Electrochemical Cells (College Board AP® Chemistry): Study Guide

Electrochemical Cells

• Redox reactions involve simultaneous oxidation and reduction as electrons flow from the reducing agent to the oxidising agent

• Which way electrons flow depends on the reactivity of the species involved

• Redox chemistry has very important applications in electrochemical cells, which come in two types:

  • Voltaic or galvanic cells

  • Electrolytic cells 

• Oxidation takes place at the anode and reduction takes place at the cathode regardless of the type of cell

  • But depending on the type of cell, the polarity changes 

    • In voltaic cells, the anode is negative

    • In electrolytic cells, the anode is positive 

Comparing Voltaic & Electrolytic Cells Summary Table

Galvanic (Voltaic) & Electrolytic Cells

Voltaic cells

• A voltaic cell generates a potential difference known as an electromotive force or EMF

  • The EMF is also called the cell potential and is given the symbol E

• The absolute value of a cell potential cannot be determined, only the difference between one cell and another

  • This is analogous to arm-wrestling: you cannot determine the strength of an arm-wrestler unless you compare them to the other competitors

• Voltaic (or Galvanic) cells generate electricity from spontaneous redox reactions, e.g.

Zn (s) + CuSO₄ (aq) → Cu (s) + ZnSO₄ (aq)

• Instead of electrons being transferred directly from the zinc to the copper ions, a cell is built which separates the two redox processes

• Each part of the cell is called a half-cell

• If a rod of metal is dipped into a solution of its own ions, an equilibrium is set up

  • For example

Zn (s)  ⇌  Zn²⁺ (aq) + 2e^– 

Zinc metal in a solution of zinc sulfate

When a metal is dipped into a solution containing its ions, an equilibrium is established between the metal and its ions

• This is a half-cell and the strip of metal is an electrode

• The position of the equilibrium determines the potential difference between the metal strip and the solution of metal

• The Zn atoms on the rod can deposit two electrons on the rod and move into solution as Zn²⁺ ions:

                   Zn (s) ⇌ Zn²⁺(aq) + 2e^– 

• This process would result in an accumulation of negative charge on the zinc rod

• Alternatively, the Zn²⁺ ions in solution could accept two electrons from the rod and move onto the rod to become Zn atoms:

                  Zn²⁺(aq) + 2e^– ⇌ Zn (s)

• This process would result in an accumulation of positive charge on the zinc rod

• In both cases, a potential difference is set up between the rod and the solution

  • This is known as an electrode potential

• A similar electrode potential is set up if a copper rod is immersed in a solution containing copper ions (eg CuSO₄), due to the following processes:

Cu²⁺(aq) + 2e^– ⇌ Cu (s)  – reduction (rod becomes positive)

Cu (s) ⇌ Cu²⁺(aq) + 2e^– – oxidation (rod becomes negative)

• Note that a chemical reaction is not taking place – there is simply a potential difference between the rod and the solution

Creating an EMF

• If two different electrodes are connected, the potential difference between the two electrodes will cause a current to flow between them

  • Thus an electromotive force (EMF) is established and the system can generate electrical energy

• A typical electrochemical cell can be made by combining a zinc electrode in a solution of zinc sulfate with a copper electrode in a solution of copper sulfate

Electrochemical cell 

The zinc-copper voltaic cell (also known as the Daniell Cell)

• The circuit must be completed by allowing ions to flow from one solution to the other

• This is achieved using a salt bridge

  • This is often a piece of filter paper saturated with a solution of an inert electrolyte such as KNO₃ (aq)

• The EMF can be measured using a voltmeter

  • Voltmeters have a high resistance so that they do not divert much current from the main circuit

• The two half cells are said to be in series as the same current is flowing through both cells

• The combination of two electrodes in this way is known as a voltaic cell and can be used to generate electricity

Conventional Representation of Cells

• Chemists use a type of shorthand convention to represent electrochemical cells

• In this convention:

  • A solid vertical (or slanted) line shows a phase boundary, which is an interface between a solid and a solution

  • A double vertical line (sometimes shown as dashed vertical lines) represents a salt bridge

    • A salt bridge has mobile ions that complete the circuit

    • Potassium chloride and potassium nitrate are commonly used to make the salt bridge as chlorides and nitrates are usually soluble

    • This should ensure that no precipitates form which can affect the equilibrium position of the half-cells

• The substance with the highest oxidation state in each half-cell is drawn next to the salt bridge

• The cell potential difference is shown with the polarity of the right-hand electrode

• The cell convention for the zinc and copper cell would be

Zn (s) ∣ Zn²⁺ (aq) ∥ Cu²⁺ (aq) ∣ Cu (s)   E _cell = +1.10 V

• This tells us the copper half-cell is more positive than the zinc half-cell so that electrons would flow from the zinc to the copper half-cell

• The same cell can be written as:

Cu (s) ∣ Cu²⁺ (aq) ∥ Zn²⁺ (aq) ∣ Zn (s)   E _cell = -1.10 V

• The polarity of the right-hand half-cell is negative, so we can still tell that electrons flow from the zinc to the copper half-cell

Electrolytic cells

• Electrolytic cells can be constructed using a beaker or crucible as the cell depending on whether the ionic compound is in solution or molten

Electrolysis of a molten ionic compound

Positive ions move to the negative electrode and negative ions move to the positive electrode

• In electrolytic cells, the substance that the current passes through and splits up is called the electrolyte

  • The electrolyte contains positive and negative ions

What happens to these ions during electrolysis?

• Negative ions move to the anode and lose electrons - this is oxidation

• Positive ions move to the cathode and gain electrons - this is reduction

• Electrically neutral atoms or molecules are released

Electrolysis of molten lead bromide

• The reactions which take place at the electrodes can be shown by half equations

• When the positive lead ions move to the cathode, they gain electrons in a reduction reaction:

Pb²⁺(aq) + 2e^-  ⇌ Pb (s)

• Similarly, when the negative bromide ions move to the anode they lose electrons in an oxidation reaction: 

2Br^- (l) - 2e^-  ⇌ Br₂ (l) 

• Sometimes oxidation reactions are written with '+2e^-' on the right of the arrow instead of '-2e' on the left

• In this case, the alternative half equation is:

2Br^- (l) ⇌ Br₂ (l) + 2e^- 

• Since metals are always cations and non-metal are always anions, it is easy to predict the products of electrolysis of molten salts:

  • Metals will always be formed at the cathode and non-metals at the anode