3.6_Aqueous_AS91392

=3.6 AS91392 Demonstrate understanding of equilibrium principles in aqueous systems (5 credits)=

__ SOLUBILITY CONSTANT __

 * [[image:ksquestion.jpg width="286" height="129" link="@http://wps.prenhall.com/wps/media/objects/4679/4792059/ch16_10.htm"]] || [[image:phetsolubility.jpg width="259" height="198" link="@https://phet.colorado.edu/en/simulation/soluble-salts"]] || [[image:ksgraph.jpg width="352" height="132" link="@http://employees.oneonta.edu/viningwj/sims/solubility_product_constant_s.html"]] ||

In an experiment, a saturated solution was made by dissolving 1.44 × 10–3 g of Ag2CrO4 in water, and making it up to a volume of 50.0 mL.
====M (Ag2CrO4) = 332 g mol–1. Calculate the solubility of Ag2CrO4(s), and hence give the [Ag+] and [CrO42–] in the solution.Determine the Ks(Ag2CrO4). ====

[[image:NCEA2009.jpg]]
====Solid sodium chloride is added to 5.00 L of 0.100 mol L–1 silver nitrate solution. Calculate the minimum mass of sodium chloride that would be needed to produce a saturated solution of AgCl. Assume that there is no change in volume when the sodium chloride is added. M(NaCl) = 58.5 g mol–1 ====

a) Describe what is meant by the term ‘solubility’.
====b) The chloride ion concentration in sea water can be determined by titrating a sample with aqueous silver nitrate (AgNO3) using potassium chromate (K2CrO4) as the indicator. As the silver nitrate is added, a precipitate of silver chloride, (AgCl) forms. When most of the AgCl has precipitated, the Ag+(aq) concentration becomes high enough for a red precipitate of Ag2CrO4 to form. Show that the solubility of Ag2CrO4 in pure water at 25°C is higher than that of AgCl. Ks(AgCl) = 1.56 × 10–10 Ks(Ag2CrO4) = 1.30 × 10–12 ==== ====c) If the concentration of chromate ions is 6.30 × 10–3 mol L–1 at the point when the Ag2CrO4 starts to precipitate, calculate the concentration of Ag+ ions in the solution. ====



NCEA PAST EXAM QUESTION: Solubility of solids in solutions with a common ion
  ==== A sample of seawater has a chloride ion concentration of 0.440 mol L –1. Determine whether a precipitate of lead(II) chloride will form when a 2.00 g sample of lead(II) nitrate is added to 500 mL of the seawater. //K// s (PbCl 2 ) = 1.70 × 10 –5 // M // (Pb(NO 3 ) 2 ) = 331 g mol –1 ====  ==== Determine whether a precipitate of iron(III) hydroxide, Fe(OH) 3, will form when Fe(NO3)3 is dissolved in water. [Fe(NO3)3] = 1.05 × 10–4 mol L–1. Assume the pH of the water is 7. //K//s(Fe(OH)3) = 2.00 × 10–39 ====

Discuss how the solubility of Ag 2 CrO  4  will change if it is dissolved in 0.1 mol L  –1  K  2  CrO  4  // No calculations // //are necessary.//
==== Sea water contains many dissolved salts. The chloride ion concentration in a sample of sea water is 0.440 mol L–1. Determine whether a precipitate of lead(II)chloride will form when a 1.00 g sample of lead(II) nitrate is added to 500 mL of the sea water. Your answer must be clearly justified. // M // (Pb(NO3)2) = 331 g mol–1 ==== ====Sea-water contains appreciable amounts of ions other than Na+ and Cl–. Evaporating the sea-water to dryness would produce a mixture of salts including NaCl. However, precipitation of NaCl occurs if concentrated hydrochloric acid is added to a saturated NaCl solution. Explain why this precipitation occurs. ====

