Acid and Bases Ib
Acids And BAses
Acids And BAses
8.1 8.2 8.3 8.4 18.1 18.2 18.3 18.4 18.5 Theories of acids and bases Properties of acids and bases Strong and weak acids & bases The pH scale Calculations involving acids and bases (AHL) Buffer solutions (AHL) Salt hydrolysis (AHL) Acid-base titrations (AHL) Indicators (AHL)
8
8.1 THeORies OF Acids And BAses
8.1.1 Define acids and bases according to the Brønsted–Lowry and Lewis theories. 8.1.2 Deduce whether or not a species could act as a Brønsted–Lowry and/or a Lewis acid or base. 8.1.3 Deduce the formula of the conjugate acid (or base) of any Brønsted–Lowry base (or acid). © IBO 2007 and oxonium ion). In these terms the above equation becomes
HX (aq) + H2O (l)
H3O+ (aq) + X– (aq)
One of the first theories to explain the fact that all acids had similar reactions, was that of Arrhenius. This proposed that in aqueous solution all acids, to some extent (dependent on the strength of the acid), split up to form a hydrogen ion and an anion, i.e. for an acid HX:
This also emphasises the fact that water is not an inert solvent, but is necessary for acid–base activity. Indeed solutions of acids in many non–aqueous solvents do not show acidic properties. For example a solution of hydrogen chloride in methylbenzene does not dissociate and hence, for example, it will not react with magnesium. Invoking the hydronium ion is useful in discussing some aspects of acid–base theory, such as conjugate acid–base pairs, but apart from this the simpler terminology of the hydrated proton/hydrogen ion, H+ (aq), will be adopted in this book. The similar reactions of acids can be explained as all being reactions of the hydrogen ion and it is perhaps more accurate to write them as ionic equations, for example the reaction of an aqueous acid with magnesium can be written as:
HX (aq)
H+ (aq) + X– (aq)
The hydrogen ion is hydrated, like all ions in aqueous solution, but some chemists prefer to show this reaction more explicitly
Acids And BAses
8.1 8.2 8.3 8.4 18.1 18.2 18.3 18.4 18.5 Theories of acids and bases Properties of acids and bases Strong and weak acids & bases The pH scale Calculations involving acids and bases (AHL) Buffer solutions (AHL) Salt hydrolysis (AHL) Acid-base titrations (AHL) Indicators (AHL)
8
8.1 THeORies OF Acids And BAses
8.1.1 Define acids and bases according to the Brønsted–Lowry and Lewis theories. 8.1.2 Deduce whether or not a species could act as a Brønsted–Lowry and/or a Lewis acid or base. 8.1.3 Deduce the formula of the conjugate acid (or base) of any Brønsted–Lowry base (or acid). © IBO 2007 and oxonium ion). In these terms the above equation becomes
HX (aq) + H2O (l)
H3O+ (aq) + X– (aq)
One of the first theories to explain the fact that all acids had similar reactions, was that of Arrhenius. This proposed that in aqueous solution all acids, to some extent (dependent on the strength of the acid), split up to form a hydrogen ion and an anion, i.e. for an acid HX:
This also emphasises the fact that water is not an inert solvent, but is necessary for acid–base activity. Indeed solutions of acids in many non–aqueous solvents do not show acidic properties. For example a solution of hydrogen chloride in methylbenzene does not dissociate and hence, for example, it will not react with magnesium. Invoking the hydronium ion is useful in discussing some aspects of acid–base theory, such as conjugate acid–base pairs, but apart from this the simpler terminology of the hydrated proton/hydrogen ion, H+ (aq), will be adopted in this book. The similar reactions of acids can be explained as all being reactions of the hydrogen ion and it is perhaps more accurate to write them as ionic equations, for example the reaction of an aqueous acid with magnesium can be written as:
HX (aq)
H+ (aq) + X– (aq)
The hydrogen ion is hydrated, like all ions in aqueous solution, but some chemists prefer to show this reaction more explicitly