5.1 - Voltaic cells

Question 1

Define electrochemistry?

Answer 1

The study of the redox processes by which chemical energy is converted to electrical energy and vice versa.

Question 2

Define an electrochemical cell?

Answer 2

An electrochemical cell is an apparatus that uses a redox reaction to produce electrical energy or uses electrical energy to cause a chemical reaction.

Question 3

Define a voltaic cell?

Answer 3

A voltaic or galvanic cell is an electrochemical cell that produces electrical energy from a spontaneous redox reaction.

Question 4

What are the components of a voltaic cell?

Answer 4

1 - Oxidation half-cell and reduction half-cell, each contains an electrode and a solution of ions.
2- An electrode is an electrically conductive material, usually a metallic strip or graphite that conducts electrons into and out of the solution in the half cell.

• Anode is the electrode where oxidation takes place.
• Cathode is the electrode where reduction takes place

3- Salt bridge: an inverted U-shaped tube filled with an electrolytic solution such as KCl, NaCl or NaNO3. It is placed on the cell, allowing the ions to flow from one half-cell to the other. The flow of ions prevents the accumulation of positive charge in the anode’s compartment. It also closes the circuit which allows producing the electric current in the external wire

4- External wire that connects the two electrodes allowing the electric current to flow

Question 5

What is the function of the salt bridge?

Answer 5

It is placed on the cell, allowing the ions to flow from one half-cell to the other. The flow of ions prevents the accumulation of positive charge in the anode’s compartment. It also closes the circuit which allows producing the electric current in the external wire.

Question 6

What do electrons flow through in the voltaic cell?

Answer 6

Electrons flow through the electrical wire.

Question 7

In terms of atoms, ions, and electrons, explain why the mass decreased at one electrode and increased at the other?

Answer 7

• The mass of the anode decreases because the atoms are oxidized ( dissolved ) forming ions.
• The mass of the cathode increases because the ions will be captured by the electrons forming atoms.

Question 8

1- If you made a new voltaic cell with Zn and Ag electrodes, what metal would be the anode, and which would be the cathode? why?
2- In this new cell, what electrode would be oxidized and which will be reduced?

3- In this new cell, what direction would electrons flow?

4- Write the half-reaction that occurs at the anode and cathode.

Answer 8

1) Zn > Anode Ag > Cathode Because zinc is more reactive than silver, therefore, it is more capable of giving electrons.
2) Zinc will be oxidized and silver ions (Ag⁺) will be reduced.

3) Zinc electrode > Silver electrode.

4)

Oxidation half-reaction: Zn(s)⟶ Zn²⁺(aq) + 2e-

Reduction half-reaction: 1e- + Ag⁺(aq) ⟶ Ag(s)

Question 9

What are some characteristics of oxidation half-cell?

Answer 9

1- The electrode, called the anode, is made of metal with higher reactivity

2- Oxidation takes place

3- The anode becomes negatively charged

4- The anode loses mass

5- Cations build up in the compartment

6- Anions from the salt bridge move to the anode

Question 10

What are some characteristics of reduction half-cell?

Answer 10

1- The electrode, called the cathode, is made of metal with lower reactivity

2- Reduction takes place

3- The cathode becomes positively charged

4- The cathode gains mass

5- Anions build up in the compartment

6- Cations from the salt bridge move to the cathode

Question 11

Explain what happens in a voltaic cell as a spontaneous redox reaction occurs?

Answer 11

The spontaneous redox reaction starts when the oxidation and reduction half-cells are connected (using a wire to connect the two electrodes and a salt bridge to connect the solutions) In the oxidation half-cell above, metal X (which is more reactive than metal Y) undergoes oxidation. Atoms of metal X in the anode lose electrons to form positive ions and electrons. The Electrons flow through the wire from the anode to the cathode and the positive ions (cations) go into the solution. As metal X atoms change into cations, the anode started losing mass In the reduction half-cell above, metal Y (which is less reactive than metal X) undergoes reduction. Positive ions of metal Y around the cathode gain the electrons transferred through the wire and form neutral metal Y atoms. As atoms of metal Y deposit on the cathode, its mass increases As the reaction proceeds, metal X cations are produced in the anode half-cell and metal Y cations are consumed in the cathode half-cell. To maintain charge balance, inert ions from the salt bridge flow into the solutions. Increasing concentrations of metal X cations in the anode half-cell are balanced by an influx of anions from the salt bridge, while a flow of cations into the cathode half-cell compensates for the decreasing metal Y cations concentration

Question 12

How is electricity produced in a voltaic cell?

