18. Aqueous Equilibrium

Buffer

Henderson-Hasselbalch Equation

You are asked to prepare a pH = 3.00 buffer solution starting 
from 1.25 L of a 1.00 M solution of hydrofluoric acid 
(HF) and any amount you need of sodium fluoride (NaF). 
(b) How many grams of sodium 
fluoride should be added to prepare the buffer solution? 
Neglect the small volume change that occurs when the sodium 
fluoride is added.

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?

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?

HC2H3O2 / NH3

A buffer consists of a conjugate acid-base pair.

Correct, F&#8722; is the conjugate base of HF.

 A buffer contains significant amounts of a weak acid and its conjugate base, or a weak base and its conjugate acid. 

Which of the following would produce a buffer?

Buffers Concept 1

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?

Method 1 for Creating a Buffer

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

Calculate the pH of the solution that results from each
mixture.
a. 50.0 mL of 0.15 M HCHO2 with 75.0 mL of 0.13 M NaCHO2

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+?

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

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?

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

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.

Consider the curve shown here for the titration of a weak
monoprotic acid with a strong base and answer each
question.
<img src="https://lightcat-files.s3.amazonaws.com/problem_images/15606c70bd6936be-1660313345820.jpg" alt="" />
c. At what volume of added base does pH = pKa?

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] 
<img src="https://lightcat-files.s3.amazonaws.com/problem_images/9ac6871141185b09-1667309337158.jpg" alt="" />

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.

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

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?

Referring to Table 17.1, pick an indicator for use in the titration
of each base with a strong acid.
a. CH3NH2<img src="https://lightcat-files.s3.amazonaws.com/problem_images/75d87aca5d4cdb38-1661871866750.jpg" alt="" />

If a strong base&#160;is added then the buffer resists a pH change by having the weak acid neutralize it.&#160;

Buffers

HC2H3O2 and NH3

CH3CH2NH2 and CH3CH2NH3+

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

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

Adding 0.003 mol HC2H3O2

Adding 0.007 mol Ca(C2H3O2)2

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?

50 mL of 0.10 M HF with 50 mL of 0.10 M NaOH.

50 mL of 0.10 M HNO2 with 25mL of 0.10 M Ca(OH)2

50 mL of 0.10 M CH3CO2H with 60 mL of 0.10 M NaOH

50 mL of 0.10 M HF with 30 mL of 0.10 M NaOH.

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

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

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?

Use the Henderson–Hasselbalch equation to calculate the pH of 
each solution in Problem 29.

Henderson-Hasselbalch Equation Example 2

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

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).

(b) What is the ratio of HCO3 - to H2CO3 in an exhausted 
marathon runner whose blood pH is 7.1?

Calculate the ratio of NaF to HF required to create a buffer with 
pH = 4.00.

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

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

Consider a galvanic cell that utilizes the following half-reactions:
 
<QUESTION REFERS TO APPENDIX B> 
(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?

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

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

This would be the pH before the KOH was added.

The addition of KOH will force the equilibrium to the right increasing the concentration of the conjugate base and decreasing the acid concentration.

Correct calculation of the pH after the equilibrium shift to the right.

The ratio of conjugate base to acid, is the ratio of [F&#8722;] to [HF] after the equilibrium shift rather the vice versa.

The ratio of conjugate base to acid, is the ratio of [F&#8722;] to [HF] rather than vice versa and the equilibrium shifts to the right rather the left.

 KOH reacts with a weak acid stoichiometrically to produce the conjugate base of the weak acid. 

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. Ka for HF = 3.5 &#215; 10&#8722;4.

Buffers resist drastic changes to the pH if a strong acid or strong base is added.&#160;

Correct, this amount of HCl is more than enough to consume all of the NaC2H3O2.

0.50 moles HC2H3O2

This addition would increase the ability of the buffer to react with a base.

20.0 g NaC2H3O2

This addition would increase the ability of the buffer to react with acid.

There would still be 0.30 moles of HC2H3O2 left to react with more base.

Only one of the additions depletes the buffer.

 You must completely eliminate all of one of the species in the buffer to destroy it. 

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?

The third method in creating a buffer is mixing a strong base with a weak acid. In this case since we have a strong species mixing with a weak species then we must make sure the weak species is higher in amount.&#160;

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

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.) <img src="https://lightcat-files.s3.amazonaws.com/problem_images/4dc2bc6ff1263b90-1675370487427.jpg" alt="" />
(b) Which solution has the greatest buffer capacity?

