Redox Equivalent Concept & Titration flash cards

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

   Oxidation = loss of electrons (increase in oxidation number). Reduction = gain of electrons (decrease in oxidation number). Mnemonic: OIL RIG (Oxidation Is Loss, Reduction Is Gain).

   Hint: Think about electron bookkeeping.

2. 2.What is a redox reaction?

   A reaction in which oxidation and reduction occur simultaneously — electrons lost by one species are gained by another, so oxidation number changes of at least two atoms are coupled.

   Hint: Two half-reactions, one process.

3. 3.Define oxidizing agent and reducing agent.

   Oxidizing agent (oxidant): gets reduced, gains electrons, oxidizes the other species. Reducing agent (reductant): gets oxidized, loses electrons, reduces the other species.

   Hint: The agent does the opposite to itself.

4. 4.In classical (oxygen/hydrogen) terms, what is oxidation?

   Oxidation = addition of oxygen or an electronegative element, OR removal of hydrogen or an electropositive element. Reduction is the reverse.

   Hint: Pre-electronic definitions.

5. 5.State the oxidation number rule for free elements and diatomic molecules.

   Every atom in a free element or in a homonuclear molecule has oxidation number $0$. Examples: $O_2$, $H_2$, $P_4$, $S_8$, $Na$, $Cl_2$ all have O.N. $=0$.

   Hint: Uncombined = zero.

6. 6.What is the oxidation number of a monatomic ion?

   Equal to its charge. E.g. $Na^+ = +1$, $Ca^{2+} = +2$, $Cl^- = -1$, $S^{2-} = -2$.

   Hint: Charge = O.N. for lone ions.

7. 7.Give the usual oxidation numbers of H and O in compounds, plus exceptions.

   H $=+1$ (but $-1$ in metal hydrides like $NaH$). O $=-2$ (but $-1$ in peroxides, $-1/2$ in superoxides, $+2$ in $OF_2$, $+1$ in $O_2F_2$).

   Hint: Watch hydrides, peroxides, fluorides.

8. 8.What oxidation number does fluorine always have in compounds?

   Always $-1$ (fluorine is the most electronegative element, so it can never be positive in any compound).

   Hint: Most electronegative → always negative.

9. 9.How do you find the oxidation number of an atom in a neutral compound or polyatomic ion?

   Sum of all oxidation numbers = $0$ for a neutral compound, or = net charge for a polyatomic ion. Solve algebraically for the unknown.

   Hint: Set the sum equal to overall charge.

10. 10.Can oxidation numbers be fractional? Explain with an example.

    Yes — the O.N. is an average. In $Fe_3O_4$, Fe averages $+8/3$; in $Na_2S_4O_6$ (tetrathionate) S averages $+2.5$. Fractions arise when identical atoms are in different environments.

    Hint: Average over non-equivalent atoms.

11. 11.What is the oxidation state of S in $H_2SO_4$?

    $+6$. Solve: $2(+1) + S + 4(-2) = 0 \Rightarrow S = +6$.

    Hint: H is +1, O is -2, sum to zero.

12. 12.What is the oxidation state of Mn in $KMnO_4$?

    $+7$. Solve: $(+1) + Mn + 4(-2) = 0 \Rightarrow Mn = +7$.

    Hint: K = +1, O = -2.

13. 13.What is the oxidation state of Cr in $K_2Cr_2O_7$?

    $+6$ for each Cr. Solve: $2(+1) + 2Cr + 7(-2) = 0 \Rightarrow 2Cr = +12 \Rightarrow Cr = +6$.

    Hint: Two Cr share +12 total.

14. 14.What is a disproportionation reaction?

    A redox reaction in which the same element in a single oxidation state is simultaneously oxidized and reduced to two different oxidation states. Requires an intermediate O.N. for that element.

    Hint: One species, two fates.

15. 15.Give a classic example of disproportionation.

    $Cl_2 + 2OH^- \rightarrow Cl^- + ClO^- + H_2O$. Chlorine ($0$) goes to $-1$ (reduced) and $+1$ (oxidized). Also $2H_2O_2 \rightarrow 2H_2O + O_2$.

    Hint: Cold dilute alkali on chlorine.

16. 16.What is comproportionation (the reverse of disproportionation)?

