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18. Aqueous Equilibrium

Intro to Buffers

18. Aqueous Equilibrium

Intro to Buffers

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Acid-Base Buffers

Acid-Base Buffers

Buffer Creation

Buffer Creation

Intro to Buffers Example

Intro to Buffers Example

Buffer Capacity

Buffer Capacity

Intro to Buffers Example

Intro to Buffers Example

Intro to Buffers Example

Intro to Buffers Example

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Calculating the pH of a Buffer

Calculating the pH of a Buffer

Practice Problem: Henderson-Hasselbalch Equation Calculations

Practice Problem: Henderson-Hasselbalch Equation Calculations

Professor Dave Explains

Derive the Henderson-Hasselblach Equation

Derive the Henderson-Hasselblach Equation

Henderson-Hasselbalch Equation

Henderson-Hasselbalch Equation

Chemistry with Dr. T

Buffer Solutions

Buffer Solutions

The Organic Chemistry Tutor

Buffer solution pH calculations | Chemistry | Khan Academy

Buffer solution pH calculations | Chemistry | Khan Academy

Khan Academy Organic Chemistry

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Buffers, the Acid Rain Slayer: Crash Course Chemistry #31

Buffers, the Acid Rain Slayer: Crash Course Chemistry #31

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Practice this topic

Multiple Choice

Which one of the following combinations does not create a buffer?

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Multiple Choice

Which of the following combinations can result in the formation of a buffer?

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Multiple Choice

Which of the following combinations can result in the formation of a buffer?

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Multiple Choice

A buffer solution is comprised of 50.0 mL of a 0.100 M HC₂H₃O₂ and 60.0 mL of a 0.100 M NaC₂H₃O₂. Which of the following actions would completely destroy the buffer?

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Multiple Choice

Calculate the pH of a solution formed by mixing 200 mL of a 0.400 M C₂H₅NH₂ solution with 350 mL of a 0.450 M C₂H₅NH₃⁺solution. (K_b of C₂H₅NH₂is 5.6 x 10 ^-4).

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Multiple Choice

Determine how many grams of sodium acetate, NaCH₃CO₂ (MW:82.05 g/mol), you would mix into enough 0.065 M acetic acid CH₃CO₂H (MW:60.05 g/mol) to prepare 3.2 L of a buffer with a pH of 4.58. The K_a is 1.8 x 10^-5.

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Multiple Choice

A buffer solution is made by combining a weak acid with its conjugate salt. What will happen to the pH if the solution is diluted to one-fourth of its original concentration?

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Multiple Choice

Which of the following statements is true about a solution prepared with equal amounts of HNO₂ and NaNO₂?

Multiple Choice

The addition of which of the following would destroy a 1.0 L buffer that is 0.50 M HC₂H₃O₂ and 0.40 M NaC₂H₃O₂?

Multiple Choice

Which of the following would produce a buffer?

Multiple Choice

What is the pH of a 1.0 L buffer solution containing 0.500 M HF and 0.300 M NaF. K_a for HF = 3.5 × 10⁻⁴.

Multiple Choice

Determine the pH of a 1.0 L buffer prepared by adding 0.100 moles of NaOH to 0.250 moles of HF. K_a for HF = 3.5 × 10⁻⁴.

Multiple Choice

What mass of NaC₂H₃O₂ (molar mass = 82.03 g/mol) would need to be added to 0.450 moles of HC₂H₃O₂ to prepare a 1.0 L buffer with a pH = 4.95. K_a for HC₂H₃O₂ = 1.8 × 10⁻⁵.

Multiple Choice

1.0 L of a buffer is prepared to produce a solution that is 0.400 M in HF and 0.250 M in LiF. What is the pH after the addition of 2.5 g of KOH? Assume no volume change.
K_a for HF = 3.5 × 10⁻⁴.

