CHEMISTRY 30 - EQUILIBRIUM, ACIDS & BASES



UNIT 3 CHEMISTRY 30 - EQUILIBRIUM, ACIDS & BASES

1. For the reaction, 2 X(g) + 3 Y(g) Z(g) the equilibrium constant expression is

A. [Z] B. [Z] C. [X]2 [Y]3 D. [Z]___

[X] [Y] [2X] [3Y] [Z] [X]2 [Y]3

2. For the reaction, H2(g) + I2(g) 2 HI(g) the equilibrium constant at a particular temperature is 30. At the same temperature, what is the equilibrium constant for 2 H2(g) + 2 I2(g) 4 HI(g)

A. 30 x 2 B. 1/30 C. 30 D. 30 X 30

3. In the equilibrium, N2(g) + O2(g) 2 NO(g) the initial concentrations (before reaction) of N2(g), O2(g), and NO(g) were respectively 0.5 mol/L, 0.2 mol/L and 0.0 mol/L. If the equilibrium concentration of NO(g) is x mol/L, the equilibrium concentrations in mol/L of N2(g) and O2(g) respectively are

A. (0.5 - x) and (0.2 - x) B. (0.5 - x/2) and (0.2 - x/2)

C. (0.5 - 2x) and (0.2 - 2x) D. (0.25 - x) and (0.1 - x)

4. In the equilibrium 2 SO2(g) + O2(g) 2 SO3(g) the initial concentrations of the reactants before the reaction started was 0.1 mol/L for SO2(g) and 0.02 mol/L for O2(g). There was no SO3(g) present initially. At equilibrium, the concentration of SO3(g) produced is 2x mol/L. The equilibrium concentrations of SO2(g) and O2(g) in mol/L are respectively

A. (0.1 - 2x) and (0.02 - x) B. (0.1 - 2x) and (0.02 - 2x)

C. (0.2 - 2x) and (0.04 - x) D. (0.1 - x) and (0.02 - x)

5. In the equilibrium FeO(s) + CO(g) Fe(s) + CO2(g) there was initially no Fe(s) or CO2(g) present. In order to produce 2 mol of CO2(g) in a 10 L container at equilibrium, what would the equilibrium concentration of CO(g) have to be in mol/L? (The initial concentration of CO(g) is x and excess FeO(s) is available.

A. x - 2 B. x - 0.2 C. 2 - x D. 0.2 - x

6. The equilibrium H2(g) + I2(g) 2 HI(g) the HI(g) is established in a 10 L container. In order to have 5 mol of HI(g) present at equilibrium, what will be the equilibrium concentration in mol/L for H2(g)? Initially, there is no HI(g) and x mol/L each of H2(g) and I2(g).

A. X - 5 B. x - 0.5 C. x - 0.25 D. x - 1.0

7. In the equilibrium A(g) + B(g) 2 C(g) the equilibrium concentrations of A(g), B(g) and 2 C(g) are respectively 0.1 - x mol/L, 0.1 - x mol/L and x mol/L. If Kc = 16, the actual equilibrium concentration, in mol/L for C(g) is

A. 8 B. 1.6 C. 0.16 D. 0.08

8. In the equilibrium A(g) + B(g) C(g) + D(g) the equilibrium concentrations of A(g), B(g), and C(g) and D(g) are respectively 0.11 - x mol/L, 0.11 - x mol/L, x mol/L and x mol/L. The Kc = 16, for the reaction is 100. The actual equilibrium concentration, in mol/L for C(g) is

A. 0.10 B. 0.01 C. 11 D. 1.1

9. For the reaction A(g) B(g) the equilibrium constant, Kc is 4.0. If originally, before equilibrium, there were only 0.2 mol/L of A(g) and no B(g), what is the equilibrium concentration of B(g)?

A. 0.8 B. 0.05 C. 0.16 D. 0.2

10. For the reaction A(g) B(g) the equilibrium constant, Kc is 9.0. If originally, before equilibrium, there were only 0.50 mol/L of A(g) and no B(g), what is the equilibrium concentration of B(g)?

A. 18 B. 0.062 C. 4.5 D. 0.45

11. For the equilibrium 2 HCl(g) H2(g) + Cl2(g), Kc = 0.020, in order to have 0.020 mol of chlorine at equilibrium, how many moles of HCl(g) should be initially placed in an empty 1.00 L container?

