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According to ChemGuide,

The Brønsted-Lowry theory doesn't go against the Arrhenius theory in any way - it just adds to it.

According to Study.com:

All Arrhenius acids and bases are also Brønsted-Lowry acids and bases.

This makes sense for Arrhenius acids, such as hydrochloric acid. Hydrochloric acid produces hydrogen ions when dissolved in water. Additionally, it donates a proton to water.

However, sodium hydroxide is an Arrhenius base, since it produces hydroxide ions when dissolved in water. However, sodium hydroxide cannot accept a proton to form $\ce{HNaOH+}$. So, how can it be a Brønsted-Lowry base? Similar reasoning can be applied to other Arrhenius bases, such as potassium hydroxide.

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Sure it can. You're just incorrect about the products when $\ce{NaOH}$ accepts an $\ce{H+}$ in aqueous solution. The result is $\ce{H2O(l) + Na+(aq)}$ instead of $\ce{HNaOH+(aq)}$. That is, the $\ce{NaOH}$ doesn't stay together.

Now mind you, most people would point out to you that the $\ce{NaOH}$ doesn't stay together as soon as it's dissolved, meaning $\ce{NaOH(s) -> Na+(aq) + OH-(aq)}$ happens right away, and then as soon as you have some $\ce{H+(aq)}$ around it reacts only with the $\ce{OH-(aq)}$ and the $\ce{Na+(aq)}$ doesn't get involved, probably isn't even anywhere near. But this doesn't change the main point, which is that if you add $\ce{NaOH}$ to water, you have a solution that can react with $\ce{H+(aq)}$, which means you added a Brønsted-Lowry base.

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  • $\begingroup$ Quoting from Peter Atkins, NaOH is merely the provider of the Bronsted-Lowry base, the hydroxide ion. $\endgroup$ – Tan Yong Boon Jun 10 '18 at 4:47
  • $\begingroup$ Sure, that's a point of view @TanYongBoon. The only complication is that you could really say the same thing about strong acids. Since HCl(g) + H2O(l) -> H3O+(aq) + Cl-(aq) happens as soon as you put the HCl in, and any addition acid-base reaction is between the H3O+ and base, you could equally well say HCl is merely the provider of the Bronsted-Lowry acid instead of a BL acid itself. But we don't. There is definitely a mild logical lacuna here in our definitions, over which reasonable men could and probably will disagree. $\endgroup$ – Christopher Grayce Jun 10 '18 at 20:17
  • $\begingroup$ @ChristopherGrayce Is H2O (l) + Na+(aq) considered to be the conjugate acid of NaOH (aq)? Also, does NaOH (aq) have a Kb value? $\endgroup$ – user62238 Jun 11 '18 at 1:56
  • $\begingroup$ @user62238 yes, if we're asked we might say Na+(aq) is the conjugate acid of NaOH(aq), although the idea of a conjugate is a little strange for strong acids and bases, since the conjugates have no acid/base activity. Na+(aq) for example will not react with bases. Similarly the conjugate base of HCl is Cl-(aq) but Cl-(aq) will not react with acids. It's not of great use to attach Ka or Kb values to strong acids and bases (at least in aqueous solution) like HCl or NaOH, because of what's called the "leveling effect." $\endgroup$ – Christopher Grayce Jun 11 '18 at 2:32
  • $\begingroup$ In essence, every acid stronger than H3O+(aq) and every base strong than OH-(aq) gets immediately converted to these species on dissolution, and then all further acid/base chemistry happens with those species. So in aqueous solution, there's no point in ranking how strong an acid or base is, once it reacts completely on dissolution. (This is not true in organic solvents, by the way.) $\endgroup$ – Christopher Grayce Jun 11 '18 at 2:34

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