HC2H3O2 and KC2H3O2

H2SO3 and NaHSO3

H3PO4 and NaH2PO4

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

Consider the curve shown here for the titration of a weak base 
with a strong acid and answer each question. 
<img src="https://lightcat-files.s3.amazonaws.com/problem_images/0d29558418e96476-1660314271378.jpg" alt="" /> 
c. At what volume of added acid does pH = 14 - pKb?

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

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),

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.

You are asked to prepare a pH = 4.00 buffer starting 
from 1.50 L of 0.0200 M solution of benzoic acid 
1C6H5COOH2 and any amount you need of sodium benzoate 
1C6H5COONa2. (b) How many 
grams of sodium benzoate should be added to prepare the 
buffer? Neglect the small volume change that occurs when 
the sodium benzoate is added.

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?

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

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

Calculating the pH of a Buffer

The sodium hydroxide consumes a stoichiometric amount of HF reducing its concentration.

Correct - the [conjugate base] is 0.100 M and [conjugate acid] is 0.150 M.

This value is the pKa of HF. The ratio of conjugate base to acid must be considered.

This would be the pH if the concentrations of the conjugate acid and base were reversed from the given data.

 NaOH reacts with a weak acid stoichiometrically to produce the conjugate base of the acid. 

Determine the pH of a 1.0 L buffer prepared by adding 0.100 moles of NaOH to 0.250 moles of HF. Ka for HF = 3.5 &#215; 10&#8722;4.

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

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-

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.) <img src="https://lightcat-files.s3.amazonaws.com/problem_images/266427f4dcf565b6-1673020533427.jpg" alt="" />
(a) Which of the solutions are buffer solutions?

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)

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?

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)

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.

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

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

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?

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

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

Buffer Solutions

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?<img src="https://lightcat-files.s3.amazonaws.com/problem_images/d1b0622f1bee628c-1673038324370.jpg" alt="" />

HF is not a strong acid. Use the Henderson-Hasselbalch equation in which the [conjugate base] is 0.300 M and [conjugate acid] is 0.500 M.

Use the Henderson-Hasselbalch equation in which the [conjugate base] is 0.300 M and [conjugate acid] is 0.500 M.

Correct application of the Henderson-Hasselbalch equation.

This value is the pKa of HF. The ratio of conjugate base to acid must be considered.

 The use of an ICE table is helpful in buffer calculations. 

What is the pH of a 1.0 L buffer solution containing 0.500 M HF and 0.300 M NaF. Ka for HF = 3.5 &#215; 10&#8722;4.

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

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?
<img src="https://lightcat-files.s3.amazonaws.com/problem_images/1df33168aec488a4-1667307541814.jpg" alt="" />

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?

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

Practice Problem: Henderson-Hasselbalch Equation Calculations

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.

Calculate the ratio of CH3NH2 to CH3NH3Cl concentration 
required to create a buffer with pH = 10.24.

Derive the Henderson-Hasselblach Equation

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?

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]
<img src="https://lightcat-files.s3.amazonaws.com/problem_images/e03dbf7139c7f107-1667308554950.jpg" alt="" />

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.<img src="https://lightcat-files.s3.amazonaws.com/problem_images/fa212ffef7fbc863-1667306866391.jpg" alt="" />

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?

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]-

Henderson-Hasselbalch Equation Example 1

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

A buffer contains 0.10 mol of acetic acid and 0.13 mol of
sodium acetate in 1.00 L. (c) What is the pH of the buffer after the addition of 0.02 mol of HNO3?

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.)

A buffer contains 0.10 mol of acetic acid and 0.13 mol of
sodium acetate in 1.00 L. (b) What is the pH of the buffer after the addition of 0.02 mol of KOH?

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?

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

The first and most obvious way to create a buffer is to simply combine a weak acid and its conjugate base. In this case, a buffer is most ideal when both components are highly concentrated and equal to one another.&#160;

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

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

(b) What is the most significant difference between the sulfides precipitated in group 2 and those precipitated in group 3?<img src="https://lightcat-files.s3.amazonaws.com/problem_images/79f8c1482d0b3382-1667319865320.jpg" alt="" />

What mass of sodium benzoate should you add to 150.0 mL of a 
0.15 M benzoic acid solution to obtain a buffer with a pH of 
4.25? (Assume no volume change.)