A. 0.18 B. 0.14 C. 0.04 D. 0.02

12. A very large equilibrium constant indicates

A. nothing about the rates of reaction

B. the rate of the forward reaction is faster than the rate of the reverse reaction

C. the rate of the reverse reaction is faster than the rate of the forward reaction

D. the rates of the forward and reverse reaction are equal and very fast

13. If the forward reaction and the reverse reaction of a system at equilibrium occurs at a very fast rate

A. the equilibrium constant will be very large

B. equilibrium can be established quickly

C. the equilibrium constant will be very large and equilibrium can be established quickly

D. the equilibrium constant will be very small but equilibrium can be established quickly

CHEMICAL EQUILIBRIUM

1. For each of the following, write the chemical reaction equation with appropriate equilibrium arrows.

a) pH measurements indicate that acetic acid in vinegar is approximately 1% ionized into hydrogen ions and acetate ions.

b) Quantitative analysis of the reaction of sodium sulfate and calcium chloride solutions shows that the products are favored.

c) Aluminum sulfate solution reacts quantitatively with a sodium hydroxide

solution.

2. Chlorine and carbon monoxide gases are mixed in a 1.00 L container and the following equilibrium is established.

CO(g) + Cl2(g) COCl2(g)

Initially, 1.5 mol of chlorine was present with excess carbon monoxide. At equilibrium, 0.80 mol of COCl2(g) was found.

a) Calculate the percent reaction.

b) Write the equilibrium law for this reaction.

c) At equilibrium, 1.75 mol of carbon monoxide and 0.70 mol of chlorine were present. Calculate the equilibrium constant.

3. Write the equilibrium law for each of the following chemical reaction equations.

a) 2 SO2(g) + O2(g) 2 SO3(g)

b) 2 NO2(g) 2 NO(g) + O2(g)

c) N2(g) + 3 H2(g) 2 NH3(g)

4. In an experiment at a high temperature, 0.500 mol/L of hydrogen bromide gas is placed into a sealed container and decomposes into hydrogen and bromine gases.

a) Write the equilibrium concentrations and law for this reaction.

b) The equilibrium concentrations in this system are [HBr(g)] = 0.240 mol/L,

[H2(g)][Br2(g)] = 0.130 mol/L. Calculate the equilibrium constant.

EQUILIBRIUM, ACIDS & BASES

1. Consider the gaseous equilibrium PCl3(g) + Cl2(g) PCl5(g). In which direction will this reaction move if the reaction pressure is increased by the addition of an inert gas?

A. In the forward direction.

B. In the reverse direction.

C. There will be no change in the position of equilibrium.

2. A catalyst is added to the equilibrium reaction NH4Cl(s) NH3(g) + Cl2(g). The reaction will

A. produce more of the products B. produce more reactant

C. go to completion D. reach equilibrium faster

3. Consider the gaseous equilibrium S8(s) + 8 O2(g) SO2(s). In which direction will this reaction move if more solid sulfur is added to the reaction vessel?

A. In the forward direction.

B. In the reverse direction.

C. There will be no change in the position of equilibrium.

4. Which change listed below would shift the following reaction to the right?

4 HCl(g) + O2(g) 2 Cl2(g) + 2 H2O(g)

A. addition of Cl2(g)

B. removal of O2(g)

C. increase of pressure, by volume decrease

D. addition of inert helium gas

5. In an endothermic reaction at equilibrium, what is the effect of lowering the temperature? The reaction

A. makes more products

B. makes more reactants

C. is unchanged

6. If more reactant is added to a gaseous reaction at equilibrium, what will happen to the value of the equilibrium constant? It will

A. increase B. decrease C. remain the same

7. If the temperature of a reaction increases, what will happen to the value of the

equilibrium constant? It

A. will increase B. will decrease

C. remain the same D. can either increase or decrease

8. If the temperature of an exothermic reaction increases, what will happen to the value of the equilibrium constant? It

A. will increase B. will decrease

C. remain the same D. can either increase or decrease

9. What major industrial chemical is produced in the Haber process?

A. Nitric acid B. Sulfuric acid

C. Ammonia D. Sodium hydroxide

10. In which of the following is ammonia NOT used?

A. Fertilizers B. Refrigerants

C. Explosives D. Fuels

11. Using Le Chatelier's Principle, we would predict that the exothermic reaction known as the Haber Process should be operated at

