lecture 9. electrical conductivity of electrolyte’s solutions prepared by phd halina falfushynska

Lecture 9. Electrical conductivity of electrolyte’s solutions

Prepared by PhD Halina Falfushynska

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• Electrolytes and Nonelectrolytes

– Electrolyte: substance that dissolved in water produces a solution that conducts electricity

• Contains ions

– Nonelectrolyte: substance that dissolved in water produces a solution that does not conduct electricity

• Does not contain ions

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• Dissociation - ionic compounds separate into constituent ions when dissolved in solution

• Ionization - formation of ions by molecular compounds when dissolved

4

• Strong and weak electrolytes– Strong Electrolyte: 100% dissociation

• All water soluble ionic compounds, strong acids and strong bases

– Weak electrolytes• Partially ionized in solution• Exist mostly as the molecular form in

solution • Weak acids and weak bases

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• Examples of weak electrolytes

– Weak acids

HC2H3O2(aq) C2H3O2 (aq) + H+ (aq)

– Weak bases

NH3 (aq) + H2O(l) NH4+ (aq) + OH (aq)

(Note: double arrows indicate a reaction that occurs in both directions - a state of dynamic equilibrium exists)

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Method to Distinguish Types of Electrolytes

•nonelectrolyte •weak electrolyte •strong electrolyte

Arrhenius theory of ionization

•(i) In aqueous solution, the molecules of an electrolyte undergo spontaneous dissociation to form positive and negative ions.•(ii) Degree of ionization (α)•= Number of dissociate d molecules / Total number of molecules of electrolyte before dissociation•(iii) At moderate concentrations, there exists an equilibrium between the ions and undissociated molecules, such as, NaOH Na+ + OH–

•This equilibrium state is called ionic equilibrium.•(iv) Each ion behaves osmotically as a molecule.

•(i) At normal dilution, value of α is nearly 1 for strong electrolytes, while it is very less than 1 for weak electrolytes.•(ii) Higher the dielectric constant of a solvent more is its ionising power. Water is the most powerful ionising solvent as its dielectric constant is highest.•(iii) α 1/Con. of solution 1/wt.of solution∝ ∝• ∝ Dilution of solution Amount of solvent∝•(iv) Degree of ionisation of an electrolyte in solution increases with rise in temperature.•(v) Presence of common ion: The degree of ionisation of an electrolyte decreases in the presence of a strong electrolyte having a common ion.

FACTORS AFFECTING DEGREE OF IONIZATION

DISSOCIATION CONSTANT

Debye–Hückel theory

•The dissolved electrolyte is completely dissociated; it is a strong electrolyte.•Ions are spherical and are not polarized by the surrounding electric field. Solvation of ions is ignored except insofar as it determines the effective sizes of the ions.•The solvent plays no role other than providing a medium of constant relative permittivity (dielectric constant).•There is no Electrostriction.•Individual ions surrounding a "central" ion can be represented by a statistically averaged cloud of continuous charge density, with a minimum distance of closest approach.

Activity coefficientа = γ·С

Ion strengthI=1/2(C1z1

2 + C2z22 +

…)

Izz 51.0lg

ELECTROCHEMICAL CELLS

• Constitute ：

•

• Two conductors (called electrodes), Electrolyte solution , Salt bridge.

• Types of electrochemical cells:• （ 1 ） Galvanic (or voltaic) cells store electrical energy.

Batteries.

• （ 2 ） Electrolytic cell requires an external source of

• electrical energy for operation.

• Cathodes and Anodes （ galvanic or electrolytic cell ）• The cathode ： reduction

• The anode ： oxidation

•What is electrode? First type conductor, Second type conductor.

Electrochemical cells• Zn ZnSO4(0.0200M) CuSO4(0.0200M) Cu

• Ecell=Eright-Eleft

•Calomel electrode

•Ag/AgCl electrode

ELECTRODE POTENTIALS

•The Standard Hydrogen Reference Electrode (SHE)

The Standard Hydrogen Electrode (abbreviated SHE) is the universal

reference for reporting relative half-cell potentials. It is a type of gas

electrode and was widely used in early studies as a reference

electrode. The SHE is also called the “ Normal Hydrogen Electrode

”

•What is Reference Electrode ( RE)?

•Figure 6-3

•To provide the reference

•standard for measuring

•potential.

• The platinum electrode (Pt) is made of a small square of platinum foil which is platinized (known as platinum black). Hydrogen gas, at a pressure of 1 atmosphere, is bubbled around the platinum electrode. The platinum black serves as a large surface area for the reaction to take place, and the stream of hydrogen keeps the solution saturated at the electrode site with respect to the gas.

•The potential of the standard hydrogen electrode is assigned a value of 0.000V at all temperatures.

•How to make SHRE?

Electrode Potentials

Standard Electrode Potential (E0)

Pt, H2 (1.00 atm) | H+(a=1.00 ） || Ag+(a=1.00) | Ag

or SHE || Ag+(a=1.00) | Ag

When the silver ion activity is 1.00, the cell potential E is

the standard electrode potential of the Ag+/Ag.

E0Ag

+/ Ag =0.799V

Electrode Potential of Silver Electrode:

Ecell=Eright-Eleft=EAg-ESHE

= EAg-0.000= EAg


•Figure 6-4

Electrochemical methods - electrodes

• Electrodes are conductors in contact with an electrolyte. It would be better to speak about half-cells, because they are “halves” of galvanic cells. We already know that certain (equilibrium) voltage arises on them.

• Electrodes of the 1st kind: exchange of ions and electrons between the solution and the electrode takes place. They can be cationic (metallic or gaseous hydrogen electrode) with equilibrium between neutral atoms and cations released into solution. Anionic electrodes are possible too. A typical electrode of the 1st kind is the copper electrode immersed in a solution of Cu2+ ions.

• Electrodes of 2nd kind consist of three parts. The metal is covered by a layer of its poorly soluble salt or hydroxide, and immersed into an electrolyte containing the same anion as the salt or hydroxide. Example: calomel electrode (Hg/Hg2Cl2/KCl) and silver chloride electrode (Ag / AgCl / KCl).

Electrodes

• Oxidoreduction electrodes are formed by a noble metal conductor (gold or platinum), immersed in a solution containing reduced as well as oxidised form of a substance.

• Ion-selective electrodes are formed by membranes permeable to given ions, and their potential depends on the activity of these ions present in solution. The most important ion selective electrode is the glass electrode, specific for H3O+ ions.

• Enzyme electrodes are a special kind of ion-selective electrodes. They contain enzyme splitting substrate the concentration of which should be determined. The reaction product must be of ionic character, to be determined by the respective ion selective electrode.

• Ion selective and enzyme electrodes are important for biosensor technologies.

Hydrogen electrode• The standard hydrogen electrode is considered as the standard electrode,

with a potential conventionally equal to zero. The potential of any other electrode is defined as the voltage of the galvanic cell formed by the electrode and the standard hydrogen electrode. It is made of platinum covered by platinum black, immersed in a solution of hydrogen ions, and saturated by gaseous hydrogen (bubbling around the electrode and absorbed by the platinum black). The potential of the hydrogen electrode depends on the activity (concentration) of hydrogen ions and equals zero at unit activity of these ions. However, this electrode is not utilised to measure pH in practice because of its difficult preparation. We can write:

where pH = -logaH+

•http://www.chemguide.co.uk/physical/redoxeqia/introduction.html

pHF

TRa

F

TRa

F

TREE HHoHHH

303.2lnln2

Calomel electrode

• The calomel electrode is together with the silver chloride electrode the most important electrode of the 2nd kind. It is used as reference electrode in the determination of potentials of other electrodes. It is made of mercury covered by the calomel layer (Hg2Cl2) and KCl solution. The potential of this electrode is given by the equilibrium concentration of Cl- anions in the electrode reaction:

• Hg2Cl2(s) + 2 e- = 2 Hg(l) + 2 Cl-

• This equilibrium is also influenced by concentration of KCl. Saturated calomel electrode is usually prepared – solution of KCl is saturated. It is easy to prepare and its potential is reproducible and very stable.

•http://www.resonancepub.com/electrochem.htm

Glass electrode• The glass electrode is an ion selective electrode used in the

determination of pH. Its main part is a silver chloride electrode (4) placed in medium of known pH, e.g. in solution of NaCl (2). This solution is separated from a solution with unknown pH by a thin glass membrane (1). It forms a concentration cell the potential of which is given by the activities (concentrations) of hydrogen ions on either side of the membrane, and is partly influenced by alkaline ions present both in the glass and measured solution. For the surface potential of the glass membrane we can write:

• E = Eo - 0,059 pH [V],• where Eo is a characteristic electrode constant. The voltage on

the glass electrode is measured by electronic voltmeters which display directly the pH values. These instruments are called pH-meters. As a reference electrode (6), the silver chloride or calomel electrode surrounded by 0.1 M HCl solution is usually used. Both electrodes often form an integral immersion body (5). (7) is a porous junction to the measured solution. Modified pH-electrodes can be used directly for pH measurement in blood, gastric juice etc. Microelectrodes can be used directly for pH measurement inside cells.

•http://commons.wikimedia.org/wiki/Image:Glass_electrode_scheme.jpg

Measurement of pH

• pH meters use electrochemical reactions.

• Ion selective probes: respond to the presence of a specific ion. pH probes are sensitive to H+.

• Specific reactions:

•Hg2Cl2(s) + 2e- 2Hg(l) + 2Cl-(aq) E°1/2 = 0.27 V

•Hg2Cl2(s) + H2(g) 2Hg(l) + 2H+(aq) + 2Cl-(aq)

•H2(g) 2H+(aq) + 2e- •E°1/2 = 0.0 V

•E°cell = 0.27 V

Measurement of pH (cont.)

•Ecell = E°cell - (0.0591)log[H+] + constant

•• Ecell is directly proportional to log [H+]

•electrode

•Walther Nernst(1864-1941) received the 1920 Nobel Prize in chemistry for his numerous contributions to the field of chemical thermodynamics. Nernst (far left) is see here with Albert Einstein, Max Planck, Robert A. Millikan, and Max von Laue in 1982.

The Nernst Equation

0 0.059E=E + log

o

Rn

0E=E + lno

R

RT

nF

•Ox ＋ ne- Red

•E0: Standard Electrode Potential

•R: ideal gas constant, 8.314J K -1 mol-1

•T: temperature, K; n: number of moles of electrons

•F: the faraday=96,485C ln: natural logarithm=2.303log

The Nernst Equation

Concentration and Ecell

• Consider the following redox reaction:

•Zn(s) + 2H+ (aq) Zn2+(aq) + H2(g) •E°cell = 0.76 V

G°= -nFE°cell < 0 •(spontaneous)

• What if [H+] = 2 M?

•Expect driving force for product formation to increase.

•Therefore G decreases, and Ecell increases

•How does Ecell dependend on concentration?

Concentration and Ecell (cont.)

• Recall, in general:G = G° + RTln(Q)

• However:

G = -nFEcell

•-nFEcell = -nFE°cell + RTln(Q)

•Ecell = E°cell - (RT/nF)ln(Q)

•Ecell = E°cell - (0.0591/n)log(Q)

•The Nernst Equation

Concentration and Ecell (cont.)• With the Nernst Eq., we can determine the effect of concentration on cell

potentials.


• Example. Calculate the cell potential for the following:

•Fe(s) + Cu2+(aq) Fe2+(aq) + Cu(s)

•Where [Cu2+] = 0.3 M and [Fe2+] = 0.1 M



•• First, need to identify the 1/2 cells

•Cu2+(aq) + 2e- Cu(s) •E°1/2 = 0.34 V

•Fe2+(aq) + 2e- Fe(s) •E°1/2 = -0.44 V

•Fe(s) Fe 2+(aq) + 2e- •E°1/2 = +0.44 V

•Fe(s) + Cu2+(aq) Fe2+(aq) + Cu(s) •E°cell = +0.78 V

Concentration and Ecell (cont.)•• Now, calculate Ecell



QFe2 Cu2

(0.1)

(0.3)0.33

•Ecell = 0.78 V - (0.0591 /2)log(0.33)

•Ecell = 0.78 V - (-0.014 V) = 0.794 V

Concentration and Ecell (cont.)•• If [Cu2+] = 0.3 M, what [Fe2+] is needed so that • Ecell = 0.76 V?



•0.76 V = 0.78 V - (0.0591/2)log(Q)

•0.02 V = (0.0591/2)log(Q)

•0.676 = log(Q)

•4.7 = Q



•4.7 = Q

QFe2 Cu2

4.7

QFe2 0.3

4.7

•[Fe2+] = 1.4 M

•Calculation redox-equilibrium constants

• For example: biological redox systems

The scheme Cytochrome c

Applications of Standard Electrode Potentials

Concentration cell

• The concentration cell is formed by two electrodes made of the same metal which are immersed in solution of respective ions of different activity (concentration) a1 and a2. Considering the Nernst equation, the standard voltage U° is equal to zero and the second term is simplified (the activities of metals are identical). Then:

UR T

F

a

a

ln 2

1

Conductometry (coulometry)

Conductometry (coulometry) is measurement of conductance or conductivity of electrolytes. Electric resistance of a conductor is given by:

where is resistivity, l – length of the conductor, and A its cross-section area. The reciprocal value of resistance is called the conductance, G = 1/R [-1 = siemens, S]. The conductivity is the reciprocal of the resistivity ( = 1/). C is the resistance constant of the conductometric vessel. The quantities l and A are difficult to measure in most cases. In practice, the resistance constant C is determined from experimentally measured resistance or conductance of an electrolyte with known conductivity.

CA

l

A

lR

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Conductometry (coulometry)We can also write:

G = /C, = G.C and C = .R

The conductivity of electrolytes depends on concentration of ions and their mobility, which is of practical importance.

The quantity of conductivity called molar conductivity (lambda) is defined:

= /c,

where c is the concentration of the electrolyte.

Conductometers (coulometers)

• Conductometers can consist of a common instrument for resistance measurement in a circuit of low-voltage alternating current with a frequency of e.g. 1kHz. The direct current cannot be used, because it causes polarization of electrodes and electrolysis of the solution. The pair of measuring electrodes is made of platinum. The instrument scale is calibrated directly in units of conductance.

• Conductometry is used to check purity of distilled water, to check for the quality of potable water, for the measurement of water content in food or soil, etc. Chemists use this method in conductometric titration (see practical exercises).

Polarography and voltametry

• Polarography and voltametry are electrochemical analytical methods, which utilise electrolytic processes on polarizable electrodes. Principle of polarography was discovered by Jaroslav Heyrovský (1890-1967) in 1922 (Nobel award for chemistry in 1959).

Polarography

• Polarography is based on the measurement of the dependence of electric current on the voltage across the mercury dropping electrode (cathode). This voltage usually does not exceed -2 V. Drops of mercury are formed in short regular intervals at the end of the immersed capillary and fall to the bottom of measuring vessel. This means that the mercury surface is renewed after each drop fall.

• On the mercury surface, cations are reduced and deposited at the characteristic so-called half-wave potentials which can be read in polarographic curves (polarograms). Reduction of individual cations manifests itself near ‘half-wave’ potentials, as increase in electric current, which is proportional to the concentration of given ions in solution.

Classical setup of polarographyhttp://www.chem.ntnu.edu.tw/changijy/secondyear/teachingcontent.files/image054.jpg

Example of a polarogram. U1, U2, U3 are so called half/wave potentials of different cations present on the solution. I is the height of the polarographic half-wave proportional to the concentration of the respective cation.

Polarography

Modifications of polarography• The sensitivity of polarography was increased by several

modifications (the detection limit ranges from tens to hundreds of nM concentrations of ions). We can measure using the hanging mercury drop electrode (not falling) so that the analysed ions are collected on the electrode surface during linearly increasing voltage.

• A modern version of polarography is the differential pulse polarography. The voltage increases linearly but small voltage pulses (e.g. 50 mV) are superimposed.

• In oscillographic polarography, alternating voltage is applied. The electrode process is then given not only by faradic currents (the exchange of electrons between the electrode and the ions) but also by capacity currents (the electrode surface behaves like a capacitor). The surface capacity depends on the way of deposition of adsorbed substances. So we can study also the substances which cause no faradic currents, such as nucleic acids and their components. This kind of polarography is sometimes called tensametry.

Potentiometry Devices

• Electrochemical devices generally denoted as potentiometry devices, are used for the determination of ion concentrations based on measurement of potential of the respective electrodes.

• The most important potentiometric measurement is the measurement of pH.

• Except of pH-metry, we can often encounter potentiometric determination of potassium, sodium or calcium ions.

• The measuring system always consists of a measuring electrode, reference electrode, and a sensitive voltmeter.

ELECTROLYTES

• Na+: most abundant electrolyte in the body • K+: essential for normal membrane excitability for

nerve impulse• Cl-: regulates osmotic pressure and assists in

regulating acid-base balance• Ca2+: usually combined with phosphorus to form

the mineral salts of bones and teeth, promotes nerve impulse and muscle contraction/relaxation

• Mg2+: plays role in carbohydrate and protein metabolism, storage and use of intracellular energy and neural transmission. Important in the functioning of the heart, nerves, and muscles

MAJOR ELECTROLYTE IMBALANCES

• Hyponatremia (sodium deficit < 130mEq/L)

• Hypernatremia (sodium excess >145mEq/L)

• Hypokalemia (potassium deficit <3.5mEq/L)

• Hyperkalemia (potassium excess >5.1mEq/L)

• Chloride imbalance (<98mEq/L or >107mEq/L)

• Magnesium imbalance (<1.5mEq/L or >2.5mEq/L)

lecture 9. electrical conductivity of electrolyte’s solutions prepared by phd halina falfushynska

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