Free Printable Electrochemistry: Fuel Cells and Batteries

Worksheets

Electrochemistry: Fuel Cells and Batteries

Total questions: 31

Worksheet time: 3hrs 35mins

Name

Class

Date

What is an anode?

It is the electrode where oxidation takes place.

It is the electrode where reduction takes place.

It is an aqueous solution containing an electrode.

It is the salt bridge that connects half-cells.

What is a cathode?

It is the electrode where oxidation takes place.

It is the electrode where reduction takes place.

It is an aqueous solution containing an electrode.

It is the salt bridge that connects half-cells.

Given their standard reduction potentials, which of the species is going to be oxidized?

Cu²⁺/Cu = 0.34V

Zn²⁺/Zn = -0.76V

CuSO4

ZnSO4

What occurs to the mass of copper electrode in the following reaction?

Zn/Zn²⁺ // Cu²⁺/Cu

increases

decreases

remains the same

What reaction occurs at the anode?

Ag⁺/Ag = 0.80V

Ni²⁺/Ni = -0.25V

Ag⁺ + e^- →Ag

Ag → Ag⁺ + e^-

Ni²⁺ + 2e^- → Ni

Ni → Ni²⁺ + 2e^-

What would be the theoretical cell potential of the previous electrochemical cell?

Ag⁺/Ag = 0.80V

Ni²⁺/Ni = -0.25V

1.05V

-1.05V

0.55V

-0.55V

What reaction occurs at the anode?

Ag⁺ + e^- →Ag

Ag → Ag⁺ + e^-

Ni²⁺ + 2e^- → Ni

Ni → Ni²⁺ + 2e^-

Which statement best describes how a salt bridge maintains electrical neutrality in the half-cells of an electrochemical cell?

It prevents the migration of electrons.

It prevents the reaction from occurring spontaneously.

It permits the migration of ions.

It allows for the reaction from occurring spontaneously.

When an electrochemical cell is operating, it is

approaching equilibrium

using external energy

undergoing oxidation, only

undergoing reduction, only

10.

What is one problem with fuel cells

They are inefficient

They produce harmful chemicals

It is difficult to supply them with fuel

They can't help make petrol

11.

Which one below is not a major reason to develop automotive fuel cell technology?

Efficiency

Low capacitance

Low or zero emissions

Local source production

12.

The half-equation for the reaction occurring at the (+) electrode in a hydrogen-oxygen fuel which has an alkaline electrolyte would be

H₂(g) + 2OH^-(aq)→ 2H₂O(l) + 2e^-

H₂(g) → 2H⁺(g) + 2e^-

O₂(g) + 4H⁺(aq) + 4e^- → 2H₂O(l)

O₂(g) + 2H₂O(l) + 4e^- → 4OH^- (aq)

13.

Methane gas is used as a fuel in an acidic fuel cell. The half-equation occurring at the anode will be

O₂(g) + 4H⁺(aq) + 4e^- →2H₂O(g)

CH₄(g) + 2H₂O(g) → CO₂(g) + 8H⁺ (aq) + 8e^-

CH₄(g) + 4H⁺(g)→ CO₂(g) + 4H₂O(l)

CO₂(g) + 8H⁺(aq) + 8e^- → CH₄(g) + 2H₂O(g)

14.

In a fuel cell based on the oxidation of methane, the equation for the anode half-reaction is:
CH₄(g) + 2H₂O(l) → CO₂(g) + 8H⁺(aq) + 8e^-
The corresponding equation for the half-reaction at the cathode is

2H₂O(l) + 4e^- → 4H⁺(aq) + O₂(g)

4H⁺(aq) + O₂(g) → 2H₂O(l)

2H₂O(l) → 4H⁺(aq) + O₂ (g) + 4e^-

4H⁺(aq) + O₂(g) + 4e^- → 2H₂O(l)

15.

In the image is an acidic hydrogen fuel cell. The half reaction occurring at the identified electrode is...

O₂ (g) + 4H⁺ (aq) + 4e^- → 2H₂O (l)

O₂ (g) + 2H₂O (l) + 4e^- → 4OH^- (aq)

H₂ (g) → 2H⁺ (aq) + 2e^-

H₂ (g) + 2OH^- (aq) → 2H₂O (l) + 2e^-

16.

In the image is an acidic hydrogen fuel cell. The half reaction occurring at the identified electrode is...

O₂ (g) + 4H⁺ (aq) + 4e^- → 2H₂O (l)

O₂ (g) + 2H₂O (l) + 4e^- → 4OH^- (aq)

H₂ (g) → 2H⁺ (aq) + 2e^-

H₂ (g) + 2OH^- (aq) → 2H₂O (l) + 2e^-

17.

What is the main reason why hydrogen is more difficult to store than diesel?

Hydrogen is a gas at room temperature but diesel is a liquid

Hydrogen can be liquefied by cooling it

Hydrogen can be stored under pressure

18.

What is a strength of using hydrogen-oxygen fuel cells in manned spacecraft, compared to using chemical cells?

They produce electricity

They produce water that the astronauts can drink

They produce a voltage

19.

What does the potential difference of a cell depend on?

The anode and cathode.