A. high pressures and high temperatures

B. high pressures and low temperatures

C. low pressures and high temperatures

D. low pressures and low temperatures

Le Châtelier's Principle

1. Nitrogen monoxide, a major air pollutant, is formed in automobile engines from the endothermic reaction of nitrogen and oxygen gases.

a) Write the equilibrium reaction equation including the term "energy" in the equation.

b) Describe the direction of the equilibrium shift if the concentration of oxygen is increased.

c) Describe the direction of the equilibrium shift if the pressure is increased.

d) Gasoline burns better at higher temperatures. What are some disadvantages of the operation of automobile engines at higher temperatures?

2. In a sealed container, nitrogen monoxide and oxygen gases are in equilibrium with nitrogen dioxide gas. The reaction of nitrogen monoxide and oxygen is exothermic.

2 NO(g) + O2(g) 2 NO2(g) + energy

Predict the equilibrium shift when the following changes are made.

a) The temperature is decreased.

b) The concentration of NO(g) is decreased.

c) The concentration of NO2(g) is increased.

d) The volume of the system is decreased.

3. The equilibrium of the iron(III) thiocyanate system (Lab Exercise 14A) is convenient to study. Thiocyanate ions are colorless and, at very low concentrations, iron(III) ions are essentially colorless. However, the FeSCN2+(aq) complex is highly colored even at low concentrations. Predict the color change in the equilibrium mixture when each of the following changes is made.

Fe3+(aq) + SCN-(aq) FeSCN2+(aq)

yellow colorless red

a) A crystal of KSCN(s) is added to the system.

b) A crystal of FeCl3(s) is added to the system.

c) A crystal of NaOH(s) is added to the system.

EQUILIBRIUM, ACIDS & BASES

Calculate the pH of each of the following:

1. 0.0040 mol/L HCl(aq)

2. 0.0090 mol/L HNO3(aq)

3. 0.060 mol/L HClO4(aq)

4. 0.015 mol/L NaOH(aq)

5. 0.00040 mol/L KOH(aq)

6. 0.725 mol/L HBr(aq)

Kw, pH, and pOH Calculations

1. Calculate the [OH-(aq)] in limes which have a [H+(aq)] of 1.3 x 10-2 mol/L.

2. Calculate the [H+(aq)] in lemons which have a [OH-(aq)] of 2.0 x 10-12 mol/L.

3. A sodium hydroxide solution is prepared by dissolving 2.50 g to make 2.00 L of solution. Calculate the hydroxide and hydrogen ion concentrations.

4. A 0.728 g sample of hydrogen chloride gas is dissolved in 200 mL of solution. Calculate the hydrogen and hydroxide ion concentrations.

5. A vinegar solution has a hydrogen ion concentration of 1.5 x 10-3 mol/L. Calculate the pH.

6. An ammonia solution has a pOH of 2.92. What is the concentration of hydroxide ions in the solution?

7. Calculate the pOH and pH of a solution made by dissolving 7.50 g of strontium hydroxide to make 500 mL of solution.

Complete the following table.

| |Substance |[H+(aq)] (mol/L) |pH |[OH-(aq)] (mol/L) |pOH |Acidic, Basic, or Neutral |

|8. |milk | | |3.2 x 10-8 | | |

|9. |pure water | |7.0 | | | |

|10. |blood |4.0 x 10-8 | | | | |

|11. |cleaner | | | |3.20 | |

EQUILIBRIUM, ACID & BASE QUESTIONS

1. List 5 properties of acids (operational definitions).

1.

2.

3.

4.

5.

2. List 5 properties of bases (operational definitions).

1.

2.

3.

4.

5.

3. When 60 g of NaOH(s) are dissolved in 250 mL of water, find:

a) [H3O+]

b) [OH-]

c) pH

EQUILIBRIUM, ACID & BASE QUESTIONS

1. Calculate the pH of a 0.10 mol/L solution of nitrous acid.

(2.13 NO QUAD 1.68 QUAD)

2. What is the [OH-(aq)] of the above solution?

(4.8 X 10 -13)

3. Calculate the pH of 0.010 mol/L Ca(OH)2(aq).

(12.30 )

4. A 0.10 mol/L solution of an unknown acid has a pH of 3.60. Determine the Ka and % reaction of this solution with water.

