CHEMICAL CHANGE: ACIDS AND BASES Flashcards
Acid-base indicator
A substance that can act as either an acid or a base.
Arrhenius theory
An acid is a substance that produces hydrogen ions (H+)/ hydronium ions (H3O+) when it dissolves in water. A base is a substance that produces hydroxide ions (OH-) when it dissolves in water.
Amphiprotic substance/ampholyte
A substance that can act as either an acid or a base.
Auto-ionisation of water
A reaction in which water reacts with itself to form ions (hydronium ions and hydroxide ions).
Concentrated acids/bases
Contain a large amount (number of moles) of acid/base in proportion to the volume of water.
Conjugate acid-base pair
A pair of compounds or ions that differ by the presence of one H+ ion.
Conjugate acid and base
A conjugate acid has one H+ ion more than its conjugate base.
Dilute acids/bases
Contain a small amount (number of moles) of acid/base in proportion to the volume of water.
Diprotic acid
An acid that can donate two protons. Example: H2SO4
Dissociation
The process in which ionic compounds split into ions.
Endpoint
The point in a titration where the indicator changes colour.
Equivalence point
The point in a reaction where equivalent amounts of acid and base have reacted completely.
Hydrolysis
The reaction of a salt with water. OR The reaction of an ion with water to produce the conjugate acid and a hydroxide ion or the conjugate base and a hydronium ion.
Ionisation
The process in which ions are formed during a chemical reaction.
Ion product of water
The product of the ions formed during auto-ionisation of water i.e. [H3O+][OH–] at 25 °C.
Ionisation constant of water (Kw)
The equilibrium value of the ion product [H3O+][OH–] at 25 °C.
Ka value
Ionisation constant for an acid.
Kb value
Dissociation or ionisation constant for a base.
Lowry-Brønsted theory
An acid is a proton (H+ ion) donor.
A base is a proton (H+ ion) acceptor.
Monoprotic acid
An acid that can donate one proton. Example: HCℓ
Neutralisation
The reaction of an acid with a base to form a salt (ionic compound) and water.
pH
The negative of the logarithm of the hydronium ion concentration in mol·dm-3.
In symbols: pH = -log[H3O+]
Unit: None
pH scale
A scale from 0 – 14 used as a measure of the acidity and basicity of solutions where
pH = 7 is neutral,
pH > 7 is basic and
pH < 7 is acidic.
Salt
The ionic compound that is the product of a neutralisation reaction.
Standard Solution
A solution of precisely known concentration.
Strong bases
Dissociate COMPLETELY in water to form a high concentration of OH- ions. Examples: sodium hydroxide (NaOH) and potassium hydroxide (KOH)
Strong acids
Ionise completely in water to form a high concentration of H3O+ ions. Examples: hydrochloric acid (HCℓ), sulphuric acid (H2SO4) and nitric acid (HNO3)
Titration
The procedure for determining the amount of acid (or base) in a solution by determining the volume of base (or acid) of known concentration that will completely react with it.
Weak acids
Ionise INCOMPLETELY in water to form a low concentration of H3O+ ions. Examples: ethanoic acid (CH3COOH) and oxalic acid (COOH)2
Weak bases
Dissociate/ionise incompletely in water to form a low concentration of OH- ions. Examples: ammonia (NH3), sodium hydrogen carbonate (NaHCO3), sodium carbonate (Na2CO3), potassium carbonate (K2CO3), calcium carbonate (CaCO3)
