Acids and Bases

Common household acids Citric acid Ethanoic acid Lactic acid Stearic acid Acetylsailicylic Acid

Common laboratory acids Hydrochloric acid -HCl Hydrochloric acid -HCl Nitric acid -HNO 3 Nitric acid -HNO 3 Sulfuric acid-H 2 SO 4 Sulfuric acid-H 2 SO 4 Phosphoric acid-H 3 PO 4 Phosphoric acid-H 3 PO 4

Arrhenius theory of acid Arrhenius was a Sweedish Arrhenius was a Sweedishchemist Put forward a theory of Put forward a theory of acids in the 1880’s Stated that: Stated that: An acid is a substance that dissociates in water to form H + ions.

Arrhenius theory of acid For example: when HCl is added to water: HCl H + + Cl - In general: HA H + + A -

Acids HCl and HNO 3 are monobasic acids as they donate one H + ion. HNO 3 H + + NO 3 - H 2 SO 4 is a dibasic acid as it donates two H + ions. H 2 SO 4 2H + + SO 4 2- H 3 PO 4 is a tribasic acid as it donates three H + ions. H 3 PO 4 3H + + PO 4 3-

A strong acid is one which dissociates fully in water Example: HCl, H 2 SO 4, HNO 3 HCl + H 2 O H 3 O + + Cl - A weak acid is one which does not fully dissociate in water Example: CH 3 COOH (ethanoic acid) CH 3 COOH + H 2 O H 3 O + + CH 3 COO -

Common household bases Magnesium hydroxide Ammonia Sodium hydroxide Sodium hydrogen carbonate Calcium hydroxide

Common laboratory bases Sodium hydroxide-NaOH Sodium hydroxide-NaOH Calcium hydroxide-Ca(OH) 2 Calcium hydroxide-Ca(OH) 2 Ammonia-NH 3 Ammonia-NH 3 Sodium carbonate-Na 2 CO 3 Sodium carbonate-Na 2 CO 3

Arrhenius theory of bases Arrhenius defined a base as: A substance that dissociates in water to produce OH - ions. For example: when NaOH is added to water: NaOH Na + + OH - In general: XOH X + + OH -

A strong base is one which dissociates fully in water Example: NaOH A weak base is one which does not fully dissociate in water Example: Mg(OH) 2

Arrhenius theory Combining: Combining: HA H + + A - XOH X + + OH - we get: HA + XOH AX + H 2 O acid + base salt + water

Limitations of Arrhenius theory The acids and bases must be in aqueous solutions (i.e. water). This prevents the use of other solvents benzene Not all acid – base reactions are in solution, e.g. ammonia gas and hydrogen chloride gas produce ammonium chloride According to Arrhenius, the salt produced should not be acidic or basic. This is not always the case, for example in the above reaction ammonium chloride is slightly acidic

Hydronium Ion Arrhenius thought that an acid gives off H + ions in solution. H + ions are protons and can not exist independently. When the acid dissociates, the H + ions react with water molecules: H + + H 2 O H 3 O + The H 3 O + ion is called the hydronium ion. This is another limitation of the Arrhenius theory.

Brønsted-Lowry Theory In 1923, Johannes Brønsted (a Danish chemist) and Thomas Lowry (an English chemist) proposed new definitions of acids and bases. Brønsted Lowry

Brønsted-Lowry Theory Brønsted and Lowry had worked independently of each other but they both arrived at the same definitions: An acid is a substance that donates protons (hydrogen ions). A base is a substance that accepts protons.

Acid = Proton Donor HCl + H 2 O H 3 O + + Cl - The HCl donates a proton and so is an acid The H 2 O, in this case, accepts a proton and so is a base Remember: Proton = H + Donates a Proton Accepts a Proton

Likewise: HNO 3 + H 2 O H 3 O + + NO 3 - and H 2 SO 4 + H 2 O H 3 O + + HSO 4 - HSO H 2 O H 3 O + + SO 4 -2

Base = Proton Acceptor NH 3 + H 2 O NH OH - The NH 3 accepts a proton and so is a base. The H 2 O, in this case, donates a proton and so is an acid. Accepts a proton Donates a proton

Amphoteric As can be seen from the previous two examples, water is capable of acting as both and acid and a base. Any substance that can act as both an acid and a base is said to be amphoteric.

Acid – Base Reaction HCl + NH 3 Cl - + NH 4 + Acid – Donates Protons Base – Accepts Protons

Neutralisation The reaction between an acid and a base to produce a salt and water A salt is formed when the hydrogen of an acid is replaced by a metal (or ammonium ion)

Neutralisation Acid + Base Salt + Water HCl + NaOH NaCl + H 2 O but since the acid and base dissociate in water we can write: H + + Cl - + Na + + OH - Na + + Cl - + H 2 O we can cancel the Na + and Cl - on both sides leaving: H + + OH - H 2 O

Everyday Examples of Neutralisation Indigestion remedies are bases that neutralise excess stomach acid Lime is a base that neutralises acid in soil Toothpaste is a base that neutralises acid in the mouth

Wasp stings are basic They can be neutralised with vinegar or lemon juice Nettle, bee and ant stings are acidic They can be neutralised with baking soda

Conjugate Acid-Base Pairs Acids and bases exist in pairs called conjugate acid-base pairs. Every time an acid donates/loses a proton, it becomes its conjugate base. Example: CH 3 COOH + H 2 O CH 3 COO - + H3O+H3O+ Acid Conjugate Base

Likewise: When a base accepts a proton, it becomes its conjugate acid. Example: NH 3 + H 2 O NH OH - Base Conjugate Acid

Examples: H 2 SO 4 + H 2 O HSO H 3 O + NH 3 + H 2 S NH HS - Acid Conjugate Base Conjugate Acid Base Conjugate Base Conjugate Acid