic}\correct)[/latex]
Because this reaction is exothermic, we tin can write information technology with heat as a product.
[latex]{\text{H}}_{two}\left(thousand\right)+{\text{I}}_{2}\left(g\right)\rightleftharpoons2\text{Howdy}\left(g\right)+\text{heat}[/latex]
Increasing the temperature of the reaction increases the internal free energy of the arrangement. Thus, increasing the temperature has the issue of increasing the corporeality of one of the products of this reaction. The reaction shifts to the left to relieve the stress, and there is an increase in the concentration of Htwo and Iii and a reduction in the concentration of How-do-you-do. Lowering the temperature of this system reduces the amount of energy present, favors the production of rut, and favors the formation of hydrogen iodide.
When nosotros change the temperature of a system at equilibrium, the equilibrium constant for the reaction changes. Lowering the temperature in the HI system increases the equilibrium abiding: At the new equilibrium the concentration of HI has increased and the concentrations of Htwo and I2 decreased. Raising the temperature decreases the value of the equilibrium constant, from 67.v at 357 °C to l.0 at 400 °C.
Temperature affects the equilibrium between NOii and NorthiiO4 in this reaction
[latex]{\text{N}}_{2}{\text{O}}_{4}\left(g\right)\rightleftharpoons2{\text{NO}}_{ii}\left(grand\right)\Delta H=57.twenty\text{kJ}[/latex]
The positive ΔH value tells the states that the reaction is endothermic and could be written
[latex]\text{rut}+{\text{N}}_{2}{\text{O}}_{four}\left(g\right)\rightleftharpoons2{\text{NO}}_{2}\left(g\right)[/latex]
At higher temperatures, the gas mixture has a deep brownish color, indicative of a significant corporeality of brown NO2 molecules. If, yet, nosotros put a stress on the system by cooling the mixture (withdrawing energy), the equilibrium shifts to the left to supply some of the free energy lost by cooling. The concentration of colorless Due northiiO4 increases, and the concentration of brown NO2 decreases, causing the brown colour to fade.
This interactive animation allows y'all to employ Le Châtelier'due south principle to predict the effects of changes in concentration, pressure, and temperature on reactant and product concentrations.
Catalysts Do Not Bear on Equilibrium
Every bit we learned during our study of kinetics, a catalyst can speed up the charge per unit of a reaction. Though this increase in reaction rate may cause a system to attain equilibrium more than quickly (past speeding up the forrard and reverse reactions), a catalyst has no effect on the value of an equilibrium constant nor on equilibrium concentrations.
The interplay of changes in concentration or pressure, temperature, and the lack of an influence of a catalyst on a chemic equilibrium is illustrated in the industrial synthesis of ammonia from nitrogen and hydrogen according to the equation
[latex]{\text{Due north}}_{2}\left(k\correct)+3{\text{H}}_{two}\left(g\right)\rightleftharpoons2{\text{NH}}_{3}\left(g\right)[/latex]
A large quantity of ammonia is manufactured by this reaction. Each year, ammonia is among the top x chemicals, by mass, manufactured in the world. Most 2 billion pounds are manufactured in the United States each year.
Ammonia plays a vital role in our global economy. It is used in the product of fertilizers and is, itself, an important fertilizer for the growth of corn, cotton, and other crops. Large quantities of ammonia are converted to nitric acrid, which plays an important function in the production of fertilizers, explosives, plastics, dyes, and fibers, and is too used in the steel manufacture.
Fritz Haber
Figure two. The work of Nobel Prize recipient Fritz Haber revolutionized agricultural practices in the early on 20th century. His work also affected wartime strategies, calculation chemical weapons to the arms.
In the early 20th century, German language chemist Fritz Haber (Figure 2) adult a practical process for converting diatomic nitrogen, which cannot be used past plants as a nutrient, to ammonia, a course of nitrogen that is easiest for plants to absorb.
[latex]{\text{N}}_{2}\left(k\right)+three{\text{H}}_{two}\left(g\right)\rightleftharpoons2{\text{NH}}_{three}\left(yard\right)[/latex]
The availability of nitrogen is a strong limiting factor to the growth of plants. Despite accounting for 78% of air, diatomic nitrogen (N2) is nutritionally unavailable due the tremendous stability of the nitrogen-nitrogen triple bond. For plants to use atmospheric nitrogen, the nitrogen must exist converted to a more bioavailable class (this conversion is chosen nitrogen fixation).