NCEA PAST EXAM QUESTION: Solubility of solids in solutions forming a complex ion
  ====In another experiment, 0.0100 g of Ag2CrO4in beaker A was made up to a volume of 50.0 mL with water. In beaker B, 0.0100 g of Ag2CrO4was made up to a volume of 50.0 mL with 0.100 mol L–1ammonia solution. ====

<span style="font-family: Arial,Helvetica,sans-serif;">Compare and contrast the solubility of Ag2CrO4in beaker A and beaker B. No calculations are necessary.
====<span style="color: #f95725; font-family: Arial,Helvetica,sans-serif;"> Solid sodium chloride is added to 5.00 L of 0.100 mol L –1 silver nitrate solution. Calculate the minimum mass of sodium chloride that would be needed to produce a saturated solution of AgCl. Assume that there is no change in volume when the sodium chloride is added. //M//(NaCl) = 58.5 g mol –1 ====

<span style="color: #f95725; font-family: Arial,Helvetica,sans-serif;">NCEA PAST EXAM QUESTION: Solubility of solids in solutions with a change in pH
<span style="color: #f95725; font-family: Arial,Helvetica,sans-serif;"> <span style="color: #f95725; font-family: Arial,Helvetica,sans-serif;"> ==== The solubility of zinc hydroxide, Zn(OH) 2, can be altered by changes in pH. Some changes in pH may lead to the formation of complex ions, such as the zincate ion, [Zn(OH) 4 ] 2– Use equilibrium principles to explain why the solubility of zinc hydroxide increases when the pH is less than 4 or greater than 10. ====

====<span style="font-family: Arial,Helvetica,sans-serif;">1) Some sulfides have very low solubility products. When hydrogen sulfide gas is bubbled through solutions of these ions, these ions separate from a mixture of ions. i) In a saturated solution of hydrogen sulfide [H3O+]2[S2–] = 1.10 x 10–23Calculate the sulfide ion concentration when the pH of the solution is 4.20. ====

<span style="font-family: Arial,Helvetica,sans-serif;">ii) Calculate the solubility of FeS in this solution, in mol L–1. Ks(FeS) = 4.90 x 10–18
====<span style="font-family: Arial,Helvetica,sans-serif;">2) A solution contains a mixture of the two metal ions Cu2+ and Zn2+, both of the same concentration. The solution is saturated with hydrogen sulfide and adding hydrochloric acid lowers the pH of the solution. Ks(CuS) = 6.30 x 10–36 Ks(ZnS) = 1.6 x 10–24 ==== ====<span style="font-family: Arial,Helvetica,sans-serif;">Account for the fact that at a pH close to 7 all the metal sulfides will precipitate whereas only the most insoluble sulfides precipitate out at a lower pH. ====

<span style="font-family: Arial,Helvetica,sans-serif;">Discuss the effect of decreasing the pH of the water on the solubility of Fe(OH)3.
====<span style="font-family: Arial,Helvetica,sans-serif;">A saturated solution of zinc hydroxide, Zn(OH)2 contains a small amount of solid Zn(OH)2 at the bottom of the container. The pH of the solution is ==== ====<span style="font-family: Arial,Helvetica,sans-serif;">increased. Discuss the effect of increasing the pH on the amount of solid present, and also on the nature and concentration of the species present in the ====

<span style="font-family: Arial,Helvetica,sans-serif;">Discuss how the solubility of Ag2CrO4will change if it is dissolved in 0.1 mol L–1NH3No calculations are necessary.
<span style="font-family: Arial,Helvetica,sans-serif;"> ==== The //K// s of aluminium hydroxide, Al(OH) 3 , at 25°C, is 3 × 10 –34 , indicating that it has very low solubility. The solubility may be altered by changes in pH (due to acidic or basic properties) and formation of complex ions such as the aluminate ion, [Al(OH) 4  ]  –. Discuss why aluminium hydroxide becomes more soluble in aqueous solutions that have a pH less than 4, or a pH greater than 10. In your answer include: ====