(6.32 X 10 -7 0.251%))

5. Which of the following solutions would have the lowest pH? Explain.

A. boric acid

B. hydrofluoric acid

C. oxalic acid

D. acetic acid

6. A student has a 10.0 mL sample of HNO3(aq). What happens to the pH if distilled water is added to the solution?

7. Calculate the [H+(aq)], [OH-(aq)] and pH for a 0.025 mol/L solution of HF(aq).

(2.67 X 10-12 (2.52 X 10-12 NON Q) )

8. Calculate the [H+(aq)], [OH-(aq)] and pH for a 0.0100 mol/L solution of H2S(aq).

(2.98 X 10 -5 M)

9. Calculate the pH for each of the following:

a) 0.0050 mol/L KHSO4(aq)

(2.44)

b) 0.010 mol/L KHSO3(aq)

(4.60)

c) 1.0 mol/L KHSO3(aq)

(3.60)

d) 0.500 mol/L KOOCCOOH(aq)

(2.066)

10. Find the pH of a solution made by dissolving 6.0 g of acetic acid in 200 mL of water.

(2.5)

11. An acid, HX(aq), is found to have a pH = 4.25 in a 0.33 mol/L solution. Find:

a) [H+(aq)]

(5.62 X 10-5 M)

b) % reaction

(0.0017%)

c) Ka

(9.57 X 10 -9)

12. A 0.0133 mol/L solution of acid has pH = 3.00. Find the % dissociation for this solution. What is the significance of this value?

(7.5%)

13. Calculate the pH of 0.050 mol/L H2S(aq), neglecting the second dissociation.

(4.13)

14. A 0.50 mol/L phosphoric acid is 12% ionized. Calculate the pH of the solution.

(1.22)

15. Calculate the pH of 0.150 mol/L HCN(aq).

(5.016)

16. If the pH goes from 4 to 3, what has happened to the [H+(aq)]?

17. The pH of 0.20 mol/L butyric acid is 4.66. What is the Ka for this acid? What is the significance of this number?

(2.39 X 10 -9 M)

Strengths of Acids

1. List three empirical properties that may be measured to distinguish among acids of different strengths.

2. Calculate the hydrogen ion concentration and pH of a 0.10 mol/L solution of nitrous

acid.

(8.1 X 10 -3 M)

3. Calculate the hydrogen ion concentration and pH of a solution prepared by dissolving 10.70 g of ammonium chloride to make 2.00 L of solution.

(7.6 X 10-6 M)

4. Use the ka value to determine the mass of sodium hydrogen sulfate required to prepare 500 mL of solution with a pH of 1.57.

(6.0 G)

5. A 0.80 mol/L solution of an unknown acid, HX(aq), has a pH of 3.75.

a) Calculate the percent reaction.

(2.2 X 10 -2% )

b) Calculate the acid ionization constant.

(4.0 X 10 -8)

6. Calculate the pH of a solution containing 0.25 mol/L of an acid with an acid ionization constant of 3.2 x 10-6 mol/L.

(3.05)

EQUILIBRIUM, ACID & BASE QUESTIONS

Label each species in the following equations as Brønsted-Lowry acid or base.

1. HSO3-(aq) + H2O(l) H3O+(aq) + SO32-(aq)

________ _________

2. NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)

________ _________

3. HF(aq) + HSO3-(aq) F-(aq) + H2SO3(aq)

________ _________

4. H2SO3(aq) + HS- (aq) HSO3-(aq) + H2S(aq)

________ _________

Write in the products for each reaction that follows and label each species in each equation as an acid or a base. Assume the transfer of only one proton. Charge must be conserved in the balanced equation.

5. HNO2(aq) + Cl-(aq) ________ + _________

________ _________

6. CH3COOH(aq) + SO42-(aq) ________ + _________

________ _________

7. H2S(aq) + NO2-(aq) ________ + _________

________ _________

8. HCO3-(aq) + S2-(aq) ________ + _________

________ _________

9. CO32-(aq) + H2O(l) ________ + _________

________ _________

10. List all species in the nine reactions above that have appeared both as an acid and a base.

11. Do Brønsted-Lowry acids (like Arrhenius acids) have to form acidic aqueous solutions? Explain.

EQUILIBRIUM, ACIDS & BASES

A. Write Brønsted-Lowry acid-base reactions (net ionic equations) for each of the following. Include states and label acids and bases. Predict equilibrium using equilibrium arrows.

