Describe LeChatelier's principle.

(a)  

Consider the following reaction at equilibrium:

BaSO₄ $$\leftrightarrow$$ Ba²⁺(aq) +SO₄^2-(aq); ΔH > 0

In which direction will equilibrium shift if...

a. H₂SO₄ is added: ​ (a)  

b. BaCl₂ is added: ​ (b)  

c. NaCl is added: ​ (c)  

d. Heat is added: ​ (d)  

Consider the following reaction at equilibrium:

BaCO₃(s) $$\leftrightarrow$$ Ba²⁺(aq) + CO₃^2-(aq); ΔH > 0

In which direction will equilibrium shift if...

a. BaCl₂ (aq) is added: ​ (a)  

b. Li₂CO₃ (aq) is added: ​ (b)  

c. KNO₃ (aq) is added: ​ (c)  

d. BaCO₃ (s) is added: ​​ (d)  

e. HNO₃ (aq) is added: ​ (e)  

Consider the following weak acid reaction at equilibrium:

HNO₂(aq)+H₂O(l) $$\leftrightarrow$$ H₃O⁺(aq) + NO₂^-

In which direction will equilibrium shift if...

a. NaOH is added: ​​ (a)  

b. NaNO₂ is added: ​​ (b)  

c. HCl is added: ​​ (c)  

d. The acid is made more dilute: ​ (d)  

Consider the following weak acid reaction at equilibrium:

Ag⁺ (aq) + 2NH₃ (aq) $$\leftrightarrow$$ [Ag(NH₃)₂]⁺ (aq)

In which direction will equilibrium shift if...

a. NH₃ (aq) is added: ​​ (a)  

b. HNO₃ (aq) is added: ​​ (b)  

c. AgC₂H₃O₂ (aq) is added: ​​ ​ (c)  

Consider the following equilibrium.

2HI (g) $$\leftrightarrow$$ H₂ (g) + I₂ (g)

colorless $$\leftrightarrow$$ colorless + purple

When the reaction is heated, purple color faded to colorless. Is the reaction endothermic or exothermic? Explain.

The pH at the half-equivalence point of a weak acid titrated with a strong base is 5.289.
Calculate the acid-dissociation constant (Ka).

During a titration, the pH at the half-equivalence point for a weak acid is measured to be 3.752.
What is the Ka value?

In a titration of a weak acid with a strong base, the half-equivalence point is reached and the pH is found to be 6.105.
Determine the acid-dissociation constant (Ka).

Calculate the acid-dissociation constant (Ka) using the titration curve shown below. The titration is of a weak monoprotic acid with a strong base.

• At the half-equivalence point, 12.5 mL of NaOH has been added, and the pH is 5.25.
  (Assume the volume at the equivalence point is about 25.0 mL.)

Using the titration curve provided, determine the acid-dissociation constant (Ka) for a weak monoprotic acid titrated with a strong base.

• At the half-equivalence point, 15.0 mL of NaOH has been added, and the pH is 3.98.
  (The full equivalence occurs at about 30.0 mL.)

A weak acid is titrated with a strong base. Based on the titration curve, find the Ka for the acid.

• At the half-equivalence point, 10.0 mL of NaOH has been added, and the pH is 4.75.
  (The equivalence point is reached at 20.0 mL.)

If 25.15 mL of 0.1880 M NaOH are required to titrate 20.00 mL of a monoprotic acid HA to the equivalence point, calculate the original molarity of HA.

A student titrates 30.00 mL of a weak acid HA solution with 0.1500 M NaOH. It takes 35.45 mL of the NaOH solution to reach the equivalence point. What was the original molarity of the HA solution?

Round to 4 sig figs

During a titration, 18.67 mL of 0.2500 M NaOH are needed to titrate 15.00 mL of a monoprotic acid HA to the equivalence point. Determine the initial molarity of the HA solution.

A solution is made by mixing 25.0 mL of 0.300 M formic acid (HA) and 25.0 mL of 0.200 M sodium formate (A⁻).
The pKa\text{pKa}pKa of formic acid is 3.75.
Find the pH of the resulting buffer solution.

You create a buffer using 0.500 M ammonia (NH₃, the base) and 0.250 M ammonium chloride (NH₄⁺, the conjugate acid).
The pKa\text{pKa}pKa for ammonium ion is 9.25.
What is the pH of the buffer?

You have a 25.00 mL solution of 0.100 M acetic acid (HA). The pKa of acetic acid is 4.76. Calculate the initial pH of the solution before any NaOH is added.

In a titration of 50.00 mL of 0.200 M formic acid (HCOOH) with 0.100 M NaOH, 20.00 mL of NaOH has been added. Calculate the pH of the solution at this point. The pKa of formic acid is 3.75.

25.00 mL of 0.150 M propionic acid (CH₃CH₂COOH, pKa = 4.87) is titrated with 0.150 M NaOH. Calculate the pH at the equivalence point.

Balance the following reaction occurring in acidic solution:

MnO4−​(aq)+Fe2+(aq)→Mn2+(aq)+Fe3+(aq)

​ (a)   Fe2+ + ​ (b)   H+ + ​ (c)   MnO4− ​→ ​ (d)   Fe3+ + Mn2+ + ​ (e)   H2​O

Balance this redox reaction in acidic solution:

Cr2​O₇^-2−​(aq)+Cl−(aq)→Cr3+(aq)+Cl2​(g)

​ (a)   Cl− + ​ (b)   H+ + Cr2​O72− ​→ ​ (c)   Cl2​ + ​ (d)   Cr3+ + ​ (e)   H2​O

Balance this redox reaction in acidic solution:

NO3−​(aq)+Cu(s)→NO(g)+Cu2+(aq)

​ (a)   Cu+ ​ (b)   H+ +​ (c)   NO3−​ → ​ (d)   Cu2+ +​ (e)   NO + 4H2​O

• Given the balanced equation between KMnO₄ and Na₂C₂O₄.

