Water & Solutions Forum – Water Science, Processing and Maintenance



Basics of oxidation and reduction


 (1) pHと電位
 (2) 標準酸化還元電位
 (3) 酸化還元電位のまとめ
Table of contents

1. Electron transfer between substances
2. Free energy change of the reaction
3. What is redox potential?
 (1) pH and potential
 (2) Standard redox potential
 (3) Summary of redox potential
4. Specific examples of redox potential
 (1) Redox potential of metal
 (2) Redox potential of each substance
5. Rate of oxidation-reduction reaction
6. Example of water quality measurement
7. Example of water purification
8. Function and name of electrode



4Fe2+ + O2 + 4H+ = 4Fe3+ + 2H2O  (1)

4Fe2+ = 4Fe3+ + 4e  (1a)
O2 + 4H+ + 4e = 2H2O  (1b)


1. Transfer of electron between substances

 Substances, such as simple molecules, ions and complex compounds, maintain a stable state by exchanging electrons that form the atoms in the substance with each other. In the electron transfer reaction, there is always a substance that emits Red (electron donor or reducing agent) and a substance that receives Ox (electron acceptor or oxidizing agent): Red-reductant, Ox-oxidant; Redox reaction-Red and Ox reaction.
 Apparently, the electron transfer reaction between substances cannot be observed, but the electron transfer reaction can be understood by decomposing the reaction process involving the redox reaction.
 For example, a reaction in which Fe2+ is oxidized to Fe3+ by dissolved oxygen O2 in an acidic aqueous solution is represented by the following formula (1).

4Fe2+ + O2 + 4H+ = 4Fe3+ + 2H2O  (1)

 The above reaction process is decomposed into the following equation.
4Fe2+ = 4Fe3+ + 4e  (1a)
O2 + 4H+ + 4e = 2H2O  (1b)

 As shown in the formula (1a), Fe2+ emits an electron e and changes to Fe3+. Meanwhile, O2 receives 4e and combines with 4H+, resulting in that two molecules of H2O is produced.
 Incidentally, the reaction rate of the formula (1) in acidic is very slow, the reaction rate will increase exponentially with increase in pH, but the reaction systems of Fe2+ and Fe3+ are extremely complicated because of formations of hydroxide compounds in neutral and alkaline water. Therefore, in order to simplify the explanation, we will deal with acidic reaction systems.

2.反応の自由エネルギー変化 ΔG


GX = μX   μX = μ°X + RT ln aX  (2)

ΔG = Gright – Gleft  (3)
= 4GFe3+ + 2GH2O – 4GFe2+ – GO2 – 4GH+  (3a)
= ΔG’°+ RTln K’
= ΔG°+ RTln K  (3b)
ΔG’° = 4μ°Fe3+ + 2μ°H2O – 4μ°Fe2+ – μ°O2– 4μ°H+
K’ = (aFe3+)4(aH2O)2/((aFe2+)4(aO2)(aH+)4)
ΔG° = ΔG’° – 2RTln aH2O
K = (aFe3+)4/((aFe2+)4(PO2)(aH+)4)  (3c)
pK = 4paFe3+ – 4paFe2+pPO2 – 4paH+
pK = 4pFe3+ – 4pFe2+pPO2 – 4pH  (3d)
(K:平衡定数、pX = –log[X])
(希薄溶液:aX = fX[X] = [X]、fX = 1)

(a) Gleft > Gright すなわち 4pFe2+ + pPO2 + 4pH < pFe3+ のとき、右方向へ反応が進行する(例えば、Fe2+濃度・酸素分圧PO2の増加やpHの低下)。
(b) Gleft < Gright すなわち 4pFe2+ + pPO2 + 4pH > pFe3+ のとき、左方向へ反応が進行する(例えば、無酸素状態やpHの上昇)。
(b)において、他の酸化剤を投入するか、白金などの不活性な一対の電極を浸して電圧を加えるなどでは、Gleftに外部エネルギーGoutが加わり、Gleft + Gout > Grightとなれば、右方向へ反応が進行することとなる。

2. Change in free energy of reaction ΔG

 Free energy GX of substance X is shown by the following equation.

