Question: Give two examples of units derived from

> Go to http://dx.doi.org/10.18434/T43G6C, where you will find a link to the NIST Statistical Reference Datasets site. Find the Dataset Archives, locate the Univariate Summary Statistics section, and select the Mavro data set. This actual data set is the r

> Calculate a pooled estimate of s from the following spectrophotometric analysis for NTA (nitrilotriacetic acid) in water from the Ohio River:

> Nine samples of illicit heroin preparations were analyzed in duplicate by a gas chromatographic method. The samples can be assumed to have been drawn randomly from the same population. Pool the following data to establish an estimate of s for the procedu

> Describe the steps in a sampling operation.

> State quantitatively the null hypothesis H0 and the alternative hypothesis Ha for the following situations, and describe the type I and type II errors. If these hypotheses were to be tested statistically, comment on whether a one- or two-tailed test woul

> In a volumetric determination of an analyte A, the data obtained and their standard deviations are as follows: From the data, find the coefficient of variation of the final result for the %A that is obtained by using the equation that follows and assumin

> Six bottles of wine of the same variety were analyzed for residual sugar content with the following results: / Evaluate the standard deviation s for each set of data. Pool the data to obtain an absolute standard deviation for t

> Analysis of several plant-food preparations for potassium ion yielded the following data: The preparations were randomly drawn from the same population. Find the mean and standard deviation Ñ• for each sample. Obtain the pooled

> Chapter 22 shows that quantitative molecular absorption spectrometry is based on Beer’s law, which can be written as -log T = εbcX where T is the transmittance of a solution of an analyte X, b is the thickness of the absorbing solution, cX is the molar c

> Calculate the molar concentration of a 20.0% solution (w/w) of KCl that has

> Identify the base on the left and its conjugate acid on the right in the equations for Problem 7-4.

> Chapter 26 describes inductively coupled plasma atomic emission spectrometry. In that method, the number of atoms excited to a particular energy level is a strong function of temperature. For an element of excitation energy E in joules (J), the measured

> The standard deviation in measuring the diameter d of a sphere is ±0.02 cm. What is the standard deviation in the calculated volume V of the sphere if d = 2.35 cm?

> Define a. Sample standard deviation. b. Coefficient of variation. c. Variance. d. Standard error of the mean.

> Calculate the absolute standard deviation and the coefficient of variation for the results of the following calculations. Round each result to include only significant figures. The numbers in parentheses are absolute standard deviations. a. y = log[2.00(

> The color change of a chemical indicator requires an overtitration of 0.03 mL. Calculate the percent relative error if the total volume of titrant is

> What is the object of the sampling step in an analysis?

> Estimate the absolute deviation and the coefficient of variation for the results of the following calculations. Round each result so that it contains only significant digits. The numbers in parentheses are absolute standard deviations. a. y = 3.95(±0.03)

> Name three types of systematic errors.

> Suggest two sources of systematic error and two sources of random error in measuring the length of a 3-m table with a 1-m metal rule.

> Calculate the molar concentration of a solution that is 50.0% NaOH (w/w) and has a specific gravity of 1.52.

> Explain the difference between a. random and systematic error. b. constant and proportional error. c. absolute and relative error. d. mean and median.

> Find the mean and median of each of the following sets of data. Determine the deviation from the mean for each data point within the sets, and find the mean deviation for each set. Use a spreadsheet if it is convenient. 6

> Simplify the following quantities using a unit with an appropriate prefix: a. 5.8 × 108 Hz. b. 4.37 × 1027 g. c. 9.31 × 107 mmol. d. 8.3 ×1010 s. e. 3.96 × 106 nm. f. 53,000 g.

> A loss of 0.4 mg of Zn occurs in the course of an analysis for that element. Calculate the percent relative error due to this loss if the mass of Zn in the sample is a 30 mg. b 100 mg. c 300 mg. d 500 mg.

> The color change of a chemical indicator requires an overtitration of 0.03 mL. Calculate the percent relative error if the total volume of titrant is a 50.00 mL. b 10.0 mL. c 25.0 mL. d 30.0 mL

> The method described in Problem 3-8 is to be used for the analysis of ores that assay about 1.2% gold. What minimum sample mass should be taken if the relative error resulting from a 0.4-mg loss is not to exceed a 20.1%? b 20.4%? c 20.8%? d 21.1%?

> A method of analysis yields masses of gold that are low by 0.4 mg. Calculate the percent relative error caused by this result if the mass of gold in the sample is a. 500 mg. b. 250 mg. c. 125 mg. d. 60 mg.

