Acid/Base Definitions and Concepts

Arrhenius Definition

The Arrhenius definition of acids and bases focuses on the behavior of substances in aqueous solutions:

Acid: A substance that increases the concentration of hydrogen ions (\H^+) when dissolved in water. For example, hydrochloric acid (\HCl) dissociates into $H^+$ and $Cl^-$ ions in water, thereby increasing the concentration of $H^+$ ions.
Base: A substance that increases the concentration of hydroxide ions (\OH^−) when dissolved in water. For example, sodium hydroxide (\NaOH) dissociates into $Na^+$ and $OH^−$ ions in water, increasing the concentration of $OH^−$ ions.

Bronsted-Lowry Definition

The Bronsted-Lowry definition broadens the scope of acids and bases beyond aqueous solutions:

Acid: A substance that donates a proton (\H^+). It is also known as a proton donor.
Base: A substance that accepts a proton (\H^+). It is also known as a proton acceptor.
In this context, an acid-base reaction involves the transfer of a proton from an acid to a base. For example, in the reaction of ammonia (\NH3) with water ( $\H2O$ ):

$NH3 + H2O \rightleftharpoons NH4^+ + OH^-$ $\H2O$ acts as an acid by donating a proton to \NH_3, which acts as a base by accepting the proton.

Lewis Definition

The Lewis definition is the most inclusive, focusing on the transfer of electron pairs:

Acid: A substance that accepts an electron pair. It is also known as an electron pair acceptor.
Base: A substance that donates an electron pair. It is also known as an electron pair donor.
For instance, in the reaction between ammonia (\NH3) and boron trifluoride (\BF3):

$NH3 + BF3 \rightarrow NH3BF3$

\BF3 acts as a Lewis acid by accepting an electron pair from \NH3, which acts as a Lewis base by donating the electron pair.

pH and pOH

Definitions of pH and pOH

pH: A measure of the concentration of hydrogen ions (\H^+) in a solution. It is defined as the negative logarithm (base 10) of the hydrogen ion concentration:

\pH = -\log_{10} [H^+]

pOH: A measure of the concentration of hydroxide ions (\OH^−) in a solution. It is defined as the negative logarithm (base 10) of the hydroxide ion concentration:

\pOH = -\log_{10} [OH^-]

Relationship between pH and pOH

In aqueous solutions at 25°C, the sum of pH and pOH is always 14:

\pH + pOH = 14

pH Scale

The pH scale typically ranges from 0 to 14:

Acidic Solutions: pH 10 [H^+] $</li></ul>From pH to [H+]$ [H^+] = 10^{-\pH} $From [OH-] to pOH$ \pOH = -\log_{10} [OH^-] $From pOH to [OH-]$ [OH^-] = 10^{-\pOH} $From pH to pOH$ \pOH = 14 - pH $From pOH to pH$ \pH = 14 - pOH $<h4 collapsed="false" seolevelmigrated="true">Conjugate Acids and Bases</h4><h5 collapsed="false" seolevelmigrated="true">Identifying Conjugate Pairs</h5>In the context of the Bronsted-Lowry definition:<ul><li>Acid: A species that donates a proton.</li><li>Base: A species that accepts a proton.</li><li>Conjugate Acid: The species formed when a base accepts a proton.</li><li>Conjugate Base: The species formed when an acid donates a proton. For example, consider the reaction: $ \HA + H2O \rightleftharpoons H3O^+ + A^- $</li><li>$ \HA $is the acid because it donates a proton.</li><li>$ \H_2O $is the base because it accepts a proton.</li><li>$ \H3O^+ $is the conjugate acid of$ \H2O $because it is formed when$ \H_2O $accepts a proton.</li><li>$ \A^− $is the conjugate base of$ \HA $because it is formed when$ \HA $donates a proton.</li></ul><h4 collapsed="false" seolevelmigrated="true">Strong Acids and Strong Bases</h4><h5 collapsed="false" seolevelmigrated="true">O Strong Acids</h5>Strong acids are acids that completely dissociate into ions when dissolved in water. This means that for a strong acid$ \HA $, the dissociation reaction is:$ \HA \rightarrow H^+ + A^- $ Common strong acids include:<ul><li>Hydrochloric acid ($ \HCl $)</li><li>Sulfuric acid ($ \H2SO4 $)</li><li>Nitric acid ($ \HNO_3$$)
Hydro