Redox Reaction flash cards

1. 1.Define a redox reaction.

   A reaction in which oxidation and reduction occur simultaneously. Electrons lost by one species (oxidised) are gained by another (reduced), so the two processes are always coupled.

   Hint: Reduction + Oxidation = Redox.

2. 2.Define oxidation in terms of oxygen and hydrogen (classical view).

   Oxidation = addition of oxygen or removal of hydrogen. Example: $2Mg + O_2 \rightarrow 2MgO$ (Mg oxidised); $H_2S + Cl_2 \rightarrow 2HCl + S$ ($H_2S$ loses H, oxidised).

   Hint: Gain O or lose H.

3. 3.Define reduction in terms of oxygen and hydrogen (classical view).

   Reduction = removal of oxygen or addition of hydrogen. Example: $CuO + H_2 \rightarrow Cu + H_2O$ (CuO reduced); $N_2 + 3H_2 \rightarrow 2NH_3$ ($N_2$ gains H, reduced).

   Hint: Lose O or gain H.

4. 4.Define oxidation and reduction in terms of electron transfer.

   Oxidation = loss of electrons; Reduction = gain of electrons. This is the most general definition and applies even when no oxygen or hydrogen is involved.

   Hint: OIL RIG.

5. 5.What does the mnemonic OIL RIG stand for?

   Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons).

   Hint: Two three-letter words about electrons.

6. 6.In $Zn + Cu^{2+} \rightarrow Zn^{2+} + Cu$, which species is oxidised and which is reduced?

   $Zn$ loses 2 electrons ($Zn \rightarrow Zn^{2+} + 2e^-$) so it is oxidised. $Cu^{2+}$ gains 2 electrons ($Cu^{2+} + 2e^- \rightarrow Cu$) so it is reduced.

   Hint: Track the electrons on each half.

7. 7.What is an oxidising agent (oxidant)?

   A substance that oxidises another (accepts electrons) and is itself reduced in the process. Its oxidation number decreases.

   Hint: It takes electrons; it goes down in ON.

8. 8.What is a reducing agent (reductant)?

   A substance that reduces another (donates electrons) and is itself oxidised. Its oxidation number increases.

   Hint: It gives electrons; it goes up in ON.

9. 9.Define oxidation number (oxidation state).

   The charge an atom would carry if all bonds were assumed 100% ionic — i.e. shared electrons are assigned to the more electronegative atom. It may be positive, negative, zero, or fractional.

   Hint: Hypothetical charge from ionic assumption.

10. 10.What is the oxidation number of any element in its free (elemental) state?

    Zero. Examples: $O_2$, $H_2$, $P_4$, $S_8$, $Cl_2$, metallic $Na$, $O_3$ all have each atom at ON $= 0$.

    Hint: Uncombined = 0.

11. 11.What is the oxidation number rule for a monatomic ion?

    It equals the charge of the ion. Example: $Na^+ = +1$, $Mg^{2+} = +2$, $Cl^- = -1$, $S^{2-} = -2$.

    Hint: ON = ionic charge.

12. 12.State the usual oxidation number of hydrogen and its exception.

    Hydrogen is usually $+1$ (with non-metals). In metal hydrides (e.g. $NaH$, $CaH_2$) it is $-1$.

    Hint: +1 normally, −1 with metals.

13. 13.State the usual oxidation number of oxygen and its main exceptions.

    Oxygen is usually $-2$. Exceptions: peroxides ($-1$, e.g. $H_2O_2$), superoxides ($-\tfrac{1}{2}$, e.g. $KO_2$), and in $OF_2$ it is $+2$ (F is more electronegative).

    Hint: −2 usually; peroxide −1; $OF_2$ +2.

14. 14.What oxidation number do alkali (Group 1) and alkaline earth (Group 2) metals always show in compounds?

    Group 1 metals are always $+1$; Group 2 metals are always $+2$.

    Hint: Match the group number.

15. 15.What is the oxidation number of fluorine in all its compounds?

    Always $-1$, because fluorine is the most electronegative element and never carries a positive oxidation state.

    Hint: Most electronegative → always −1.

16. 16.What is the sum rule for oxidation numbers in a neutral molecule and in a polyatomic ion?

    For a neutral molecule the sum of all oxidation numbers $= 0$. For a polyatomic ion the sum equals the charge on the ion.

    Hint: Sum = 0 or = ion charge.

17. 17.QUESTION: Find the oxidation number of S in $H_2SO_4$.

    Let ON of S $= x$. $2(+1) + x + 4(-2) = 0 \Rightarrow 2 + x - 8 = 0 \Rightarrow x = +6$. So S is $+6$.

    Hint: H = +1, O = −2, sum = 0.

18. 18.QUESTION: Find the oxidation number of Mn in $KMnO_4$.

    $K = +1$, each $O = -2$. $+1 + x + 4(-2) = 0 \Rightarrow x - 7 = 0 \Rightarrow x = +7$. Mn is $+7$.

    Hint: K = +1, O = −2, total = 0.

19. 19.QUESTION: Find the oxidation number of Cr in $Cr_2O_7^{2-}$.

    $2x + 7(-2) = -2 \Rightarrow 2x - 14 = -2 \Rightarrow 2x = 12 \Rightarrow x = +6$. Each Cr is $+6$.

    Hint: Sum equals ion charge, −2.

20. 20.QUESTION: Find the oxidation number of N in $HNO_3$.

    $+1 + x + 3(-2) = 0 \Rightarrow x - 5 = 0 \Rightarrow x = +5$. N is $+5$.

    Hint: H = +1, O = −2.

21. 21.QUESTION: Find the oxidation number of C in $CO_2$ and in $CH_4$.

    In $CO_2$: $x + 2(-2) = 0 \Rightarrow x = +4$. In $CH_4$: $x + 4(+1) = 0 \Rightarrow x = -4$. Carbon ranges widely.

    Hint: O drags it up, H drags it down.

22. 22.QUESTION: Find the average oxidation number of S in $Na_2S_2O_3$ (sodium thiosulphate).

    $2(+1) + 2x + 3(-2) = 0 \Rightarrow 2 + 2x - 6 = 0 \Rightarrow 2x = 4 \Rightarrow x = +2$. Average ON of S $= +2$ (the two S atoms are actually inequivalent).

    Hint: Na = +1, O = −2; it's an average.

23. 23.QUESTION: Find the oxidation number of Fe in $Fe_3O_4$.

    $3x + 4(-2) = 0 \Rightarrow 3x = 8 \Rightarrow x = +\tfrac{8}{3}$. The average is $+\tfrac{8}{3}$ (it is a mix of $Fe^{2+}$ and $2\,Fe^{3+}$).

    Hint: Fractional — it's an average of +2 and +3.

24. 24.QUESTION: Find the oxidation number of Cl in $HClO_4$ (perchloric acid).

    $+1 + x + 4(-2) = 0 \Rightarrow x - 7 = 0 \Rightarrow x = +7$. Chlorine is $+7$, its highest state.

    Hint: H = +1, O = −2.

25. 25.QUESTION: Find the oxidation number of S in $H_2SO_3$.

    $2(+1) + x + 3(-2) = 0 \Rightarrow 2 + x - 6 = 0 \Rightarrow x = +4$. Sulphur is $+4$.

    Hint: Compare with +6 in $H_2SO_4$.

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