Multiple Choice

Which pairs of compounds are capable of making a buffer? Select all that apply.

a) 1.3 M LiOH and 1.7 M HCOOH c) 0.35 M CH₃CO₂H and 0.35 M NaOH

b) 0.784 M NH₄⁺ and 0.800 M HClO₄ d) 0.80 HNO₃ and 0.15 MgO

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Multiple Choice

Determine which of the following actions will destroy a buffer composed of 0.50 L of 1.44 M H₃PO₄ and 0.60 L of 1.25 M NaH₂PO₄⁻.

a) Addition of 1.45 moles of KH₂PO₄

b) Addition of 0.85 moles of HCl

c) Addition of water

d) Addition of 0.30 moles of Ca(OH)₂

e) Addition of 0.70 moles of HIO₄

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Textbook Question

The beaker on the right contains 0.1 M acetic acid solution with methyl orange as an indicator. The beaker on the left contains a mixture of 0.1 M acetic acid and 0.1 M sodium acetate with methyl orange. (b) Which solution is better able to maintain its pH when small amounts of NaOH are added? Explain. [Sections 17.1 and 17.2]

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Textbook Question

The beaker on the right contains 0.1 M acetic acid solution with methyl orange as an indicator. The beaker on the left contains a mixture of 0.1 M acetic acid and 0.1 M sodium acetate with methyl orange. (a) Using Figures 16.8 and 16.9, which solution has a higher pH?

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Textbook Question

The indicator methyl orange has been added to both of the following solutions. Based on the colors, classify each statement as true or false: (a) The pH of solution A is definitely less than 7.00.

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Textbook Question

The following diagram represents a buffer composed of equal concentrations of a weak acid, HA, and its conjugate base, A-. The heights of the columns are proportional to the concentrations of the components of the buffer. (c) Which of the three represents a situation that cannot arise from the addition of either an acid or a base? [Section 17.2]

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Textbook Question

The following diagram represents a buffer composed of equal concentrations of a weak acid, HA, and its conjugate base, A-. The heights of the columns are proportional to the concentrations of the components of the buffer. (a) Which of the three drawings, (1), (2), or (3), represents the buffer after the addition of a strong acid? [Section 17.2]

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Textbook Question

What is the pH of a buffer solution prepared by dissolving 0.250 mol of NaH2PO4 and 0.075 mol of NaOH in enough water to make 1.00 L of solution? (Ka (H2PO4-) = 6.2 X 10^-8) (a) 6.32 (b) 6.83 (c) 7.21 (d) 7.71

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Textbook Question

Consider the following table of standard electrode potentials for a series of hypothetical reactions in an aqueous solution: reduction half-reaction E °(V) (c) Which substance(s) can oxidize C2+?

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Textbook Question

Consider the following table of standard reduction potentials:. Which substance(s) can be reduced by C-? (a) D and B (b) A- (c) D3+ and B2+ (d) A

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Textbook Question

Bars of iron are put into each of the three beakers as shown here. In which beaker—A, B, or C—would you expect the iron to show the most corrosion ? [Section 20.8]

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Textbook Question

You are titrating an acidic solution with a basic one, and just realized you forgot to add the indicator that tells you when the equivalence point is reached. In this titration, the indicator turns blue at the equivalence point from an initially colorless solution. You quickly grab a bottle of indicator and add some to your titration beaker, and the whole solution turns dark blue. What do you do now?

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Textbook Question

In excess of NH3(aq), Zn2+ forms a complex ion, [Zn(NH3)4]2+ which has a formation constant Kf = 7.8 x 10^8. Calculate the concentration of Zn2+ in a solution prepared by adding 1.00 x 10^-2 mol Zn(NO3)2 to 1.00 L of 0.250 M NH3. (a) 7.9 x 10^-4 M (b) 2.8 x 10^-6 M (c) 3.9 x 10^-9 M (d) 6.4 x 10^-11 M

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Textbook Question

What is the molar solubility of AgI in 0.20 M NaCN? (a) 6.2 x 10^-4 M (b) 1.0 x 10^-1 M (c) 7.6 x 10^-2 M (d) 2.1 x 10^-3 M

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Textbook Question

A solution containing sulfide ions is added to a solution of 0.036 M Cu2+ and 0.044 M Fe2+. At what concentration of sulfide ion will a precipitate begin to form? What is the identity of the precipitate? (a) 1.4 x 10^-16 M, FeS (b) 3.6 x 10^-35 M, CuS (c) 3.6 x 10^-35 M, FeS (d) 1.4 x 10^-16 M, C