    Two species containing the same element in different oxidation states react to give a single product with that element in one intermediate state. E.g. $5Cl^- + ClO_3^- + 6H^+ \rightarrow 3Cl_2 + 3H_2O$.

    Hint: Two states merge into one.

17. 17.Why can $H_2O_2$ act as both an oxidizing and a reducing agent?

    Oxygen in $H_2O_2$ is at intermediate O.N. $-1$. It can go to $-2$ (acts as oxidant) or to $0$ in $O_2$ (acts as reductant), so it does both.

    Hint: Intermediate oxidation state = versatile.

18. 18.Why is fluorine gas a strong oxidizing agent but never a reducing agent?

    F is the most electronegative element; its lowest O.N. is $-1$ and it has no positive states, so it can only gain electrons (be reduced) — it can only oxidize others, never reduce them.

    Hint: No positive oxidation states available.

19. 19.What are the two standard methods of balancing redox equations?

    (1) Oxidation-number method (balance electron gain/loss via O.N. change). (2) Ion-electron (half-reaction) method (split into oxidation and reduction half-reactions, balance separately, combine).

    Hint: O.N. method and half-reaction method.

20. 20.List the steps of the ion-electron method in acidic medium.

    (1) Write both half-reactions. (2) Balance atoms except O, H. (3) Balance O with $H_2O$. (4) Balance H with $H^+$. (5) Balance charge with $e^-$. (6) Equalize electrons, add, cancel.

    Hint: O with water, H with protons, charge with electrons.

21. 21.How does the ion-electron method differ in basic (alkaline) medium?

    Balance as if acidic, then add to both sides as many $OH^-$ as there are $H^+$; combine $H^+ + OH^- \rightarrow H_2O$ and cancel duplicate waters.

    Hint: Neutralize the H+ afterward.

22. 22.In the oxidation-number method, what is the balancing principle?

    Total increase in oxidation number (electrons lost) must equal total decrease in oxidation number (electrons gained). Multiply species by suitable factors so these balance.

    Hint: Electrons lost = electrons gained.

23. 23.Why must the number of electrons be equal in the two half-reactions before adding them?

    Electrons are neither created nor destroyed; every electron lost in oxidation is gained in reduction. Equalizing them ensures $e^-$ cancel out and charge is conserved.

    Hint: Conservation of charge/electrons.

24. 24.Define equivalent weight of a substance (general idea).

    Equivalent weight $= \dfrac{\text{Molar mass}}{n\text{-factor}}$. It is the mass that combines with or displaces $1$ mole of electrons (or $1$ mole of $H^+$/$OH^-$ / unit charge).

    Hint: Molar mass divided by n-factor.

25. 25.Define n-factor (equivalence factor).

    The number of reactive units per formula unit: H+ replaced (acid), OH- replaced (base), electrons transferred per formula unit (redox), or total charge exchanged (salt).

    Hint: Reactive units per formula unit.

26. 26.What is the n-factor of an acid? Give examples.

    n-factor = basicity = number of replaceable $H^+$ ions. $HCl = 1$, $H_2SO_4 = 2$, $H_3PO_4 = 3$, $CH_3COOH = 1$.

    Hint: Count ionizable H+.

27. 27.What is the n-factor of a base? Give examples.

    n-factor = acidity = number of replaceable $OH^-$ ions. $NaOH = 1$, $Ca(OH)_2 = 2$, $Al(OH)_3 = 3$.

    Hint: Count ionizable OH-.

28. 28.What is the n-factor of a salt in a (non-redox) double-displacement reaction?

    n-factor = total charge on cations (= total charge on anions) per formula unit. $NaCl = 1$, $Na_2CO_3 = 2$, $Al_2(SO_4)_3 = 6$, $CaCl_2 = 2$.

    Hint: Total positive (or negative) charge.

29. 29.What is the n-factor of an oxidizing or reducing agent in a redox reaction?

    n-factor = number of electrons gained or lost per formula unit = total change in oxidation number per formula unit.

    Hint: Electrons transferred per formula unit.

30. 30.State the law of equivalence (law of chemical equivalence).

    In any reaction, equivalents of each reactant and product are equal: for A reacting with B, $\text{eq}_A = \text{eq}_B$. Equivalents = (moles) × (n-factor) = mass / equivalent weight.

    Hint: Equivalents react in a 1:1 ratio.

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