General Chemistry
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Identify the type of reaction occurring given that it is carried out in standard conditions.
Sulfur dioxide gas reacts with carbon monoxide.
H2 and Cl2 gas combine to form HCl gas.
Sodium hydride decomposes into Na metal and H2 gas.
Consider the following spontaneous cell reactions:
X + Z2+ → X2+ + Z
Y + Z2+ → Y2+ + Z
X + Y2+ → X2+ + Y
Sort the following reduction half-reactions in decreasing order of likelihood of occurrence:
(1) X2+ + 2 e– → X
(2) Y2+ + 2 e– → Y
(3) Z2+ + 2 e– → Z
Does a reaction occur when a strip of cobalt is dipped into an aqueous solution of Fe(NO3)2 under standard-state conditions? If there is, what is the reaction?
Does a reaction occur when a strip of nickel is dipped into an aqueous solution of CuNO3 under standard-state conditions? If there is, what is the reaction?
Does a reaction occur when oxygen gas is bubbled through an acidic solution of Mn(NO3)2 under standard-state conditions? If there is, what is the reaction?
Identify the best oxidizing agent and the best reducing agent.
Ferrates can be formed by heating ferric oxide with other metal oxides or carbonates. One such reaction is Fe2O3 + ZnO → Zn(FeO2)2. Identify if Fe is oxidized, reduced, or neither in this reaction.
Peracetic acid is produced by the reaction of acetaldehyde with oxygen gas: O2 + CH3CHO → CH3CO3H. Is oxygen reduced or oxidized in this reaction?
Identify the reducing agent and the oxidizing agent in the following reaction: C + 2 H2SO4 → CO2 + 2 SO2 + 2 H2O.
Identify the reducing agent and the oxidizing agent in the following reaction: 2 Sb + 2 HNO3 → Sb2O3 + 2 NO + H2O
To turn NO into NO2, a/an _ is required.
Choose the correct statement
Which of the following statements is true?
Brønsted–Lowry defined acid-base reactions as proton-transfer reactions where a weak acid will have a strong conjugate base. In terms of redox reaction, which of the following statements is true?
What type of reaction occurs between Cu and AgNO3?
Determine the part of the periodic table in the image that is easily reduced:
Select the correct statement:
For Pt(SO4)2 + 4 RbI → PtI2 + I2 + 2 Rb2SO4, give the oxidation number for each element
Which statement is true for the following reaction?
CH4 (g) + 2 O2 (g) → CO2 (g) + 2 H2O (l)
For Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s), identify the reduced and oxidized element.
Select the redox reaction from the following
Identify the reducing agent and the oxidizing agent for the reaction of gold metal with I2 and NaCl producing NaAuCl4 and NaI.
Is the given reaction a redox reaction? If it is, identify the reactant that acts as an oxidizing agent and the reactant that acts as a reducing agent.
4 Na(s) + O2(g) → 2 Na2O(s)
NaOH(aq) + HCl(aq) → H2O(l) + NaCl(aq)
Identify element being reduced and element being oxidized in the following reaction:
1. Cl2 (aq) + 2NaI (aq) → I2 (aq) + 2NaCl (aq)
2. 3 Fe(NO3)2 (aq) + 2 AI (s) → 3 Fe (s) + 2 AI(NO 3)3 (aq)
Identify the oxidizing agent and the reducing agent for the following reaction.
2 Al3+(aq) + 2 Fe(s) → 2 Al(s) + 3 Fe2+(aq)
Which of the following reactions is NOT a Redox Reaction?
Identify the following for the reaction: 3 Br2 + 2 Al → 6 Br- + 2 Al3+
oxidized species: ________ oxidizing agent: _________
reduced species: ________ reducing agent: _________
For the choices below, identify where the electrons are being transferred from and to where.
Which of the following characteristics is associated with oxidation-reduction (redox) reactions?
Sulfuric acid can be produced via a wet process:S + O2 → SO22 SO2 + O2 ⇌ 2 SO3SO3 + H2O → H2SO4
Determine the element that is oxidized and reduced for the second step of the process.
Label the following as a non-redox or redox reaction. Determine the reducing agent and the oxidizing agent for the identified redox reaction.
I. Mg (s) +Cl2 (g) → MgCl2 (s)
II. Cu (s) + 2 Ag+ (aq) → Cu2+ (aq) + Ag (s)
III. 2 K (s) + F2 (g) → 2 KF (s)
IV. N2O3 (g) + H2O (l) → HNO2 (aq)
The best reducing agents are located in the lowest-right part of the periodic table while the best oxidizing agents are found in the top-left part of the periodic table. Is this statement true or false?
Using the given reduction potentials below, identify whether the reduction of Cl2(g) by HNO2(aq) will occur spontaneously at standard-state conditions.
E°red Cl2/Cl– = 1.358 VE°red NO3–/HNO2 = 0.934 V
Using the given reduction potentials below, identify whether the oxidation of Pt(s) by Zr4+(aq) will occur spontaneously at standard-state conditions.
E°red Pt2+/Pt = 1.18 VE°red Zr4+/Zr = –1.45 V
Using the given reduction potentials below, identify whether the reduction of Ag+(aq) by Cu+(aq) will occur spontaneously at standard-state conditions.