GX = μX   μX = μ°X + RT ln aX  (2)
X : chemical potential, aX: activity)

 The free energy change ΔG of reaction(1) is
ΔG = Gright – Gleft  (3)
= 4GFe3+ + 2GH2O – 4GFe2+ – GO2 – 4GH+  (3a)
= ΔG’°+ RTln K’
= ΔG°+ RTln K  (3b)
(Substituting equation (2) into equation (3a))
ΔG’° = 4μ°Fe3+ + 2μ°H2O – 4μ°Fe2+ – μ°O2– 4μ°H+
K’ = (aFe3+)4(aH2O)2/((aFe2+)4(aO2)(aH+)4)
(Water activity aH2O is constant.)
ΔG° = ΔG’° – 2RTln aH2O
K = (aFe3+)4/((aFe2+)4(PO2)(aH+)4)  (3c)
pK = 4paFe3+ – 4paFe2+pPO2 – 4paH+
pK = 4pFe3+ – 4pFe2+pPO2 – 4pH  (3d)
(K: equilibrium constant, pX = –log[X])
(Dilute solution: aX = fX[X] = [X]、fX = 1)
(PX: partial pressure of gas (atm))

 Now, since the chemical reaction progresses in the direction of decreasing ΔG, the reaction(1) will proceed as follows from the relationship of equations(3 to 3d) for the free energy change ΔG.
(a) When Gleft > Gright, that is, 4pFe2+ + pPO2 + 4pH < pFe3+, the reaction proceeds to the right (e.g., Fe 2+ concentration and/or the oxygen partial pressure PO2 increase, and/or the pH reduces).
(b) When Gleft < Gright, taht is, 4pFe2+ + pPO2 + 4pH > pFe3+, the reaction proceeds to the left (for example, anoxic state or increase of pH).
 The reaction will proceed to the right in the cace of (b) when another oxidant is added or a pair of inert electrodes such as platinum is dipped to apply a voltage so that the external energy Gout is added to Glight to be Gleft + Gout > Gright.
 The potential-pH equilibrium diagram, which is important for the design and control of the specific electrode reaction process for water purification, is explained on another page.



aA + mH+ + ze = bB + cH2O  (4)

ΔG = -zFE → E = -ΔG/(zF)  (5)

 式(5)は、上記(2)・(3a~d)に示した各関係式と同様に扱い、E°= -ΔG°/(zF)とすると、次式で示される。(正確には、ΔG → ΔG° – c(μ°H2O + RTln aH2O) である。)
E = E°- RT/(zF) ln K  (6)
K = (aB)b/((aA)a (aH)m))

 温度25℃(T=298K)において、定数R(=8.314 J/K/mol)、F(=96,500 C/mol)、ln = 2.303logを用い、式(6)を書き直す。
E = E°- 0.0591(m/z)pH + 0.0591(1/z)(a log aA – b log aB)
= E°- 0.0591(m/z)pH – 0.0591(1/z)(apA– bpB)  (7)
(Eの単位:V(volt) = J/C)

3. What is redox potential?

 It is shown that the above reaction(1) process can be understood by electron transfer between substances when decomposed into ion reactions (1a) and (1b) including electrons. This decomposed reaction is also called a half reaction .
 In an aqueous solution, the half reaction involving substance A, electron e and hydrogen ion H+ can be generalized by the following equation.

aA + mH+ + ze = bB + cH2O  (4)

 In reaction (4), the change of free energy is
ΔG = -zFE → E = -ΔG/(zF)  (5)

 Expression (5) is treated in the same manner as the relational expressions shown in (2) and (3a to 3d) above, and if E° = −ΔG°/(zF), the following expression is given.
(To be exact, ΔG → ΔG° – c(μ°H2O + RTln) aH2O).)
E = E°- RT/(zF) ln K  (6)
K = (aB)b/((aA)a (aH)m))

