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Acids and Bases

Acids and Bases. Lactic acid. Citric acid. Stearic acid. Ethanoic acid. Acetylsailicylic Acid. Common household acids. Common laboratory acids. Hydrochloric acid - HCl Nitric acid - HNO 3 Sulfuric acid - H 2 SO 4 Phosphoric acid - H 3 PO 4. Arrhenius theory of acid.

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Acids and Bases

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  1. Acids and Bases

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

  3. Common laboratory acids Hydrochloric acid - HCl Nitric acid - HNO3 Sulfuric acid - H2SO4 Phosphoric acid - H3PO4

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

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

  6. Acids • HCl and HNO3 are monobasic acids as they donate one H+ ion. HNO3 H+ + NO3- • H2SO4 is a dibasic acid as it donates two H+ ions. H2SO4 2H+ + SO42- • H3PO4 is a tribasic acid as it donates three H+ ions. H3PO4 3H+ + PO43-

  7. A strong acid is one which dissociates fully in water Example: HCl, H2SO4, HNO3 HCl + H2O H3O+ + Cl- • A weak acid is one which does not fully dissociate in water Example: CH3COOH (ethanoic acid) CH3COOH + H2O H3O+ + CH3COO-

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

  9. Common laboratory bases Sodium hydroxide - NaOH Calcium hydroxide - Ca(OH)2 Ammonia - NH3 Sodium carbonate - Na2CO3

  10. 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-

  11. 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

  12. Arrhenius theory • Combining: HA H+ + A- XOH X+ + OH- we get: HA + XOH AX + H2O acid + base salt + water

  13. 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

  14. 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+ + H2O H3O+ • The H3O+ ion is called the hydronium ion. • This is another limitation of the Arrhenius theory.

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

  16. 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.

  17. Donates a Proton Accepts a Proton Acid = Proton Donor HCl + H2O H3O+ + Cl- • The HCl donates a proton and so is an acid • The H2O, in this case, accepts a proton and so is a base Remember: Proton = H+

  18. Likewise: • HNO3 + H2O H3O+ + NO3- and • H2SO4 + H2O H3O+ + HSO4- • HSO4- + H2O H3O+ + SO4-2

  19. Accepts a proton Donates a proton Base = Proton Acceptor NH3 + H2O NH4+ + OH- • The NH3 accepts a proton and so is a base. • The H2O, in this case, donates a proton and so is an acid.

  20. 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.

  21. Acid – Donates Protons Base – Accepts Protons Acid – Base Reaction HCl + NH3 Cl- + NH4+

  22. 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)

  23. Neutralisation Acid + Base Salt + Water HCl + NaOH NaCl + H2O but since the acid and base dissociate in water we can write: H+ + Cl- + Na+ + OH- Na+ + Cl- + H2O we can cancel the Na+ and Cl- on both sides leaving: H+ + OH- H2O

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

  25. 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

  26. 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: CH3COOH + H2O CH3COO- + H3O+ Acid Conjugate Base

  27. Likewise: • When a base accepts a proton, it becomes its conjugate acid. Example: NH3 + H2O NH4+ + OH- Base Conjugate Acid

  28. Examples: Conjugate Base Acid H2SO4 + H2O HSO4- + H3O+ NH3 + H2S NH4+ + HS- Conjugate Acid Base Conjugate Base Acid Base Conjugate Acid

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