Answer 12

To produce electricity in a voltaic cell, the oxidation and reduction half-reactions of the redox reaction are conducted in separate compartments called half-cells, so that electrons transferred between the two halves of the reaction can flow through an external wire from one compartment to another

Question 13

How does an electrochemical cell work?

Answer 13

By converting chemical energy to electrical energy.

Question 14

In a zinc and copper electrochemical cell, cations travel to which half-cell? Why?

Answer 14

Cations travel to the cathode half-cell. Because the cations in the cathode half-cell are changing to Cu atoms to form more copper, therefore, we need more positively charged ions (cations) in the cathode half-cell in order for the solution to remain neutralized.

Question 15

In a zinc and copper electrochemical cell, anions travel to which half-cell? Why?

Answer 15

Anions travel to the anode half-cell. Because the zinc electrode gives electrons and Zn²⁺ ions. therefore, the number of positively charged ions (Zn²⁺) increases in the solution, so, we need anions from the salt bridge (Cl ˉ ) to keep the solution neutralized.

Question 16

Example: On the following diagram, label the anode, charge on the anode, cathode, charge on the cathode, oxidation half-cell, reduction half-cell, and the direction of electrons flow. See figure in: (Instructional guide, Page 79)

Answer 16

See answer in: (Instructional guide, Page 79)

Question 17

Define the reduction potential of an electrode? And how is it used?

Answer 17

The reduction potential of an electrode is the tendency of the electrode to gain electrons. The reduction potential of an electrode is measured against the standard hydrogen electrode.

Question 18

Define the standard hydrogen electrode (SHE cell)?

Answer 18

The standard hydrogen electrode consists of a small sheet of platinum immersed in a hydrochloric acid (HCl) solution that has a hydrogen ion concentration of 1M. Hydrogen gas at a pressure of 1 atm is bubbled in and the temperature is maintained at 25℃. The reduction potential of this standard hydrogen electrode is defined as 0.000 V.

Question 19

Describe the importance of the standard hydrogen electrode (SHE cell)?

Answer 19

The standard hydrogen electrode can act as an oxidation-half reaction or a reduction-half reaction, depending on the half-cell to which it is connected.

Question 20

What is the oxidation half-reaction of SHE cell?

Answer 20

H₂(g) ⇌ 2H⁺(aq) + 2e⁻

Question 21

What is the reduction half-reaction of SHE cell?

Answer 21

2H⁺(aq) + 2e⁻ ⇌ H₂(g)

Question 22

Define standard reduction potential (E°) of an electrode?

Answer 22

The standard reduction potential of an electrode is the potential developed by an electrode (immersed in 1 M solution of its ions at 25℃) relative to the potential of standard hydrogen electrode.

Question 23

Standard reduction potential is expressed in what?

Answer 23

Standard reduction potential is expressed in volts (V).

Question 24

How can the standard reduction potential (E°) of an electrode be determined?

Answer 24

1) Standard reduction potential of a copper electrode A copper electrode is immersed in 1 M solution of Cu(NO3)2 at 25℃ and connected to a standard hydrogen electrode. Since copper is less reactive than hydrogen (see the reactivity series up), it undergoes reduction and hydrogen undergoes oxidation. The potential of the cell, E°, is measured as + 0.34 V Oxidation half reaction: H₂(g) ⟶ 2H⁺(aq) + 2e⁻ Reduction half-reaction: 2e- + Cu²⁺(aq) ⟶ Cu(s) 2) Standard reduction potential of a zinc electrode A zinc electrode is immersed in 1 M solution of Zn(NO3)² at 25℃ and connected to a standard hydrogen electrode. Since zinc is more reactive than hydrogen (see the reactivity series up), it undergoes oxidation and hydrogen undergoes reduction. The potential of the cell, E°, is measured as − 0.76 V Oxidation half reaction: Zn(s) ⟶ Zn²⁺(aq) + 2e- Reduction half-reaction: 2H⁺(aq) + 2e⁻ ⟶ H₂(g)

Question 25

If a reduction half-reaction is reversed to change it into oxidation, the sign of E° _________.
a) Is reversed

b) Remains the same
c) Is decreased
d) Is increased

Answer 25

a) Reversed

Question 26

How can the standard cell potential of a voltaic cell be calculated, given the standard reduction potentials?

Answer 26

E° = E° cathode − E° anode

E° reduction + E° oxidation E° = The overall standard cell potential

E° cathode = the standard reduction potential of the cathode (standard reduction half-reaction)

E° anode = the standard reduction potential of the anode (standard oxidation half-reaction)

Question 27

E° of an electrode is always given as ________________.

