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Defining Acids and Bases
The Arrhenius Definition (Outdated):
Acids donate hydrogen ions (H⁺) into aqueous solutions, and bases donate hydroxide ions (OH⁻) into aqueous solutions.
This model describes a large portion of acid-base chemistry accurately, but the definitions themselves aren't accurate. It is because the model only considers hydroxide as a base and applies exclusively to aqueous systems.
The Brønsted–Lowry Definition:
Acids donate protons (H⁺), while bases accept protons.
This definition is accurate for describing proton transfer, as it defines the fundamental behavior of acids and bases during reactions.
Conjugates of the Brønsted–Lowry Model:
Conjugate Base (of an Acid): What is left of the compound after giving away a proton. Conjugate bases can accept protons to revert back into an acid.
Conjugate Acid (of a Base): The compound after it accepts a proton. Conjugate acids can give away protons
MODEL EXAMPLE: HA + B ⇌ A⁻ + HB⁺
The acid is HA, as it gave away a proton during the reaction.
The base is B, as it accepted the proton during the reaction.
A- is the conjugate base.
HB+ is the conjugate acid.
Amphoterism: The ability to react as either an acid or a base.
Example: 2H₂O (l) ⇌ H₃O⁺ (aq) + OH⁻ (aq)
During the self-ionization of water, one water molecule donates a proton, while another water molecule accepts a proton.
The Lewis Definition (Not Relevant Here):
Bases donate electron pairs to form covalent bonds. Anions and neutral chemicals with lone pairs are typically capable to donate electron pairs.
Acids accept electron pairs. Cations, atoms capable of hypervalency, and atoms with incomplete octets are typically capable to accept electron pairs.
This definition focuses on electron interactions rather than proton transfer and is commonly applied in higher-level chemistry.
Strong and Weak Acids/Bases:
Strong acids and bases dissociate completely due to their weaker bonds and more chemically-stable conjugates. Equations involving strong acids/bases are typically written with a forward arrow (→) because the reactions are shifted heavily rightwards.
Example of Strong Acid Dissociation: HNO₃ (aq) → H⁺ (aq) + NO₃⁻ (aq)
Example of Strong Base Dissociation: NaOH (aq) → Na⁺ (aq) + OH⁻ (aq)
Weak acids do not dissociate completely, and their dissociation process is reversible. Most weak bases do not dissociate, and instead directly react with water molecules to take their protons. Equations involving weak acids/bases are typically written with an equilibrium arrow (⇌).
Example of Weak Acid Dissociation: CH₃COOH (aq) ⇌ H⁺ (aq) + CH₃COO⁻ (aq)
Example of Weak Base Dissociation: NH₃ (aq) + H₂O (l) ⇌ NH₄⁺ (aq) + OH⁻ (aq)
AP Chemistry Standards:
A list of strong acids and bases that the curriculum requires you to know.
Strong Acids:
HCl (Hydrochloric acid)
HBr (Hydrobromic acid)
HI (Hydroiodic acid)
HNO₃ (Nitric acid)
H₂SO₄ (Sulfuric acid – only the first proton is strong)
HClO₄ (Perchloric acid)
Strong Bases:
Group 1 Hydroxides:
LiOH (Lithium hydroxide)
NaOH (Sodium hydroxide)
KOH (Potassium hydroxide)
RbOH (Rubidium hydroxide)
CsOH (Cesium hydroxide)
Group 2 Hydroxides:
Ca(OH)₂ (Calcium hydroxide)
Sr(OH)₂ (Strontium hydroxide)
Ba(OH)₂ (Barium hydroxide)
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