Primary Cells (DP IB Chemistry)

Primary Cells

• We have seen previously that 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 cells

  • 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 by means of 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

• 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

Fuel cells

• A fuel cell is an electrochemical cell in which a fuel donates electrons at one electrode and oxygen gains electrons at the other electrode

• These cells are becoming more common in the automotive industry to replace petrol or diesel engines

• As the fuel enters the cell it becomes oxidised which sets up a potential difference or voltage within the cell

• Different electrolytes and fuels can be used to set up different types of fuel cells

• An important cell is the hydrogen-oxygen fuel cell which combines both elements to release energy and water

The hydrogen-oxygen fuel cell

A hydrogen-oxygen fuel cell combines both elements to release energy and water

• The fuel cell consists of

  • A reaction chamber with separate inlets for hydrogen and oxygen gas

  • An outlet for the product - water

  • An electrolyte of aqueous sodium hydroxide 

  • A semi-permeable membrane that separates the hydrogen and oxygen gases

• The half equations are

2H₂ (g) + 4OH^– (aq)  →  4H₂O (l) +  4e^–                     E^θ = -0.83 V 

O₂ (g) +  2H₂O  +  4e^– →  4OH^– (aq)                      E^θ = +0.40 V 

• The overall reaction is found by combining the two half equations and cancelling the common terms:

2H₂ (g) + 4OH^– (aq) + O₂ (g) +  2H₂O   +  4e^– →   4H₂O (l) +  4e^– + 4OH^– (aq)

2H₂ (g) + O₂ (g)  →   2H₂O (l)           E^θ = +1.23 V

Benefits of fuel cells

• Water is the only reaction product, so fuel cells present obvious environmental advantages over other types of cells

• The reaction is the same as hydrogen combusting in oxygen, but since the reaction takes place at room temperature without combustion, all the bond energy is converted into electrical energy instead of heat and light

• There are no harmful oxides of nitrogen produced, which are usually formed in high-temperature combustion reactions where air is present

• Fuel cells have been used on spacecraft, where the product can be used as drinking water for astronauts

Risks and problems of fuel cells

• Hydrogen is a highly flammable gas and the production and storage of hydrogen carries safety hazards

• Very thick walled cylinders and pipes are needed to store hydrogen which has economic impacts

• The production of hydrogen is a by-product of the crude oil industry, which means it relies on a non-renewable, finite resource

• Until a cheap way is found to make hydrogen, its widespread use in fuel cells will be limited

• Hydrogen has high energy density, that is, the amount of energy contained in 1g of the fuel is high compared to other fuels, but because it is a gas, its energy density per unit volume is low which means larger containers are needed compared to liquid fuels

Redox in Voltaic Cells

• Electrochemical cells can be either voltaic (galvanic) or electrolytic cells

  • Voltaic cells generate electricity from chemical reactions

    • This is a spontaneous reaction which drives electrons around a circuit

  • Electrolytic cells drive chemical reactions using electrical energy

    • An electric current reverses the normal directions of chemical change and this is non-spontaneous

    • Stable compounds such as sodium chloride or lead bromide can be broken down into their elements

• Oxidation takes place at the anode and reduction will occur at the cathode

  • 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