The choice of platinum as the electrode material in the standard hydrogen electrode (SHE) is due to its unique properties. Platinum is chemically inert, meaning it does not react with the hydrogen or the ions in the solution. This inertness ensures that the electrode does not participate in the chemical reactions occurring in the cell, thereby not affecting the standard electrode potential being measured. Platinum also has a high catalytic activity, which is crucial for the hydrogen electrode reaction, where hydrogen gas is involved. The platinum surface catalyses the dissociation of hydrogen molecules (H₂) into hydrogen ions (H⁺) and electrons, facilitating the establishment of equilibrium necessary for the measurement of electrode potentials. Moreover, platinum's conductivity ensures efficient electron transfer, crucial for accurate and stable electrode potential readings. These properties make platinum an ideal choice for the electrode material in the SHE.

The standard electrode potential provides information about the thermodynamic feasibility of a redox reaction, but it does not predict the reaction rate. The electrode potential indicates the tendency of a species to gain or lose electrons, thereby determining the direction in which a reaction is thermodynamically favoured. However, the rate of the reaction depends on kinetic factors, which are not accounted for by the electrode potential. Factors such as the activation energy, temperature, and the presence of catalysts influence the reaction rate. A reaction with a positive standard cell potential (indicating spontaneity) may still proceed slowly if it has a high activation energy or occurs at a temperature that does not provide sufficient energy to the reactants. Therefore, while the standard electrode potential is crucial in understanding the spontaneity and direction of redox reactions, additional kinetic considerations are necessary to predict the reaction rate.

Using standard conditions when measuring electrode potentials is crucial for consistency and comparability of results. Standard conditions include a temperature of 298 K, a pressure of 1 atmosphere, and a concentration of 1 molar for solutions. These conditions ensure that the electrode potentials measured for different substances can be compared directly. If non-standard conditions were used, the potentials measured would be influenced by varying temperatures, pressures, or concentrations, making it difficult to ascertain the intrinsic electrochemical properties of the substances being tested. This comparability is vital for constructing an electrochemical series and for applications in electrochemistry, such as battery design and corrosion prediction. Additionally, standard conditions provide a reference point against which changes in electrode potentials under different conditions can be measured and understood, aiding in the study of electrochemical processes in various environmental and industrial contexts.

In environmental chemistry, the standard electrode potential is used to understand and manage redox reactions in natural and contaminated environments. For instance, it helps in predicting the behaviour of heavy metals in water bodies, essential for assessing their mobility, bioavailability, and potential toxicity. By knowing the standard electrode potentials of various metal ions, scientists can predict which metals are likely to be reduced or oxidised in a particular environment, influencing their solubility and interaction with organisms. Furthermore, this knowledge is applied in designing processes for water treatment and pollution remediation. Techniques such as electrochemical reduction or oxidation, which rely on applying potentials to drive desired redox reactions, are developed based on the standard electrode potentials of pollutants. These methods are effective in removing or detoxifying contaminants like heavy metals, organic pollutants, and inorganic anions from water and soil.

Temperature can significantly impact the standard electrode potential. The standard conditions for measuring electrode potentials include a temperature of 298 K. If the temperature changes, the equilibrium position of the half-cell reaction can shift, leading to a change in the electrode potential. This change is due to the effect of temperature on the reaction kinetics and equilibrium constants. As temperature increases, reaction rates generally increase, which can either enhance the reduction potential of a reductant or decrease it, depending on the reaction specifics. Additionally, the Nernst equation, which relates electrode potential to temperature, ion concentration, and the number of electrons transferred, indicates that potential varies with temperature changes. In practical applications, temperature control is essential for accurate measurements of electrode potentials, as deviations can lead to incorrect interpretations of a substance's electrochemical properties.