Answer 27

Reduction half-reactions

Question 28

The electrode with higher reduction potential undergoes _________ (Reduction or Oxidation) and the electrode with lower reduction potential undergoes _________ (Reduction or Oxidation).

Answer 28

1) Reduction 2) Oxidation

Question 29

The electrode with ________ (Higher or Lower) reduction potential undergoes reduction and the electrode with ________ (Higher or Lower) reduction potential undergoes oxidation.

Answer 29

1) Higher 2) Lower

Question 30

If a half-reaction is multiplied by a coefficient to balance the redox equation, E° ____________. a) Is reversed b) Remains the same c) Is decreased d) Is increased

Answer 30

b) Remains the same

Question 31

A voltaic cell is composed of a salt bridge, a silver electrode, and a magnesium electrode.

a) Write a balanced equation for the redox reaction taking place
b) Draw the voltaic cell, identifying the anode, cathode, salt bridge, and the direction of electron flow
c) Identify the reaction taking place at the cathode and the reaction taking place at the anode
d) Write the cell notation
e) Determine the standard cell potential

Answer 31

a) Looking at the standard reduction potentials, magnesium has a lower reduction potential and so undergoes oxidation and acts as an anode. Silver, which has a higher reduction potential than magnesium, undergoes reduction and acts as a cathode. Mg²⁺ + 2e- → Mg E° = -2.372V Ag⁺ + e- → Ag E° = +0.8000 V (Reverse the magnesium reduction reaction to change it into oxidation. Multiply the reduction half-reaction of silver by 2 and add the two half-reactions to get the overall equation for the reaction taking place) Oxidation half-reaction: Mg → Mg²⁺ + 2e” Reduction half-reaction: Ag⁺ + e- → Ag The balanced equation for the reaction taking place: Mg + 2Ag⁺ → Mg²⁺ + 2Ag
b) Answer on page 84 of the instructional guide.
c) Reaction at the anode: Mg → Mg²⁺ + 2e-

Reaction at the cathode: Ag⁺ + e- → Ag

d) Mg I Mg²⁺ II Ag⁺ I Ag
e) E° = E° cathode — E° anode E° = (+0.800) — (-2.372) E° = +3.172 V

Question 32

Using the cell potential, when is a redox reaction is spontaneous and when is it nonspontaneous?

Answer 32

A redox reaction is spontaneous if the cell potential is > 0 and nonspontaneous if the cell potential is < 0.

Question 33

Calculate the cell potential to determine if the following redox reaction spontaneous or not? a) Sn(s) + Cu²⁺(aq) → Sn²⁺(aq) + Cu(s)

Answer 33

E° = E° cathode − E° anode E° = (+0.342) – (-0.138) E° = +0.480 V, spontaneous

Question 34

Calculate the cell potential to determine if the following redox reaction spontaneous or not? b) Mg(s) + Pb²⁺(aq) → Pb(s) + Mg²⁺(aq)

Answer 34

E° = E° cathode − E° anode E° = (-0.126) – (-2.372) E° = +2.246 V, spontaneous

Question 35

Calculate the cell potential to determine if the following redox reaction spontaneous or not? c) 2SO₄²ˉ(aq) + Co²⁺(aq) → Co(s) +S₂O₈²ˉ(aq)

Answer 35

E° = E° cathode − E° anode E° = (-0.28) – (+2.010) E° = -2.29 V, nonspontaneous

Question 36

The electrode, called the ________, is made of metal with lower reactivity than the electrode called the _______.

Answer 36

1) Cathode 2) Anode

Question 37

What is the general cell notation for a redox reaction occurring in a voltaic cell?

Answer 37

Question 38

Write the cell notation for the following cell:

Answer 38

Question 39

A redox reaction is ___________ if the cell potential is > 0 and ____________ if the cell potential is < 0

Answer 39

1) spontaneous
2) nonspontaneous

Question 40

Calculate the cell potential to determine if each of the following redox reactions is spontaneous as written:

a) Sn_(s)+ Cu²⁺_(aq)→ Sn²⁺_(aq)+ Cu_(s)
b) Mg_(s)+ Pb²⁺_(aq)→ Pb_(s)+ Mg²⁺_(aq)
c) 2SO₄^2-_(aq)+ Co²⁺_(aq)→ Co_(s)+ S₂O₈^2-_(aq)

Answer 40

a) E°= E° cathode −E° anode

E°= (+0.342) –(-0.138)

E°= +0.480V, spontaneous

b) E°= E° cathode −E° anode

E°= (-0.126) –(-2.372)

E°= +2.246V, spontaneous

c) E°= E° cathode −E° anode

E°= (-0.28) –(+2.010)

E°= -2.29V, nonspontaneous