Yes, a substance can act as both an oxidising and reducing agent in different reactions, or even in the same reaction under certain conditions. This dual role is typically exhibited by substances that have a medium range of electronegativity and can both gain and lose electrons easily. A classic example is hydrogen peroxide (H₂O₂). In reactions where it interacts with substances more readily oxidised than itself, like in its reaction with potassium iodide (KI), it acts as an oxidising agent. However, when it reacts with more powerful oxidising agents, like potassium permanganate (KMnO₄), it serves as a reducing agent. This versatility is what makes hydrogen peroxide a widely used substance in various chemical reactions and applications, ranging from industrial processes to antiseptic formulations.

The oxidation state of an element is a key indicator of its role as an oxidising or reducing agent in a redox reaction. The oxidation state, essentially a hypothetical charge an atom would have if electrons were transferred completely, helps in understanding the electron transfer in a reaction. An increase in the oxidation state of an element during a reaction signifies that it has lost electrons, indicating its role as a reducing agent. Conversely, a decrease in the oxidation state implies that the element has gained electrons, characterising it as an oxidising agent. By examining the changes in oxidation states of the elements in a reaction, one can decipher the direction of electron flow and thereby identify the oxidising and reducing agents. This approach is particularly useful in complex reactions where the electron transfer might not be immediately apparent.

The concentration and physical state of a substance greatly influence its effectiveness as an oxidising or reducing agent. A higher concentration of an oxidising or reducing agent generally increases the rate of the redox reaction, as there are more reactive molecules or ions available to undergo electron transfer. For example, a concentrated solution of an oxidising agent like potassium permanganate will react more vigorously than a diluted one. Similarly, the physical state (solid, liquid, or gas) affects the surface area available for the reaction and the ease with which the molecules or ions can interact. Gases and liquids generally react more quickly than solids due to greater molecular mobility and larger surface areas for interaction. Therefore, understanding the influence of concentration and physical state is crucial in predicting and controlling the outcomes of redox reactions.

Environmental factors such as temperature, pH, and the presence of catalysts can significantly influence the role and effectiveness of a substance as an oxidising or reducing agent. Temperature, for instance, often increases the rate of redox reactions by providing the necessary energy for electron transfer. A higher temperature can enhance the reactivity of oxidising and reducing agents. The pH of the environment can also affect the redox potential of substances. For example, acidic or alkaline conditions can alter the oxidation state of certain elements, influencing their ability to act as oxidising or reducing agents. Catalysts play a pivotal role too, as they can lower the activation energy of a reaction, making it easier for the oxidising and reducing agents to react. These environmental factors, therefore, must be considered when analysing or predicting the behaviour of substances in redox reactions.

Electronegativity, which measures an atom's ability to attract and bond with electrons, is a crucial factor in identifying oxidising and reducing agents. In general, elements with high electronegativity, such as oxygen, fluorine, and chlorine, are more likely to act as oxidising agents because they tend to attract electrons towards themselves. This attribute enables them to gain electrons and undergo reduction. Conversely, elements with low electronegativity, such as metals like sodium, magnesium, and calcium, often serve as reducing agents. Their lower electronegativity means they readily lose electrons and get oxidised. Therefore, by assessing the relative electronegativities of elements in a reaction, one can predict which substances are more likely to be oxidising or reducing agents. For example, in a reaction between chlorine (high electronegativity) and sodium (low electronegativity), chlorine would act as the oxidising agent, and sodium as the reducing agent.