15.3 Multiple Equilibria

15.3 Multiple Equilibria

  • Check the work.
  • The silver ion concentration can be calculated by dissolving AgNO3 and KCN in enough water to make a solution.
  • Sometimes one equilibrium reaction doesn't adequately describe the system being studied.
    • Sometimes we have more than one type of equilibrium occurring at the same time.
  • Carbon dioxide in the air forms carbonic acid in the water.
  • Reefbuilding coral can be negatively impacted by ocean acidification as they cannot absorb the calcium carbonate they need to grow and maintain their skeletons.
    • The local bio system depends on the health of the reefs for its survival.
    • The disruptions to sea life can be felt around the world if enough local reefs are also affected.
    • The world's oceans are in the midst of a period of intense acidification, believed to have begun in the mid-nineteenth century, and which is now speeding up faster than any change to oceanic pH in the last 20 million years.
  • When coral are unable to adequately build and maintain their calcium carbonite skeletons because of excess ocean acidification, the reef is only capable of hosting a small fraction of the species as before, and the local food chain starts to collapse.
  • Learn how the ocean affects other marine creatures.
  • Coral reefs are specifically affected by ocean acidification.
  • Solids derived from weak acids are usually dissolved in strong acids.
  • The shift of the equilibrium concentrations to the right is a result of the decrease in the concentration of the anion.
  • The dissolution of Ca5(PO4)3OH in acid is relevant to us.
    • The principal mineral in the teeth is a biological form of hydroxylapatite, which is a subclass of apatites.
  • The crystal of the mineral is shown here.
    • The sample of pure apatite is colored because of the presence of impurities.
  • The reactions decrease the ion concentrations and the additional hydroxylapatite is dissolved in an acidic solution in accord with Le Chatelier's principle.
    • When acid waste is produced on our teeth, it causes the hydroxylapatite of the tooth to break down.
    • NaF, or stannous fluoride, is found in toothpastes.
    • Ca5(PO4)3F: NaF + Ca5(PO4)3OH is the result of replacing the OH- ion in hydroxylapatite with F- ion.
    • The fluorapatite is more resistant to attack by acids than the hydroxylapatite.
    • The chemistry in everyday life feature has information on the role of fluoride in preventing tooth decay.
  • Ca5(PO4)3F is formed by reacting with hydroxylapatite and is used to protect our teeth.
    • Since it doesn't have a hydroxide ion, fluorapatite is more resistant to attacks by acids in our mouths and protects our teeth.
    • Scientists discovered that naturally fluorinated water could be beneficial to your teeth, and so it became common practice to add fluoride to drinking water.
  • It is found in many toothpastes and helps prevent tooth decay.
  • Excess fluoride can affect its advantages.
    • Natural sources of drinking water in various parts of the world have varying concentrations of fluoride, and places where that concentration is high are prone to certain health risks when there is no other source of drinking water.
    • The most serious side effect of excess fluoride is bone disease.
    • The bones can stiffen due to excess fluoride in the body.
    • Mobility can be severely impacted by it.
  • Over two million people suffer from Skeletal fluorosis.
    • The US Environmental Protection Agency sets a maximum level of fluoride for drinking water in the US.
    • Dehydration is a problem in areas with high levels of fluoride in the water.
  • The acid attack on limestone and marble is similar to the acid attack on hydroxylapatite.
  • If enough acid is present, the concentration of carbonate ion is reduced to such a low level that the reaction quotient for the dissolution of calcium carbonate remains less than the solubility product of calcium carbonate, even after all of the calcium carbonate has dissolved.

  • The only significant reaction that can occur when a solid is dissolved is the formation of its ion.
  • The silver bromide reaction can be found in the OpenStax book at http://cnx.org/content/col11760/1.9 sp.
    • Adding S2 O3 causes Reaction (2) to shift to the right.
  • Unreacted S2O3 is present in Step 3.
  • The solution was prepared from 1.9 g Na2S2O3 and dissolved 1.0 g of AgBr.

  • A solution with NH3 dissolved 2.0 g of AgCl was prepared.
  • Le Chatelier's principle can be used to determine how to shift the concentration of ion in the equilibrium between a slightly soluble solid and a solution of its ion.
  • This could be important to a laundry because clothing washed in water with a high concentration of manganese may be stained by it.
    • The concentration of hydroxide ion can be increased to reduce the concentration of manganese.
    • We could add a small amount of NaOH or some other base such as the silicates found in laundry detergents.
    • As the concentration of OH- ion increases, the equilibrium responds by shifting to the left and decreasing the amount of Mn2+ ion in the equilibrium mixture.
  • At the new equilibrium, [OH-] is less than in the solution of Mg(OH)2 in pure water.
  • There is more solid Mg(OH)2.
  • At the new equilibrium, [OH-] is greater and [Mg2+] is less than in pure water.
    • There is more solid Mg(OH)2.
  • This OpenStax book is available for free at http://cnx.org/content/col11760/1.9 to relieve the stress of less OH- ion.
    • At the new equilibrium, [OH-] is less than in the solution of Mg(OH)2 in pure water.
    • More Mg(OH)2 is dissolved.