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Textbook Question

Which of the following solutions is a buffer? (a) A solution made by mixing 100 mL of 0.100 M CH3COOH and 50 mL of 0.100 M NaOH, (b) a solution made by mixing 100 mL of 0.100 M CH3COOH and 500 mL of 0.100 M NaOH, (c) A solution made by mixing 100 mL of 0.100 M CH3COOH and 50 mL of 0.100 M HCl, (d) A solution made by mixing 100 mL of 0.100 M CH3COOK and 50 mL of 0.100 M KCl.

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Textbook Question

A buffer contains 0.10 mol of acetic acid and 0.13 mol of sodium acetate in 1.00 L. (a) What is the pH of this buffer?

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Textbook Question

You have to prepare a pH = 3.50 buffer, and you have the following 0.10 M solutions available: HCOOH, CH3COOH, H3PO4, HCOONa, CH3COONa, and NaH2PO4. Which solutions would you use?

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Textbook Question

You have to prepare a pH = 5.00 buffer, and you have the following 0.10 M solutions available: HCOOH, HCOONa, CH3COOH, CH3COONa, HCN, and NaCN. How many milliliters of each solution would you use to make approximately 1 L of the buffer?

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Textbook Question

You have to prepare a pH = 5.00 buffer, and you have the following 0.10 M solutions available: HCOOH, HCOONa, CH3COOH, CH3COONa, HCN, and NaCN. Which solutions would you use?

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Textbook Question

A buffer contains significant amounts of acetic acid and sodium acetate. Write equations showing how this buffer neutralizes added acid and added base.

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Textbook Question

Consider the following table of standard reduction potentials:

(b) Which substances can be oxidized by B2+? Which can be reduced by D?

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Textbook Question

A buffer contains significant amounts of ammonia and ammonium chloride. Write equations showing how this buffer neutralizes added acid and added base.

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Textbook Question

Addition of the indicator methyl orange to an unknown solution leads to a yellow color. The addition of bromthymol blue to the same solution also leads to a yellow color. (b) What is the range (in whole numbers) of possible pH values for the solution?

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Textbook Question

The following pictures represent solutions that contain one or more of the compounds H2A, NaHA, and Na2A, where H2A is a weak diprotic acid. (Na+ ions and solvent water molecules have been omitted for clarity.)

(b) Which solution has the greatest buffer capacity?

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Textbook Question

The following pictures represent solutions that contain one or more of the compounds H2A, NaHA, and Na2A, where H2A is a weak diprotic acid. (Na+ ions and solvent water molecules have been omitted for clarity.)

(a) Which of the solutions are buffer solutions?

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Textbook Question

Assume that 30.0 mL of a 0.10 M solution of a weak base B that accepts one proton is titrated with a 0.10 M solution of the monoprotic strong acid HA. (b) What is the predominant form of B at the equivalence point?

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Textbook Question

Addition of phenolphthalein to an unknown colorless solution does not cause a color change. The addition of bromthymol blue to the same solution leads to a yellow color. (c) What other indicator or indicators would you want to use to determine the pH of the solution more precisely?

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Textbook Question

From each of the following pairs of substances, use data in Appendix E to choose the one that is the stronger reducing agent: (d) BrO3-1aq2 or IO3-1aq2

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Textbook Question

From each of the following pairs of substances, use data in Appendix E to choose the one that is the stronger reducing agent: (a) Fe(s) or Mg(s)

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Textbook Question

By using the data in Appendix E, determine whether each of the following substances is likely to serve as an oxidant or a reductant: (d) Zn(s)

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Textbook Question

By using the data in Appendix E, determine whether each of the following substances is likely to serve as an oxidant or a reductant: (b) MnO4- (aq, acidic solution),

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Textbook Question

Is each of the following substances likely to serve as an oxidant or a reductant: (d) N2O51g2?

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Textbook Question

(a) Assuming standard conditions, arrange the following in order of increasing strength as oxidizing agents in acidic solution: Cr2O72-, H2O2, Cu2+, Cl2, O2.