E°red Ag+/Ag = 0.800 VE°red Cu2+/Cu+ = 0.153 V
The following reaction is used by a voltaic cell:
2 Cu2+(aq) + H2(g) → 2 H+(aq) + 2 Cu+(aq)
When [Cu2+] = 2.50 M, PH2 = 0.79 atm, [Cu+] = 0.0014 M, and the pH in both half-cells is 4.50, what is the emf for this cell?
2 H+(aq) + 2 e− → H2(g) E°red = 0.00 V
Cu2+(aq) + e− → Cu+(aq) E°red = 0.16 V
Provide the balanced equation, standard cell potential, ΔG° and K at 298 K for the following reaction: The oxidation of chromium(II) ion in an acidic solution to chromium(III) ion by hypochlorous acid.
Cr3+(aq) + e− → Cr2+(aq) E°red = —0.50 V
2 HClO(aq) + 2 H+(aq) + 2 e− → Cl2(g) + 2 H2O(l) E°red = 1.61 V
A cell with the same half-reactions at the anode and the cathode was constructed with 0.125 M Ag+ in both electrode compartments. A stoichiometric amount of a cyanide ion-generating solution is carefully added to one of the two compartments. Handling the solution very carefully, the cell potential was measured and recorded for future reference. The measured cell potential is 0.380 V. Assuming that there is no change in volume, calculate the formation constant Kf for Ag(CN)2−.
A certain voltaic cell uses the half-reactions shown below:
Anode: Zr(s) + 2 H2O(l) → ZrO2(s) + 4 H+(aq) + 4 e–Cathode: Mn2+(aq) + 2 e– → Mn(s)
Using the thermodynamic values below at 25.0°C, calculate the cell potential (E°) and the equilibrium constant K for the overall cell reaction.
A rechargeable battery is classified as a secondary battery. Nickel-iron battery is an example of a secondary battery. Its half-reactions occur as follows:
Cathode: NiO(OH)(s) + H2O(l) + e− → Ni(OH)2(s) + OH−(aq)
Anode: Fe(s) + 2 OH−(aq) → Fe(OH)2(s) + 2e−
Given the following data, determine the standard cell potential for the nickel-iron battery:
NiO(OH)(s) + H2O(l) + e− → Ni(OH)2(s) + OH−(aq)
E°red,NiO(OH) = 0.490 V
E°red,Fe2+ = −0.450 V
Refer to the following half-reactions and their E° values:
Bi+(aq) + e– → Bi(s) E° = 0.240 VW3+(aq) + 3 e– → W(s) E° = 0.100 VGa3+(aq) + 3 e– → Ga(s) E° = –0.549 V
For the combination of reactions that produces the largest voltage at 25°C, provide the overall cell reaction and determine the cell potential (E°), ΔG° (in kJ), and equilibrium constant (K).
Consider the galvanic cell shown below with a potential of 1.369 V at 25°C
Pb(s) | PbCl2(s) | Cl–(0.20 M) || Cr2O72–(0.20 M), Cr3+(0.20 M), H+(0.20 M) | Pt(s)
Determine the Ksp value for PbCl2(s).
Cr2O72–(aq) + 14 H+(aq) + 6 e– → 2 Cr3+(aq) + 7 H2O(l) E°red = 1.232 VPb2+(aq) + 2 e– → Pb(s) E°red = –0.126 V
Illustrate a cell representation of the electrolysis of a concentrated aqueous solution of hydrochloric acid. Use inert electrodes. Label which are the anode, cathode, positive and negative electrodes, and the direction of ion and electron flow. Show the balanced equations for the important reactions involved.
Due to an unexpected storm, the power in your resthouse has been knocked out. You want to know the latest news, but your tablet is dead. You want to create a battery to charge it. From the materials found in your resthouse, the following allows you to make a battery:
1.00 M Cu2+ solution from herbicide
1.00 M Zn2+ solution from calamine lotion
Salt bridge from filter paper soaked in a KNO3 solution derived from toothpaste
Calculate the voltage that can be generated.
Molten carbonate fuel cells (MCFC) are newly developed fuel cells that commonly use molten lithium and potassium carbonates as electrolytes. The half-reactions in the cell are shown below:
Anode: H2(g) + CO32–(l) → H2O(l) + CO2(g) + 2 e–
Cathode: 1/2 O2(g) + CO2(g) + 2 e– → CO32–(l)
Using the thermodynamic values below, determine the cell potential (E°) and the equilibrium constant (K) for the overall reaction at 25°C. Identify whether the values of E° and K will increase, decrease or remain unchanged as the temperature decreases.
