Experiment 21 Hard Water Analysis Deposits of hardening ions (generally calcium carbonate deposits) can reduce the flow of water in plumbing • To learn the cause and effects of hard water • To determine the hardness of a water sample Objectives The following techniques are used in the Experimental Procedure: Techniques Hardening ions present in natural waters are the result of slightly acidic rainwater owing over mineral deposits of varying compositions; the acidic rainwater1 reacts with the very slightly soluble carbonate salts of calcium and magnesium and with various ironcontaining rocks A partial dissolution of these salts releases the ions into the water supply, which may be surface water or groundwater Introduction Ϫ CO2(aq) ϩ H2O(l) ϩ CaCO3(s) l Ca (aq) ϩ HCO3 (aq) 2ϩ 2ϩ 2ϩ 2ϩ (21.1) Surface water: water that is collected from a watershed—for example, lakes, rivers, and streams ϩ Hardening ions such as Ca , Mg , and Fe (and other divalent, , ions) form insoluble compounds with soaps and cause many detergents to be less effective Soaps, which are sodium salts of fatty acids such as sodium stearate, C17H35CO2ϪNaϩ, are very effective cleansing agents so long as they remain soluble; the presence of the hardening ions however causes the formation of a gray, insoluble soap scum such as (C17H35CO2)2Ca: C17H35CO2ϪNaϩ(aq) ϩ Ca2ϩ(aq) l (C17H35CO2)2Ca(s) ϩ Naϩ(aq) (21.2) This gray precipitate appears as a bathtub ring and also clings to clothes, causing white clothes to appear gray Dishes and glasses may have spots, shower stalls and lavatories may have a sticky film, clothes may feel rough and scratchy, hair may be dull and unmanageable, and your skin may be irritated and sticky because of hard water Hard water is also responsible for the appearance and undesirable formation of “boiler scale” on tea kettles and pots used for heating water The boiler scale is a poor conductor of heat and thus reduces the ef ciency of transferring heat Boiler scale also builds on the inside of hot water pipes, causing a decrease in the ow of water (see opening photo); in extreme cases, this buildup causes the pipe to burst Boiler scale consists primarily of the carbonate salts of the hardening ions and is formed according to ⌬ Ca2ϩ(aq) ϩ HCO3Ϫ(aq) l CaCO3(s) ϩ CO2(g) ϩ H2O(l) (21.3) CO2 dissolved in rainwater makes rainwater slightly acidic: CO2(g) ϩ H2O( l ) l H3Oϩ(aq) ϩ HCO3Ϫ(aq) The greater the CO2(g) levels in the atmosphere due to fossil fuel combustion, the more acidic will be the rainwater Experiment 21 249 Table 21.1 Hardness Classification of Water* Hardness ( ppm CaCO3) Classi cation Ͻ17.1 ppm 17.1 ppm–60 ppm 60 ppm–120 ppm 120 ppm–180 ppm Ͼ180 ppm Soft water Slightly hard water Moderately hard water Hard water Very hard water *U.S Department of Interior and the Water Quality Association Figure 21.1 Stalactite and stalagmite formations are present in regions having large deposits of limestone, a major contributor of hardening ions Colored formations are often due to trace amounts of Fe2ϩ, Mn2ϩ, or Sr2ϩ, also hardening ions Theory of Analysis Complex ion: generally a cation of a metal ion to which is bonded a number of molecules or anions (see Experiment 36) Titrant: the solution placed in the buret in a titrimetric analysis Analyte: the solution containing the substance being analyzed, generally in the receiving flask in a titration setup _ _ Na+ O O O Na+ C H2C C N O CH2 CH2 CH2 H2C O C N CH2 C O OH OH Na2H2⌼ Notice that this reaction is just the reverse of the reaction for the formation of hard water (equation 21.1) The same two reactions are also key to the formation of stalactites and stalagmites for caves located in regions with large limestone deposits (Figure 21.1) Because of the relatively large natural abundance of limestone deposits and other calcium minerals, such as gypsum, CaSO4•2H2O, it is not surprising that Ca2ϩ ion, in conjunction with Mg2ϩ, is a major component of the dissolved solids in hard water Hard water, however, is not a health hazard In fact, the presence of Ca2ϩ and Mg2ϩ in hard water can be considered dietary supplements to the point of providing their daily recommended allowance (RDA) Some research studies (though