Rewrite equation (6) using constants R (= 8.314 J/K/mol), F (= 96,500 C/mol), and ln = 2.303 log at a temperature of 25°C (T = 298K ).
E = E°- 0.0591(m/z)pH + 0.0591(1/z)(a log aA – b log aB)
= E°- 0.0591(m/z)pH – 0.0591(1/z)(apA– bpB)  (7)
(Unit of E:V(volt) = J/C)



 このケースでは、酸化還元電位Eは、 r = apA – bpB = p(Aa/Bb) によって変化する。pHは関与しない。


 このケースでは、酸化還元電位Eは、前記 r が一定であれば、pHが±1ほど‘増減’すれば、0.059(m/z)[V]ほど‘減増’することとなる。

(1) pH and potential

(a) m = 0

 In this case, the redox potential E changes with r = apA – bpB = p(Aa/Bb). pH is not involved.

(b) m ≠ 0

 In this case, the oxidation-reduction potential E is “increased” by 0.059 (m/z) [V] when the pH is “decreased” by ±1, if the r above is constant.

(2)標準酸化還元電位 E°

 一般に数表に表示されている酸化還元電位は、反応式(4)において、関与する全ての物質Xの活量aX = 1 としたときの値でpX = 0となり、式(7)はE = E°となる。
 なお、酸化還元電位の絶対値を求めることは不可能で、基準となる半反応の電位を‘0’と定め、これに対して目的とする半反応の電位との電位差を測定し、その値を目的とする半反応の酸化還元電位としている。一般的には、標準水素電極SHE(standard hydrogen electrode)の電位を基準としている。

2H+ + 2e = H2  (8)
pH = 1.0、PH2 = 1.0atm)

(2) Standard redox potential E°

 Generally, the redox potential shown in the table is pX = 0 when the activity aX = 1 of all involved substances X in the reaction(4), and the formula (7) is E = E°. The potential when the temperature of 25°C and the activity of the substances involved is set to ‘1’ is called the standard redox potential. In addition, it is impossible to obtain the absolute value of the redox potential, and the potential of the reference half-reaction is set to ‘0’, and the potential difference from the potential of the desired half-reaction is measured, and the difference value is determined and defined as the redox potential of the desired half reaction. Generally, the potential of a standard hydrogen electrode SHE is used as a reference.

2H+ + 2e = H2  (8)
pH = 1.0、PH2 = 1.0atm)


(a) 酸化還元電位の基準
(b) 物質Xの活量
aX(s) = 1 → log aX(s)= 0; aH2O = 1 → log aH2O = 0
物質Xの活量aXは、水溶液中の濃度[X]を用いて表すと、aX = fX[X]の関係があり、fX活量係数(0<fX≦1)といい、無限希釈([X] → 0)のとき、fX → 1となる。したがって、希薄溶液では、aX = [X]として扱ってもよいこととなる。
(c) 標準酸化還元電位E°
 半反応(4)の標準酸化還元電位は、温度25℃、反応に関わる全物質の活量aXi = 1のときの、式(6)で示す酸化還元電位Eである。
 電解質を加えない純水中で物質Xを無限に希釈して、電位Eを測定すれば、aX = [X]となり、ΔGから求めた式(5)と実験的に式(6)から求めた電位Eは一致することとなろうが、このような実験系での電位の測定は不可能である。

