Understanding Redox Reactions

In a redox reaction, if a species is reduced, what occurs in terms of electron transfer and what happens to its oxidation number?

Consider the polyatomic ion $SO_4^{2-}$. What is the overall oxidation number of this ion, and what principle dictates this value?

During a redox reaction, a substance undergoes oxidation. Which of the following statements accurately describes what happens to this substance?

When balancing redox half-equations, why is it crucial to balance the oxygen atoms by adding $H_2O$ to the appropriate side of the equation?

In the compound $OF_2$, what is the oxidation number of oxygen, and why does it differ from its usual oxidation number?

What is the primary reason for balancing redox equations, beyond simply ensuring the same number of atoms of each element on both sides?

Consider a redox reaction where a metal atom, M, is oxidized to $M^{2+}$. What specific changes occur at the atomic level during this oxidation process?

If a species is described as being reduced in a redox reaction, what directly measurable effect does it have on another species involved in the same reaction?

For a compound such as $H_2O$, how do you confirm that the sum of the oxidation numbers of all atoms equals zero, and what does this principle indicate about charge distribution?

What is the critical role of balancing 'charge with e-' in the final step of balancing redox half-equations, and what chemical reality does it ensure?

In balancing redox reactions, why is it necessary to balance hydrogen atoms ($H$) by adding $H^+$ ions, and what specific conditions does this assumption typically imply about the reaction environment?

In an electrolytic process, if a metal ion $M^{3+}$ gains three electrons to deposit as a solid metal on the cathode, how does this reduction affect the local electrochemical environment?

Consider a scenario where hydrogen is bonded to a metal, such as in sodium hydride ($NaH$). How does this affect the oxidation number of hydrogen, and why?

Given the unbalanced redox reaction: $MnO_4^- + Fe^{2+} ightarrow Mn^{2+} + Fe^{3+}$ in an acidic solution, what is the stoichiometric coefficient for $Fe^{2+}$ once the equation is fully balanced?

If chlorine gas ($Cl_2$) is bubbled into a solution containing bromide ions ($Br^−$), what observable change would indicate a redox reaction has occurred, and what is the underlying chemical process?

In the context of oxidation numbers, how does the concept of electronegativity govern the assignment of oxidation states in covalent compounds?

What is the implication of the statement 'the species with the e- is getting oxidized' in terms of electron availability during a chemical reaction?

Given a voltaic cell constructed with zinc and copper electrodes, predict what occurs at the zinc electrode as the voltaic cell discharges, and justify your prediction based on redox principles.

Consider a galvanic cell where the half-reactions involve complex ion formation. How does the formation of a stable complex ion affect the reduction potential of a metal ion in the half-cell?