disputed) have also indicated a positive correlation between water hardness and decreased heart disease The concentration of the hardening ions in a water sample is commonly expressed as though the hardness is due exclusively to CaCO3 Hardness is commonly expressed as mg CaCO3/L, which is also ppm CaCO3,2—or grains per gallon, gpg CaCO3, where gpg CaCO3 ϭ 17.1 mg CaCO3/L A general classi cation of hard waters is listed in Table 21.1 In this experiment, a titration technique is used to measure the combined hardening divalent ion concentrations (primarily Ca2ϩ and Mg2ϩ) in a water sample The titrant is the disodium salt of ethylenediaminetetraacetic acid (abbreviated Na2H2Y).3 In aqueous solution, Na2H2Y dissociates into Naϩ and H2Y2Ϫ ions The H2Y2Ϫ ion reacts with the hardening ions, Ca2ϩ and Mg2ϩ, to form very stable complex ions, especially in a solution buffered at a pH of about 10 An ammonia–ammonium ion buffer is often used for this pH adjustment in the analysis As H2Y2Ϫ titrant is added to the analyte, it complexes with the “free” Ca2ϩ and 2ϩ Mg of the water sample to form the respective complex ions: (21.4a) Ca2ϩ(aq) ϩ H2Y2Ϫ(aq) l [CaY]2Ϫ(aq) ϩ Hϩ(aq) 2ϩ 2Ϫ 2Ϫ ϩ (21.4b) Mg (aq) ϩ H2Y (aq) l [MgY] (aq) ϩ H (aq) From the balanced equations, it is apparent that once the molar concentration of the Na2H2Y solution is known, the moles of hardening ions in a water sample can be calculated, a 1Ϻ1 stoichiometric ratio: volume H2Y2Ϫ ϫ molar concentration of H2Y2Ϫ ϭ moles H2Y2Ϫ ϭ moles hardening ions (21.5) The hardening ions, for reporting purposes, are assumed to be exclusively Ca2ϩ from the dissolving of CaCO3 Since one mole of Ca2ϩ forms from one mole of CaCO3, the hardness of the water sample expressed as mg CaCO3 per liter of sample is (21.6) moles hardening ions ϭ moles Ca2ϩ ϭ moles of CaCO3 mg CaCO3 mol CaCO3 100.1 g CaCO3 mg ppm CaCO3 ϭ (21.7) ϫ ϫ Ϫ3 L sample L sample mol 10 g ppm means “parts per million”—1 mg of CaCO3 in 1,000,000 mg (or kg) solution is ppm CaCO3 Assuming the density of the solution is g/mL (or kg/L), then 1,000,000 mg solution ϭ L solution Therefore, mg/L is an expression of ppm Ethylenediaminetetraacetic acid is often simply referred to as EDTA with an abbreviated formula of H4Y 250 Hard Water Analysis A special indicator is used to detect the endpoint in the titration Called Eriochrome Black T (EBT),4 it forms complex ions with the Ca2ϩ and Mg2ϩ ions, but binds more strongly to Mg2ϩ ions Because only a small amount of EBT is added, only Mg2ϩ complexes; no Ca2ϩ ion complexes to EBT—therefore, most all of the hardening ions remain “free” in solution The EBT indicator is sky blue in solution but forms a wine-red complex with Mg2ϩ: (21.8) Mg2ϩ(aq) ϩ EBT(aq) [Mg-EBT]2ϩ(aq) sky blue wine red Therefore, before any H2Y2Ϫ titrant is added for the analysis, the analyte is wine red because of the [Mg-EBT]2ϩ complex ion As the H2Y2Ϫ titrant is added, all of the “free” Ca2ϩ and Mg2ϩ ions in the water sample become complexed just prior to the endpoint; thereafter, the H2Y2Ϫ removes the trace amount of Mg2ϩ from the wine-red [Mg-EBT]2ϩ complex At this point, the solution changes from the wine-red color back to the original sky-blue color of the EBT indicator to reach the endpoint All hardening ions have been complexed with H2Y2Ϫ: [Mg2ϩ -EBT]2ϩ(aq) ϩ H2Y2Ϫ(aq) l [MgY]2Ϫ(aq) ϩ Hϩ(aq) ϩ EBT(aq) (21.9) wine red sky blue 2ϩ Therefore, the presence of Mg in the sample is a must in order for the color change from wine red to sky blue to be observed To ensure the appearance of the endpoint, oftentimes a small amount of Mg2ϩ as [MgY]2Ϫ is initially added to the analyte along with the EBT indicator to form the wine-red color of [Mg-EBT]2ϩ The Indicator for the Analysis HO N N HO _ SO3 NO2 Eriochrome Black T The mechanism for the process of adding both [MgY]2Ϫ and EBT is as follows: The [MgY]2Ϫ dissociates in the analyte because the Y4Ϫ (as H2Y2Ϫ in water) is more strongly bonded to the Ca2ϩ of the sample; the “freed” Mg2ϩ then combines with the EBT to form the wine-red color (equation 21.8) The complexing of the “free” Ca2ϩ and Mg2ϩ with the H2Y2Ϫ titrant continues until both are depleted At that point, the H2Y2Ϫ reacts with the [Mg-EBT]2ϩ in the sample until the endpoint is reached (equation 21.9) Because Mg2ϩ and Y4Ϫ (as H2Y2Ϫ) are freed initially from the added [MgY]2Ϫ, but later consumed at the endpoint, no additional H2Y2Ϫ titrant is required for the analysis of hardness in the water sample The standardization of a Na2H2Y solution is determined by its reaction with a known amount of calcium ion in a (primary) standard Ca2ϩ solution (equation 21.4a) The measured aliquot of the standard Ca2ϩ solution is buffered to a pH of 10 and titrated with the Na2H2Y solution to the Eriochrome Black T sky blue endpoint (equation 21.9) To achieve the endpoint, a small amount of Mg2ϩ in the form of [MgY]2Ϫ is added to the standard Ca2ϩ solution Note that the standardization of the Na2H2Y solution with a standard Ca2ϩ solution in Part A is reversed in Part B, where the (now) standardized Na2H2Y solution is used to determine the concentration of Ca2ϩ (and other hardening ions) in a sample Procedure Overview: A (primary) standard solution of Ca2ϩ is used to standardize a prepared ϳ0.01 M Na2H2Y solution The (secondary) standardized Na2H2Y solution is subsequently used to titrate the hardening ions of a water sample to the Eriochrome Black T (or calmagite) indicator endpoint The standardized Na2H2Y solution may have already been prepared by stockroom personnel If so, obtain 100 mL of the solution and proceed to Part B Consult with your instructor Three trials are to be completed for the standardization of the ϳ0.01 M Na2H2Y solution Initially prepare three clean 125-mL Erlenmeyer asks for Part A.3 A Standard Na H Y Solution Experimental Procedure A A Standard 0.01 M Disodium Ethylenediaminetetraacetate, Na H Y, Solution Calmagite may be substituted for Eriochrome Black T as an indicator The same wine-red to sky-blue endpoint is observed Ask your instructor Experiment 21 251 Measure of the mass of the Na2H2Y solution Calculate the mass of Na2H2Y•2H2O (molar mass ϭ 372.24 g/mol) required to prepare 250 mL of a 0.01 M Na2H2Y solution See Prelaboratory Assignment question and show this calculation on the Report Sheet Measure this mass on weighing paper, transfer it to a 250-mL volumetric ask containing 100 mL of deionized water, swirl to dissolve, and dilute to the mark (slight heating may be required) Prepare a buret for titration Rinse a clean buret with the Na2H2Y solution several times and then ll Record the volume of the titrant using all certain digits plus one uncertain digit Prepare the standard Ca2؉ solution Obtain ϳ80 mL of a standard Ca2ϩ solution and record its exact molar concentration (ϳ0.01 M) Pipet 25.0 mL of the standard Ca2ϩ solution into a 125-mL Erlenmeyer ask, add mL of buffer (pH ϭ 10) solution, and drops of EBT indicator (containing a small amount of [MgY]2Ϫ) Titrate the standard Ca2؉ solution Titrate the standard Ca2ϩ solution with the Na2H2Y titrant; swirl continuously Near the endpoint, slow the rate of addition to drops; the last few drops should be added at 3–5-second intervals The solution changes from wine red to purple to sky blue—no tinge of the wine-red color should remain; the solution is blue at the endpoint Record the nal volume in the buret Repeat the titration with the standard Ca2؉ solution Repeat the titrations on the remaining two samples Calculate the molar concentration of the Na2H2Y solution Save the standard Ca2ϩ solution for Part B Read and record the volume in the buret to the correct number of significant figures B Analysis of Water Sample Complete three trials for your analysis The rst trial is an indication of the hardness of your water sample You may want to adjust the volume of water for the analysis of the second and third trials Obtain the water sample for analysis a Obtain about 100 mL of a water sample from your instructor You may use your own water sample or simply the tap water in the laboratory b If the water sample is from a lake, stream, or ocean, you will need to gravity lter the sample before the analysis c If your sample is acidic, add M NH3 until it is basic to litmus (or pH paper) Prepare the water sample for analysis Pipet 25.0 mL of your ( ltered, if necessary) water sample5 into a 125-mL Erlenmeyer ask, add mL