Haber was born in Breslau, Prussia (presently Wroclaw, Poland) in December 1868. He went on to study chemical science and, while at the University of Karlsruhe, he developed what would afterward be known every bit the Haber process: the catalytic formation of ammonia from hydrogen and atmospheric nitrogen under loftier temperatures and pressures. For this work, Haber was awarded the 1918 Nobel Prize in Chemistry for synthesis of ammonia from its elements. The Haber procedure was a boon to agriculture, as it allowed the product of fertilizers to no longer be dependent on mined feed stocks such as sodium nitrate. Currently, the annual production of constructed nitrogen fertilizers exceeds 100 million tons and synthetic fertilizer production has increased the number of humans that arable state can support from 1.9 persons per hectare in 1908 to iv.iii in 2008.
In addition to his work in ammonia production, Haber is besides remembered past history as ane of the fathers of chemic warfare. During World War I, he played a major office in the development of poisonous gases used for trench warfare. Regarding his role in these developments, Haber said, "During peace fourth dimension a scientist belongs to the World, simply during war time he belongs to his country."[1] Haber defended the utilise of gas warfare against accusations that it was inhumane, saying that expiry was decease, past any ways it was inflicted. He stands as an example of the upstanding dilemmas that face scientists in times of state of war and the double-edged nature of the sword of science.
Like Haber, the products fabricated from ammonia can be multifaceted. In addition to their value for agronomics, nitrogen compounds tin also be used to accomplish destructive ends. Ammonium nitrate has also been used in explosives, including improvised explosive devices. Ammonium nitrate was i of the components of the bomb used in the attack on the Alfred P. Murrah Federal Building in downtown Oklahoma City on April 19, 1995.
It has long been known that nitrogen and hydrogen react to grade ammonia. However, it became possible to manufacture ammonia in useful quantities by the reaction of nitrogen and hydrogen only in the early 20th century after the factors that influence its equilibrium were understood.
To exist applied, an industrial procedure must give a large yield of product relatively speedily. One way to increase the yield of ammonia is to increase the pressure on the organization in which Northward2, H2, and NH3 are at equilibrium or are coming to equilibrium.
[latex]{\text{North}}_{ii}\left(g\right)+iii{\text{H}}_{ii}\left(g\right)\rightleftharpoons2{\text{NH}}_{3}\left(thousand\right)[/latex]
The formation of boosted amounts of ammonia reduces the full pressure exerted by the system and somewhat reduces the stress of the increased pressure.
Although increasing the pressure of a mixture of N2, H2, and NH3 will increase the yield of ammonia, at low temperatures, the charge per unit of formation of ammonia is slow. At room temperature, for instance, the reaction is so slow that if we prepared a mixture of Nii and H2, no detectable amount of ammonia would form during our lifetime. The germination of ammonia from hydrogen and nitrogen is an exothermic process:
[latex]{\text{N}}_{2}\left(g\right)+iii{\text{H}}_{2}\left(chiliad\right)\rightarrow 2{\text{NH}}_{3}\left(g\correct)\Delta H=-92.2\text{kJ}[/latex]
Thus, increasing the temperature to increase the rate lowers the yield. If we lower the temperature to shift the equilibrium to favor the germination of more ammonia, equilibrium is reached more slowly because of the large decrease of reaction rate with decreasing temperature.
Role of the rate of formation lost past operating at lower temperatures tin can be recovered by using a catalyst. The net outcome of the goad on the reaction is to crusade equilibrium to be reached more rapidly.
In the commercial production of ammonia, conditions of nearly 500 °C, 150–900 atm, and the presence of a catalyst are used to give the all-time compromise among charge per unit, yield, and the cost of the equipment necessary to produce and contain high-force per unit area gases at loftier temperatures (Figure iii).
Effigy 3. Commercial production of ammonia requires heavy equipment to handle the loftier temperatures and pressures required. This schematic outlines the blueprint of an ammonia institute.