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Textbook Question

A 100.0-mL buffer solution is 0.175 M in HClO and 0.150 M in NaClO. b. What is the pH after addition of 150.0 mg of HBr?

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Textbook Question

A 100.0-mL buffer solution is 0.175 M in HClO and 0.150 M in NaClO. a. What is the initial pH of this solution?

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Textbook Question

Determine whether or not the mixing of each pair of solutions results in a buffer. e. 125.0 mL of 0.15 M NH3; 150.0 mL of 0.20 M NaOH

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Textbook Question

Determine whether or not the mixing of each pair of solutions results in a buffer. d. 175.0 mL of 0.10 M NH3; 150.0 mL of 0.12 M NaOH

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Textbook Question

Determine whether or not the mixing of each pair of solutions results in a buffer. c. 50.0 mL of 0.15 M HF; 20.0 mL of 0.15 M NaOH

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Textbook Question

Determine whether or not the mixing of each pair of solutions results in a buffer. b. 50.0 mL of 0.10 M HCl; 35.0 mL of 0.150 M NaOH

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Textbook Question

Determine whether or not the mixing of each pair of solutions results in a buffer. a. 100.0 mL of 0.10 M NH3; 100.0 mL of 0.15 M NH4Cl

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Textbook Question

Determine whether or not the mixing of each pair of solutions results in a buffer. e. 105.0 mL of 0.15 M CH3NH2; 95.0 mL of 0.10 M HCl

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Textbook Question

Determine whether or not the mixing of each pair of solutions results in a buffer. c. 165.0 mL of 0.10 M HF; 135.0 mL of 0.050 M KOH

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Textbook Question

Determine whether or not the mixing of each pair of solutions results in a buffer. b. 150.0 mL of 0.10 M HF; 135.0 mL of 0.175 M HCl

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Textbook Question

Which metal could you use to reduce Mn2+ ions but not Mg2+ ions?

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Textbook Question

Which metal can be oxidized with an Sn2+ solution but not with an Fe2+ solution?

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Textbook Question

The fluids within cells are buffered by H2PO4- and HPO42 - . b. Could a buffer system employing H3PO4 as the weak acid and H2PO4- as the weak base be used as a buffer system within cells? Explain.

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Textbook Question

A 500.0-mL buffer solution is 0.100 M in HNO2 and 0.150 M in KNO2. Determine if each addition would exceed the capacity of the buffer to neutralize it. d. 1.35 g HI

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Textbook Question

A 500.0-mL buffer solution is 0.100 M in HNO2 and 0.150 M in KNO2. Determine if each addition would exceed the capacity of the buffer to neutralize it. b. 350 mg KOH

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Textbook Question

A 500.0-mL buffer solution is 0.100 M in HNO2 and 0.150 M in KNO2. Determine if each addition would exceed the capacity of the buffer to neutralize it. c. 1.25 g HBr

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Textbook Question

A 500.0-mL buffer solution is 0.100 M in HNO2 and 0.150 M in KNO2. Determine if each addition would exceed the capacity of the buffer to neutralize it. a. 250 mg NaOH

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Textbook Question

Which metal cation is the best oxidizing agent? a. Pb2+ b. Cr3+ c. Fe2+ d. Sn2+

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Textbook Question

Which of the following gives a buffer solution when equal volumes of the two solutions are mixed? (a) 0.10 M NH3 and 0.10 M HCl (b) 0.20 M NH3 and 0.10 M HCl (c) 0.10 M NH4Cl and 0.10 M NH3 (d) 0.20 M NH4Cl and 0.10 M NaOH

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Textbook Question

Which of the following solutions has the greater buffer capacity: 100 mL of 0.30 M HNO2-0.30 M NaNO2 or 100 mL of 0.10 M HNO2-0.10 M NaNO2? Explain.

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Textbook Question

From the value of Kf listed in Table 17.1, calculate the concentration of Ni2 +1aq2 and Ni1NH326 2+ that are present at equilibrium after dissolving 1.25 g NiCl2 in 100.0 mL of 0.20 M NH31aq2.