A galvanic cell is made up of a chromium electrode dipped in a 0.015 M Cr(NO3)3 and a lanthanum electrode dipped in a solution of La(NO3)3. What is the concentration of La3+ in the La(NO3)3 solution if the cell potential measured at 25 °C is 1.63 V? (E°red,La3+ = −2.38 V; E°red,Cr3+ = −0.73 V)
Consider a galvanic cell with the following reaction:
2 Au3+(aq) + 3 Pb(s) → 3 Pb2+(aq) + 2 Au(s)
What is the potential of the cell at 25.0 °C if the concentrations of Pb2+(aq) and Au3+(aq) are 0.250 M and 0.0150 M, respectively? (E°red,Pb2+ = −0.13 V, E°red,Au3+ = 1.50 V)
Consider the following cell:
Co(s) | Co2+ (2.0 M) || H+(? M) | H2( 1.0 atm) | Pt(s)
Calculate the pH of the solution in the cathode compartment if the observed potential of the cell is 0.14 V at 25 °C. (E°Co2+ = − 0.28 V)
A photographer was on a trip to an exotic place to capture stunning landscapes when the electrical system of his RV was knocked out. He tried to locate himself on the road using the maps app on his tablet but the tablet's battery was dead. In his RV, he found safety matches that he could use to make a 1.0 M KClO3 solution. He bought his RV from an old photographer, so he was able to find a small amount of silver nitrate (AgNO3) and made a 1.0 M solution. He was wearing a silver ring and had a $100 American Platinum Eagle coin in his pocket. He used a paper towel dipped in salt solution as a salt bridge.The battery in the tablet requires 5.0 V for charging. Can this battery charge his tablet?
What are the cathode, anode and overall cell reaction for the electrolysis, in a cell with inert electrodes, of the aqueous solution of a KClO2 salt?
ClO2−(aq) → ClO2(g) + e− E° = −0.95 V
K+(aq) + e− → K(s) E° = −2.92 V
Increasing the pH of the solution will favor the conversion of iron to rust. Is this statement true or false?
Magnesium can be acquired by the electrolysis of molten MgCl2. Calculate the required current to acquire 650 g of Mg in a span of 2 hours.
The electrochemical process of anodizing is used to add color and also a corrosion-resistant coat to metals like aluminum and titanium. Unlike aluminum anodizing, titanium anodizing requires a significantly thinner oxide layer. This thin layer of titanium oxide works by interfering with the wavelength of the incident light which results in different colors. The reactions involved in this process are:
Cathode (reduction): 4 H+(aq) + 4 e– → 2 H2(g)
Anode (oxidation): Ti(s) + 2 H2O(l) → TiO2(s) + 4 H+(aq) + 4 e–
The thickness and the various color of the oxide layer can be controlled by changing the voltage. The higher the voltage applied, the higher the current generated, and the thicker the oxide layer formed. Calculate the time it will take to produce a dark green titanium oxide film with a thickness of 0.13 µm on a square sheet of titanium metal with an edge length of 20.0 cm. The density of titanium oxide is 4.23 g/cm3 and the current used is 0.25 A.
An aqueous metal nitrate [M(NO3)3] solution was subjected to electrolysis using a current of 25.0 A. If 408 g of the metal is produced in 400 min, what is the identity of the metal ion (M3+)?
What are the cathode, anode and overall cell reactions for the electrolysis, in a cell with inert electrodes, of the aqueous solution of a NiMnO4 salt?
What are the cathode, anode and overall cell reaction for the electrolysis, in a cell with inert electrodes, of the aqueous solution of a CrI2 salt?
Which of the following are the end products of the overall process that takes place when an aqueous BaI2 solution is electrolyzed?
How many kilowatt-hours of electricity are needed to obtain 1.0×103 kg (1 metric ton) of chromium through the electrolysis of Cr3+, given that this process has a 40% efficiency and the applied voltage is 5.10 V?
(a) Assuming that the electrolytic cell has 78% efficiency, what is the mass of Cs formed from the electrolysis of molten CsCl using a current of 2.5×103 A for a period of 12 h? (b) Determine the minimum voltage needed to drive the reaction.
Cs+ + e– → Cs E°red = –3.03 VCl2 + 2 e– → 2 Cl– E°red = –1.36 V
An ammeter, which measures electrical current, is created by a student using the electrolysis of water into hydrogen and oxygen gases. If the student wants to collect 15.5 mL of water-saturated H2(g), how long should an electrical current of 0.670 amperes be run through the device? The temperature of the system is 27.0 °C, and the atmospheric pressure is 850 torr.
Active metals such as Mn are obtained via the electrolysis of molten salts. Why can't we use aqueous solutions to obtain active metals via electrolysis?
True or False. We can obtain potassium metal from the electrolysis of an aqueous solution of KF because K+ ions are reduced to K(s)
What happens at the anode when molten KF is electrolyzed?
The MBr3 is a metal bromide salt. What is the molar mass of the metal if 2.56 g of the metal is deposited when MBr3 undergoes electrolysis for 18.4 minutes with a 3.29 A current?
In the electrolysis of an aqueous solution of Pb(IO3)2, what are the half-reaction equations that occur at the anode and cathode?
For the electrolysis of a molten mixture of LiF and LiCl, determine the products formed.
For the electrolysis of pure molten KBr, determine the products formed.
A 250 mL of a 0.250 M NaCl solution with an initial pH of 7.00 was subjected to electrolysis. A 15.0 mL portion of the solution was set for analysis after 20 minutes of the process. 32.6 mL of a 0.150 M HCI solution was used to reach the endpoint.