The type of electrolyte and fuel

Whether it's alkali or not

The type of electrode and the electrolyte.

20.

In non-rechargeable cells, why do the chemical reactions stop over time?

One of the reactants becomes used up.

The energy has depleted.

The products are oxidised.

The electrodes become reactive.

21.

Complete the equation which occurs at the negative electrode of a hydrogen fuel cell 2H₂ + 4OH⁻ →

H₂O + 4e⁻

4H₂O + e⁻

4H₂O + 4e⁻

3H₂O + 2e⁻

22.

Why can some cells not be recharged?

Because the reaction is not reversible

Because the reaction is reversible

Because the reactant is used up

Because the products are used up

23.

A galvanic cell was set up using the Cu²⁺/Cu and Fe³⁺/Fe half-cells. The standard redox potentials for these half-cells are:
Cu²⁺(aq) + 2e^- → Cu(s); E⁰ = 0.34 V
Fe³⁺(aq) + e^- → Fe²⁺(aq); E⁰ = 0.77 V
The oxidant and reductant reacting in this galvanic cell are, respectively:

Fe³⁺; Cu

Fe²⁺; Cu

Cu²⁺; Fe³⁺

Cu²⁺; Fe²⁺

24.

A galvanic cell was set up from a Cl₂/Cl^- half-cell and a standard hydrogen electrode (SHE). It was found that the voltmeter reading was 1.36 V and the chlorine half-cell was the positively-charged electrode. Which statement about this galvanic cell is INCORRECT?

Both of the electrodes should be made of platinum.

The chlorine half-cell is the anode.

As the cell operates, the pH of the solution in the SHE decreases.

The standard redox potential for the chlorine half-cell must be +1.36 V.

25.

In the ethanol fuel cell, ethanol is pumped in at one electrode and oxygen is pumped in at the other. The only products of the cell reaction are the same as those that would be produced if ethanol was burnt directly in excess air. What is the correct equation for the reaction at the anode, if an acidic electrolyte is used?

O₂(g) + 4H⁺(aq) + 4e^-→ 2H₂O(l)

O₂(g)+ 2H⁺(aq) + 2e^- → 2H₂O₂(aq)

C₂H₅OH(aq) + 3H₂O(l) → 2CO₂(g) + 12H⁺(aq) + 12e^-

C₂H₅OH(aq) + 3H₂O(l) → 2CO₂(g) + 11H⁺(aq) + 11e^-

26.

A galvanic cell was set up using the Pb⁴⁺/Pb²⁺ and Pb²⁺/Pb half-cells. The standard redox potentials for these half-cells are:
Pb⁴⁺(aq) + 2e^- → Pb²⁺(aq); E⁰ = 1.69 V
Pb²⁺(aq) + e^- → Pb(s); E⁰ = -0.13 V
What metal should be used for the electrode for the Pb⁴⁺/Pb²⁺ half-cell and what would be the potential difference across the cell?

Lead; 1.56 V

Lead; 1.82 V

Platinum; 1.56 V

Platinum; 1.82 V

27.

In a methane–oxygen fuel cell, the net cell reaction is:
CH₄(g) + O₂(g) → CO₂(g) + 2H₂O(l)
If an alkaline electrolyte is used, the equation for the reaction at the positive electrode is:

O₂(g) + 2H₂O(l) + 4e^-→ 4OH^-(aq)

O₂(g) + 4H⁺(aq) + 4e^-→ 2H₂O(l)

CH₄(g) + 2H₂O(l) →CO₂(g) + 8H⁺(aq) + 8e^-

CH₄(g) + 8OH^-(aq) →CO₂(g) + 6H₂O(l) + 8e^-

28.

When the current is flowing in a primary galvanic cell (such as the alkaline cell used in a torch):

electrons travel through the porous barrier that separates the two half-cells.

the electrode marked with a (+) sign is the anode.

a reduction reaction occurs at the negative electrode.

the anode starts to corrode.

29.

The cell reaction occurring in a particular button cell was:
Zn(s) + Ag₂O(s) + H₂O(l) → 2Ag(s)+ Zn(OH)₂(s)
This cell delivered 1.50 V. The half-equation for the silver half-cell is:
Ag₂O(s)+ H₂O(l)+ 2e^-→ 2Ag(s)+ 2OH^-(aq); E⁰ = 0.34 V
The standard redox potential for the zinc half-cell must be:

-1.16 V

+1.16 V

-1.84 V

+1.84 V

30.

What is the key difference between a galvanic cell and a fuel cell?

Galvanic cells deliver direct current; fuel cells deliver alternating current.

Galvanic cells are essentially for energy storage; fuel cells work continuously.

Galvanic cells do not use gases as reactants; fuel cells only use gases.

Galvanic cells are not rechargeable; fuel cells are constantly recharged.

31.

How do fuel cells produce energy (look at the image)

The hydrogen gives of oxygen to create energy then recombine to create water.

Oxygen enters then gives of a electron then combines with hydrogen to create water.

Hydrogen gives of an electron then recombined with the electron and oxygen to create water.

All of the above

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Electrochemistry: Fuel Cells and Batteries

Total questions: 31

Worksheet time: 3hrs 35mins

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