of the buffer (pH ϭ 10) solution, and drops of EBT indicator Titrate the water sample Titrate the water sample with the standardized Na2H2Y until the blue endpoint appears (as described in Part A.4) Repeat (twice) the analysis of the water sample to determine its hardness Disposal: Dispose of the analyzed solutions in the Waste EDTA container The Next Step (1) Because hardness of a water source varies with temperature, rainfall, seasons, water treatment, and so on design a systematic study of the hardness of a water source as a function of one or more variables (2) Compare the incoming versus the outgoing water hardness of a continuous water supply (3) Compare the water hardness of drinking water for adjacent city and county water supplies and account for the differences If your water is known to have a high hardness, decrease the volume of the water proportionally until it takes about 15 mL of Na2H2Y titrant for your second and third trials Similarly, if your water sample is known to have a low hardness, increase the volume of the water proportionally 252 Hard Water Analysis Experiment 21 Prelaboratory Assignment Hard Water Analysis Date Lab Sec Name Desk No What cations are responsible for water hardness? Experimental Procedure, Part A.1 Calculate the mass of disodium ethylenediaminetetraacetate (molar mass ϭ 372.24 g/mol) required to prepare 250 mL of a 0.010 M solution Show the calculation here and on the Report Sheet Express the mass to the correct number of signi cant gures Experimental Procedure, Part A.3 A 25.7-mL volume of a prepared Na2H2Y solution titrates 25.0 mL of a standard 0.0107 M Ca2ϩ solution to the Eriochrome Black T endpoint What is the molar concentration of the Na2H2Y solution? a Which hardening ion, Ca2ϩ or Mg2ϩ, binds more tightly to (forms a stronger complex ion with) the Eriochrome Black T indicator used for today’s analysis? b What is the color change at the endpoint? Experiment 21 253 A 50.0-mL water sample requires 16.33 mL of 0.0109 M Na2H2Y to reach the Eriochrome Black T endpoint a Calculate the moles of hardening ions in the water sample b Assuming the hardness is due exclusively to CaCO3, express the hardness concentration in mg CaCO3/L sample See equation 21.7 c What is this hardness concentration expressed in ppm CaCO3? d Classify the hardness of this water according to Table 21.1 a Determine the number of moles of hardening ions present in a 100-mL volume sample that has a hardness of 58 ppm CaCO3 See equations 21.6 and 21.7 b What volume of 0.100 M Na2H2Y is needed to reach the Eriochrome Black T endpoint for the analysis of the solution See equation 21.5 c Water hardness is also commonly expressed in units of grains/gallon, where grain/gallon equals 17.1 ppm CaCO3 Express the hardness of this “slightly hard” water sample in grains/gallon 254 Hard Water Analysis Experiment 21 Report Sheet Hard Water Analysis Date Lab Sec Name Desk No A A Standard 0.01 M Disodium Ethylenediaminetetraacetate, Na2H2Y, Solution Calculate the mass of Na2H2Y•2H2O required to prepare 250 mL of a 0.01 M Na2H2Y solution Volume of standard Ca2ϩ solution (mL) Trial Trial Trial 25.0 _ 25.0 _ 25.0 _ Concentration of standard Ca2ϩ solution (mol/L) _ Mol Ca2ϩ ϭ mol Na2H2Y (mol) _ _ _ Buret reading, initial (mL) _ _ _ Buret reading, nal (mL) _ _ _ Volume of Na2H2Y titrant (mL) _ _ _ Molar concentration of Na2H2Y solution (mol/L) _ _ _ Average molar concentration of Na2H2Y solution (mol/L) _ B Analysis of Water Sample Trial Trial Trial Total volume of water sample (mL) _ _ _ Buret reading, initial (mL) _ _ _ Buret reading, nal (mL) _ _ _ Volume of Na2H2Y titrant (mL) _ _ _ Experiment 21 255 Mol Na2H2Y ϭ mol hardening ions, Ca2ϩ and Mg2ϩ (mol) _ _ _ Mass of equivalent CaCO3 (g) _ _ _ ppm CaCO3 (mg CaCO3 /L sample) _ _ _ Average ppm CaCO3 _ Average gpg CaCO3 _ 10 Standard deviation of ppm CaCO3 _ Appendix B 11 Relative standard deviation of ppm CaCO3 (%RSD) _ Appendix B Laboratory Questions Circle the questions that have been assigned Part A.3 State the purpose for the mL of buffer (pH ϭ 10) being added to the standard Ca2ϩ solution Part A.3 The Eriochrome Black T indicator is mistakenly omitted What is the color of the analyte (standard Ca2ϩ solution)? Describe the appearance of the analyte with the continued addition of the Na2H2Y solution Explain *3 Part A.3 The buffer solution is omitted from the titration procedure, the Eriochrome Black T indicator and a small amount of Mg2ϩ are added, and the standard Ca2ϩ solution is acidic a What is the color of the solution? Explain b The Na2H2Y solution is dispensed from the buret What color changes are observed? Explain Part A.4 Deionized water from the wash bottle is used to wash the side of the Erlenmeyer ask How does this affect the reported molar concentration of the Na2H2Y solution—too high, too low, or unaffected? Explain Part A.4 The dispensing of the Na2H2Y solution from the buret is discontinued when the solution turns purple Because of this technique error, will the reported molar concentration of the Na2H2Y solution be too high, too low, or unaffected? Explain Part B.3 The dispensing of the Na2H2Y solution from the buret is discontinued when the solution turns purple Because of this technique error, will the reported hardness of the water sample be too high, too low, or unaffected? Explain Part A.4 and Part B.3 The dispensing of the Na2H2Y solution from the buret is discontinued when the solution turns purple However in Part B.3, the standardized Na2H2Y solution is then used to titrate a water sample to the (correct) blue endpoint Will the reported hardness of the water sample be too high, too low, or unaffected? Explain *8 Washing soda, Na2CO3•10H2O (molar mass ϭ 286 g/mol), is often used to “soften” hard water—that is, to remove hardening ions Assuming hardness is due to Ca2ϩ, the CO32Ϫ ion precipitates the Ca2ϩ: Ca2ϩ(aq) ϩ CO32Ϫ(aq) l CaCO3(s) How many grams and pounds of washing soda are needed to remove the hardness from 500 gallons of water having a hardness of 200 ppm CaCO3 (see Appendix A for conversion factors)? 256 Hard Water Analysis Experiment 22 Molar Solubility, Common-Ion Effect Silver oxide forms a brown mudlike precipitate from a mixture of silver nitrate and sodium hydroxide solutions • To determine the molar solubility and the solubility constant of calcium hydroxide • To study the effect of a common ion on the molar solubility of calcium hydroxide Objectives The following techniques are used in the Experimental Procedure: Techniques Salts that have a very limited solubility in water are called slightly soluble (or “insoluble”) salts A saturated solution of a slightly soluble salt is a result of a dynamic equilibrium between the solid salt and its ions in solution; however, because the salt is only slightly soluble, the concentrations of the ions in solution are low For example, in a saturated silver sulfate, Ag2SO4, solution, the dynamic equilibrium between solid Ag2SO4 and the Agϩ and SO42Ϫ ions in solution lies far to the left because of the low solubility of silver sulfate: Introduction Ag2SO4(s) Agϩ(aq) ϩ SO42Ϫ(aq) Slightly soluble salt: a qualitative term that reflects the very low solubility of a salt Dynamic equilibrium: the rate of the forward reaction equals the rate of the reverse reaction (22.1) The mass action expression for this system is [Agϩ]2[SO42Ϫ] (22.2) As Ag2SO4 is a solid, its concentration is constant and therefore does not appear in the mass action expression At equilibrium, the mass action expression equals Ksp, called the solubility product or, more simply, the equilibrium constant for this slightly soluble salt The molar solubility of Ag2SO4, determined experimentally, is 1.4 ϫ 10Ϫ2 mol/L This means that in 1.0 L of a saturated Ag2SO4 solution, only 1.4 ϫ 10Ϫ2 mol of silver sulfate dissolves, forming 2.8 ϫ 10Ϫ2 mol of Agϩ and 1.4 ϫ 10Ϫ2 mol of SO42Ϫ The solubility product of silver sulfate equals the product of the molar concentrations of the ions, each raised to the power of its coef cient in the balanced equation: Ksp ϭ [Agϩ]2[SO42Ϫ] ϭ [2.8 ϫ 10Ϫ2]2[1.4 ϫ 10Ϫ2] ϭ 1.1 ϫ 10Ϫ5 Molar solubility: the number of moles of salt that dissolve per liter of (aqueous) solution (22.3) What happens to the molar solubility of a salt when an ion, common to the salt, is added to the saturated solution? According to LeChâtelier’s principle (Experiment 16), Experiment 22 257