1) An oven cleaner spill (NaOH) is neutralized by household vinegar.

2) Excess stomach acid (HCl) is treated with Milk of Magnesia (Mg(OH)2).

3) A solution of potassium hydrogen carbonate is mixed with a sodium sulfate solution.

4) Three solutions: nitrous acid, ethanoic acid, and sodium acetate are mixed.

B. 1) Write the conjugate acid of:

a. NO2- ___________________

b. HPO42- ___________________

c. HSO4- ___________________

d. OH- ___________________

2) Write the conjugate base of:

a. HF ___________________

b. NH4+ ___________________

c. HPO42- ___________________

d. HCO3- ___________________

3) Predict whether the equilibrium favors reactants or products.

a. NaHSO4 + NaNO2 ___________________

b. NaHCO3 + NaHSO3 ___________________

c. CH3COO- + H3PO4 ___________________

4) NaH2PO4(aq), NH4Cl(aq), KHSO4((aq), NaHCO3(aq) (All 0.10 mol/L.)

a. Which has the highest pH? ___________________

b. Which has the weakest bond with hydrogen? ________________

c. Which is the strongest base? ___________________

EQUILIBRIUM, ACIDS & BASES

1. Given: HX(aq) + B-(aq) HB(aq) + X-(aq)

a. Which species is the stronger acid?

b. Which species is the stronger base?

2. NaX(aq) has a higher pH than NaY(aq). Which is a stronger base X-(aq) or Y-(aq)? Explain.

3. Predict the most likely reaction:

a. Hydrofluoric acid is added to potassium sulfate solution.

b. Nitric acid is added to a calcium fluoride solution.

c. Hydroiodic acid is added to an ammonium acetate solution.

d. Carbonic acid and OCl-(aq).

e. Hydrofluoric acid and ammonia.

4. Given: HA + C-(aq) ——> HC + A-(aq)

HB + D-(aq) ——> HD + B-(aq)

HC + D-(aq) HA + D-(aq)

List the bases in order from strongest to weakest.

5. Choose the acids in each of the following reactions.

a. HPO42-(aq) + HSO4-(aq) H2PO4-(aq) + SO42-(aq)

b. HSO3-(aq) + H2O(l) H3O+(aq) + SO32-(aq)

c. CH3COOH(aq) + SO42-(aq) CH3COO-(aq) + HSO4-(aq)

Use appropriate sets of equilibrium arrows to indicate the equilibrium situation for the following acid/base reactions.

1. hydroiodic acid and sulfate ion

2. hydroiodic acid and sulfite ion

3. bromothymol blue and fluoride ion

4. hydronium ion and oxalate ion

5. ammonium ion and chloride ion

6. OH- and NH4+

7. H3O+ and NH3

8. H2CO3 and HSO3-

9. OH- and ethanoic acid

10. H3O+ and methyl orange ion

11. HPO42- and nitrite ion

12. cyanide ion and HOOCCOO-

13. H3O+ and OH-

14. phosphoric acid and C6H5COO-

15. H2S and NO2

Predicting Acid-Base Equilibria

1. What is the generalization used to predict the position of an acid-base equilibrium?

Predict the acid-base reaction for each of the following questions. Communicate your answer using the five-step method.

2. A nitrous acid spill is neutralized with sodium hydrogen carbonate (baking soda).

3. In a chemical analysis, a sample of methanoic acid is titrated in a quantitative reaction with standardized sodium hydroxide.

4. Could hypochlorous acid in a bleach solution be neutralized with baking soda?

5. Small quantities of the poisonous hydrocyanic acid can be produced in a research laboratory by quantitatively reacting sodium cyanide and hydrochloric acid.

6. A student mixes solutions of sodium acetate and ammonium chloride to test the predictive power of the Brønsted-Lowry concept.

7. Some excess hydrofluoric acid remained after its use in glass etching. To dispose of this hazardous acid, a technician decided to neutralize the solution.

a) What readily available base would she choose to ensure a complete neutralization?

b) Predict the neutralization reaction assuming a quantitative reaction.