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Textbook Question

Consider the curve shown here for the titration of a weak monoprotic acid with a strong base and answer each question.

c. At what volume of added base does pH = pKa?

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Textbook Question

Consider the curve shown here for the titration of a weak base with a strong acid and answer each question.

c. At what volume of added acid does pH = 14 - pKb?

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Textbook Question

Calculate the pH of a buffer solution that is 0.20 M in HCN and 0.12 M in NaCN. Will the pH change if the solution is diluted by a factor of 2? Explain.

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Textbook Question

Use the Henderson–Hasselbalch equation to calculate the pH of a buffer solution that is 0.25 M in formic acid (HCO2H) and 0.50 M in sodium formate (HCO2Na).

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Textbook Question

A 30.0-mL sample of 0.165 M propanoic acid is titrated with 0.300 M KOH. Calculate the pH at each volume of added base: one-half equivalence point.

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Textbook Question

Arrange the following oxidizing agents in order of increasing strength under standard-state conditions: Br2(aq), MnO4+(aq), Sn4+(aq).

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Textbook Question

The ratio of HCO3- to H2CO3 in blood is called the 'bicarb number' and is used as a measure of blood pH in hospital emergency rooms. A newly diagnosed diabetic patient is admitted to the emergency room with ketoacidosis and a bicarb number of 10. Calculate the blood pH. Ka for carbonic acid at room temperature (37 degrees Celsius) os 7.9 x 10^-7).

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Textbook Question

A solution of Na2SO4 is added dropwise to a solution that is 0.010 M in Ba2 +1aq2 and 0.010 M in Sr2 +1aq2. (a) What concentration of SO42 - is necessary to begin precipitation? (Neglect volume changes. BaSO4: Ksp = 1.1 * 10-10; SrSO4: Ksp = 3.2 * 10-7.)

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Textbook Question

A solution of Na2SO4 is added dropwise to a solution that is 0.010 M in Ba2 +1aq2 and 0.010 M in Sr2 +1aq2. (c) What is the concentration of SO42 - 1aq2 when the second cation begins to precipitate?

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Textbook Question

A 25.0-mL sample of 0.125 M pyridine is titrated with 0.100 M HCl. Calculate the pH at each volume of added acid: one-half equivalence.

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Textbook Question

In what volume ratio should you mix 1.0 M solutions of NH4Cl and NH3 to produce a buffer solution having pH = 9.80?

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Textbook Question

Consider the following substances: I2(s), Fe2+(aq), Cr2O72-(aq). Which is the strongest oxidizing agent? Which is the weakest oxidizing agent?

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Textbook Question

Give a recipe for preparing a CH3CO2H-CH3CO2Na buffer solution that has pH = 4.44.

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Textbook Question

Consider the following substances: Fe2+(aq), Sn2+(aq), I-(aq). Identify the strongest reducing agent and the weakest reducing agent.

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Textbook Question

You need a buffer solution that has pH = 7.00. Which of the following buffer systems should you choose? Explain. (a) H3PO4 and H2PO4 - (b) H2PO4- and HPO42- (c) HPO42- and PO43-

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Textbook Question

Consider the following substances: Fe(s), PbO2(s), H+(aq), Al(s), Ag(s), Cr2O72-(aq). (d) Which substances can be oxidized by Cu2+(aq)? Which can be reduced by H2O2(aq)?

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Textbook Question

Consider a buffer solution that contains equal concentrations of H2PO4- and HPO42-. Will the pH increase, decrease, or remain the same when each of the following substances is added? (a) Na2HPO4 (b) HBr (c) KOH (d) KI (e) H3PO4 (f) Na3PO4

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Textbook Question

Which of the indicators given in Figure 16.5, methyl violet, bromcresol green, phenol red, or thymolphthalein, would be most appropriate to detect a pH change from: (c) 2 to 0?

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Textbook Question

Which of the indicators given in Figure 16.5, methyl violet, bromcresol green, phenol red, or thymolphthalein, would be most appropriate to detect a pH change from: (b) 8 to 10?

