Lewis acids and pH

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Anushi Patel 1J
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Joined: Fri Sep 28, 2018 12:19 am
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Lewis acids and pH

Postby Anushi Patel 1J » Sat Dec 08, 2018 10:04 pm

How do Lewis bases/acids affect the pH of a solution? If they are bases, they should cause the OH- concentration to increase or the H+ concentration to decrease. If they are acids, the opposite. However, I don't see how they can do that if they are simply electron acceptors/donors. Can someone explain the mechanism?

Akhil Paladugu 3G
Posts: 30
Joined: Fri Sep 28, 2018 12:26 am

Re: Lewis acids and pH

Postby Akhil Paladugu 3G » Sat Dec 08, 2018 10:19 pm

Many Lewis acids are water reactive and will hydrolyze in water. Lewis bases will usually act as Bronsted bases as well and protonate in water thus raising the pH

Iona Pay
Posts: 43
Joined: Fri Sep 28, 2018 12:15 am

Re: Lewis acids and pH

Postby Iona Pay » Sat Dec 08, 2018 10:24 pm

Lewis acids and bases are generally defined as electron pair acceptors and donators respectively. When a Lewis acid and base are combined, they form a Lewis adduct (for our purposes here, a single product of two distinct molecules).

So, to understand how Lewis acid/bases affect pH:

Think of a proton (H+) as a Lewis acid - it has a positive charge, and could accept a pair of electrons. When adding H+ to a solution, we either raise or lower a pH. For example, protonating ammonia (H+ + NH3 → NH4+) lowers the pH of the solution because it produces a basic Lewis adduct. If we protonate hydroxide (H+ + OH− → H2O), we produce water, which is neutral.

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