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List five common uses of acids and bases in our daily lives.

Chemistry

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Answer

$ClO_4^-, HSO_4^-, NO_3^-

HClO_4, HSO_4^-, HNO_3$

Step-by-step solution

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  1. Lewis acid definition

    A Lewis acid is a substance (atom, ion, or molecule) that can accept a pair of electrons to form a coordinate covalent bond. A Lewis base is a substance that can donate a pair of electrons.

  2. Common inorganic and organic Lewis acids

    Inorganic Lewis acids: BF3BF_3, AlCl3AlCl_3, BeCl2BeCl_2, Fe3+Fe^{3+}, H+H^+, SO3SO_3.
    Organic Lewis acids: carbocations like CH3+CH_3^+, (CH3)3C+(CH_3)_3C^+, acyl chlorides like CH3COClCH_3COCl.

  3. Lewis acids and bases in daily life

    Lewis acids: CO2CO_2 (in carbonated drinks, accepts electrons from water), SO2SO_2 (in pollution, from burning coal).
    Lewis bases: NH3NH_3 (in household cleaners), OH−OH^- (in soaps).

  4. Lewis structures for the acids

    HNO3_3: Central N bonded to one OH (single bond), one =O (double bond), one O−^- (single bond, resonance). N has 8 electrons (one from charge).

    H–O–N∣∣O\begin{align*} &\text{H--O--N} \\ &\qquad || \\ &\qquad \text{O} \end{align*}

    (with lone pairs: O in OH has 2, =O has 2, terminal O has 3; resonance forms).

    H2_2SO4_4: Central S bonded to two OH (single), two =O (double). S has 12 electrons (expanded octet).

     HO–S–OH∣∣∣∣OO\begin{align*} &\text{ HO--S--OH} \\ &\qquad || \quad || \\ &\qquad \text{O} \qquad \text{O} \end{align*}

    HClO: H--O--Cl, O has 2 lone pairs, Cl has 3 lone pairs.

    H–O–Cl(: on O and Cl)\text{H--O--Cl} \quad (: \text{ on O and Cl})

    HBr: H--Br, Br has 3 lone pairs.

    H–Br:\text{H--Br:}

    HClO4_4: Central Cl bonded to three =O, one OH. Cl has 12 electrons.

    HO–Cl∣∣O∣O\begin{align*} &\text{HO--Cl} \\ &\qquad || \\ &\qquad \text{O} \\ &\qquad | \\ &\qquad \text{O} \end{align*}

    (resonance, all O equivalent).

  5. Lewis structures for NH4+_4^+ and H3_3O+^+

    NH4+_4^+: Central N with 4 N--H single bonds (8 electrons on N).

    [\text{H--N--H} \atop | \atop \text{H--} \text{--H}]^+

    H3_3O+^+: Central O with 3 O--H bonds, 1 lone pair on O.

    [\text{H--O--H} \atop | \atop \text{H}]^+ \quad (: \text{ on O})

  6. Formulas for conjugate bases

    Formula first: Conjugate base = acid −- H+^+.
    HClO4→_4 \to ClO4−_4^-.

    \ceHClO4−>[−H+]ClO4−\ce{HClO4 ->[ -H^+ ] ClO4^-}

    H2_2SO4→_4 \to HSO4−_4^-.

    \ceH2SO4−>[−H+]HSO4−\ce{H2SO4 ->[ -H^+ ] HSO4^-}

    HNO3→_3 \to NO3−_3^-.

    \ceHNO3−>[−H+]NO3−\ce{HNO3 ->[ -H^+ ] NO3^-}

  7. Formulas for conjugate acids

    Formula first: Conjugate acid = base ++ H+^+.
    ClO4−→_4^- \to HClO4_4.

    \ceClO4−−>[+H+]HClO4\ce{ClO4^- ->[ +H^+ ] HClO4}

    SO42−→_4^{2-} \to HSO4−_4^-.

    \ceSO42−−>[+H+]HSO4−\ce{SO4^2- ->[ +H^+ ] HSO4^-}

    NO3−→_3^- \to HNO3_3.

    \ceNO3−−>[+H+]HNO3\ce{NO3^- ->[ +H^+ ] HNO3}

  8. Amphoteric and amphiprotic substances

    Amphoteric (act as both Lewis acid and base): \ceAl(OH)3\ce{Al(OH)3}, \ceZn(OH)2\ce{Zn(OH)2}, \ceSn(OH)2\ce{Sn(OH)2}.
    Example: \ceAl(OH)3+OH−−>Al(OH)4−\ce{Al(OH)3 + OH^- -> Al(OH)4^-} (accepts e−^- pair, Lewis acid); \ceAl(OH)3+3H+−>Al3++3H2O\ce{Al(OH)3 + 3H^+ -> Al^3+ + 3H2O} (Lewis base).

    Amphiprotic (act as both Brønsted acid and base): \ceH2O\ce{H2O}, \ceHCO3−\ce{HCO3^-}, \ceHSO4−\ce{HSO4^-}.
    Example: \ceH2O+HCl−>H3O++Cl−\ce{H2O + HCl -> H3O^+ + Cl^-} (acid); \ceH2O+NH3−>NH4++OH−\ce{H2O + NH3 -> NH4^+ + OH^-} (base).

  9. Sulfur atom in SO2_2 as Lewis acid

    Lewis structure of SO2_2: Central S with 1 lone pair, double bond to one O, single bond to other O (resonance, 10 electrons on S).

    O∣∣O=S\begin{align*} & \qquad \text{O} \\ & \quad || \\ \text{O=S} \end{align*}

    S has incomplete octet (can accept e−^- pair), so SO2_2 is Lewis acid (e.g., forms adduct with amines).
    This explains corrosion in pipes: SO2_2 from pollution accepts electrons from metal oxides, breaking layers.

    ClO_4^-, HSO_4^-, NO_3^-

    HClO_4, HSO_4^-, HNO_3

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