(3) Summary of redox potential

(a) Redox potential reference
 ① The potential SHE of the standard hydrogen electrode shown in reaction (8) is set to ‘0’.
 ② The potential E of the half reaction (4) is indicated by the potential difference [V] with respect to SHE.
(b) Activity of substance X
 ① Each substance X involved in the redox reaction is shown by activity aX.
 ② The activity of solid X(s) and the activity of water H2O shall be ‘1’.
aX(s) = 1 → log aX(s) = 0; aH2O = 1 → log aH2O = 0
 ③ When the activity aX of the substance X is expressed using the concentration [X] in the aqueous solution, there is a relationship of aX = fX[X], and fX is called an activity coefficient (0<fX≦1), which is infinite dilution (when [X] → 0, fX → 1). Therefore, in dilute solution, it can be treated as aX = [X].
(c) the standard redox potential E°
 The standard redox potential of the half reaction (4) is the redox potential E shown in equation (6), when the temperature is 25°C and the activity of all substances involved in the reaction is aXi = 1.
 In general, the standard redox potential E° is obtained from the free energy change ΔG shown in equation(5). Actually, when the potential E of the reaction(4) is measured with respect to SHE using the electrochemical method, it does not match the potential E calculated from the equation(6). Since an electrolyte is added in the electrochemical measurement, the activity of the substance X involved in the reaction is different due to the interaction with the electrolyte.
 If the substance X is diluted infinitely in pure water without adding any electrolyte and the potential E is measured, aX = [X], the E-value obtained theoretically from the equation(5) using ΔG and the E-value experimentally from the equation(6) would be the same, but it is impossible to measure the potentials in such an experimental system abve.




① Redox電位の低い金属(電子を放出する傾向が極めて高い)
② Redox電位が中程度の金属(電子を放出する傾向が中程度)
 Mg~Pbは常温では空気中の酸素と反応しにくいが(Mgは比較的反応しやすい。)、放置すると表面に酸化皮膜をつくる。高温では空気中の酸素とよく反応し、MgやAlの粉末を空気中で熱すると、強い光を出して燃え、酸化物 MgOやAl2O3となる。
③ Redox電位の高い金属(電子を放出する傾向が低い)

4. Specific example of redox potential

(1) Metal

(a) Redox potential and general properties of metals

Table 1 shows standard redox potentials (standard electrode potentials) of typical hydrated metal ions. A metal having a low redox potential has a high outer shell electron energy, and has a strong tendency to emit electrons to be ionized (reducing power).
① Metals with low redox potential (very high tendency to emit electrons)
 Metal of Li to Na react rapidly with oxygen in the air even at room temperature to form oxides. These metals react with water at room temperature to generate H2 gas to become ions and dissolve in water.
② Metals with a medium redox potential (moderate tendency to emit electrons)
 Metals of Mg to Pb do not easily react with oxygen in the air at room temperature (Mg reacts relatively easily), but if left in the air, oxide films are formed on these surfaces. At high temperatures, these react well with oxygen in the air, and when Mg and Al powders are heated in the air, they emit intense light and burn to form oxides MgO and Al2O3.
 Mg reacts with hot water to form hydroxide Mg(OH)2. Zn and Fe react with high temperature steam to generate H2 gas, and become ZnO and Fe3O4 oxides.
 Ni, Sn, and Pb dissolve in dilute acid to generate H2, but the reaction with oxygen in the air is unlikely to occur. Metals with a higher Redox potential than Ni hardly react with water. Pb is practically insoluble in HCl and H2SO4 because of the formations of PbCl2 and PbSO4 which are insoluble in water.
③ Metals with high Redox potential (low tendency to emit electrons)
 Metals of Cu to Au, which have higher redox potentials than hydrogen H2, do not easily react with oxygen in the air and do not react with dilute acid. Cu and Hg are oxidized to oxides when heated in the air. Metals of Cu to Ag react with oxidizing acids (dilute/concentrated nitric acid and hot concentrated sulfuric acid) to respectively generate NOx and SO2 gases, to ionize and dissolve in acids.
 Pt and Au are ionized and dissolved only in aqua regia (concentrated hydrochloric acid(1) + concentrated nitric acid(3)). Pt and Au are stable even in the air.

表1 金属の酸化還元電位と酸素・水との反応
Table 1 Redox potentials of metals and their reactions with oxygen and water.