Calculate the current (in A) running through the cell
The molecular view of the process taking place at the anode in a voltaic cell is shown in the diagram below. Why are the spheres used to depict the atoms in the anode larger than those in the solution?
It matters to be specific as to whether a molten or an aqueous solution will be subjected to electrolysis since the products vary for each case. List the products that will be formed when aqueous NaCl is electrolyzed in a cell with inert electrodes and briefly account for the difference.
Draw a cell setup for the electrolysis of molten ZnCl2 and indicate the direction of the flow of electrons.
Molten sodium chloride can be reduced to sodium metal by passing an electric current through the molten salt to split it into its elements. Illustrated below is a simple diagram of the cell. Show the balanced equations for the anode, cathode, and overall cell reactions.
Draw a cell setup for the electrolysis of molten AlBr3 and indicate the direction of the flow of electrons.
The cell shown in the image below is used to commercially produce Sr, the element, via electrolysis from a molten salt (the "electrolyte"), such as SrCl2. Remember that in an electrolytic cell, which is the opposite of what we see in batteries, the anode is given the - sign and the cathode is given the + sign. What is the half-reaction at the anode of the electrolytic cell?
An experiment is conducted on an alternative source for the production of methanol (CH3OH). One way to determine if methanol is produced is through titration with acidic potassium dichromate. The presence of methanol can be determined since Cr2O72− is converted to Cr3+ and indicated by a color change from orange to green. During titration, methanol is oxidized to formaldehyde (CH2O) as described in the following chemical reaction:
3CH3OH(aq) + Cr2O72−(aq) + 8H+(aq) → 3CH2O(aq) + 2Cr3+(aq) + 7H2O(l)
If the standard half-cell potential for the reduction of formaldehyde to methanol is 0.110 V, what is E° for the reaction?
During the weekend, you stayed in your vacation villa in the countryside when a blackout happened. You desire to create a battery to charge your phone. You brought a liquid foundation with you which can be used to make a 1.00 M Mg2+ solution. You have bathroom cleaners in your villa that can produce a 1.00 M H2O2 solution. You also made a salt bridge from tissue paper soaked in a KNO3 solution made from makeup primer.
Determine the voltage that can be produced. (E°red,Mg2+ = −2.37 V; E°red,H2O2 = 1.78 V)
Provide the balanced equation and calculate the value of E° for the formation of O2 and H2O from H2O2 given the following half-reactions and corresponding E° values:
H2O2(aq) + 2 H+(aq) + 2e− → 2 H2O(l) E°H2O2→H2O = 1.78 V
O2(g) + 2 H+(aq) + 2e− → H2O2(aq) E°O2→H2O2 = 0.70 V
Determine if the reaction is spontaneous under standard-state conditions.
In an experiment, Hg2Cl2(s) is converted into Hg(l) and Cl2(g). Is the overall reaction a galvanic or electrolytic cell?
Determine E° for the following reaction:
F2(g) + 2Na(s) → 2Na+(aq) + 2F−(aq)
Under standard-state conditions, is the reaction spontaneous?
Among the following spontaneous cell reactions,
Y+ + W → Y + W+
W+ + Z → W + Z+
Y+ + Z → Y + Z+
X+ + Z → X + Z+
Y+ + X → Y + X+
Identify the reaction that produces the highest voltage.
True or False. In the hydrogen fuel cell, the reaction of hydrogen gas with oxygen that produces water is spontaneous while the reverse reaction is not spontaneous.
True or False. The reduction of Co3+ to Co2+ by H2O2 can occur spontaneously in an acidic solution under standard conditions.
The overall reaction for a zinc-bromine battery is: Br2(l) + Zn(s) → 2 Br–(aq) + Zn2+(aq). Using the standard reduction potentials below, determine the standard emf generated by the battery.
Zn2+(aq) + 2 e– → Zn(s) E°red = –0.76 V
Br2(l) + 2 e– → 2 Br–(aq) E°red = +1.07 V
Everyday household items such as flashlights and radios use alkaline batteries. The overall reaction in the battery is as follows: Zn(s) + 2 MnO2(s) ⇌ ZnO(s) + Mn2O3(s). Using a reliable source, give the two half-reactions that occur in the battery and calculate the standard emf generated by the voltaic cell.
Consider the following half-reactions with their standard reduction potentials.
(a) Au+(aq) + e– → Au(s) E°red = 1.692 V(b) Cr2+(aq) + 2 e– → Cr(s) E°red = –0.913 V(c) Zn2+(aq) + 2 e– → Zn(s) E°red = –0.763 V(d) Bi3+(aq) + 3 e– → Bi(s) E°red = 0.308 V
Identify the combination of these half-cell reactions that will results in an overall reaction with the smallest positive E°cell, then provide the value of E°cell.
Determine the standard cell potential for the overall reaction: 6 HClO(aq) + 6 H+(aq) + 2 Au(s) → 3 Cl2(g) + 6 H2O(l) + 2 Au3+(aq).
