Acids + Bases

• Theories of acids and bases include Arrhenius (grade 10 de initiation), Bronsted-Lowry (IUPAC de initiation), and Lewis.
• In Arrhenius theory, an acid is a molecular compound that ionizes in water, for example, binary acids like HCl and H2SO4.
• In Bronsted-Lowry theory, a base is an ionic compound that dissociates in water, for example, metal hydrides like NaOH and Ca(OH)2.
• In Lewis theory, an acid is a substance that accepts electrons and a base is a substance that donates electrons.
• Some acids and bases may have more than one hydrogen ion.
• If only one acid, H+, is generated, the acid is monoprotic and can produce more than one acid.
• If only one base, OH-, is generated, the base is monobasic and can produce more than one base.
• Acids and Bases are defined as substances that donate or accept a proton.
• Acids are substances that donate a proton.
• Bases are substances that accept a proton.
• Neutralization is the process where an acid and a base react to form a water molecule and a salt.
• pH is a measure of the acidity or basicity of a solution.
• Periodicity 3 is the subject of this chapter.
• pH is measured on a scale from 0 to 14.
• A pH of 0 is the most acidic, while a pH of 14 is the most basic.
• The pH of a solution can be adjusted by adding or removing acid or base.
• Examples of weak acids include organic acids such as CH3COOH, H2CO3.
• ● E x : o r g a n i c a c i d s ( e x : C H 3 C O O H ) , H 2 C O 3 ( a q )
• Examples of strong acids include HCl, H2SO4, HNO3.
• Weak acids partially dissociate into ions HA (aq) + H2O (l) ⇌ H3O+ (aq) + A− (aq).
• Strong acids completely dissociate into ions HA (l) + H2O (l) → H3O+ (aq) + A− (aq).
• Acids and bases can be classified as strong or weak based on their dissociation constant.
• To determine the [H+] and [OH-] of the solution, use the equation pH = -log[H+], where pH is the negative logarithm of [H+].
• The reaction between an acid and a base is exothermic, meaning that the heat is released and the change in energy is negative.
• An aqueous solution has a pH of 3.20 at 298 K.
• The pH of an aqueous solution can be calculated using the formula pH = -log[H+], where pH is the negative logarithm of [H+].
• The ionic product of water is the result of auto-ionisation of water, where water molecules dissociate to a very small extent H2O (l) ⇌ H+ (aq) + OH– (aq).
• The ionic product of water is represented by the equation Kw = [H+][OH-], and at 298 K, Kw = 1.00x10^-14.
• Salt is produced in the reaction between an acid and a base.
• General equation: alkali plus acid results in salt and water, represented by the equation: Alkali + Acid → Salt + Water.
• The products of the reaction between a hydrochloric acid and lithium hydroxide are hydrochloric acid and sodium carbonate.
• In a neutralisation reaction, an acid reacts with a base to produce a salt and water.
• Ionic equation: Net ionic equation: Neutralisation reactions involve an acid reacting with a base to produce a salt and water.
• The products of the reaction between an acid and a base depend on the parent acid and base.
• The products of the reaction between a nitric acid and calcium hydroxide are nitric acid and calcium nitrate.
• Strong acids have a dissociation constant greater than 1, while strong bases have a dissociation constant less than 1.
• Weak acids have a dissociation constant between 1 and 10, while weak bases have a dissociation constant between 1 and 10.
• pH = -log[H+ (aq)].
• [OH- (aq)] = 10^-pOH.
• pOH = -log[OH- (aq)].