Zn2+ + 2e ← Zn E = -0.76 + RT/(2F)ln[Zn2+]  (9)
Cu2+ + 2e → Cu E = +0.34 + RT/(2F)ln[Cu2+]  (10)

 外殻電子のエネルギーは、e(in Zn) > e(in Cu)であるので、電子はZn→Cuへ移動して安定化する(図2に示すように、電子は高いエネルギー状態から低いエネルギー状態へ移動する。)。すなわち、式(9)で左方向へ、式(10)で右方向へ反応が進行し、理論的にはlog[Zn2+]/[Cu2+]=37(式(9)と式(10)のEが等しくなるとき)に達したときに反応が停止する。
図1(B)は、Zn|Zn2+‖Cu2+|Cu 、と簡略表記される。式(9)と(10)に示すように、イオン濃度により電位が変化するので、条件によっては電子の移動方向(電流は逆方向)は異なるが、ECu > EZn の場合には、(-)Zn|Zn2+‖Cu2+|Cu(+)、と表記することもある。’|’は電子移動反応が発生する界面、’‖’は隔膜や塩橋(電解質イオンは透過するが、溶液は混合しない)を示す。Pt(不活性で電子の授受のみに関与)を用いた水素電極SHEとCu電極を塩橋で接続した電池は、Cu|Cu2+‖H+|Pt,H2(1atm)、のように表記される。

(b) Reaction of metal with other metal ions

 As shown in Fig.1(A), when a zinc plate Zn is immersed in an aqueous solution of Cu(II) sulfate CuSO4, Zn dissolves as Zn2+ and Cu deposits on the Zn plate to form a film. Further, as shown in Fig.1(B), both chambers separated by a porous partition wall (unbaked plate) are filled with respective aqueous solutions of copper(II) sulfate and zinc(II) sulfate (which are acidified with dilute sulfuric acid), When a Cu plate and a Zn plate are dipped in each, and both plates are connected by a lead wire, an electric current flows, Cu is deposited on the Cu plate, and Zn2+ is eluted from the Zn plate (Daniel battery).
 These electron transfer reactions are explained as follows. For simplicity, the activity coefficient of each ion is set to ‘1’.

Zn2+ + 2e ← Zn E = -0.76 + RT/(2F)ln[Zn2+]  (9)
Cu2+ + 2e → Cu E = +0.34 + RT/(2F)ln[Cu2+]  (10)

 Since the energy of the outer shell electron is e–(in Zn) > e–(in Cu), the electron moves from Zn to Cu and stabilizes (As shown in Fig.2, eletron moves to a lower energy state.). That is, the reaction proceeds to the left in Eq.(9) and to the right in Eq.(10), and the reaction stops theoretically when log[Zn2+]/[Cu2+]=37 (E-values in Eqs.(9) and (10) are equal).
 The energy released by this reaction is converted into the temperature rise of the solution in Fig.1(A), and is converted into the heat generation energy of the external resistor in Fig.1(B). Strictly speaking in Fig.1(B), some energy is also consumed in moving ions in the solution with a slight increase in the temperature of the solution.
Battery notation
Figure 1(B) is abbreviated as Zn|Zn2+‖Cu2+|Cu. As shown in Eqs. (9) and (10), the potential changes depending on the ion concentration, so the electron transfer direction (current is opposite direction) differs depending on the conditions, but in the case of ECu > EZn, it may also be expressed as (-)Zn|Zn2+|Cu2+|Cu(+). “|” indicates an interface where an electron transfer reaction occurs, and “‖” indicates a diaphragm or a salt bridge (electrolyte ions permeate but a solution does not mix). A battery in which the hydrogen electrode SHE, using Pt (inactive and involved only in the transfer of electrons), and the Cu electrode are connected by a salt bridge is described as Cu|Cu2+‖H+|Pt,H2(1atm).
 What are described on this page are equilibrium reactions (time unit is not included), and the velocities of electron transfers are not mentioned. The magnitude of electric current (in the amount of electricity flowing per unit time or the amount of change in substance) depends on various factors such as the surface area of ​​the electrode metal, the concentrations of reactive ions and their migration rates, and the activation energies at the water-metal interface are involved, so they are described on another page.

