Electrochemical Cells (Edexcel IAL Chemistry)

Core Practical 12: Investigating Electrochemical Cells

Spec point — 16.6

Measuring the EMF of a cell

• To measure a cell EMF you will need

  • Two small beakers, around 75 cm³ capacity
  • Strips of suitable metals such as copper, zinc, iron and silver
  • 1.0 mol dm^-3 solutions of the metal ions (nitrates, chlorides or sulfates depending on their solubility)
  • A high resistance voltmeter (usually a digital multimeter has this)
  • Two sets of wires with crocodile clips
  • A salt bridge consisting of a strip of filter paper soaked in saturated potassium nitrate

The experimental set up for measuring the EMF of a cell made of two metal / metal ion half cells

The procedure

• The strips of metals need to be freshly cleaned to remove any oxide coatings

  • This can be done with a piece of sandpaper

• To support the metals, it is easiest to have long strips that can be folded over the side of the beaker and held in place with the crocodile clips

• Fill up the beakers to about two thirds of the way with the metal ion solutions

• Using tongs, dip a strip of filter paper into a beaker of saturated potassium nitrate solution and then place it between the two beakers making sure the ends of the strip are well immersed in the solutions

• Connect the crocodile clips to the voltmeter, wait for a steady reading and record the measurement

Practical tips

• If you don't get a positive reading on the voltmeter swap the terminals around
• Voltmeters will have marked positive and negative terminals (usually in red and black, respectively), so when you get a positive reading this tells you the relative polarity of the metals in the cell
• Change the salt bridge each time, to prevent cross contamination of ions between half cells

Calculating the Theoretical EMF of a Cell

• The voltage you measure in your experiment can be compared to a theoretical value calculated using standard electrode potentials (E°)

  • These values are found in the electrochemical series in your data book.

• The EMF of a cell is the difference between the E° values of the two half-cells

• The formula is:

E°(cell) = E°(positive electrode) - E°(negative electrode)

• A simple way to remember this is "most positive E° value minus most negative E° value

Worked Example Calculate the theoretical EMF for a cell made of zinc and iron.

1. Find the E° values from a data book:

• Fe²⁺(aq) + 2e⁻ ⇌ Fe(s) has an E° = -0.44 V
• Zn²⁺(aq) + 2e⁻ ⇌ Zn(s) has an E° = -0.76 V

1. Identify the positive and negative electrodes:

• The half-cell with the more positive (less negative) E° value will be the positive electrode

  • In this case, that is the Fe²⁺/Fe half-cell.

• The half-cell with the more negative E° value will be the negative electrode

  • This is the Zn²⁺/Zn half-cell.

1. Apply the formula:

• E°(cell) = E°(positive) - E°(negative)
• E°(cell) = (-0.44) - (-0.76)
• E°(cell) = +0.32 V

This calculation is how the theoretical values in the specimen results table are determined.

Specimen Results

• Here is a set of typical results for this experiment:

Analysis

• It is unlikely you will get very close to the theoretical results as these would be obtained under standard conditions(1.0 mol dm^-3 concentration, 298 K temperature)

  • This is because these conditions are hard to achieve in a school laboratory

• However, the relative EMF of cells you construct should match the theoretical values

• The higher the EMF, the larger the difference in reactivity ('electron pushing power') between the metals