2 HClO(aq) + 2 H+(aq) + 2 e– → Cl2(g) + 2 H2O(l) E°red = +1.63 V
Au3+(aq) + 3 e– → Au(s) E°red = +1.498 V
A certain voltaic cell has a standard cell potential of 2.516 V and is given by the following overall reaction: AuBr4–(aq) + Al(s) → Au(s) + 4 Br–(aq) + Al3+(aq). If the E°red of Al is –1.662 V, determine the E°red for the reaction associated with Au.
Which of the following images shows the correct electrolytic cell with labeled anode and cathode and indicated electron flow direction where Co is oxidized to Co2+ and Ni2+ is reduced to Ni? What are the half-reactions at the anode and the cathode? How much voltage is needed to run the reaction?
For the following electrochemical cell, calculate the standard cell potential:
MnO4–(aq) + 4 H+(aq) + 3 K(s) → MnO2(s) + 2 H2O(l) + 3 K+(aq)
For the following balanced redox reaction,
Pb2+(aq) + H2S(g) → Pb(s) + S(s) + 2H+(aq)
What is the E°cell? Is the forward direction of the redox reaction spontaneous or nonspontaneous?
Cr3+(aq) + Fe(s) → Cr(s) + Fe3+(aq)
Is the forward direction of the redox reaction spontaneous?
Ba2+(aq) + 2 Na(s) → Ba(s) + 2 Na+(aq)
Sn4+(aq) + Mg(s) → Sn2+(aq) + Mg2+ (aq)
Calculate the standard cell potential for the following reaction:
IO3- (aq) + 2 H+ (aq) + 2 Cl- (aq) → Cl2 (aq) + IO- (aq) + 2 H2O (l)
At 25°C, which of the following pairs would you predict to react spontaneously?
Ba2+(aq) + 2e- → Ba(s) E° = -2.90 V
Sn4+(aq) + 2e- → Sn2+(aq) E° = 0.15 V
Mg2+(aq) + 2e- → Mg(s) E° = -2.37 V
Zn2+(aq) + 2e- → Zn(s) E° = -0.73 V
Na+(aq) + e- → Na(s) E° = -2.71 V
Fe3+(aq) + 3e- → Fe(s) E° = -0.036 V
Cr3+(aq) + 3e- → Cr(s) E° = -0.73 V
Ni2+(aq) + 2e- → Ni(s) E° = -0.23 V
The redox titration of tin and potassium dichromate (VI) is performed in an electrochemical cell that has a platinum electrode, a calomel reference electrode made of an Hg2Cl2/Hg electrode, and a saturated KCl solution with 3.0 M Cl−. What is the cell potential of the solution if 55.0 mL of 0.020 M K2Cr2O7 is added to a 150 mL solution of 0.020 M Sn2+ in 1.75 M H2SO4? (E°red,Sn4+/Sn2+ = 0.15 V; E°red,Cr2O7 2−/Cr3+ = 1.36 V)
Consider the electrode half-reaction of an alkaline battery:
2 MnO2(s) + H2O(l) + 2 e− → Mn2O3(s) + 2 OH−(aq) E°red = 0.15 V
ZnO(s) + H2O(l) + 2 e− → Zn(s) + 2 OH−(aq) E°red = −1.25 V
How does a fivefold increase in the concentration of NaOH in the electrolyte affect the cell voltage? Explain.
Breathalyzers are used to test the amount of blood alcohol level in the exhaled breath of suspected drunk drivers. An acidic solution of potassium dichromate is usually used to oxidize alcohol (ethanol) to acetic acid. In an experiment, an acidic solution of potassium permanganate was used rather than potassium dichromate.
5 CH3CH2OH(aq) [ethanol] + 4 MnO4−(aq) + 12 H+(aq) → 5 CH3CO2H(aq) [acetic acid] + 4 Mn2+(aq) + 11 H2O(l)
Potassium permanganate changes color from a deep purple to a colorless solution after being reduced. The Breathalyzer can measure this change in color to get a reading of how much alcohol is in the blood. Determine the potential of the reaction if the concentrations of ethanol, acetic acid, MnO4−, and Mn2+ is 1.5 M and the pH of the solution is 4.50. (E°red,CH3CO2H = 0.058 V; E°red,MnO4− = 1.51 V)
Iron can reduce sulfuric acid to produce solid sulfur. It can also reduce sulfuric acid to produce sulfur dioxide:
(i) 3 Fe(s) + SO42−(aq) + 8 H+(aq) → S(s) + 4 H2O(l) + 3 Fe2+(aq) E° = 0.49 V
(ii) Fe(s) + SO42−(aq) + 4 H+(aq) → SO2(g) + 2 H2O(l) + Fe2+(aq) E° = 0.28 V
What is the potential for reaction (i) if there are 0.15 M Fe2+ and 0.20 M H2SO4? What is the potential for reaction (ii) if there are 0.15 M Fe2+, 0.20 M H2SO4, and 1.0 atm SO2 gas? Which of the two reactions has a greater thermodynamic tendency if the concentration of H2SO4 is 0.15 M? (Assume that H2SO4 fully dissociates.)
Some naturally occurring metals react spontaneously with the oxygen in the air.