図1 酸化還元電位の異なる金属での電子移動反応
Fig.1 Electron transfer reactions on metals with different redox potentials.

図2 酸化還元系における電子移動反応
Fig.2 Electron transfer reaction in redox system.



(2) Redox potential of various substances

 Standard redox potentials of various substances are shown in Tables 2 to 7. Here, Typical examples of redox reactions related to water purification will be briefly described. Note that some of the battery/electrolysis system notations listed below are not official, but are used for simplicity and simplification.

表2 難溶性塩の標準酸化還元電位
Table 2 Standard redox potentials of sparingly soluble salts.

表3 金属イオンの標準酸化還元電位
Table 3 Standard redox potentials of metal ions.

表4 金属錯イオンの標準酸化還元電位
Table 4 Standard redox potentials of metal complex ions.

表5 イオウおよびハロゲンの標準酸化還元電位
Table 5 Standard redox potentials of sulfur and halogen.

表6 無機物質の標準酸化還元電位
Table 6 Standard redox potentials of inorganic substances.
表7 有機化合物の標準酸化還元電位
Table 7 Standard redox potentials of organic compounds.



5. Kinetics of redox reactions

 These are described on another page.


(a) pHの測定

 ガラス膜センサーの両面(検水側 pHx と標準液側 pHs = 7)に発生する電位差ΔEを参照電極を用いて測定し、検水のpHxを測定する。電位差は、高入力インピーダンス電圧計に、つぎに示すAg(+)極およびAg(-)極をリード線で接続して測定する。
 検水(t=25℃)のpHx = 7のときΔE = 0となるように電圧計を調整すると、pHx = ・・、6、7、8、・・のとき、ΔE[mV] = ・・、59.1、0.00、-59.1、・・となる。

6. Examples of water quality measurements

(a) pH measurement

 Measure the pHx of the test water by measuring the potential difference ΔE generated on both sides (pHx on the test water side and pHs = 7 on the standard solution side) of the glass membrane sensor using the reference electrode. The potential difference is measured by connecting the following Ag(+) and Ag(-) electrodes with lead wires to a high input impedance voltmeter.
 If the voltmeter is adjusted so that ΔE = 0 at pHx = 7 of the test water (t=25°C), ΔE[mV] = …, 59.1, 0.00, -59.1,…, when pHx = …, 6, 7, 8, …
 As the reference electrode, a silver-silver chloride electrode (Ag|AgCl/Cl–) (Table 2-19) is widely used because it has excellent potential reproducibility and accuracy and is easy to handle. The potential generation mechanism of the glass membrane is omitted.

Ag(+)|AgCl/Cl (reference electrode) ‖Test water (pHx)|Glass membrane|Standard solution (pHs)‖Cl/AgCl|Ag(-) (reference electrode)
ΔE = -RT/(F)(pHx – pHs)  (9)

(b) DO(溶存酸素)の測定


(b) DO (dissolved oxygen) measurement

 DO measurement by diaphragm electrode method includes polarographic method (electrolysis) and galvanic cell method. The polarographic method requires an external power source to apply a voltage to the electrodes, whereas the galvanic battery method does not require an applied voltage because the measurement system itself forms a battery.
 In both cases, the measurement electrode (working electrode) (inert metal such as Pt, Au, or Ag) and reference electrode (counter electrode) (Ag|AgCl/Cl, Ag|Ag2O/OH) immersed in electrolyte (KCl), Pb|PbO/OH, etc.), which is partitioned by a membrane that selectively permeates oxygen.
 Connect the Pt(+) and Ag(-) electrodes to the ammeter with lead wires. When oxygen in the test water permeates the diaphragm and enters the above-mentioned electrolytic system, the next electrolytic reaction occurs, electrons move from the Ag electrode to the Pt electrode, and a current proportional to the oxygen concentration (the opposite direction of the electrons) flows and convert this value to a DO value. The relationship between DO and voltage (which is converted from current) is calibrated with an oxygen-free and saturated DO aqueous solution.