EQUILIBRIUM, ACIDS & BASES

Questions

For the following equations:

a. label the acid, base, conjugate acid and conjugate base.

b. place the equilibrium arrows.

1. Aqueous solutions of hydroiodic acid and sodium sulfate are mixed.

2. Magnesium hydroxide is neutralized with acetic acid.

3. Aqueous solutions of nitric acid and ammonium carbonate are mixed.

4. Solutions of phosphoric acid and silver acetate are mixed.

5. Methyl orange ions were added to a solution of sodium hydrogen phosphate.

Brønsted-Lowry Definitions and Indicators

1. According to the Brønsted-Lowry concept, define an acid and a base in terms of an acid-base reaction.

2. a) What is meant by an "acidic" solution?

b) Does a Brønsted-Lowry acid have to form an acidic solution?

3. There are many species that are classified as amphiprotic.

a) What does amphiprotic mean?

b) What general type of species is amphiprotic?

4. HOOCCOO-(aq) + PO43-(aq) HPO42-(aq) + OOCCOO2-(aq)

a) Label each species as A or B for both forward and reverse reactions.

b) Identify all conjugate acid-base pairs.

c) According to the Brønsted-Lowry concept, what determines the position of this equilibrium?

5. What is the generalization about the strength of an acid relative to its conjugate base?

6. Use the following pH scale to label the colors for bromothymol blue over the 0-14 pH range. Identify the form of the indicator for each distinct color using conventional symbols.

| | | | | | | | | | | | | | |

0 7 14

7. Problem What is the approximate pH of an unknown solution?

Evidence Separate samples of the unknown solution turned blue litmus to red, congo red to blue, and orange IV to yellow.

Analysis

8. Three unknown solutions in unlabelled beakers have pH values of 5.8, 7.8 and 9.8. Write two diagnostic tests using indicators to identify the pH of each solution.

EQUILIBRIUM, ACIDS & BASES

1. Using the modified theory, complete the following

a. SO32-(aq) and H2O(l) ——>

b. SO2(aq) and H2O(l) ——>

c. HPO42-(aq) and H2O(l) ——>

2. All but which of the following is amphoteric?

H2O(l) HSO3-(aq) H2PO4-((aq) NO3-(aq)

3. a. Complete the following single-step reaction:

CO32-(aq) + HS-(l) ——>

b. If some K2S(s) is added to the above system, what will be the effect on equilibrium, according to Le Chatelier's Principle?

4. a. Complete the following reaction:

SO42-(aq) + HSO3-(aq) ——> __________ + __________

b. __________ __________ _________ ______________

Label each of the above species as acid or base in the spaces provided.

5. For each of the following, write the reaction equation. Label each species as an acid or a base and indicate the equilibrium arrows.

a. A sodium bicarbonate solution is mixed with a potassium hydroxide solution.

b. A lithium bisulfate solution is reacted with water.

c. HPO42-(aq) + HSO41-((aq)

d. Potassium carbonate and hydrochloric acid are reacted.

EQUILIBRIUM, ACID & BASE QUESTIONS

1. What volume of 0.0220 mol/L HNO3(aq) is required to neutralize 75.0 mL of 0.0340 mol/L KOH(aq)? If the neutral solution was boiled to dryness, what would be the formula of the melting salt?

2. A 50.0 mL sample of NaOH(aq) solution of unknown concentration was titrated to an endpoint using 68.4 mL of 1.45 mol/L HNO3(aq). Find [NaOH(aq)].

3. When 35.0 mL of HBr(aq) was titrated to an endpoint using 0.452 mol/L KOH(aq) it was found that 28.7 mL of the KOH(aq) had been used. What was the [HBr(aq)]?

4. What mass of KOH(s) is required to neutralize 45.0 mL of 0.560 mol/L HCl(aq)?

5. What volume of 0.230 mol/L Ca(OH)2(aq) is required to neutralize 50.0 mL of 0.560 mol/L HCl(aq)?

6. A 30.0 mL sample of LiOH(aq) solution was titrated to an endpoint using 14.9 mL of 0.654 mol/L HBr(aq). Find [LiOH(aq)].

7. 25.0 mL of an HCl(aq) solution [unknown] was titrated to an endpoint using 21.6 mL of 0.275 mol/L NaOH(aq). Find [HCl(aq)].