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Textbook Question

Use the data in Appendix D to predict whether the following reactions can occur under standard-state conditions. (a) Oxidation of Sn2+(aq) by Br2(aq)

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Textbook Question

(b) What is the most significant difference between the sulfides precipitated in group 2 and those precipitated in group 3?

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Textbook Question

Methyl red has a pKa of 5.0 and is red in its acid form and yellow in its basic form. If several drops of this indicator are placed in a 25.0-mL sample of 0.100 M HCl, what color will the solution appear? If 0.100 M NaOH is slowly added to the HCl sample, in what pH range will the indicator change color?

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Textbook Question

Derive an equation similar to the Henderson–Hasselbalch equation relating the pOH of a buffer to the pKb of its base component.

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Textbook Question

Referring to Table 17.1, pick an indicator for use in the titration of each acid with a strong base. a. HF

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Textbook Question

Referring to Table 17.1, pick an indicator for use in the titration of each base with a strong acid. a. CH3NH2

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Textbook Question

The acid–base indicator bromcresol green is a weak acid. The yellow acid and blue base forms of the indicator are present in equal concentrations in a solution when the pH is 4.68. What is the pKa for bromcresol green?

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Textbook Question

Two buffers are prepared by adding an equal number of moles of formic acid (HCOOH) and sodium formate (HCOONa) to enough water to make 1.00 L of solution. Buffer A is prepared using 1.00 mol each of formic acid and sodium formate. Buffer B is prepared by using 0.010 mol of each. (b) Which buffer will have the greater buffer capacity?

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Textbook Question

Mathematically prove that the pH at the halfway point of a titration of a weak acid with a strong base (where the volume of added base is half of that needed to reach the equivalence point) is equal to pKa for the acid.

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Textbook Question

A mixture of copper and gold metals that is subjected to electrorefining contains tellurium as an impurity. The standard reduction potential between tellurium and its lowest common oxidation state, Te4+, is Te4+1aq2 + 4 e- ¡ Te1s2 E°red = 0.57 V Given this information, describe the probable fate of tellurium impurities during electrorefining. Do the impurities fall to the bottom of the refining bath, unchanged, as copper is oxidized, or do they go into solution as ions? If they go into solution, do they plate out on the cathode?

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Textbook Question

A solution is 0.010 M in Ba2 + and 0.020 M in Ca2 + . b. What is the remaining concentration of the cation that precipitates first, when the other cation begins to precipitate?

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Textbook Question

A solution is 0.010 M in Ba2 + and 0.020 M in Ca2 + . a. If sodium sulfate is used to selectively precipitate one of the cations while leaving the other cation in solution, which cation will precipitate first? What minimum concentration of Na2SO4 will trigger the precipitation of the cation that precipitates first?

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Textbook Question

The solubility product for Zn1OH22 is 3.0 * 10-16. The formation constant for the hydroxo complex, Zn1OH242 - , is 4.6 * 1017. What concentration of OH- is required to dissolve 0.015 mol of Zn1OH22 in a liter of solution?

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Textbook Question

The value of Ksp for Cd1OH22 is 2.5 * 10-14. (b) The solubility of Cd1OH22 can be increased through formation of the complex ion CdBr42 - 1Kf = 5 * 1032. If solid Cd1OH22 is added to a NaBr solution, what is the initial concentration of NaBr needed to increase the molar solubility of Cd1OH22 to 1.0 * 10-3 mol/L?

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Textbook Question

A solution is made 1.1 * 10 - 3 M in Zn(NO3)2 and 0.150 M in NH3. After the solution reaches equilibrium, what concentration of Zn2 + (aq) remains?

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Textbook Question

(b) As the acid is titrated, the pH of the solution after the addition of 11.05 mL of the base is 4.89. What is the Ka for the acid?