Key Concepts and Summary
Systems at equilibrium can be disturbed by changes to temperature, concentration, and, in some cases, volume and pressure level; volume and force per unit area changes will disturb equilibrium if the number of moles of gas is different on the reactant and product sides of the reaction. The system's response to these disturbances is described by Le Châtelier's principle: The organisation volition answer in a way that counteracts the disturbance. Not all changes to the arrangement consequence in a disturbance of the equilibrium. Adding a goad affects the rates of the reactions but does not modify the equilibrium, and irresolute pressure or book volition non significantly disturb systems with no gases or with equal numbers of moles of gas on the reactant and production side.
| Table 1. Furnishings of Disturbances of Equilibrium and 1000 | |||
|---|---|---|---|
| Disturbance | Observed Change equally Equilibrium is Restored | Direction of Shift | Effect on M |
| reactant added | added reactant is partially consumed | toward products | none |
| product added | added product is partially consumed | toward reactants | none |
| subtract in book/increase in gas pressure | pressure decreases | toward side with fewer moles of gas | none |
| increase in book/subtract in gas pressure | pressure increases | toward side with fewer moles of gas | none |
| temperature increase | estrus is absorbed | toward products for endothermic, toward reactants for exothermic | changes |
| temperature subtract | heat is given off | toward reactants for endothermic, toward products for exothermic | changes |
Exercises
- The following equation represents a reversible decomposition: [latex]{\text{CaCO}}_{three}\left(s\right)\rightleftharpoons\text{CaO}\left(southward\right)+{\text{CO}}_{2}\left(g\right)[/latex]
Under what conditions volition decomposition in a closed container proceed to completion so that no CaCOiii remains? - Explain how to recognize the conditions under which changes in pressure would touch on systems at equilibrium.
- What property of a reaction can we use to predict the effect of a modify in temperature on the value of an equilibrium constant?
- What would happen to the color of the solution in part (b) of Effigy 1 if a pocket-sized amount of NaOH were added and Fe(OH)3 precipitated? Explain your respond.
- The following reaction occurs when a burner on a gas stove is lit: [latex]{\text{CH}}_{4}\left(chiliad\right)+two{\text{O}}_{two}\left(thousand\right)\rightleftharpoons{\text{CO}}_{two}\left(k\right)+ii{\text{H}}_{2}\text{O}\left(1000\correct)[/latex]
Is an equilibrium among CHiv, O2, COtwo, and HtwoO established under these conditions? Explicate your respond. - A necessary step in the manufacture of sulfuric acrid is the germination of sulfur trioxide, SO3, from sulfur dioxide, SOii, and oxygen, O2, shown below. At loftier temperatures, the rate of formation of SO3 is higher, but the equilibrium amount (concentration or partial pressure) of Then3 is lower than information technology would be at lower temperatures. [latex]2{\text{Then}}_{2}\left(k\right)+{\text{O}}_{2}\left(chiliad\right)\rightarrow ii{\text{SO}}_{3}\left(g\right)[/latex]
- Does the equilibrium constant for the reaction increase, decrease, or remain about the same every bit the temperature increases?
- Is the reaction endothermic or exothermic?
- Propose four means in which the concentration of hydrazine, N2H4, could be increased in an equilibrium described past the following equation: [latex]{\text{N}}_{two}\left(thou\right)+2{\text{H}}_{2}\left(g\right)\rightleftharpoons{\text{N}}_{2}{\text{H}}_{4}\left(g\correct)\Delta H=95\text{kJ}[/latex]
- Propose four ways in which the concentration of PHthree could be increased in an equilibrium described by the following equation: [latex]{\text{P}}_{4}\left(g\right)+half dozen{\text{H}}_{two}\left(yard\correct)\rightleftharpoons4{\text{PH}}_{3}\left(g\correct)\Delta H=110.v\text{kJ}[/latex]
- How volition an increment in temperature affect each of the following equilibria? How will a decrease in the book of the reaction vessel impact each?
- [latex]2{\text{NH}}_{3}\left(g\right)\rightleftharpoons{\text{North}}_{two}\left(g\correct)+3{\text{H}}_{ii}\left(thousand\correct)\Delta H=92\text{kJ}[/latex]
- [latex]{\text{N}}_{2}\left(g\correct)+{\text{O}}_{2}\left(m\right)\rightleftharpoons2\text{NO}\left(g\right)\Delta H=181\text{kJ}[/latex]
- [latex]ii{\text{O}}_{3}\left(yard\right)\rightleftharpoons3{\text{O}}_{2}\left(g\right)\Delta H=-285\text{kJ}[/latex]
- [latex]\text{CaO}\left(s\correct)+{\text{CO}}_{2}\left(1000\right)\rightleftharpoons{\text{CaCO}}_{3}\left(s\right)\Delta H=-176\text{kJ}[/latex]
- How will an increment in temperature affect each of the following equilibria? How will a decrease in the volume of the reaction vessel affect each?