8. What volume of 0.200 mol/L NaOH(aq) is required to neutralize 50.0 mL of 0.350 mol/L HCl(aq)?

9. What volume of 0.156 mol/L LiOH(aq) is required to neutralize 4.50 L of 0.205 mol/L HClO4(aq)?

10. What volume of 2.00 mol/L HCl(aq) is needed to neutralize 1.20 g of dissolved NaOH(s)?

11. What volume of 3.00 mol/L HNO3(aq) is needed to neutralize 450 mL of 0.100 mol/L Sr(OH)2(aq)?

12. In a titration of sulfuric acid with a solution of 0.20 mol/L sodium hydroxide the following buret readings were obtained. Use the information to find [H2SO4(aq)].

ACID BASE

initial volume 4.55 mL initial volume 6.05 mL

final volume 28.60 mL final volume 35.40 mL

13. Find the volume of base required when 50.0 mL of 2.00 mol/L hydrochloric acid was titrated against a 1.50 mol/L potassium hydroxide solution.

14. If 30.0 mL of hydrochloric acid reacts completely with a 0.500 g sample of Na2CO3(s), calculate the [acid].

15. A 30.0 mL sample of HCl(aq), pH = 2.17, is reacted completely with 20.0 mL of NaOH(aq). What is the pH of the base?

16. A 20 000 L tank car that is filled with 20 mol/L H2SO4(aq) spills its contents into a trackside pit. A chemist hired by the railway company recommends that tank cars of 10 mol/L NaOH(aq) be brought to the scene to neutralize the spill. What volume of base is required?

17. It takes 28.7 mL of KOH(aq) to reach the second endpoint (ie. 2 H+ removed) with 19.3 mL of 0.150 mol/L H2SO4(aq). Calculate [KOH(aq)].

18. Calculate the volume of 0.0230 mol/L Ca(OH)2(aq) required to reach the second endpoint with 10.0 mL of 0.355 mol/L oxalic acid.

19. In a titration, 12.4 mL of HCl(aq) is required to react with 15.0 mL of 0.250 mol/L NH3(aq). What is the pH of the HCl(aq)?

20. A pellet of Ba(OH)2(aq) requires 25.7 mL of 2.25 mol/L HCl(aq) to neutralize. What is the mass of the pellet?

21. A small pellet of solid acid (3.20 g) required 32.0 mL of 0.30 mol/L NaOH(aq) to neutralize. Determine:

a. moles of H3O+

b. moles of OH-

c. molar mass of the acid

Acid-Base Stoichiometry

1. In a chemical analysis, 25.0 mL of sulfuric acid solution was titrated to the second endpoint with 0.358 mol/L KOH(aq). In the titration, an average volume of 18.2 mL was required. Calculate the molar concentration of the sulfuric acid.

2. Several 10.0 mL vinegar samples were titrated with a standardized 0.582 mol/L solution of sodium hydroxide. An average volume of 13.8 mL of sodium hydroxide was required to reach the phenolphthalein endpoint. What is the concentration of the vinegar solution?

3. A sodium borate solution was titrated to the second endpoint with 0.241 mol/L hydrobromic acid. An average volume of 15.2 mL of hydrobromic acid was required to react with 20.0 mL samples of sodium borate. Calculate the molar concentration of sodium borate.

4. Problem: What is the molar concentration of a hydrochloric acid solution?

Experimental Design: 100.0 mL of a standard solution of sodium oxalate was prepared using 1.85 g of the dry solid. Using the second endpoint, 10.0 mL samples were titrated with hydrochloric acid.

Evidence:

|TITRATION OF 10.0 mL OF Na2OOCCOO(aq) WITH HCl(aq) |

|Trial 1 2 3 4 |

|Final buret reading (mL) 16.1 31.5 46.9 16.9 |

|Initial buret reading (mL) 0.3 16.1 31.5 1.5 |

|Comment on endpoint poor good good good |

Analysis:

(ENRICHMENT) EXCESS PROBLEMS

1. Titrations involve a procedure where the reaction is stopped when it is quantitative and complete. A pH meter or indicator is used for the purpose of knowing when to stop an acid-base titration. Other contexts, such as plotting pH curves, involve adding an excess of one chemical or not adding enough of one chemical. In both cases the reaction may be stoichiometric and quantitative, but is not

complete.

