Hydrogen half-reaction

In the discussion of the Daniell cell, we indicated that this cell produces a voltage of 1.10 V. This voltage is really the difference in potential between the two half-cells. The cell potential (really the half-cell potentials) is dependent upon concentration and temperature, but initially we ll simply look at the half-cell potentials at the standard state of 298 K (25°C) and all components in their standard states (1M concentration of all solutions, 1 atm pressure for any gases and pure solid electrodes). Half-cell potentials appear in tables as the reduction potentials, that is, the potentials associated with the reduction reaction. We define the hydrogen half-reaction (2H+(aq) + 2e - H2(g)) as the standard and has been given a value of exactly 0.00 V. We measure all the other half-reactions relative to it some are positive and some are negative. Find the table of standard reduction potentials in your textbook. 

Near the middle of the list, you will see 0 volts arbitrarily assigned to the standard hydrogen electrode all other potentials are relative to the hydrogen half-reaction. The voltages are given signs appropriate for a reduction reaction. For oxidation, the sign is reversed thus, the oxidation half-reaction,... 

Many reactions that occur in living cells are oxidation-reduction reactions. Appendix IX lists several compounds of biological importance and shows their relative tendencies to gain electrons ai 25°C and pH 7 under standard conditions. The numerical values of Ho reflect the reduction potentials relative to the 2H + 2e" H2 half-reaction which is taken as — 0.414 volt at pH 7. The value for the hydrogen half-reaction at pH 7 was calculated from the arbitrarily assigned value (Ho) of 0.00 volt under true standard-state conditions (1 M H and 1 atm Hs). For those few halfreactions of biological importance that do not involve as a reactant, the Ho and Ho values are essentially identical. 

E° in this reaction is defined as zero when H2 gas is at 1.0 atm, [H+j is 1.0 M (i.e., pH 0), and the temperature is 25°C, thus fixing the scale of E° value for other reactions. If it is easier to remove electrons from a material, E° will be negative if it is harder, E° will be positive. E° for the hydrogen half-reaction is used as the zero point even when reference is made to E° values. Thus,... 

Dihydrogen is a reducing agent. The H2/H couple at [H ] = I M defines the zero of the potential scale. At pH = 7 the hydrogen half-reaction... 

Figure 21.8 Determining an unknown Ehait-ceii with the standard reference (hydrogen) eiectrode. A voltaic cell has the 2n halfreaction in one half-cell and the hydrogen reference half-reaction in the other. The magnified view of the hydrogen half-reaction shows two H3O+ ions being reduced to two H2O molecules and an H2 molecule,...

To see which metals reduce (referred to as displacing H2 ) from acids, choose a metal, write its half-reaction as an oxidation, combine this half-reaction with the hydrogen half-reaction, and see if °en is positive. What you find is that the metals Li through Pb, those that lie below the standard hydrogen (reference) half-reaction in Appendix D, give a positive °en when reducing H. Iron, for example, reduces H from an acid to H2 ... 

All electrochemical cells involve two half-reactions an oxidation half-reaction in which electrons are released, and a reduction half-reaction in which electrons are taken up. The net voltage of the cell, the only quantity that can be measured experimentally, is the algebraic sum of the potentials for the two half-reactions. Each potential is a measure of the relative ability of a given half-reaction to occur. However, since the potential for a half-reaction cannot be measured directly, numerical values for half-reaction potentials are arbitrary and must be based on a reference potential. The hydrogen half-reaction serves as the reference for all electrochemical potentials. 

Metals that can displace H2from acid. The standard hydrogen half-reaction represents the reduction of H+ ions from an acid to H2 ... 

Using experimental values on the exchange current densities and symmetry coefficients and employing the Butler-Volmer equation, calculate the current density of the hydrogen half-reaction, 2H+(aq) + 2e = H2(g), on Pt electrode in 1 mol L" H2SO4 (aq) at an anodic polarization of 0.1 V. 

Calculate the limiting current for the hydrogen half-reaction, H+(aq) + e =... 

Half-reaction (i) means that Co(II) in aqueous solution cannot be oxidised to Co(III) by adding ammonia to obtain the complexes in (ii), oxidation is readily achieved by, for example, air. Similarly, by adding cyanide, the hexacyanocobaltate(II) complex becomes a sufficiently strong reducing agent to produce hydrogen from water ... 