- [latex]2{\text{H}}_{two}\text{O}\left(g\right)\rightleftharpoons2{\text{H}}_{two}\left(one thousand\right)+{\text{O}}_{2}\left(g\right)\Delta H=484\text{kJ}[/latex]
- [latex]{\text{N}}_{2}\left(g\right)+3{\text{H}}_{2}\left(1000\correct)\rightleftharpoons2{\text{NH}}_{3}\left(g\right)\Delta H=-92.2\text{kJ}[/latex]
- [latex]2\text{Br}\left(g\right)\rightleftharpoons{\text{Br}}_{2}\left(g\right)\Delta H=-224\text{kJ}[/latex]
- [latex]{\text{H}}_{two}\left(g\correct)+{\text{I}}_{2}\left(s\right)\rightleftharpoons2\text{HI}\left(chiliad\right)\Delta H=53\text{kJ}[/latex]
- Water gas is a 1:1 mixture of carbon monoxide and hydrogen gas and is called h2o gas because it is formed from steam and hot carbon in the following reaction: [latex]{\text{H}}_{two}\text{O}\left(m\correct)+\text{C}\left(southward\right)\rightleftharpoons{\text{H}}_{2}\left(grand\right)+\text{CO}\left(yard\right)\text{.}[/latex] Methanol, a liquid fuel that could possibly replace gasoline, can be prepared from h2o gas and hydrogen at high temperature and pressure in the presence of a suitable catalyst.
- Write the expression for the equilibrium constant (Thousandc ) for the reversible reaction
[latex]ii{\text{H}}_{two}\left(g\correct)+\text{CO}\left(yard\right)\rightleftharpoons{\text{CH}}_{3}\text{OH}\left(g\right)\Delta H=-xc.2\text{kJ}[/latex] - What will happen to the concentrations of Htwo, CO, and CH3OH at equilibrium if more H2 is added?
- What will happen to the concentrations of Hii, CO, and CH3OH at equilibrium if CO is removed?
- What will happen to the concentrations of H2, CO, and CH3OH at equilibrium if CHiiiOH is added?
- What volition happen to the concentrations of Hii, CO, and CH3OH at equilibrium if the temperature of the system is increased?
- What will happen to the concentrations of H2, CO, and CH3OH at equilibrium if more goad is added?
- Write the expression for the equilibrium constant (Thousandc ) for the reversible reaction
- Nitrogen and oxygen react at high temperatures.
- Write the expression for the equilibrium abiding (Chiliadc ) for the reversible reaction
[latex]{\text{N}}_{2}\left(g\right)+{\text{O}}_{ii}\left(thou\right)\rightleftharpoons2\text{NO}\left(k\correct)\Delta H=181\text{kJ}[/latex] - What will happen to the concentrations of Nii, O2, and NO at equilibrium if (i) more O2 is added?
- What will happen to the concentrations of N2, O2, and NO at equilibrium if N2 is removed?
- What will happen to the concentrations of Northii, Otwo, and NO at equilibrium if NO is added?
- What will happen to the concentrations of Ntwo, Otwo, and NO at equilibrium if the pressure on the arrangement is increased by reducing the volume of the reaction vessel?
- What will happen to the concentrations of Ntwo, Otwo, and NO at equilibrium if the temperature of the system is increased?
- What will happen to the concentrations of N2, O2, and NO at equilibrium if a goad is added?
- Write the expression for the equilibrium abiding (Chiliadc ) for the reversible reaction
- Water gas, a mixture of H2 and CO, is an important industrial fuel produced past the reaction of steam with reddish hot coke, substantially pure carbon.
- Write the expression for the equilibrium constant for the reversible reaction
[latex]\text{C}\left(s\right)+{\text{H}}_{2}\text{O}\left(g\right)\rightleftharpoons\text{CO}\left(g\right)+{\text{H}}_{2}\left(chiliad\right)\Delta H=131.30\text{kJ}[/latex] - What will happen to the concentration of each reactant and product at equilibrium if more C is added?
- What will happen to the concentration of each reactant and product at equilibrium if H2O is removed?
- What will happen to the concentration of each reactant and product at equilibrium if CO is added?
- What will happen to the concentration of each reactant and product at equilibrium if the temperature of the system is increased?
- Write the expression for the equilibrium constant for the reversible reaction
- Pure iron metallic tin be produced by the reduction of fe(III) oxide with hydrogen gas.
