Chapter 5: Galvanic and Fuel Cells Flashcards
Electrochemical cell
- Device that converts chemical energy into electrical energy (or vice versa)
Galvanic cell
- Type of electrochemical cell in which chemical energy is converted into electrical energy
- Reactions that occur in them are spontaneous and exothermic
NOTE: Galvanic cells are also known as voltaic cells.
Battery
- Combination of several cells in series to obtain a higher potential difference or voltage
Reactions in each half cell
- The anode (-) is where oxidation occcurs
- The cathode (+) is where reduction occurs
TIP: AN OIL RIG CAT = anode, oxidation is loss & reduction is gain, cathode.
Salt bridge purpose
- Often made from filter paper soaked in a relatively unreactive electrolyte (e.g. KNO₃)
- Allow cells to produce electricity by allowing ions to move between the two half-cells (balances charges)
- Cations (e.g K⁺) move toward the cathode and anions (e.g. NO₃-) move toward the anode
What occurs when there is no salt bridge?
- The reaction would not proceed due to buildup of charge
- One half cell would accumulate a negative charge and the other would accumulate a positive charge
Electrode materials
- Redox pair with a metal – given metal is used
- E.g. Ag⁺(aq)/Ag(s): silver electrode
- Redox pair without a metal – platinum or graphite is used
- E.g. Fe³⁺(aq)/Fe²⁺(aq): platinum or graphite electrode
Strong vs weak reductants and oxidants
-
Strong reductants donate electrons more readily than weak reductants
- They have weak conjugate oxidising agents
-
Strong oxidants accept electrons more readily than weak oxidants
- They have weak conjugate reducing agents
Predicting cell reactions
-
Reduction – higher in the electrochemical series
- Goes forward →
-
Oxidation– lower in the electrochemical series
- Reversed ←
TIP: ‘Clockwise’
Potential difference
= higher half-cell E⁰ (oxidant) – lower half-cell E⁰ (reductant)
- Measures the tendency to push electrons into the external circuit
- It is the electromotive force between two points in a circuit
Standard electrode potential (E⁰)
- The voltage measured when a half-cell is connected to a standard hydrogen half-cell at standard conditions
- Exists when one half-cell has a greater tendency to push electrons into the external circuit than the other cell
NOTE: It is also known as electromotive force (emf) or voltage.
Limitations of predictions using standard electrode potentials
- The E⁰ values in the electrochemical series are only under standard conditions (values and order on the series vary under different conditions)
- Does not provide information about the rate of reaction
Fuel cells
- Type of galvanic cell that generate electricity from redox reactions
- Chemical energy is converted directly into electrical energy
- Reactants are not stored and must be continuously supplied
- Two types: acidic and alkaline
Fuel cell design
-
Electrodes – must be conducting and porous to allow hydrogen and oxygen to come into contact with the ions in the electrolyte
- Size determines the size of the current drawn from cell
-
Electrolyte – allows ions from anode and cathode to come into contact (does not allow transfer of electrons)
- E.g. KOH
- Catalysts – often coat electrodes to increase the reaction rate and current produced by the cell
Advantages and disadvantages of fuel cells
-
Advantages
- Efficiency of 40-60% due to direct energy conversion
- Reduced greenhouse gas emissions
- Can use a variety of fuels
- Generate electricity as long as fuel is supplied (conventional batteries need to be recharged/replaced)
-
Disadvantages
- Require a constant fuel supply
- Expensive
- Storage and safety concerns