> A 0.005-g sample of a rock is to be analyzed, and iron is to be determined at the ppm level. Determine the type of analysis and type of constituent.

> What kind of systematic errors are detected by varying the sample size?

> Describe how systematic method errors may be detected.

> What mass of solute in grams is contained in a. 450.0 mL of 0.0986 M H2O2? b. 26.4 mL of 9.36 3 1024 M benzoic acid (122.1 g/mol)? c. 2.50 L of a solution that contains 23.4 ppm SnCl2? d. 21.7 mL of 0.0214 M KBrO3?

> Describe at least three ways in which a systematic error might occur while using a pipet to transfer a known volume of liquid.

> Describe at least three systematic errors that might occur while weighing a solid on an analytical balance.

> Exactly 750.0 mL of a solution that contains 500.0 ppm of Ba(NO3)2 is mixed with 200.0 mL of a solution that was 0.04100 M in Al2(SO4)3. a. What mass of solid BaSO4 is formed? b. What is the molar concentration of the unreacted reagent [Al2(SO4)3 or Ba(N

> What volume of 0.01000 M AgNO3 is required to precipitate all of the I- in 150 mL of a solution that contains 22.50 ppt KI?

> Identify the acid on the left and its conjugate base on the right in the following equations:

> What mass of MgNH4PO4 precipitates when 200.0 mL of a 1.000% (w/v) solution of MgCl2 is treated with 40.0 mL of 0.1753 M Na3PO4 and an excess of NH4+? What is the molar concentration of the excess reagent (Na3PO4 or MgCl2) after the precipitation is comp

> Exactly 75.00 mL of a 0.3132 M solution of Na2SO3 is treated with 150.0 mL of 0.4025 M HClO4 and boiled to remove the SO2 formed. a. What is the mass in grams of SO2 that is evolved? b. What is the concentration of the unreacted reagent (Na2SO3 or HClO4)

> Exactly 25.0 mL of a 0.3757 M solution of Na3PO4 is mixed with 100.00 mL of 0.5151 M HgNO3. a. What mass of solid Hg3PO4 is formed? b. What is the molar concentration of the unreacted species (Na3PO4 or HgNO3) after the reaction is complete?

> Find the number of K+ ions in 1.43 mole of K2HPO4?

> Exactly 0.118 g of pure Na2CO3 is dissolved in 100.0 mL of 0.0731 M HCl. a. What mass in grams of CO2 is evolved? b. What is the molar concentration of the excess reactant (HCl or Na2CO3)?

> Find the number of millimoles of solute in a. 386 mL of 0.210 M HClO4. b. 25.0 L of 8.05 × 10-3 M K2CrO4. c. 4.50 L of an aqueous solution that contains 6.95 ppm of AgNO3. d. 537 mL of 0.0200 M KOH.

> What mass of solid PbCl2 (278.10 g/mol) is formed when 200 mL of 0.125 M Pb2+ is mixed with 400 mL of 0.175 M Cl?

> What mass of solid La(IO3)3 (663.6 g/mol) is formed when 50.0 mL of 0.250 M La3+ is mixed with 75.0 mL of 0.302 M IO3-?

> Describe the preparation of a. 2.50 L of 0.0250 M KMnO4 from the solid reagent. b. 4.00 L of 0.250 M HClO4, starting with an 8.00 M solution of the reagent. c. 500 mL of a solution that is 0.0200 M in I-, starting with MgI2. d. 200 mL of 1.00% (w/v) aque

> Describe the preparation of a. 500 mL of 0.1000 M AgNO3 from the solid reagent. b. 1.00 L of 0.1000 M HCl, starting with a 6.00 M solution of the reagent. c. 250 mL of a solution that is 0.0810 M in K+, starting with solid K4Fe(CN)6. d. 500 mL of 3.00% (

> Is the following statement true or false or both? Define your answer with equations, examples, or graphs. “A buffer maintains the pH of a solution constant.”

> Describe the preparation of 750 mL of 3.00 M HNO3 from the commercial reagent that is 70.5% HNO3 (w/w) and has a specific gravity of 1.42.

> What volume of 2.00 M NaOH must be added to 200.0 mL of 1.00 M glycolic acid to produce a buffer solution having a pH of 4.15?

> Describe the preparation of 500 mL of 3.00 M H3PO4from the commercial reagent that is 86% H3PO4(w/w) and has a specific gravity of 1.71.