Pt(+): O2 + 4H+ + 4e- → 2H2O  (e-の流入とO2への供与)
Ag(-): 4Ag + 4Cl → 4AgCl + 4e-  (e-の放出)

①電気分解: 正極=アノード、負極=カソード、②電池: 正極=カソード、負極=アノード。
Polarity of electrode
  In electrochemistry (electrolysis, battery) and electronic circuits, the electrode where the current i/electron e inflows/outflows from/to the outside is called anode, and the electrode where the current i/electron e outflows/inflowsis called a cathode. On the other hand, when defining the polarity of the electrode by the potential, the one with the higher potential is called the positive electrode (+) and the one with the lower potential is called the negative electrode (-). The current flows from high to low, and the electrons move from low to high.
 Note that the directions of electrons (or current) flow in the battery system (discharging) and the electrolysis system (charging) are opposite.
 ① Electrolysis: positive electrode = anode, negative electrode = cathode, ② battery: positive electrode = cathode, negative electrode = anode.
  “Function and name of electrode” is described in detail at the bottom of this page.

(c) COD(化学的酸素要求量)の測定


(c) Measurement of COD (chemical oxygen demand)

 For COD, potassium permanganate K2MnO4 (Table 6-9) or potassium dichromate K2Cr2O7 (Table 6-8) is added to the test water as an oxidizing agent under certain conditions (addition of acid, heating for a predetermined time). After reacting the organic substance with the oxidant added and quantifying the remaining oxidant after the reaction, the amount of consumed oxidant is calculated, and this consumed amount is converted to the amount of oxygen [mg-O2/L].
 A reducing agent such as potassium oxalate K2C2O4 (Table 7-9) is used to quantify the remaining oxidizing agent.


(a) 電解鉄とリン除去


7. Examples of Water purification processes

(a) Iron electrodes and phosphorus removal

 The iron electrodes method is widely used for small septic tank of phosphorus removal type. When a pair of iron plates are connected to a power source, electrolyzed Fe2+ is immediately oxygen-oxidized and converted to Fe3+ which reacts with phosphate ion to form the insoluble salt Fe2PO4.

(b) オゾンと上水・下水


(b) Ozone and tap-water/sewage

Ozone O3 has a very strong oxidizing power (Table 6-13) and is easily soluble in water, so it oxidizes substances in water and is widely used for decolorization, deodorization and sterilization. It is a safe method as it does not generate trihalomethanes and spontaneously decomposes in water.

(c) 生物による酸化還元反応


(c) Redox reactions by living organisms

 As shown in Table 7, the redox potencials of the general organic compounds are not within the range where theoretical electrolysis of water (hydrogen generation: 0.00 to oxygen generation: 1.23V), herefore, the redox ractions in their aqueous soution do not proceed under normal conditions. As mentioned in the above COD measurement, the extremly strong oxidizing organic substances such as potassium dichromate K2CrO7 are able almost to completetly oxidize and decopose the organic substances to CO2 only in boiling sulfuric acid aqueous solution with reflux (105°C) for an hour. However, even under the same conditions, potassium permanganate KMnO4 cannot completely oxidatively decompose organic substances to CO2.
 However, as seen in oxidative decomposition reactions by aerobic organisms and methane fermentation by anaerobic organisms (symbiotic reactions by complex bacterial groups), redox reactions of organic substances proceed at room temperature. In the photosynthetic reaction by plants, carbon dioxide is reduced at room temperature to synthesize organic substances. In addition, there are bacteria that use hydrogen sulfide (Table 6-2) as an energy source at hot springs and deep-sea hot water outlets, and at the deep-sea hot water outlets, there are special organisms and bacteria that prey on bacteria. Inhabited by organisms that attach to/hold organs and ingest organic substances. These reactions are composed of complicated biochemical reactions, the mechanism is extremely complicated, and life is amazing.