- Write the expression for the equilibrium constant (Kc ) for the reversible reaction
[latex]{\text{Fe}}_{2}{\text{O}}_{3}\left(due south\right)+three{\text{H}}_{ii}\left(g\right)\rightleftharpoons2\text{Fe}\left(south\correct)+three{\text{H}}_{2}\text{O}\left(g\right)\Delta H=98.seven\text{kJ}[/latex] - What will happen to the concentration of each reactant and product at equilibrium if more Atomic number 26 is added?
- What will happen to the concentration of each reactant and product at equilibrium if HiiO is removed?
- What will happen to the concentration of each reactant and product at equilibrium if H2 is added?
- What will happen to the concentration of each reactant and production at equilibrium if the pressure on the system is increased by reducing the book of the reaction vessel?
- What will happen to the concentration of each reactant and production at equilibrium if the temperature of the system is increased?
- Write the expression for the equilibrium constant (Kc ) for the reversible reaction
- Ammonia is a weak base that reacts with h2o co-ordinate to this equation: [latex]{\text{NH}}_{3}\left(aq\right)+{\text{H}}_{two}\text{O}\left(l\right)\rightleftharpoons{\text{NH}}_{4}{}^{+}\left(aq\right)+{\text{OH}}^{-}\left(aq\right)[/latex]
Will any of the following increase the percent of ammonia that is converted to the ammonium ion in water?- Addition of NaOH
- Addition of HCl
- Addition of NHivCl
- Acetic acid is a weak acid that reacts with water according to this equation: [latex]{\text{CH}}_{three}{\text{CO}}_{2}\text{H}\left(aq\right)+{\text{H}}_{2}\text{O}\left(aq\right)\rightleftharpoons{\text{H}}_{3}{\text{O}}^{+}\left(aq\correct)+{\text{CH}}_{three}{\text{CO}}_{ii}{}^{-}\left(aq\right)[/latex]
Will any of the following increase the percent of acetic acid that reacts and produces [latex]{\text{CH}}_{3}{\text{CO}}_{2}{}^{-}[/latex] ion?- Addition of HCl
- Addition of NaOH
- Addition of NaCH3COii
- Propose two ways in which the equilibrium concentration of Ag+ can exist reduced in a solution of Na+, Cl–, Ag+, and [latex]{\text{NO}}_{three}{}^{\text{-}}[/latex], in contact with solid AgCl.
[latex]{\text{Na}}^{+}\left(aq\correct)+{\text{Cl}}^{-}\left(aq\right)+{\text{Ag}}^{+}\left(aq\right)+{\text{NO}}_{three}{}^{-}\left(aq\right)\rightleftharpoons\text{AgCl}\left(south\right)+{\text{Na}}^{+}\left(aq\right)+{\text{NO}}_{three}{}^{-}\left(aq\right)[/latex]
[latex]\Delta H=-65.9\text{kJ}[/latex] - How can the force per unit area of water vapor be increased in the following equilibrium? [latex]{\text{H}}_{ii}\text{O}\left(l\right)\rightleftharpoons{\text{H}}_{2}\text{O}\left(k\correct)\Delta H=41\text{kJ}[/latex]
- Additional solid silvery sulfate, a slightly soluble solid, is added to a solution of silver ion and sulfate ion at equilibrium with solid silver sulfate: [latex]two{\text{Ag}}^{+}\left(aq\right)+{\text{And so}}_{four}{}^{two-}\left(aq\correct)\rightleftharpoons{\text{Ag}}_{ii}{\text{SO}}_{iv}\left(s\right)[/latex]
Which of the following will occur?- Ag+ or [latex]{\text{SO}}_{4}{}^{2-}[/latex] concentrations volition not modify.
- The added silver sulfate will dissolve.
- Boosted argent sulfate volition course and precipitate from solution as Ag+ ions and [latex]{\text{SO}}_{4}{}^{2-}[/latex] ions combine.
- The Ag+ ion concentration will increase and the [latex]{\text{So}}_{four}{}^{2-}[/latex] ion concentration will decrease.
Glossary
Le Châtelier'southward principle: when a chemical system at equilibrium is disturbed, information technology returns to equilibrium by counteracting the disturbance
position of equilibrium: concentrations or partial pressures of components of a reaction at equilibrium (commonly used to depict conditions before a disturbance)
stress: change to a reaction's conditions that may cause a shift in the equilibrium
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