> The following table gives the sample means and standard deviations for six measurements each day of the purity of a polymer in a process. The purity is monitored for 24 days. Determine the overall mean and standard deviation of the measurements, and cons

> What mass of solute in milligrams is contained in a. 26.0 mL of 0.250 M sucrose (342 g/mol)? b. 2.92 L of 5.23 3 1024 M H2O2? c. 673 mL of a solution that contains 5.76 ppm Pb1NO322 1331.20 g>mol2? d. 6.75 mL of 0.0426 M KNO3?

> What mass of KIO3 is needed to convert the copper in 0.1570 g of CuSO4 . 5H2O to Cu1IO322?

> What volume of 0.200 M HCl must be added to 500.0 mL of 0.300 M sodium mandelate to produce a buffer solution with a pH of 3.25?

> Describe the preparation of a. 1.50 L of 21.0% (w/v) aqueous glycerol (C3H8O3, 92.1 g/mol). b. 1.50 kg of 21.0% (w/w) aqueous glycerol. c. 1.50 L of 21.0% (v/v) aqueous glycerol.

> What mass of sodium glycolate should be added to 400.0 mL of 1.00 M glycolic acid to produce a buffer solution with a pH of 4.25?

> Describe the preparation of a. 500 mL of 5.25% (w/v) aqueous ethanol (C2H5OH, 46.1 g/mol). b. 500 g of 5.25% (w/w) aqueous ethanol. c. 500 mL of 5.25% (v/v) aqueous ethanol.

> What mass of sodium formate must be added to 500.0 mL of 1.00 M formic acid to produce a buffer solution that has a pH of 3.75?

> Why are Ce4+ solutions never used for the titration of reductants in basic solutions?

> Why are Ce4+ solutions never used for the titration of reductants in basic solutions?

> Why are standard solutions of reductants less often used for titrations than standard solutions of oxidants?

> Write a balanced net ionic equation for the reduction of UO22+ in a Walden reductor.

> Why is a Walden reductor always used with solutions that contain appreciable concentrations of HCl?

> How many millimoles of solute are contained in a. 2.95 mL of 0.0789 M KH2PO4? b. 0.2011 L of 0.0564 M HgCl2? c. 2.56 L of a 47.5 ppm solution of Mg1NO322? d. 79.8 mL of 0.1379 M NH4VO3 1116.98 g>mol2?

> The following data represent measurements made on a process for 30 days. One measurement was made each day. Assuming that 30 measurements are enough that x S m and s S s, find the mean of the values, the standard deviation, and the upper and lower contro

> Write balanced net ionic equations to describe

> Use a spreadsheet to do the calculations and plot the titration curves for the following titrations. Calculate potentials after the addition of titrant corresponding to 10%, 20%, 30%, 40%, 50%, 60%, 70%, 80%, 90%, 95%, 99%, 99.9%, 100%, 101%, 105%, 110%,

> A 11.4% (w/w) NiCl2 (129.61 g/mol) solution has a density of 1.149 g/mL. Calculate a. The molar concentration of NiCl2 in this solution. b. The molar Cl2 concentration of the solution. c. The mass in grams of NiCl2 contained in each liter of this solutio

> A gas mixture was passed at the rate of 2.50 L/min through a solution of sodium hydroxide for a total of 59.00 min. The 2 SO in the mixture was retained as sulfite ion: After acidification with HCl, the sulfite was titrated with 5.15 mL of 0.002997MKIO3

> A 2.552-g sample containing both Fe and V was dissolved under conditions that converted the elements to Fe(III) and V(V). The solution was diluted to 500.0 mL, and a 50.00-mL aliquot was passed through a Walden reductor and titrated with 17.79 mL of 4 0

> A sensitive method for I− in the presence of CI− and Br− entails oxidation of the I− to IO3−with Br2. The excess Br2 is then removed by boiling or by reduction with formate ion. The IO3− produced is determined by addition of excess I- and titration of th

> The ethyl mercaptan concentration in a mixture was determined by shaking a 1.795-g sample with 50.00 mL of 2 0 01204 MI. in a tightly stoppered flask: The excess 2I was back-titrated with 15.21 mL of 22 3 0 01437MNa S O . . Calculate thepercentage of 2 5

> An 8.13-g sample of an ant-control preparation was decomposed by wet ashing with H2SO4 and HNO3. The As in the residue was reduced to the trivalent state with hydrazine. After removal of the excess reducing agent, the As(III) required a 31.46-mL titratio

> The 3 KClO in a 0.1791-g sample of an explosive was determined by reaction with 50.00 mL of 0 0. 873 2 MFe + When the reaction was complete, the excess 2 Fe + was back-titrated with 14.95 mL of 4 0 06970MCe + . . Calculate the percentage of 3 KClO in the

> How many millimoles of solute are contained in a. 2.00 L of 2.76 3 1023 M KMnO4? b. 250.0 mL of 0.0423 M KSCN? c. 500.0 mL of a solution containing 2.97 ppm CuSO4? d. 2.50 L of 0.352 M KCl?