8. Function and name of electrode

Definition of anode/cathode

Electrode of electric circuit
 An electrode is a passive element, an active element such as a vacuum tube or a semiconductor element, an electrolysis device, a battery, or the like, which is a portion electrically connected to the object for the purpose of working or measuring an electric signal.
 In electrochemistry (electrolysis) and diodes (vacuum tubes, semiconductor devices), the electrode where the current (i) flows in from the external circuit and electrons (e-) flow out to the external circuit is called the anode.
 On the contrary, the electrode through which current flows out of the external circuit and electrons flow in from the external circuit is called the cathode.
(The terms anode and cathode are named Faraday, and are derived from the Greek words’anodos’ meaning’upstream’ and’cathodos’ meaning ‘downstream’.)
Electrolysis/battery-discharge electrodes
 An oxidation reaction occurs when the electrode that receives electrons (e-) from the active substance in the electrolytic solution (including solid) and releases the electrons to the external circuit is the anode. Electrons flow from an external circuit, and an electrode that donates electrons to the active substance in the electrolyte is the cathode, and a reduction reaction occurs.
 Note that, with an inert electrode such as Pt or C, only the electron transfer proceeds by acting with the active substance in the electrolytic solution. However, in an active anode electrode such as Fe/Pb, the electrode itself releases (donates) electrons to an active substance in the electrolytic solution and dissolves as an ion or reacts with a substance in the electrolytic solution to form an insoluble salt or oxide.


アノード = 外部回路に電子を放出する電極 = 電解液中の活性物質から電子を受容する電極 = 酸化反応が起こる電極
 電位により極性を定義する場合は、電位が高い方を正極(+)(positive electrode)、低い方を負極(-)(negative electrode)と呼ぶ。両極に抵抗等の負荷を掛けて接続したとき、外部電流は正極(+)から負極(-)へ、電子は負極(-)から正極(+)へ流れる。
  電解(充電): 正極 = アノード、負極 = カソード
  電池(放電): 正極 = カソード、負極 = アノード


Notation on this site

 On this site, “electrode functions and names” are described as follows.
Anode = Electrode that emits electrons to an external circuit = Electrode that accepts electrons from the active substance in the electrolyte = Electrode where oxidation reaction occurs
Cathode responds to the above items in reverse.
Positive and negative electrodes
 When the polarity is defined by the potential, the higher potential is called the positive electrode (+) and the lower potential is called the negative electrode (-). When a load such as a resistance is applied to both electrodes and connected, an external current flows from the positive electrode (+) to the negative electrode (-), and electrons flow from the negative electrode (-) to the positive electrode (+).
Battery and electrolysis
 The correspondence between positive/negative electrodes and anode/cathode is reversed in batteries and electrolysis (electrolysis).
  Electrolysis (charging): Positive electrode = Anode, Negative electrode = Cathode
  Battery (discharge): Positive electrode = cathode, negative electrode = anode
 This corresponds to the current flowing into the positive electrode during electrolysis (charging) and the current flowing out from the positive electrode during battery (discharging). For example, the PbO2 electrode of a lead acid battery, which is a secondary battery, is a positive electrode during charging and discharging, but becomes an anode through which current flows in during charging and is oxidized, and becomes a cathode through which current flows out during discharging and is reduced.

Anode and cathode
 The terminology for anodes and cathodes is confusing in Japan, because there have been two styles: (1) depending on the direction of current (direction of oxidation/reduction) (direct translation of anode/cathode) and (2) depending on the high or low of the potential.
 On the other hand, the terms positive electrode and negative electrode are well established as the distinction between “high and low” in potential. In Japanese high school chemistry, “positive and negative” electrodes are used for batteries and “anode and negative” electrodes for electrolysis.


Published: 2018/01/24
Updated: January 08, 2020 (Added table of contents and “8. Functions and names of electrodes”)
Updated: June 20, 2020 (English version posted)
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