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question 6 (1 point) saved
which of the following species is neither an acid or a base as defined by bronsted-lowry?
hpo₄²⁻
alcl₃
nh₃
co₃²⁻
all of these are either an acid or a base as defined by bronsted-lowry.
- Recall Bronsted - Lowry definitions: An acid is a proton ($H^+$) donor, and a base is a proton acceptor.
- Analyze each species:
- $\ce{HPO_4^{2-}}$: Can accept a proton to form $\ce{H_2PO_4^-}$ (acting as a base) or donate a proton to form $\ce{PO_4^{3-}}$ (acting as an acid), so it is an amphoteric species (both acid and base).
- $\ce{AlCl_3}$: In water, $\ce{Al^{3+}}$ can hydrolyze. The $\ce{Al^{3+}}$ ion can accept electron pairs (Lewis acid), but in terms of Bronsted - Lowry, when considering the hydrolysis reaction $\ce{Al^{3+}+3H_2O
ightleftharpoons Al(OH)_3 + 3H^+}$, the $\ce{H_2O}$ donates protons to form $\ce{H^+}$, and the $\ce{Al^{3+}}$ can be considered in the context of the overall reaction. However, more directly, if we consider the species in terms of proton donation/acceptance: $\ce{AlCl_3}$ itself does not have protons to donate, but in the hydrolysis process, the water acts as the acid. Wait, no, let's re - evaluate. Wait, actually, the Bronsted - Lowry definition is about proton ($H^+$) donation/acceptance. $\ce{AlCl_3}$ is a Lewis acid (electron pair acceptor), but does it act as a Bronsted - Lowry acid or base? A Bronsted - Lowry acid must donate $H^+$, and a base must accept $H^+$. $\ce{AlCl_3}$ has no $H$ atoms, so it cannot donate $H^+$. Can it accept $H^+$? The $Al^{3+}$ ion has a positive charge, and $H^+$ also has a positive charge, so it is not likely to accept $H^+$ (like - charges repel). Wait, but maybe I made a mistake earlier. Wait, the other species:
- $\ce{NH_3}$: Can accept a proton to form $\ce{NH_4^+}$, so it is a Bronsted - Lowry base.
- $\ce{CO_3^{2-}}$: Can accept a proton to form $\ce{HCO_3^-}$ (acting as a base).
- Wait, but the option "All of these are either an acid or a base as defined by Bronsted - Lowry" - Wait, no, earlier analysis of $\ce{AlCl_3}$ was wrong. Wait, no, let's correct. Wait, the key is that the question is asking which is neither. But according to the options, the correct answer is the last one? Wait, no, wait:
Wait, no, let's re - check:
- $\ce{HPO_4^{2-}}$: Bronsted - Lowry acid (donates $H^+$) or base (accepts $H^+$).
- $\ce{AlCl_3}$: Is it a Bronsted - Lowry acid or base? A Bronsted - Lowry acid donates $H^+$, a base accepts $H^+$. $\ce{AlCl_3}$ has no $H$ to donate. Can it accept $H^+$? The $Al^{3+}$ has a positive charge, and $H^+$ is also positive, so electrostatically, it's not favorable. But maybe in some reactions? Wait, no, the other species:
- $\ce{NH_3}$: Accepts $H^+$ (base).
- $\ce{CO_3^{2-}}$: Accepts $H^+$ (base).
Wait, but the option "All of these are either an acid or a base as defined by Bronsted - Lowry" - but $\ce{AlCl_3}$: Wait, maybe I was wrong about $\ce{AlCl_3}$. Wait, actually, in the context of aqueous solutions, when $\ce{AlCl_3}$ dissolves in water, the reaction is $\ce{AlCl_3 + 3H_2O
ightleftharpoons Al(OH)_3+3HCl}$. In this reaction, $\ce{H_2O}$ donates a proton to form $\ce{OH^-}$ (which then reacts with $\ce{Al^{3+}}$), and $\ce{HCl}$ is formed. But the $\ce{AlCl_3}$ itself: does it act as a Bronsted - Lowry acid or base? It doesn't donate $H^+$, and does it accept $H^+$? The $Al^{3+}$ ion is a Lewis acid (accepts electron pairs from $\ce{H_2O}$), but in terms of Bronsted - Lowry, the species that donates $H^+$ is $\ce{H_2O}$ (acid), and the species that accepts $H^+$? Wait, no, the $\ce{Cl^-}$ ion is a conjugate base of $\ce{HCl}$, but $\ce{AlCl_3}$ as a compound: maybe the question has a different approach. The other species:
- $\ce{HPO_4^{2-}}$: amphoteric (acid and base).
- $\ce{NH_3}$: base (accepts $H^+$).
- $\…
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All of these are either an acid or a base as defined by Bronsted - Lowry.