Classifying acids and bases

Around 1923, two scientists called Lowry and Bronsted defined acids and bases in the following way:

They defined acids as proton or H⁺ donors.
They saw bases as proton or H⁺acceptors.

Around 1938 another scientist called Lewis extended the definition of acids to include any species that accepts a lone pair of electrons. The reason he extended the definition was because the Lowry-Bronsted theory excluded substances that had no protons to donate in the first place. According to Lewis a base is a species that can donate a lone pair of electrons.

A substance that can act as either an acid or a base is called amphoteric. Examples of such substances are zinc hydroxide and aluminium hydroxide.

E.g.1

Al(OH)₃ + NaOH NaAlO₂ + 2H₂O

(acidic) (base) (sodium

aluminate)

Al(OH)₃ + 3HCl AlCl₃+ 3H₂O

(basic) (acid) (aluminium

chloride)

E.g.2

Zn(OH)₂+ 2NaOH Na₂ZnO₂ + 2H₂O

(acidic) (base) (sodium

zincate)

Zn(OH)₂ + 2HCl ZnCl₂ + 2H₂O

(basic) (acid) (zinc

chloride)

Some Common Reactions

ACID + ALKALI SALT + WATER

E.g.1: HCl_(aq) + NaOH_(aq) NaCl_(aq) + H₂O_(l)

Hydrochloric + Sodium Sodium + Water

Acid Hydroxide Chloride

E.g.2: 2HCl_(aq) + Ca(OH)_2(aq) CaCl_2(aq) + 2H₂O_(l)

Hydrochloric + Calcium Calcium + Water

Acid Hydroxide Chloride

E.g.3: HCl_(aq) + NH_3(aq) NH₄Cl_(aq)

Hydrochloric + Ammonia Ammonium

Acid Chloride

(HCl_(aq) + NH₄OH_(aq) NH₄Cl_(aq) + H₂O_(l))

(Hydrochloric + Ammonium Ammonium

Acid Hydroxide Chloride + Water)

E.g.4: HNO_3(aq) + NaOH_(aq)NaNO_3(aq) + H₂O_(l)

Nitric + Sodium Sodium + Water

Acid Hydroxide Nitrate

E.g.5: 2HNO₃ + Ca(OH)_2(aq) Ca(NO₃)_2(aq) + 2H₂O_(l)

Nitric + Calcium Calcium + Water

Acid Hydroxide Nitrate

E.g.6: HNO_3(aq) + NH_3(aq) NH₄NO_3(aq)

Nitric + Ammonia Ammonium

Acid Nitrate

(HNO_3(aq) + NH₄OH_(aq) NH₄NO_3(aq) + 2H₂0_(l))

(Nitric + Ammonium Ammonium +

Acid Hydroxide Nitrate Water)

E.g.7: H₂SO_4(aq)+ 2KOH_(aq) K₂SO_4(aq) + 2H₂0_(l)

Sulphuric + Potassium Potassium + Water

Acid Hydroxide Sulphate

E.g.8: H₂SO_4(aq)+ Ca(OH)_2(aq) CaSO_4(aq) + 2H₂O_(l)

Sulphuric + Calcium Calcium + Water

Acid Hydroxide Sulphate*

*Calcium Sulphate is only slightly soluble in water. This is not a desirable demonstration reaction.

ACID + BASE SALT + WATER

E.g.1: 2HNO_3(aq) + CuO_(s) Cu(NO₃)_2(aq)+ H₂O_(l)

Nitric + Copper Copper + Water

Acid Oxide Nitrate

E.g.2: H₂SO_4(aq) + MgO_(s) MgSO_4(aq) + H₂O_(l)

Sulphuric + Magnesium Magnesium + Water

Acid Oxide Sulphate

METAL + ACID SALT + HYDROGEN

E.g.1: Mg_(s) + 2HCl_(aq) MgCl_2(aq)+ H_2(g)

Magnesium + Hydrochloric Magnesium + Hydrogen

metal Acid Chloride

E.g.2: Ca_(s) + 2HNO_3(aq) Ca(NO₃)_2(aq) + H_2(g)

Calcium + Nitric Calcium + Hydrogen

metal Acid Nitrate

Note that metals above hydrogen in the reactivity series displace hydrogen from dilute acids.

Neutralisation

Any reaction in which an acid reacts with a base is a neutralisation reaction. A salt and water are formed.

*See above examples

H⁺_(aq) + OH^-_(aq) H₂O_(l)

Exact neutralisation reactions are carried out by a process called titration.

Acid/Base Indicators

Litmus:

‘Acid gives red,

Alkali, blue,

If you use litmus,

This rhyme is true!’

¨ When testing for acidity use blue litmus paper or solution.

¨ When testing for alkalinity use red litmus paper or solution.

Universal Indicator:

It gives us a whole range of colours that help us to see how acidic or alkaline a solution is.

pH number: 0 – 2: strongly acidic red

3 – 6: weakly acidic red-orange/orange/yellow

7: neutral green

8 – 12: weakly alkaline blue/deep-blue

12 – 14: strongly alkaline violet

By Clare O' Connell: Loreto College, Coleraine