When either hydrogen ions or hydroxide ions participate in a redox half-reaction, then clearly the redox potential is alTected by change of pH. Manganate(Vir) ions are usually used in well-acidified solution, where (as we shall see in detail later) they oxidise chlorine ions. If the pH is increased to make the solution only mildly acidic (pH = 3-6), the redox potential changes from 1.52 V to about 1.1 V, and chloride is not oxidised. This fact is of practical use in a mixture of iodide and chloride ions in mildly acid solution. manganate(VII) oxidises only iodide addition of acid causes oxidation of chloride to proceed. 

Hydrogen peroxide has both oxidising properties (when it is converted to water) and reducing properties (when it is converted to oxygen) the half-reactions are (acid solution) ... 

Standard Hydrogen Electrode The standard hydrogen electrode (SHE) is rarely used for routine analytical work, but is important because it is the reference electrode used to establish standard-state potentials for other half-reactions. The SHE consists of a Pt electrode immersed in a solution in which the hydrogen ion activity is 1.00 and in which H2 gas is bubbled at a pressure of 1 atm (Figure 11.7). A conventional salt bridge connects the SHE to the indicator half-cell. The shorthand notation for the standard hydrogen electrode is... 

For acidic solutions, balance the hydrogen in each half-reaction by adding H3O+ and H2O to opposite sides of the reaction for basic solutions, add OH and H2O to opposite sides of the reaction. 

The rate at which the corrosion of the 2iac proceeds depends on the rates of the two half reactions (eqs. 8 and 12). Equation 8, a necessary part of the desired battery reaction, fortunately represents a reaction that proceeds rather rapidly, whereas the reaction represented by equation 12 is slow. le, the generation of hydrogen on pure 2iac is a sluggish reaction and thus limits the overall corrosion reaction rate. 

Some typical half-cell reactions and their respective standard reduction potentials are listed in Table 21.1. Whenever reactions of this type are tabulated, they are uniformly written as reduction reactions, regardless of what occurs in the given half-cell. The sign of the standard reduction potential indicates which reaction really occurs when the given half-cell is combined with the reference hydrogen half-cell. Redox couples that have large positive reduction potentials... 

Pourbaix has classified the various equilibria that occur in aqueous solution into homogeneous and heterogeneous, and has subdivided them according to whether the equilibria involve electrons and/or hydrogen ions. The general equation for a half reaction is... 

This half-reaction commonly occurs when the cation in solution is very difficult to reduce. For example, electrolysis of a solution containing K+ ions (E d = —2.936 V) or Na+ ions (E = —2.714 V) yields hydrogen gas at the cathode (Table 18.4). 

Each zinc atom loses two electrons in changing to a zinc ion, therefore zinc is oxidized. Each hydrogen ion gains an electron, changing to a hydrogen atom, therefore hydrogen is reduced. (After reduction, two hydrogen atoms combine to form molecular H2.) As before, reaction (7) can be separated into two half-reactions ... 

Both of these half-reactions show production of electrons. But we know there must be an electron used for each produced, so one of the equations must be reversed. Experiment shows us it is the second because hydrogen gas is evolved from the solution. The first equation is correct as written. Lithium metal dissolves and is converted to ions. Thus,... 

Aluminum metal reacts with aqueous acidic solutions to liberate hydrogen gas. Write the two half-reactions and the net ionic reaction. 

A half-cdl consisting of a palladium rod dipping into a 1 M Pd(NOj)2 solution is connected with a standard hydrogen half-cell. The cell voltage is 0.99 volt and the platinum dectrode in the hydrogen half-cell is the anode. Determine E° for the reaction... 

H+], calculation of, 192, see also Hydrogen ion Haber, Fritz, 151 Haber process, 140, 150 Hafnium, oxidation number, 414 Haldane, J. B. S., 436 Half-cell potentials effect of concentration, 213 measuring, 210 standard, 210 table of, 211, 452 Half-cell reactions, 201 Half-life, 416 Half-reaction, 201 balancing, 218 potentials, 452 Halides... 

One combination of electrodes that could be used to determine pH is a hydrogen electrode connected through a salt bridge to a calomel electrode. The reduction half-reaction for the calomel electrode is... 

Iron(Il) hydrogen phosphite, FeHPO, is oxidized by hypochlorite ions in basic solution. The products are chloride ion, phosphate ion, and iron(lll) hydroxide. Write the balanced equation for each half-reaction and the overall equation for the reaction. 

When the two electrodes are connected, current flows from M to X in the external circuit. When the electrode corresponding to half-reaction 1 is connected to the standard hydrogen electrode (SHE), current flows from M to SHE. (a) What are the signs of ° of the two half-reactions (b) What is the standard cell potential for the cell constructed from these two electrodes ... 

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