Acids and Bases Lactic acid Citric acid Common

Acids and Bases

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

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

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

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

Acids n HCl and HNO 3 are monobasic acids as they donate one H+ ion. HNO 3 H+ + NO 3 - n H 2 SO 4 is a dibasic acid as it donates two H+ ions. H 2 SO 4 2 H+ + SO 42 - n H 3 PO 4 is a tribasic acid as it donates three H+ ions. H 3 PO 4 3 H+ + PO 43 -

n 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- n 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-

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

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

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

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

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

Limitations of Arrhenius theory 1. 2. 3. 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 n n 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 n 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 n 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 Donates a Proton HCl + H 2 O H 3 O+ + Cl- Accepts a Proton n n 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+

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

Base = Proton Acceptor Accepts a proton NH 3 + H 2 O NH 4+ + OHDonates a proton n n 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.

Amphoteric n n 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 Acid – Donates Protons HCl + NH 3 Cl- + NH 4+ 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 + Na. OH Na. Cl + H 2 O but since the acid and base dissociate in water we can write: H+ + Cl- + Na+ + OHNa+ + Cl- + H 2 O we cancel the Na+ and Cl- on both sides leaving: H+ + OHH 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 Nettle, bee and ant stings are acidic They can be neutralised with vinegar or lemon juice They can be neutralised with baking soda

Conjugate Acid-Base Pairs n n 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 Acid CH 3 COO- + H 3 O+ Conjugate Base

Likewise: n When a base accepts a proton, it becomes its conjugate acid. Example: NH 3 + H 2 O Base NH 4+ + OH- Conjugate Acid

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