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.

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