> Treatment of hydroxylamine 2 (H NOH) with an excess of Fe(III) results in the formation of N O2 and an equivalent amount of Fe(II): Calculate the molar concentration of an H N2 OH solution if the Fe(II) produced by treatment of a 25.00-mL aliquot requir

> Atomic emission measurements were made to determine sodium in a blood serum sample. The following emission intensities were obtained for standards of 5.0 and 10.0 ng/mL and for the serum sample. All emission intensities were corrected for any blank emiss

> There is 0.5690-g specimen of iron ore was dissolved and passed through a Jones reductor. Titration of the Fe(II) produced required 38.79 mL of 0.01926 M KMnO4. Express the results of this analysis in terms of (a) percent Fe and (b) percent 2 3 Fe O .

> Calculate the percentage of MnO2 in a mineral specimen if the 2I liberated by a 0.1267-g sample in the net reaction

> A 0.1853-g sample of 3 KBrO was dissolved in dilute HCl and treated with an unmeasured excess of KI. The liberated iodine required 44.36 mL of a sodium thiosulfate solution. Calculate the molar concentration of the Na S O 22 3.

> Consult Appendix 3, and pick out a suitable acid-base pair to prepare a buffer with a pH of a. 10.3. b. 6.1. c. 4.5. d. 8.1.

> A 0.2219-g sample of pure iron wire was dissolved in acid, reduced to the +2 state, and titrated with 34.65 mL of cerium (IV). Calculate the molar concentration of the 4 Ce +solution.

> How would you prepare 1.000 L of 0.05000 M KBrO3?

> How would you prepare 2.0 L of approximately 0.04 M I3 solution? Calculate the molar concentration of 4 KMnO in this solution.

> Why are KMnO4 solutions filtered before they are standardized?

> Calculations of volumetric analysis ordinarily consist of transforming the quantity of titrant used (in chemical units) to a chemically equivalent quantity of analyte (also in chemical units) through use of a stoichiometric factor. Use chemical formulas

> A solution prepared by dissolving a 0.2541-g sample of electrolytic iron wire in acid was passed through a Jones reductor. The iron(II) in the resulting solution required a 36.76-mL titration. Calculate the molar oxidant concentration if the titrant used

> In the titration of 2I solutions with Na S O 2+ 3 , the starch indicator is never added until just before chemical equivalence. Why?

> The following atomic absorption results were obtained for determinations of Zn in multivitamin tablets. All absorbance values are corrected for the appropriate reagent blank (cZn = 0.0 ng/mL). The mean value for the blank was 0.0000 with a standard devia

> Write balanced equations showing how 2 2 K Cr O could be used as a primary standard for solutions of 7Na S O 2+3.

> Suggest a way in which a solution of 3 KIO could be used as a source of known quantities of 2I .

> A standard solution of 2I increased in concentration with standing. Write a balanced net ionic equation that accounts for the increase.

> What is the primary use of standard K2Cr2O7 solutions?

> A 5.85% (w/w) Fe(NO3)3 (241.86 g/mol) solution has a density of 1.059 g/mL. Calculate a. The molar analytical concentration of Fe(NO3)3 in this solution. b. The molar NO3- concentration in the solution. c. The mass in grams of Fe(NO3)3 contained in each

> Briefly explain why there is no term in an equilibrium constant expression for water or for a pure solid, even though one (or both) appears in the balanced net ionic equation for the equilibrium.

> Why are solutions of KMnO4 and Na2S2O3 generally stored in dark reagent bottles? Answer Standard permanganate and thiosulfate solutions are generally stored in the dark because their decomposition reactions are catalyzed by light.

> Briefly explain why the concentration units of milligrams of solute per liter and parts per million can be used interchangeably for a dilute aqueous solution.

> Why are KMnO4 solutions filtered before they are standardized?

> Under what circumstance is the curve for an oxidation/reduction titration asymmetric about the equivalence point?