Excess problems are recognized in texts and tests by too much information being provided. For reactions in solution, an excess problem is recognized when the concentration and volume of two chemicals in a reaction are provided. Questions usually ask for the final concentration of an acid, base, hydronium ion, or hydroxide ion. Other questions may ask for the pH or pOH of the final mixture. The general procedure for solving excess problems is outlined below.

1. Convert both of the concentration and volume values into amounts in moles.

2. Use the mole ratio, and then subtract the amounts, to determine the identity and extent of the excess amount.

3. Add the volumes of the two solutions to determine the final volume of the mixture.

4. Calculate the concentration of the excess chemical from the excess amount and the total volume of solution.

5. Calculate the excess hydronium or hydroxide ion concentration from the percentage ionization or Ka.

6. Convert the hydronium or hydroxide ion concentration to a pH or pOH if required.

Example ------------------------------------------------------------------------------------------------------

What is the pH of a solution produced by mixing 50 mL of 0.1 mol/L sodium hydroxide with 20 mL of 0.1 mol/L aqueous hydrogen sulfide?

2 NaOH(aq) + H2S(aq) ——> 2 HOH(l) + Na2S(aq)

50 mL 20 mL

0.1 mol/L 0.1 mol/L

Step 1: nNaOH = 50 mL x 0.1 mol/ nH2S = 20 mL x 0.1 mol/L

= 5 mmol (excess reagent) = 2 mmol (limiting reagent)

Step 2: nNaOH = 2 mmol x [pic] = 4 mmol (required)

nNaOH = 5 mmol - 4 mmol = 1 mmol (in excess)

Step 3/4: CNaOH = 1 mmol = 0.014 mol/L

(50 + 20) mL

Step 5: [OH-] = CNaOH = 0.014 mol/L

[H3O+(aq)] = 1.0 x 10-14 mol2/L2 = 7.0 x 10-13 mol/L

Step 6: pH = -log (7.0 x 10-13) = 12.15

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1. What is the pH of a solution produced by mixing 100 mL of 0.25 mol/L nitric acid with 50 mL of 0.40 mol/L potassium hydroxide solution?

2. Calculate the hydrogen and hydroxide ion concentrations after 60 mL of 0.1 mol/L hydrochloric acid is reacted with 25 mL of 0.1 mol/L barium hydroxide solution.

3. Calculate the hydrogen and hydroxide ion concentrations after 100 mL of 0.20 mol/L sodium hydroxide solution is reacted with 40 mL of 0.20 mol/L sulfuric acid solution.

4. A 10 mL sample of a hydrochloric acid solution of unknown concentration was titrated with a standardized 0.0759 mol/L sodium hydroxide solution. The endpoint of 11.5 mL was overshot by 1.5 mL of NaOH(aq). According to the evidence, what is the pH of the Erlenmeyer solution after overshooting the endpoint?

pH Curves

1. Use the accompanying sketch of a pH curve for a titration to answer the following questions.

Acid-Base Reaction

(a) Does the buret contain the acid or the base?

b) Is the sample reacted an acid or a base?

(c) How many endpoints are present?

Estimate each pH endpoint.

(d) How many quantitative reactions have occurred?

e) Choose the best indicator for each endpoint.

(f) What part of the curve represents a possible

buffering region?

2. Sketch a pH curve for the reaction of sulfuric acid with sodium hydroxide solution. All reaction steps are known to be quantitative. Include reaction equations.

3. A sodium hydrogen phosphate solution is to be titrated with hydrochloric acid. Only one quantitative reaction is observed. Sketch the pH curve and write equilibrium equations.

4. What are two advantages of a color endpoint using an indicator as opposed to a pH endpoint using a pH meter?

Equilibrium, Acids & Bases

Determine the pH of the solution that results when each of the following are added together.

1. 25.0 mL of pH = 3.55 added to 25.0 mL of pH = 11.25

2. 50.0 mL of pH = 6.35 added to 100 mL of pH = 6.60

3. 100 mL of pH = 5.72 added to 250 mL of pOH = 5.04

4. 50.0 mL of pH = 5.14 added to 30.0 mL of pH = 8.94

5. 120 mL of pOH = 8.75 added to 50.0 mL of pOH = 9.02

6. 135.0 mL of pH = 8.66 added to 70.0 mL of pOH = 5.14

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