If 0.1985 g of sodium oxalate (Na₂C₂O₄) requires 28.47 mL of a KMnO₄ solution to reach the equivalence point, calculate the molarity of the KMnO₄ solution.

A titration uses 0.2567 g of sodium oxalate (Na₂C₂O₄) and 36.12 mL of a KMnO₄ solution to reach the equivalence point. Find the molarity of the KMnO₄ solution.

In an experiment, 0.1742 g of sodium oxalate (Na₂C₂O₄) is titrated with 24.75 mL of KMnO₄ solution. Calculate the molarity of the KMnO₄ solution.

Sodium oxalate is 138 g/mol.

An impure sample of sodium oxalate (Na₂C₂O₄) weighs 0.2150 g.
If 32.18 mL of 0.02050 M KMnO₄ solution are required to reach the equivalence point, calculate the percent of oxalate ion (C₂O₄²⁻) in the original sample.

A 0.1847 g sample of impure sodium oxalate is titrated, requiring 27.40 mL of 0.01876 M KMnO₄ to reach the equivalence point.
Find the percent by mass of oxalate ion (C₂O₄²⁻) in the impure sample.

A 0.2063 g impure sample of Na₂C₂O₄ is analyzed and requires 35.22 mL of 0.01745 M KMnO₄ to reach equivalence.
Calculate the percent of oxalate ion (C₂O₄²⁻) in the original sample.

What is the format for cell notation using anode and cathode & anode solution and cathode solution?

(a)  

What is the format for cell notation using anode and cathode & anode solution and cathode solution?

(a)  

How would you solve this?

• Given the standard shorthand cell notation for a voltaic cell, be able to:

  • 1. Write the balanced overall redox reaction

  • 2. Determine how many electrons were transferred in the overall reaction

  • 3. Know the following formula and be able to use it to perform calculations:                            

 E^°_cell = E^°_red(cath)  -  E^°_red(anode)

E cell =

(a)  

E cell =

(a)  

A voltaic cell has a standard cell potential of

E cell∘ = +0.89 

and transfers 2 electrons in the overall reaction.
Calculate:

• (a) ΔG∘ in KJ

• (b) K (equilibrium constant)

Calculate to the 3 sig figs.

(a)  

A voltaic cell has a standard cell potential of

E cell∘ = +1.65

and transfers 3 electrons in the overall reaction.
Calculate:

• (a) ΔG∘ in KJ

• (b) K (equilibrium constant)

Calculate to 3 sig figs

(a)  

A voltaic cell has a standard cell potential of

E cell∘ = +.24

and transfers 1 electron in the overall reaction.
Calculate:

• (a) ΔG∘ in KJ

• (b) K (equilibrium constant)

Calculate to 5 sig figs.

(a)  

Organize these options into the right categories.

• Coulombs = Amps ​ (a)   seconds

Reduction occurs at the ​ (a)   and oxidation at the ​ (b)   for an electrolytic cell.

In the electrolysis of dilute sulfuric acid (H₂SO₄), write the net ionic half-reaction that occurs at the cathode.

(a)  

In the electrolysis of dilute sulfuric acid (H₂SO₄), write the net ionic half-reaction that occurs at the cathode.

(a)  

During the electrolysis of H₂SO₄ (aq), which ion is reduced at the cathode, and what is the half-reaction?

(a)  

In the electrolysis of H₂SO₄ (aq) using inert platinum electrodes, hydrogen gas is observed to form at the cathode.
Write the balanced net ionic half-reaction for this cathode process.

(a)  

Half-reaction:

Ag+ + e− → Ag (s)

How many seconds are needed to plate 2.50 g of Ag using a constant current of 1.75 A?

Half-reaction:

Cu2+ + 2e− → Cu (s)

How many grams of copper would be deposited if a current of 3.00 A is applied for 25.0 minutes?

Half-reaction:

Al3+ + 3e− → Al (s)

A current of 5.00 A is passed through an electrolytic cell containing Al3+ ions for 2.00 hours.
How many grams of aluminum are produced?

In a lab, electrolysis of acid was performed.
The following data were collected:

• Current = 0.500 A

• Time = 3600 s

• Volume of H₂ gas collected = 250.0 mL

• Pressure of H₂ gas = 1.00 atm

• Temperature = 298 K

• R=0.08206 

Calculate the experimental value of Faraday’s constant.

During an experiment, students observed:

• Current = 0.750 A

• Time = 1800 s

• Volume of H₂ collected = 150.0 mL

• Pressure of H₂ = 0.950 atm

• Temperature = 298 K

• R=0.08206 L

Calculate the experimental Faraday’s constant.

Round to the nearest whole number.

A different lab setup gave the following:

• Current = 1.00 A

• Time = 2400 s

• Volume of H₂ gas collected = 500.0 mL

• Pressure = 0.980 atm

• Temperature = 305 K

• R=0.08206 L

Find the experimental value of Faraday's constant.

Round to the nearest whole number.