<span style="font-family: Arial,Helvetica,sans-serif;">a discussion of the equilibrium principles involved.
<span style="font-family: Arial,Helvetica,sans-serif;"> ====<span style="font-family: Arial,Helvetica,sans-serif;">An aqueous ammonia solution has a pH of 10 and when phenolphthalein indicator is added it turns pink. Solid ammonium chloride is added to this solution and the solution turns colourless due to a decrease in pH. By considering the equilibrium systems, discuss why the pH of the solution decreased. Include a relevant equation in your answer. ==== <span style="font-family: Arial,Helvetica,sans-serif;"> ====<span style="font-family: Arial,Helvetica,sans-serif;">Sea-water contains appreciable amounts of ions other than Na+ and Cl–. As part of the process for extracting table salt from sea-water, sodium hydroxide is added to the sea-water to precipitate the magnesium ions as magnesium hydroxide. The concentration of Mg2+ ions at this stage is 0.555 mol L–1. Calculate the minimum hydroxide ion concentration and hence the pH of the solution needed for precipitation to occur. ====

<span style="font-family: Arial,Helvetica,sans-serif;">Ks (Mg(OH)2) = 7.10 × 10–12
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__<span style="color: #f95725; font-family: Arial,Helvetica,sans-serif;">RELATIVE CONCENTRATIONS OF SPECIES IN SOLUTION __

 * [[image:species.jpg width="560" height="148" link="@file:species in solution.pdf"]] || [[image:l3strongweak.jpg width="408" height="164" link="@file:strongweak.pdf"]] ||


 * [[image:l3phetpH.jpg width="368" height="275" link="@https://phet.colorado.edu/en/simulation/acid-base-solutions"]] || [[image:ionssolution.jpg width="560" height="171" link="@http://cdn.wwnorton.com/college/chemistry/chem3/chemtours/chapter_04/migration_ions/Interface.swf"]] ||

<span style="color: #f95725; font-family: Arial,Helvetica,sans-serif;">NCEA PAST EXAM QUESTION: Relative concentrations of species in solution
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When chlorine gas is added to water, the equation for the reaction is:
Cl 2  (//g//) + H  2  O( l ) <span style="font-family: 'Cambria Math',serif;">⇌ HCl(//aq//) + HOCl(//aq//)

** (ii) ** List all the species present when HOCl reacts with water, in order of decreasing concentration.Justify your order.
====<span style="font-family: Arial,Helvetica,sans-serif;">1 mol of each of the following substances was placed in separate flasks, and water was added to these flasks to give a total volume of 1 L for each solution, rank these solutions in order of increasing pH. Justify your choice and include equations where appropriate. ====

<span style="font-family: Arial,Helvetica,sans-serif;">HCl
====<span style="font-family: Arial,Helvetica,sans-serif;">The conductivity of the 1 mol L–1solutions formed in (a) can be measured, rank these solutions in order of decreasing conductivity. Compare and contrast the conductivity of each of the 1 mol L–1solutions, with reference to species in solution. ====

**a) i)** Write an equation for the reaction of methanoic acid with water.
====**ii)** Methanoic acid, HCOOH, is a weak acid. A dilute aqueous solution of this acid has a pH of 2.78. List all the species in the aqueous solution of methanoic acid in order of decreasing concentration.Give reasons for you answer.====

<span style="font-family: Arial,Helvetica,sans-serif;">NH3 NaCl NH4Cl HF
====<span style="font-family: Arial,Helvetica,sans-serif;">b) Discuss the relative concentrations of the species present in each of the 0.100 mol L–1solutions of NH3 and HF. You do not need to include water. ====

<span style="font-family: Arial,Helvetica,sans-serif;">(ii) Write the equation for the ammonium ion reacting with water.
====<span style="font-family: Arial,Helvetica,sans-serif;">(b) The bar chart below shows the relative concentrations of the species (excluding water) in a solution of 0.1 mol L–1NH4Cl. (The bar chart is not drawn to scale.) ==== Identify the species **A** to **E**. Justify your answer.

