Acids and Bases
Brønsted-Lowry Theory
The Brønsted-Lowry theory defines acids and bases based on the transfer of protons ( $ H^+ $ ).
- Acid: A species that donates a proton.
- Base: A species that accepts a proton.
Key Concepts:
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Conjugate Acid-Base Pairs: An acid and its conjugate base differ by a single proton. Similarly, a base and its conjugate acid differ by a single proton. For example, in the reaction of $ HCl $ with $ H_2O $ :
$ HCl(aq) + H_2O(l) \rightleftharpoons H_3O^+(aq) + Cl^-(aq) $
$ HCl $ is the acid, $ Cl^- $ is its conjugate base, $ H_2O $ is the base, and $ H_3O^+ $ is its conjugate acid.
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Amphoteric Substances: Substances that can act as both acids and bases. Water is a classic example:
$ H_2O(l) + H_2O(l) \rightleftharpoons H_3O^+(aq) + OH^-(aq) $
In this autoionization of water, one water molecule acts as an acid (donating a proton), and the other acts as a base (accepting a proton).
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Strong vs. Weak Acids/Bases: This relates to the extent of proton donation/acceptance. Strong acids/bases completely dissociate in water, while weak acids/bases only partially dissociate. Acid and Base Strength
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** $ K_a $ and $ K_b $ :** The acid dissociation constant ( $ K_a $ ) and base dissociation constant ( $ K_b $ ) quantify the strength of weak acids and bases, respectively. A larger $ K_a $ or $ K_b $ indicates a stronger acid or base. Equilibrium Constants
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pH and pOH: These scales measure the acidity and basicity of a solution, respectively. They are related to the concentration of $ H_3O^+ $ and $ OH^- $ ions. $ pH = -\log_{10}[H_3O^+]] $ and $ pOH = -\log_{10}[OH^-]] $ . At 25°C, $ pH + pOH = 14 $ . pH and pOH Calculations
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Titrations: Titrations are used to determine the concentration of an unknown acid or base by reacting it with a solution of known concentration. Acid-Base Titrations
Example Reaction:
The reaction between ammonia ( $ NH_3 $ ) and water:
$$ NH_3(aq) + H_2O(l) \rightleftharpoons NH_4^+(aq) + OH^-(aq) $$
In this reaction, $ NH_3 $ acts as a base (accepting a proton from water), and $ H_2O $ acts as an acid (donating a proton). $ NH_4^+ $ is the conjugate acid of $ NH_3 $ , and $ OH^- $ is the conjugate base of $ H_2O $ .
Related Notes: