Electrochemical Thermodynamics and Concepts

Sensitivity of electrochemical measurements

Measurements of electrochemical processes are made by measuring electrical currents and voltages. The currents results from the flow of charged ions (positive or negative) and or electrons. Current measurement can be extremely sensitive as small as 10-15 amps (coulombs per second). Let’s get “feeling” for the sensitivity of such a measurement.

monolayer: ≈1.5 x1015 atoms/cm2 . Suppose each atom on the surface dissolves as a single charged positive ion (i.e., a cation).

corresponding to 1.6 x 10-19 C/atom of charge. Today we can easily measure at the nanoamp scale. Suppose we measure a current density of 10-6 A cm-2 for 1 s or 10-6 C. A full monolayer of metal dissolving as a +1 cation results in a charge of 1.6 x 10-19

C/atom x 1.5 x 1015 atoms/cm2 or ≈ 2.4 x 10-4 C cm-2. Then 10-6 C corresponds to only of order 0.01 ML.

M ⇒ M + + e−

+1

1 cm1 cm

A

An electrochemical reaction is a chemical reaction involving electron transfer.

An ordinary chemical reaction does not involve electron transfer

Chemical reactions: Note the mass balance

N2 +3H2 = 2NH3 synthesis of ammonia

Electrochemical Reactions: Note the mass and charge balance

2H+ +2e- = H2 Hydrogen ion reduction

Zn = Zn2+ + 2e- Zn oxidation

Since electrochemical reactions involve electron transfer we can measurethe rate of these processes by designing an appropriate electrical circuit and measuring the flow of an electric current.

Faraday�s Law relates the flow of electric current to the mass of metalreacting.

ItamnF

=

m = mass of metal reacting (gm)t = time (s, min, hours, years)I = current (Amps)a = atomic weight of metaln = number of electrons transferred in the reactionF = 96,485 C; Faraday�s constant which is equal to the

charge associated with 1 mole of electrons.

If , for example, we divide Faraday�s law by the time and area of A dissolving surface, we obtain a relation describing the rate of this process.

m iartA nF

= =

A = surface area of corroding metal (cm2)i = I/A: the current density (A/cm2)r = Rate measured as a weight per unit area per

unit time (gm/cm2-s).

Examples

Consider the case for steel or iron corrosion. Suppose we make a measurementof the current associated with the corrosion of 100 cm2 of a steel surface and find a current of 0.1A. Then dividing through by the area of the sample,

0.1 A/100 cm2 = 0.001 A/cm2 = 1 mA/cm2 current density.

m iartA nF

= =

i = 1 mA/cm2

a = 55.85 gF = 96,500 Cn = 2

2 2Fe Fe e+ -= +

27 20.001 / 55.85 2.89 10 /

2 96,500A cm g gm cm s

C-´

= ´´

We can convert this to a penetration rate (cm/s) by dividing by the density of Fe.

27 20.001 / 55.85 2.89 10 /

2 96,500A cm gr gm cm s

C-´

= = ´´

37.88 /Fe g cmr =

7 28

3

2.89 10 / 3.67 10 /7.88 /Fe

r gm cm sp cm sgm cmr

--´

= = = ´

In this notation a slash represents a phase boundary and a comma separates 2 components in the same phase.

Reference Electrodes

Faradaic and Nonfaradaic Processes

Cur

rent

Potential (V vs. NHE)

Faradaic

Faradaic

Nonfaradaic

Next consider:

Divide kJ/mole by -2F2e transfer:

H2 at 1Atm

porous barrier

{H+} =1{Zn2+} =1

Zn Pt

V

Electrochemical Cells

Pt is inert and acts as a catalyst: 2H+ + 2e- = H2

Zn is oxidized: Zn = Zn2+ +2e-

Total reaction: Zn + 2H+ = Zn2+ + H2

We measure a potential difference of - 0.762 V

2H+ + 2e- = H2 E0 = 0.000 V (defined as zero by convention)

Zn2+ +2e- = Zn E0 = - 0.762 V (NHE) sign taken to be negative

since Zn is oxidized.

H2 at 1Atm

porous barrier

{H+} =1{Cu2+} =1

Cu Pt

V

Pt is inert and acts as a catalyst: H2 = 2H+ + 2e- : hydrogen is oxidizedCu plates; Cu2+ +2e- = Cu Total reaction Cu2+ + H2 = Cu + 2H+

We measure a potential difference of 0.342 V

2H+ + 2e- = H2 E0 = 0.000 V (defined as zero by convention)

Cu2+ +2e- = Cu E0 = 0.342 V (NHE) This sign is positive sinceCu is reduced.

Electrochemical Cells

1 M ZnSO4

1 M CuSO4

+1.104 V

If we construct this �cell� we observe the following:Zn dissolves; Zn = Zn2+ +2e-; oxidation occurs at the anodeCu plates; Cu2+ +2e- =Cu reduction occurs at the cathode

If we connect a voltmeter as shown we measure a voltage of 1.104 V.

Electrochemical Cells

Daniell Cell

Zn dissolves; Zn = Zn2+ +2e-; oxidation occurs at the anodeCu plates; Cu2+ +2e- =Cu reduction occurs at the cathode

For the Cu-Zn cell we measured a voltage difference of 1.104 V.

Zn2+ +2e- = Zn E0 = - 0.762 V (NHE)

Cu2+ +2e- = Cu E0 = +0.342 V (NHE)

+1.104 V

This difference is 1.104 V

Standard Reduction Potentials: EMF Series

3+ -

+ -2 2

+ -

3+ - 2+

2

Au +3e =Au +1.498 O +4H +4e =2H 0 (pH = 0) +1.229

Ag +1e =Ag +0.798 Fe +1e =Fe +0.771 Cu + -

+ -2

2+ -

2+ -

+2e =Cu +0.3422H +2e =H +0.000

Pb +2e =Pb -0.126Sn +2e =Sn -0

2+ -

2+ -

2+ -

2+ -

.138Ni +2e =Ni -0.250Co +2e =Co -0.277Cd +2e =Cd -0.403Fe +2e =Fe

2+ -

3+ -

-0.447 Zn +2e =Zn -0.762Al +3e =Al -1.662

Noble

Active

Cathode

Anode

Volts (NHE): E0

Standard conditionsStandard states:

• For a solid; a =1: pure metal, metal oxide, etc.

• For a gas, 1 Atm pressure is taken as unit activity.

• For dilute solutes typically found in most instances of corrosion, activity is reasonably approximated by the concentration in M. The standard state is 1 M.

• Temperature is taken as 25ºC = 298 K

definition: pH = -log [H+]