Can oxygen oxidize Cr2+(aq) to Cr3+(aq) spontaneously at 25 °C under the following environmental conditions: [Cr2+] = [Cr3+] = 1×10−6 M; pH = 7.0; PO2 = 170 mmHg?
Cr3+(aq) + e− → Cr2+(aq) E°red = −0.41 V
O2(g) + 4 H+(aq) + 4 e− → 2 H2O(l) E°red = 1.23V
Determine what happens to the cell voltage when each of the following changes is made to the Daniell cell originally in standard conditions. Show the balanced equation for each reaction.
a) 2.25 M Zn(NO3)2 is added to the anode compartment
b) 2.25 M CuSO4 is added to the cathode compartment
c) 1.00 M Cu(NO3)2 is added to the cathode compartment
d) 2.25 M H2SO4 is added to the cathode compartment
The Nernst equation is applicable not just to cell reactions, but also to half-reactions. Consider the following half-reaction,
2 Hg2+(aq) + 2 e− → 2 Hg22+(aq) E° = 0.92 V
Using the Nernst equation, what is its potential at 25 °C if [Hg2+] = [Hg22+] = 0.15 M?
The following half-cell reactions are used in a certain voltaic cell:
Br2(l) + 2 e– → 2 Br–(aq) E°red = 1.07 VAg+(aq) + e– → Ag(s) E°red = 0.80 V
If the concentrations of Br–(aq) and Ag+(aq) at 298 K are 0.015 M and 0.10 M respectively, calculate the emf generated by this cell at these concentrations.
The overall reaction for a a voltaic cell is Co2+(aq) + Ni(s) → Co(s) + Ni2+(aq). The cell potential is +0.16 V when the concentration of Co2+ in the cathode is 1.5 M. Calculate the concentration of Ni2+ in the anode half-cell for this case.
Ni2+(aq) + 2 e– → Ni(s) E° = –0.26 VCo2+(aq) + 2 e– → Co(s) E° = –0.28 V
Calculate the emf for the voltaic cell with the reaction
4 Cr2+(aq) + O2(g) + 4 H+(aq) → 4 Cr3+(aq) + 2 H2O(l)
when PO2 = 0.75 atm, Cr2+(aq) = 1.7 M , Cr3+(aq) = 0.012 M, pH at the cathode half-cell = 3.25.
A voltaic cell is built and operates at 298 K using the following reaction: Fe(s) + Co2+(aq) → Co(s) + Fe2+(aq). When Fe2+(aq) = 0.750 M and Co2+ = 0.350 M, what is the emf of the cell?
A voltaic cell has the following reaction: 2 Cu+(aq) + Cd(s) → 2 Cu(s) + Cd2+(aq). Will the addition of water to the half-cell of the anode decrease, increase, or have no effect on the cell emf?
Calculate the change in cell voltage when the ion concentrations in the half-cell of the cathode is increased by a factor of 100 for the following voltaic cell.
The pH of a prepared PbO2(s)/Pb2+(aq) electrode is 12.38. At 25°C, what is the concentration of Pb2+ needed to decrease the potential of the half-cell to 0.0 V?
Consider the following half-reactions that occur in an electrochemical cell:
Ox: Mn(s) → Mn2+(aq, 2.50 M) + 2 e–
Red: O2(g, 0.150 atm) + 2 H2O(l) + 4 e– → 4 OH–(aq, 2.50 M)
Determine the cell potential at 25.0°C.
At 25°C, the voltage of a Co/Co2+ concentration cell is 0.12 V. One of the half cells has a concentration of 1.2x10-3 M. Determine the concentration of Co2+ in the other half-cell assuming that it has a smaller concentration
At 25 °C, two Zn/Zn2+ half-cells comprise a concentration cell with a cell potential of 0.32 V. Determine the proportion of Zn2+ in the two half-cells
A voltaic cell initially consists of a 0.0250 M Mg/Mg2+ half-cell and a 1.32 M Ag/Ag+ half-cell. Calculate the cell potential when the Ag+ has already dropped to 3.18 V.
A voltaic cell initially consists of a 0.0250 M Mg/Mg2+ half-cell and a 1.32 M Ag/Ag+ half-cell. Calculate the cell potential when the Ag+ has already dropped to 0.50 M.
Ox: Fe(s) → Fe3+(aq, 0.200 M) + 3 e–
Red: PbO2(s) + 4 H+(aq, 2.10 M) + SO42–(aq, 1.60 M) + 2 e– → PbSO4(s) + 2 H2O(l)
Calculate the initial cell potential of a voltaic cell consisting of a 0.0250 M Mg/Mg2+ half-cell and a 1.32 M Ag/Ag+ half-cell.
The following redox reaction occurs in a galvanic cell:
Cu2+(aq) + Ni(s) → Cu(s) + Ni2+(aq)
Determine the cell potential at 25.0°C when [Cu2+] = 2.50 M; [Ni2+] = 0.0150 M.
Calculate the pH needed for a standard hydrogen electrode to have an electrode potential of -0.153 V
The reduction of iron and the oxidation of zinc creates a theoretical battery. Calculate the initial voltage of the battery if there are 1.3 M of Zn2+ and 2.8 M of Fe2+ in a 1.0-liter half-cells initially.