<span style="color: #211d1e; font-family: Arial,Helvetica,sans-serif;">The following table lists some properties of aqueous solutions of sodium hydroxide, methylamine and methylammonium chloride.
<span style="color: #211d1e; font-family: 'Times New Roman',serif; font-size: 15.3333px;">

<span style="font-family: Arial,Helvetica,sans-serif;">(b) Justify the differences in the pH and conductivity of the three solutions.
<span style="font-family: Arial,Helvetica,sans-serif;"> ====<span style="font-family: Arial,Helvetica,sans-serif;">When bromine is added to water, it forms hypobromous acid (HOBr), a weak acid. Write an equation to show the equilibrium system that is formed with hypobromous acid and water. ====

<span style="font-family: Arial,Helvetica,sans-serif;">(b) Explain why aqueous aminomethane, CH3NH2, is a weak electrolyte.
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<span style="font-family: Arial,Helvetica,sans-serif;">**iii)** A buffer solution
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<span style="font-family: Arial,Helvetica,sans-serif;">Give reasons for arranging in this order, including equations for any reactions occurring to produce solutions that do not have a pH of 7.
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<span style="color: #f95725; font-family: Arial,Helvetica,sans-serif;">NCEA PAST EXAM QUESTION: Acid Base strength Ka (pKa)
<span style="color: #f95725; font-family: Arial,Helvetica,sans-serif;"> <span style="font-family: Arial,Helvetica,sans-serif;"> ====<span style="font-family: Arial,Helvetica,sans-serif;">1. Hypochlorous acid has a pKa of 7.53. Another weak acid, hydrofluoric acid, HF, has a pKa of 3.17. A 0.100 mol L–1 solution of each acid was prepared by dissolving it in water. Compare the pHs of these two solutions. No calculations are necessary. ====

<span style="font-family: Arial,Helvetica,sans-serif;">**2.** Aqueous methylamine, CH3NH2, solution has a pH of 11.8. Show by calculation that the concentration of this solution is 0.0912 mol L–1.
====<span style="font-family: Arial,Helvetica,sans-serif;">1. 1 mol of each of the following substances was placed in separate flasks, and water was added to these flasks to give a total volume of 1 L for each solution. In the box below, rank these solutions in order of increasing pH. Justify your choice and include equations where appropriate. ====

====<span style="font-family: Arial,Helvetica,sans-serif;">**1.** A solution prepared by dissolving hydrogen fluoride in water has a pH of 2.34. Calculate the concentration of the hydrogen fluoride in the solution. pKa (HF) = 3.17 ==== ====<span style="font-family: Arial,Helvetica,sans-serif;">**2.** Glycolic acid, HOCH2COOH, is a monoprotic acid used in various skin-care products, and can be represented as HG. Glycolic acid has a pKa value of 3.83. ====

<span style="font-family: Arial,Helvetica,sans-serif;">[[image:NCEA2010.jpg]]
====<span style="font-family: Arial,Helvetica,sans-serif;">An aqueous solution of ammonium chloride (NH4Cl) has a pH of 4.66. Calculate the concentration of NH4Cl solution. pKa(NH4+) = 9.24 Ka = 5.75 × 10 –10 ====

<span style="font-family: Arial,Helvetica,sans-serif;"> [[image:NCEA2007.jpg]]
====<span style="font-family: Arial,Helvetica,sans-serif;">1. The pH of the solution in the stomach of a patient in hospital is 2.50. As a treatment, the patient is given a small volume of sodium citrate (Na3Cit) solution. Citric acid, H3Cit, is a triprotic acid. ====

<span style="font-family: Arial,Helvetica,sans-serif;"> (ii) Explain your choice, including an appropriate equation in your answer
====<span style="font-family: Arial,Helvetica,sans-serif;">2. An aqueous ammonia solution has a pH of 10 and when phenolphthalein indicator is added it turns pink. Solid ammonium chloride is added to this solution and the solution turns colourless due to a decrease in pH. By considering the equilibrium systems, discuss why the pH of the solution decreased. Include a relevant equation in your answer. ==== <span style="color: #f95725; font-family: Arial,Helvetica,sans-serif;">