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Textbook Question

Use the appropriate values of Ksp and Kf to find the equilibrium constant for the reaction. FeS(s) + 6 CN-(aq)ΔFe(CN)64 - (aq) + S2 - (aq)

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Textbook Question

Is the solubility of Zn(OH)2 increased, decreased, or unchanged on addition of each of the following substances? Write a balanced net ionic equation for each dissolution reaction. (See Appendix C.6 for formulas of complex ions.) (c) NaCN

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Textbook Question

Is the solubility of Fe(OH)3 increased, decreased, or unchanged on addition of each of the following substances? Write a balanced net ionic equation for each dissolution reaction. (See Appendix C.6 for formulas of complex ions.) (b) NaOH(aq)

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Textbook Question

Dissolution of 5.0 x 10^-3 mol of Cr(OH)3 in 1.0 L of 1.0 M NaOH gives a solution of the complex ion [Cr(OH)4]- (Kf = 8 x10^29). What fraction of the chromium in such a solution is present as uncomplexed Cr3+?

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Textbook Question

Write a balanced net ionic equation for each of the follow-ing dissolution reactions, and use the appropriate Ksp and Kf values in Appendix C to calculate the equilibrium constant for each. (a) AgI in aqueous NaCN to form [Ag(CN)2]-

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Textbook Question

Write a balanced net ionic equation for each of the follow-ing dissolution reactions, and use the appropriate Ksp and Kf values in Appendix C to calculate the equilibrium constant for each. (b) Cu(OH)2 in aqueous NH3 to form [Cu(NH3)4]2+

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Textbook Question

Calculate the molar solubility of AgI in: (a) Pure Water (b) 0.10 M NaCN: Kf for [Ag(CN)2]- is 3.0 x 10^20

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Textbook Question

Calculate the molar solubility of Cr(OH)3 in 0.50 M NaOH; Kf for Cr(OH)4- is 8 x 10^29.

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Textbook Question

Citric acid (H3Cit) can be used as a household cleaning agent to dissolve rust stains. The rust, represented as Fe(OH)3, dissolves because the citrate ion forms a soluble complex with Fe3+ (a) Using the equilibrium constants in Appendix C and Kf = 6.3 x 10^11 for Fe(Cit), calculate the equilibrium constant K for the reaction. (b) Calculate the molar solubility of Fe(OH)3 in 0.500 M solution of H3Cit.

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Textbook Question

Calculate the solubility of silver chloride in a solution that is 0.100 M in NH3.

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Textbook Question

Calculate the solubility of CuX in a solution that is 0.150 M in NaCN. Ksp for CuX is 1.27 * 10 - 36.

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Textbook Question

Using the qualitative analysis flowchart in Figure 17.18

, tell how you could separate the following pairs of ions. (a) Ag+ and Cu2+

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Textbook Question

Give a method for seperating the following pairs of ions by the addition of no more than two substances. (a) Hg2+ and Co2+

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Textbook Question

What should the molar concentrations of benzoic acid and sodium benzoate be in a solution that is buffered at a pH of 4.55 and has a freezing point of -2.0 °C? (Assume complete dissociation of sodium benzoate and a density of 1.01 g/mL for the solution.)

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Textbook Question

Assume that you have three white solids: NaCl, KCl, and MgCl2. What tests could you do to tell which is which?

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Textbook Question

In quantitative analysis, Ag+, Hg2+, and Pb2+ are seperated from other cations by the addition of HCl. Calculate the concentration of Cl-ions required to just begin the precipitation of (a) AgCl, (b) Hg2Cl2, (c) PbCl2 in a solution hav-ing metal-ion concentrations of 0.030 M. What fraction of the Pb2+ remains in solution when the Ag+ just begins to precipitate?

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Textbook Question

A 40.0 mL sample of a mixture of HCl and H3PO4 was titrated with 0.100 M NaOH. The first equivalence point was reached after 88.0 mL of base, and the second equiva-lence point was reached after 126.4 mL of base. (f) What indicators would you select to signal the equiva-lence points?

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Textbook Question

A 1.000 L sample of HCl gas at 25°C and 732.0 mm Hg was absorbed completely in an aqueous solution that contained 6.954 g of Na2CO3 and 250.0 g of water. (a) What is the pH of the solution?

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Textbook Question

Consider a galvanic cell that utilizes the following half-reactions: (d) Will AgCl precipitate if 10.0 mL of 0.200 M NaCl is added to the solution in part (c)? Will AgBr precipitate if 10.0 mL of 0.200 M KBr is added to the resulting solution?

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