Pt (s) | H2(g, 0.839 atm) | H+ (aq, ? M) || Cr2+ (aq, 1.50 M) | Cr(s)
Calculate the H+ concentration at 25 °C given that the Ecell = -0.847 V
Use the following reduction potential values:
2 H+(aq) + 2e- → H2(g) E° = 0.00 V
Cr2+ (aq) + 2e- → Cr(s) E° = -0.913 V
Calculate the cell potential at 75 ∘C for the reaction given the following concentrations:
[Fe2+] = 1.30 M
[Mg2+] = 0.560 M
Mg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s)
Predict the sign of E° when K > 1 and determine the value of K if E° = 0 V.
For the reaction 2 I–(aq) + Hg2+(aq) → I2(s) + Hg(l) at 298 K, calculate Eºcell, K, ΔGºrxn, and fill the table below.
Consider the half-reaction below with an E° value of –2.069 V:
[AlF6]3–(aq) + 3 e– → Al(s) + 6 F–(aq)
Determine the formation constant (Kf) of [AlF6]3–. [Use E°red Al3+ = –1.662 V]
Identify whether the following statement is True or False: The value of ΔG° will be positive and K will be less than 1 when a reaction has a positive E° value.
A Nickel-Cadmium battery is rechargeable and is composed of a cadmium metal as an anode and a nickel(III) compound NiO(OH) supported on nickel metal as a cathode. This has an E° cell of 1.2 V and uses the following reactions:
Anode: Cd (s) + 2 OH− (aq) → Cd(OH)2 (s) + 2e−Cathode: 2 NiO(OH) (s) + 2 H2O (l) + 2e− → 2 Ni(OH)2 (s) + 2 OH− (aq)Overall: 2 NiO(OH) (s) + Cd (s) + 2 H2O (l) ⇌ 2 Ni(OH)2 (s) + Cd(OH)2 (s)
Find ΔG° (in kJ) and equilibrium constant K for the cell reaction at 25 °C.
The lead storage battery involves the following electrode half-reactions:
Anode: Pb (s) + HSO4− (aq) → PbSO4 (s) + H+ (aq) + 2 e− (E° = 0.296 V)
Cathode: PbO2 (s) + 3 H+ (aq) + HSO4− (aq) + 2 e− → PbSO4 (s) + 2 H2O (l) (E° = 1.680 V)
Overall: PbO2 (s) + Pb (s) + 2 H+ (aq) + 2 HSO4− (aq) → 2PbSO4 (s) + 2 H2O (l) (E° = 1.976 V)
Find the ΔG° (in kJ) and K at 25 °C.
At 298 K, the standard cell potential of a cell is +0.124 V. If there are 4 electrons transferred in the reaction, what is the equilibrium constant value for the reaction?
Given the following E°red
Cl2(g) + 2 e− → 2 Cl− (aq) E°red = 1.36 V
MnO4−(aq) + 8 H+(aq) + 5 e− → Mn2+(aq) + 4 H2O(l) E°red = 1.51 V
What is the value of the equilibrium constant for the reaction at 298 K?
10 Cl− (aq) + 2 MnO4−(aq) + 16 H+(aq) → 2 Mn2+(aq) + 8 H2O(l) + 5 Cl2(g)
What are the balanced equation, standard emf, ΔG°, and K value at 298 K for the oxidation of MnO2(s) to MnO4−(aq) by ClO−(aq) in basic solution?
What are the balanced equation, standard emf, ΔG°, and K value at 298 K for the oxidation of aqueous chloride ion by Hg22+(aq) to produce Cl2(g)?
Consider the reaction of Cu2+(aq) and Ni(s) at 25 °C. Calculate the equilibrium constant for this reaction using reduction potentials.
Determine K and ∆G°rxn at 25.0 °C using tabulated electrode potentials for the reaction of HClO(aq) with Cl-(aq) to an acidic solution producing Cl2(g).
Consider the reaction below:
Pb(s) + 2 Ag+(aq) → Pb2+(aq) + 2 Ag(s)
Calculate the equilibrium constant using the reduction potentials.
Silver is extracted from its ore via the leaching process for the cyanide ion in that ore. Shown below is the overall reaction.
4 Ag (s) + 8 CN− (aq) + O2 (g) + 2 H2O (l) → 4 [Ag(CN)2]− (aq) + 4 OH− (aq)
Use the given data to calculate ΔG° for the above reaction at 25 °C.
Ag+ (aq) + 2 CN− (aq) → Ag(CN)2− (aq); Kf = 3.00×1020
H2O (l) ⇌ H+ (aq) + OH− (aq); Kw = 1.00×10−14
O2 (g) + 4 H+ (aq) + 4 e− → 2 H2O (l); E° = 1.229 V
Ag+ (aq) + e− → Ag (s); E° = 0.800 V
For each of the Q/K ratios, identify whether Ecell and ΔG are negative, positive, or zero.
Q/K ratios: <1, =1, >1
An electrochemical is depicted with the cell notation Cu | Cu2+ || Mn3+ | Mn2+ | Pt. Draw the cell and show the flow of electrons and ions.