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<span style="color: #f95727; font-family: Arial,Helvetica,sans-serif;">TITRATION CURVES

 * [[image:RSctitrationexperiment.jpg width="387" height="290" link="@http://www.rsc.org/learn-chemistry/resources/screen-experiment/titration/experiment/2"]] || [[image:l3simulator.jpg width="318" height="320" link="@file:titrationsimulator.swf"]] ||


 * [[image:l3species.jpg width="281" height="262" link="@file:l3species.swf"]] || [[image:l3pHtitration.jpg width="487" height="228" link="@file:pHtitration.swf"]] ||

<span style="color: #f95725; font-family: Arial,Helvetica,sans-serif;">NCEA PAST EXAM QUESTION: Titration curves to represent an acid-base system
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(ii) Calculate the pH of the titration mixture after 5.00 mL of NaOH has been added.
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20.00 mL of 0.125 mol L–1 ethanoic acid is titrated with 0.125 mol L–1 sodium hydroxide solution. The equation for this reaction is:
CH3COOH(//aq//) + NaOH(//aq//) --> CH3COONa(//aq//) + H2O (//l//)

**b) i)** Show that the pH at the equivalence point for this titration is 8.78.
p//K//a(CH3COOH ) = 9.24

**iii)** Phenolphthalein is an acid-base indicator. It is a weak acid and its formula can be represented as HIn. Phenolphthalein is colourless in acidic solutions and purple in basic solutions.
p//K//a (HIn) = 9.60

<span style="font-family: Arial,Helvetica,sans-serif;">• an explanation of the significance of the pKa in selecting an indicator.
====<span style="font-family: Arial,Helvetica,sans-serif;">Below is the titration curve for 10.0 mL of 0.100 mol L–1ethanoic acid being titrated with 0.100 mol L–1sodium hydroxide. Ethanoic acid can be represented by the symbol HEt. ====

<span style="font-family: Arial,Helvetica,sans-serif;">20.00 mL of 0.160 mol L–1 ammonia is titrated with 0.230 mol L–1 hydrochloric acid. The equation for the reaction is
<span style="font-family: Arial,Helvetica,sans-serif;">NH3+ HCl → NH4++ Cl–pKa(NH4+) = 9.24, Ka = 5.75 × 10–10

25.0 mL of 0.0500 mol L –1 benzoic acid solution (C  6  H  5  COOH) is titrated with 0.0500 mol L  –1  sodium hydroxide solution. The equation for the reaction is:
C 6  H  5  COOH(//aq//) + NaOH(//aq//) → C  6  H  5  COONa(//aq//) + H  2  O(//ℓ//)

The following titration curve shows the addition of aqueous 0.100 mol L –1 sodium hydroxide to a solution of hydrazoic acid, HN  3.
p//K// a (HN 3  ) = 4.72

** b) ** The initial pH of the hydrazoic acid (HN 3 ) is 2.6. Calculate the concentration of the HN  3  solution used in the titration.
====<span style="font-family: Arial,Helvetica,sans-serif;">**7)** A 0.160 mol L–1 solution of sodium hydroxide is titrated against 50 mL of aqueous propanoic acid, HPr. 40 mL of the sodium hydroxide solution was required to exactly react with the propanoic acid. The reaction occurring can be represented as: HPr(aq) + NaOH(aq) → NaPr(aq) + H2O Ka(HPr) = 1.35 × 10–5 ====



<span style="font-family: Arial,Helvetica,sans-serif;">**ii)** Discuss the ability of the solution formed to act as a buffer. Your answer should include relevant equations.
====<span style="font-family: Arial,Helvetica,sans-serif;">**d)** The equivalence point of the titration could also be found using an acid-base indicator. Which of the following indicators would be suitable to use? Explain your choice of indicator. ====

<span style="font-family: Arial,Helvetica,sans-serif;">The graph below shows the change in pH when 40.0 mL of 0.0500 mol L–1 aqueous NH3 is titrated with 0.200 mol L–1 aqueous HCl
<span style="display: block; font-family: Arial,Helvetica,sans-serif; text-align: left;">The equation for the reaction occurring during the titration is: NH3(aq) + HCl(aq) → NH4Cl(aq) <span style="display: block; font-family: Arial,Helvetica,sans-serif; text-align: left;">**a)** Use the curve to determine pKa(NH4+) and hence calculate Ka(NH4+).

