TOPIC 8 : Acids and Bases
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This document discusses the properties and reactions of acids and bases. It states that acids have a sour taste and cause color changes in indicators like litmus. Acids react with metals to produce hydrogen gas and with carbonates to produce carbon dioxide gas. Bases have a bitter taste and feel slippery. Strong acids and bases fully dissociate in water while weak acids and bases only partially dissociate. The Bronsted-Lowry theory defines acids as proton donors and bases as proton acceptors. Water can act as both an acid and a base depending on the reaction.
TOPIC 8 : Acids and Bases
- 2. Properties of acids and bases • Acids have a sour taste : e.g vinegar owes its taste to acetic acid, and lemons and other citrus fruits contain citric acids. GENERAL PROPRTIES Acids:
- 3. • Acids cause colour changes in plant dyes • E.g change the colour of litmus from blue to red. • Acid react with certain metals e.g zinc, magnesium, and iron to produce hydrogen gas.
- 4. • E.g: reaction between hydrochloric acid and magnesium: • 2HCl (aq) + Mg (s) MgCl2 (aq) + H2 (g)
- 5. • Acids react with carbonates and bicarbonates such as Na2CO3, CaCO3, and NaHCO3 to produce carbon dioxide gas. • Example : • 2HCl (aq) + CaCO3 (s) CaCl2 (aq) + H2O (l) + CO2 (g) • HCl (aq) + NaHCO3 (s) NaCl (aq) + H2O (l) + CO2 (g) • Aqueous acids solutions conduct electricity.
- 6. • Bases have a bitter taste. • Bases feel slippery; soaps which contains bases, exhibit this property. • Bases cause colour changes in plant dyes e.g change the colour of litmus from red to blue. • Aqueous base solutions conduct electricity. Bases
- 7. • Ammonia, NH3 • Soluble carbonates, CaCO3 • Hydrogencarbonates, NaHCO3 Bases which are not hydroxide Alkalis – bases that dissolve in water
- 8. • Strong acids are all strong electrolytes that ionize completely in water (dissociate) . • Example ; hydrochloric acids, HCl nitric acids, HNO3 Perchloric acids, HClO4 Sulphuric acids, H2SO4
- 9. • HCl (aq) + H2O (l) H3O+ (aq) + Cl- (aq) • HNO3 (aq) + H2O (l) H3O+ (aq) + NO3 - (aq) • HClO4 (aq) + H2O (l) H3O+ (aq) + ClO4 - (aq) • H2SO4 (aq) + H2O (l) H3O+ (aq) + HSO4 - (aq)
- 10. • Strong bases are all strong electrolytes that ionize (dissociates) completely in solution. • Example; NaOH (aq) Na+ (aq) + OH- (aq) KOH (aq) K+ (aq) + OH- (aq) Ba(OH)2 (aq) Ba2+ (aq) + 2OH- (aq)
- 11. • Weak acids are acids that ionize only to limited extent in water (partially). • Example ; • At equilibrium, aqueous solutions of weak acids contain mixture of nonionized acid molecule, H3O+ ion and conjugate base.
- 12. • Note : the strength of acid can vary greatly due to differences in extent of ionization. • The limited ionization of weak acids is related to the equilibrium constant for ionization, Ka.
- 13. • Weak bases are bases that ionize only to a limited extent in water. • Example ; • At equilibrium, there is a mixture of nonionized , and ions.
- 17. Bronsted – Lowrey • Acids – proton donors • Bases – proton acceptor • Example ; • HCl (g) + H2O (l) H3O+ (aq) + Cl- (g) ? ?
- 18. • An acids becomes its conjugate base when it donates a proton. • A base becomes its conjugate acid when it accepts a proton.
- 19. • HCl donates its proton completely to H2O. • For a strong acid, the reverse reaction does not occur. • Thus, Cl- ion is a weak conjugate base. • So, strong acids form weak conjugate bases. • Weak acids form strong conjugate bases.
- 20. • NH3 (aq) + H2O (l) ⇌ NH4 + (aq) + OH- (aq) (base) (acid) ( ? ) (? ) NH3 – base (accepts a proton from H2O) H2O – Bronsted – Lowrey Acid EQUATION IS REVERSED ; NH 4 + acid OH- base
- 21. NH3 • Weak base • Does not accept the proton from H2O completely. NH4 + ion • strong conjugate acid H2O • A weak acid OH- • Strong conjugate base
- 22. When the equation reversed ;
- 24. • The relative strength of an acid and its conjugate base, and a base and its conjugate acid. • Explain why a strong acid or a strong base dissociates completely in aqueous solution? • Explain why weak acid or a weak base dissociates partially in aqueous solution?
- 25. • A strong acid completely breaks apart to give ions in solution (100% dissociation) whereas a weak acid only slightly dissociates in solution (perhaps less than 1%). • A strong acid, when placed in water, will almost fully ionise/dissociate straight away, producing H+ (aq) ions from water. • A weak acid will, however, only partially dissociate into ions, leaving a high percentage of unreacted molecules in the solution.
- 26. H2O amphoteric • acting as a base in the presence of an acid • acting as an acid in the presence of a base H2O + H2O ⇌ H3O+ + OH- (acid) (base)
- 27. Bronsted – Lowrey acids and bases are NOT limited to reactions with water • Example ; • HCl (g) + NH3 (g) ⇌ NH4 + + Cl- (acid) (base) (conjugate acid) (conjugate base)
- 28. Question ; Is it NaOH a Bronsted – Lowrey base ? NaOH NOT a Bronsted – Lowrey base because it does not accept a proton
- 29. • Identify Bronsted – Lowrey acid, Bronsted – Lowrey base, conjugate acid, or conjugate base, from each of the following equation. • H2CO3 + H2O ⇌ HCO3 - + H3O+ • NH4 + + H2O ⇌ NH3 + H3O+ • CH3NH2 + H2O ⇌ CH3NH3 + + OH- • CH3COOH + H2O ⇌ CH3COO- + H3O+
- 30. • Lewis acid is an atom, ion or molecule that accepts a pair of electrons to form a coordinate covalent bond. • Lewis base is an atom, ion or molecule that donates a pair of electron to form a coordinate covalent bond.
- 31. Example ; Lewis base Lewis acid
- 34. • NH3 is a Lewis base • Donates a pair of electrons to the Lewis acid H+ when it forms NH4 + ions
- 35. • OH- Lewis base – donates a pair of electrons to the Lewis acid H+ to form H2O molecule
- 36. For the following reactions, identify the Lewis acids and the Lewis base:
- 37. Lewis base Lewis acid