  • The OpenStax book is available for free at http://cnx.org/content/col11760/1.9 reaction exceeds the magnitude of the solubility product.
    • The reaction quotient is equal to the product.
  • A reagent can be added to allow one ion to leave the solution.
    • The common ion effect can play a role in precipitation reactions.
    • In the case of an ion with one of the other ion in the solution, Le Chatelier's principle applies and more precipitate comes out of solution so that the molar solubility is reduced.
  • Lewis proposed a definition for acids and bases that relied on an atom's or molecule's ability to accept or donate electron pairs.
    • A Lewis acid can accept an electron pair, whereas a Lewis base cannot.
    • Lewis acid-base adducts are examples of complex ion.
    • In a complex ion, we have a central atom, often consisting of a transition metal cation, which acts as a Lewis acid, and several neutral molecule or ion surrounding it.
    • electron pairs are shared to form bonds.
    • Formation of complex ion in solution can affect the solubility of transition metal.
  • Multiple equilibria are systems where two or more equilibria processes are happening at the same time.
    • Acid rain, fluoridation, and dissolution of carbon dioxide in sea water are some examples.
    • When looking at these systems, we need to consider each equilibrium separately and combine the individual equilibrium constants into one product or reaction quotient expression using the tools from the first equilibrium chapter.
    • Each reaction in a multiple equilibria system will shift toward reactants or products based on what is added to the initial reaction and how it affects each subsequent equilibrium reaction.
  • The following compounds havelubilities in grams per 100 mL of water according to the Handbook of Chemistry and Physics.
    • If the volume doesn't change on dissolution, you can calculate the solubility product for each compound.
  • Predict which of the salts has the most moles per liter in pure water by using solubility products.

  • Determine which of the calcium salts has the least amount of moles per liter and the most amount of grams per liter.
  • If no equilibria other than dissolution are involved, calculate the concentrations of ion in a saturated solution of each of the following.
  • The concentrations are found in a mixture of the two.

  • These calculations are for nickel carbonate.
  • Both NiCO3 and CoCO3 are dissolved in the same solution.
  • It is possible that it is necessary to control the pH.
  • A solution has 1.0 x 10-5 mol of KBr and 0.10 mol of KCl per liter.
    • Gradually AgNO3 is added to the solution.
  • A solution has 1.0 x 10-2 mol of KI and 0.10 mol of KCl per liter.
    • Gradually AgNO3 is added to the solution.
  • When a blood sample is drawn for a laboratory test, K2C2O4 is used because it removes calcium from the sample.
    • It is necessary to remove 1.0% of the Ca2+ in order to prevent coagulation.
    • The volume of the blood sample is not known.
  • 50% of the stones in the urinary calculi are composed of calcium phosphate.
    • The urine contains 0.10 g of Ca2+ per day.
    • The mid range amount of urine may be taken as 1.4 L per day.

  • The complete exercise is used.
    • Discuss how values change when you change salts.
  • Two hypothetical salts have the same H2O solubility.

  • The dissociation constant is calculated.
  • A roll of 35-mm black and white photographic film has 0.27 g of unexposed AgBr before it develops.
  • Before precipitation begins, AgNO3 is added.
  • The weak acid in the solution is HF.
    • A strong acid is pure liquid HF.
    • HNO3 is a base and accepts protons.
    • Adding one of the Lewis acids and accepting F- ion can increase the acidity of liquid HF.
    • Balanced chemical equations are needed for the reaction of pure HNO3 with pure HF.
  • H2NCH2CO2H is the simplest of the amino acids.
    • The functional groups of amino acids are an amine group, -NH2, and a carboxylic acid group, -CO2H.
    • An acid or a base can be found in an amino acid.
    • The acid strength of the carboxyl group is the same as that of acetic acid, CH3CO2H, and the base strength is slightly greater than that of ammonia, NH3.
  • Boric acid is a Lewis acid.
  • A system in equilibrium is said to be a saturated solution of a slightly soluble electrolyte and some solid electrolyte.
  • The volume should be 1.00 L.

  • NH3 has dissolved AgI and AgCl.
  • CaCO3, SrCO3 and BaCO3 are required to be dissolved by CH3CO2H.
  • Ca(OH)2 is an inexpensive base, but it is limited in its use.
  • When they reach the small intestine, magnesium hydroxide and magnesium citrate act as mild laxatives.
  • Two different substances have the same effect in the small intestine.
  • The Educational Testing Service gave permission for the following question to be used.
  • The equilibrium is represented by an equation.
  • The KF dissolved.
    • The volume change is insignificant.
  • You should give an explanation to support your conclusion.
  • List the chemical species in decreasing order of their concentrations, if you can.
  • NH3 is mixed with an equal volume of a solution.