For the given cell notation, write an overall cell reaction:
Mn(s) | Mn2+(aq) || Fe2+(aq) | Fe(s)
Draw the cell diagram, label which electrode is the anode and the cathode. Draw arrows to show the directions in which the ions and electrons flow.
A galvanic cell is made up of a Zn/Zn2+ anode and Pb/Pb2+ cathode connected together by a wire. A salt bridge also connects the two cells. Write a cell notation for this galvanic cell.
Consider the following redox reaction:
2 NH4+(aq) + MnO4−(aq) → 2 MnO2(s) + N2(g) + 4 H2O(l)
Write a shorthand notation for this reaction. In your notation, you can use an inert metal if necessary.
Ni(s) | Ni2+(aq) || Br2(l) | Br−(aq) | Pt(s)
Explain why we need an inert electrode at the cathode.
Write a shorthand notation for the following electrochemical cell:
Consider the following cell notation:
Mg(s) | Mg2+(aq) || Cu2+(aq) | Cu(s)
Determine the Mg2+:Cu2+ concentration ratio at 25.0 °C if the cell potential is measured to be 2.74 V. (E°red,Mg2+ = −2.37 V; E°red,Cu2+ = 0.34 V)
The galvanic cell, Mg(s) | Mg2+(aq) || Ce4+(aq), Ce3+(aq)| Pt(s), has a standard potential of 0.870 V. Provide the standard reduction potential for the Ce4+/Ce3+ half-cell.
For the given electrochemical cell, cell potential is dependent on the copper concentration in the cathode half-cell.Pt(s) | H2(g, 1.0 atm) | H+(aq, 1.0 M) || Cu+(aq, ? M) | Cu(s)This cell has an Ecell of 0.34 V. Determine the [Cu+] of the solution.
Calculate the E°cell of the following celll:
Fe(s) | Fe3+(aq) || Cl2(g) | Cl-(aq) | Pt(s)
Write the cell notation for the reaction below:
3 Cu2+(aq) + 2 Al(s) ↔ 3 Cu(s) + 2 Al3+(aq)
Identify and balance the reaction for the following cell:
Pb(s) | Pb2+(aq) || NO(g) | NO3-(aq),H+(aq) | Pt(s)
Identify the metals that can be used for the cathodic protection of iron: Ni, Mg, Co, Zn, Cd
(a) What is the mass of hydrogen gas (in grams) needed to generate a constant current of 12.4 A for 10.0 h?
(b) What is the volume (in liters) of the hydrogen gas if the required conditions are 650°C and 1406 mmHg?
What amount of a constant current (in amperes) is needed to generate silver at a rate of 20.0 kg/h using the smelting process?
Which of the following best describes cathodic protection?
Which of the following best explains how galvanizing prevents or minimizes rusting of steel?
3.25 g of Li is consumed at the battery's anode during the discharge of a Li-ion battery with an overall reaction of Li(s) + CoO2(s) → LiCoO2(s). What amount of electrical charge in coulombs is transferred from Li to CoO2?
When a current is passed through an HCl solution, bubbles come out of the solution as H+ becomes H2 gas. Determine the pH of a 4.18 L of 0.378 M HCl solution after passing through 27.5 A for 39 minutes.
A solution of Pb2+ is produced from dissolving 0.617 g of impure lead sample in strong acid. This solution needed 0.0263 L of 0.0568 M solution of NO3- to reach the equivalence point where NO3- is reduced to NO gas. What is the mass percent of lead in the sample if there is no other reducing agent?
Would coating iron with cadmium (Cd) metal prevent the corrosion of iron?
Would coating iron with tin (Sn) metal prevent the corrosion of iron?
At the cathode of an electrolytic cell, cadmium can be electroplated with the following half-reaction:
Cd2+(aq) + 2 e- → Cd(s)
Calculate the mass of cadmium plated onto the cathode when a 3.2 A current goes through the cell for 56 min.
Calculate the amount of time to redissolve 231 g of Ag at a charging current of 8.93 A in a rechargeable battery based on a concentration cell constructed of two Ag/Ag+ half-cells.
Calculate the amount of time needed to run a current of 7.2 A for 436 mg of Cu to be electroplated
Cu2+(aq) + 2 e– → Cu(s)
Identify the solid formed and the amount in grams when 15 A of current is applied for 10 minutes to an electrolytic cell that uses the two half-reactions below:
Mg 2+(aq) + 2e− → Mg(s) E 0 = − 2.38
Li +(aq) + e− → Li(s) E 0 = − 3.04
Using a current of 8.20 A, an aqueous solution containing Ga3+ was subjected to electrolysis. Calculate the mass of Ga(s) plated out after 3.00 days.
Ga3+(aq) + 3 e– → Ga(s) E°red = –0.549 V
A Cu/Cu2+ concentration cell has a voltage of 0.22 V at 25 °C. The concentration of Cu2+ in one of the half-cells is 1.5x10−3 M. What is the concentration of Cu2+ in the other half-cell? (Assume the concentration in the unknown cell is the lower of the two concentrations.)