The weak acid and conjugate base can be different from one another by up to a magnitude of 10. This is called the buffer range. If they are different by more than 10 then this is considered a bad buffer.

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

Whenever we have a buffer we can skip the ICE Chart and use the Henderson Hasselbalch.&#160;

Using the qualitative analysis flowchart in Figure 17.18 <img src="https://lightcat-files.s3.amazonaws.com/problem_images/24e049a90fe4fd4d-1673298887186.jpg" alt="" />, tell how you could separate the following pairs of ions. 
(a) Ag+ and Cu2+

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)

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?

Method 3 for Creating a Buffer

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.

The second method in creating a buffer is mixing a strong acid with a weak base. In this case since we have a strong species mixing with a weak species then we must make sure the weak species is higher in amount.&#160;

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)?

A 1.0-L buffer solution initially contains 0.25 mol of NH3 and 
0.25 mol of NH4Cl. In order to adjust the buffer pH to 8.75, 
should you add NaOH or HCl to the buffer mixture? What mass 
of the correct reagent should you add?

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] 
<img src="https://lightcat-files.s3.amazonaws.com/problem_images/8f011819d92c84c1-1667307874516.jpg" alt="" />

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

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

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.

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?

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

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

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

The Henderson – Hasselbalch Equation

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?

Buffer solution pH calculations | Chemistry | Khan Academy

Use the Henderson–Hasselbalch equation to calculate the pH of 
each solution. 
c. a solution that contains 10.0 g of HC2H3O2 and 10.0 g of 
NaC2H3O2 in 150.0 mL of solution

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+

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

If a strong acid&#160;is added then the buffer resists a pH change by having the conjugate base neutralize it.&#160;

Method 2 for Creating a Buffer

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

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?

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

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

Which buffer system is the best choice to create a buffer with 
pH = 9.00? For the best system, calculate the ratio of the masses 
of the buffer components required to make the buffer. 
HF/KF HNO2/KNO2 
NH3/NH4Cl HClO/KClO

Use the Henderson-Hasselbalch equation to calculate the ratio of the conjugate base to acid at the given pH.  Then use this ratio and the given concentration of the acid to determine the concentration of the conjugate base.

This value is the molar concentration of the acetate ion needed to attain the specified pH.

The Henderson-Hasselbalch equation states pH = pKa + log [base]/[acid]. The mass of NaC2H3O2 can be determined from the C2H3O2&#8722; concentration, the volume of the solution, and the molar mass.

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. Ka for HC2H3O2 = 1.8 &#215; 10&#8722;5.

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+&nbsp;solution. (Kb of&nbsp;C2H5NH2&nbsp;is 5.6 x 10&nbsp;-4).

(a) What is the ratio of HCO3 - to H2CO3 in blood of pH 7.4?

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?

Consider the following table of standard reduction potentials: <img src="https://lightcat-files.s3.amazonaws.com/problem_images/e8fc9881b39293b7-1675288357170.jpg" alt="" />
(b) Which substances can be oxidized by B2+? Which can be reduced by D?

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).

(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?

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.
(b) Which of the following can you establish about the solution: (i) A minimum pH, (ii) A maximum pH, or (iii) A specific range of pH values?

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.)

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?

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)

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

Referring to Table 17.1, pick an indicator for use in the titration
of each acid with a strong base.
a. HF<img src="https://lightcat-files.s3.amazonaws.com/problem_images/cd2d4006dbc4b99a-1661871784756.jpg" alt="" />

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.

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?

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

A buffer contains a weak acid, HA, and its conjugate
base. The weak acid has a pKa of 4.5, and the buffer has a
pH of 4.3. Without doing a calculation, state which of
these possibilities are correct at pH 4.3. (a) 3HA4 = 3A-4,
(b) 3HA4 7 3A-4, or (c) 3HA4 6 3A-4. [Section 17.2]

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

Henderson-Hasselbalch Equation Example 3

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

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.

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

As we stated earlier a buffer is composed of a weak acid and its conjugate base, but there are actually three ways to create a buffer.

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?

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?

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?

Henderson-Hasselbalch Equation Example 4

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?