Conductivity of Electrolyte Solutions
Experiment 4: Conductivity of electrolyte solutions
(Dated: October 29, 2009)
I.
INTRODUCTION
Pure water does not conduct electricity, but any solvated ionic species would contribute to conduction of electricity. An ionically conducting solution is called an electrolyte solution and the compound, which produces the ions as it dissolves, is called an electrolyte. A strong electrolyte is a compound that will completely dissociate into ions in water. Correspondingly, a weak electrolyte dissolves only partially. The conductivity of an electrolyte solution depends on concentration of the ionic species and behaves differently for strong and weak electrolytes. In this work the electric conductivity of water containing various electrolytes will be studied. The data will be extrapolated to infinitely dilute solutions and the acidity constant for a given weak electrolyte will also be determined. Additional theoretical background for electrolyte solutions can be found from Refs. [1–3].
II. THEORY
Movement of ions in water can be studied by installing a pair of electrodes into the liquid and by introducing a potential difference between the electrodes. Like metallic conducting materials, electrolyte solutions follow Ohm’s law: R= U I (1)
where R is the resistance (Ω, “ohms”), U is the potential difference (V, “Volts”), and I is the current (A, “Amperes”). Conductance G (S, Siemens or Ω−1 ) is then defined as reciprocal of the resistance: G= 1 R (2)
Conductance of a given liquid sample decreases when the distance between the electrodes increases and increases when the effective area of the electrodes increases. This is shown in the following relation: G=κ A l (3)
where κ is the conductivity (S m−1 ), A is the cross-sectional area of the electrodes (m2 ; e.g. the effective area available for conducting electrons through the liquid), and l is the distance between the electrodes (m). Molar conductivity Λm (S m2 mol−1 ) is defined as: Λm = κ c (4)
where c is the
(Dated: October 29, 2009)
I.
INTRODUCTION
Pure water does not conduct electricity, but any solvated ionic species would contribute to conduction of electricity. An ionically conducting solution is called an electrolyte solution and the compound, which produces the ions as it dissolves, is called an electrolyte. A strong electrolyte is a compound that will completely dissociate into ions in water. Correspondingly, a weak electrolyte dissolves only partially. The conductivity of an electrolyte solution depends on concentration of the ionic species and behaves differently for strong and weak electrolytes. In this work the electric conductivity of water containing various electrolytes will be studied. The data will be extrapolated to infinitely dilute solutions and the acidity constant for a given weak electrolyte will also be determined. Additional theoretical background for electrolyte solutions can be found from Refs. [1–3].
II. THEORY
Movement of ions in water can be studied by installing a pair of electrodes into the liquid and by introducing a potential difference between the electrodes. Like metallic conducting materials, electrolyte solutions follow Ohm’s law: R= U I (1)
where R is the resistance (Ω, “ohms”), U is the potential difference (V, “Volts”), and I is the current (A, “Amperes”). Conductance G (S, Siemens or Ω−1 ) is then defined as reciprocal of the resistance: G= 1 R (2)
Conductance of a given liquid sample decreases when the distance between the electrodes increases and increases when the effective area of the electrodes increases. This is shown in the following relation: G=κ A l (3)
where κ is the conductivity (S m−1 ), A is the cross-sectional area of the electrodes (m2 ; e.g. the effective area available for conducting electrons through the liquid), and l is the distance between the electrodes (m). Molar conductivity Λm (S m2 mol−1 ) is defined as: Λm = κ c (4)
where c is the
References: [1] P. W. Atkins and J. de Paula, Physical Chemistry (7th ed.) (Oxford University Press, Oxford, UK, 2002). [2] R. J. Silbey, R. A. Alberty, and M. G. Bawendi, Physical Chemistry (4th ed.) (Wiley, New York, 2004). [3] R. Chang, Physical chemistry for the chemical and biological sciences (University Science Books, Sausalito, California, 2000). 12