<span style="font-family: Arial,Helvetica,sans-serif;">**ii)** Using Ka(NH4+) determined in part (a) on the previous page, determine the pH at the equivalence point of the second titration.
<span style="font-family: 'Times New Roman','serif';"> ====<span style="font-family: Arial,Helvetica,sans-serif;">The active ingredient in many sunscreens is para-aminobenzoic acid. It is a weak monoprotic acid and can be represented as HPab, while its conjugate base is Pab–. ====

<span style="font-family: Arial,Helvetica,sans-serif;">**c)** Calculate the concentration of H3O+ in the solution.
====<span style="font-family: Arial,Helvetica,sans-serif;">**d)** The concentration of the HPab solution was determined by titration. A 20.0 mL sample of the HPab solution required 12.0 mL of 0.0500 mol L–1 NaOH to reach the equivalence point. The ====

<span style="font-family: Arial,Helvetica,sans-serif;">**ii)** Using the results from parts (c) and (d)(i), show that pKa(HPab) = 4.92.
====<span style="font-family: Arial,Helvetica,sans-serif;">**e)** Would the pH at the equivalence point of the titration of HPab with NaOH be more than 7, less than 7 or equal to 7? Give reasons and include any relevant equations that support your answer. ==== ====<span style="font-family: Arial,Helvetica,sans-serif;">**f)** Using the information above, sketch a curve showing the change in pH against the volume of sodium hydroxide added to the 20.0 mL HPab solution in the flask. ====

__<span style="color: #f95727; font-family: Arial,Helvetica,sans-serif;">BUFFER SOLUTIONS __

 * [[image:l3bufferswwnorton.jpg width="174" height="262" link="@http://cdn.wwnorton.com/college/chemistry/chem4/chemtours/chapter_16/buffers/Interface.swf"]] || [[image:l3buffersadd.jpg width="365" height="265" link="@http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/buffer12.swf"]] || [[image:l3bufferschemcollective.jpg width="278" height="267" link="@http://chemcollective.org/activities/tutorials/buffers/buffers3"]] ||

<span style="color: #f95727; font-family: Arial,Helvetica,sans-serif;">NCEA QUESTIONS ON BUFFER SOLUTIONS
==== An aqueous solution containing a mixture of HF and sodium fluoride, NaF, can act as a buffer solution. Calculate the mass of NaF that must be added to 150 mL of 0.0500 mol L–1 HF to give a buffer solution with a pH of 4.02. Assume there is no change in volume. //M//(NaF) = 42.0 g mol–1 p//K//a(HF) = 3.17 ====

[[image:NCEA2012.jpg]]
==== **a)** A mixture of aqueous solutions of NH 3 and ammonium chloride, NH 4 Cl, can act as a buffer solution. Calculate the mass of NH 4 Cl required, when added to 250 mL of a 0.150 mol L –1 NH 3 solution, to give a buffer solution with a pH of 8.60. Assume there is no change in volume. ====

// M // (NH 4 Cl) = 53.5 g mol–1 p//K//a (NH 4 + ) = 9.24
• describe the function of a buffer solution • evaluate its effectiveness when small amounts of acid or base are added • include any relevant equations.
 * b**** ) ** Discuss the ability of the NH 3 / NH 4 Cl solution to act as a buffer at a pH of 8.60. In you answer you should:

//Your answer should include relevant equations.//
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<span style="display: block; height: 1px; left: -40px; overflow: hidden; position: absolute; top: 9883px; width: 1px;"> //K//s (Mg(OH)2) = 7.10 × 10–12