1. Acids and Bases
By: Clarissa

2. THE PH SCALE
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3. pH scale
• pH is equal to -log10 of the hydrogen ion
concentration
– Equal to the power of ten of the hydrogen ion
concentration
• pH 7 is neutral
– If solution is acidic, the H+ will be >10-7 mol dm-3
 pH decrease
– Basic solutions will have a pH >7
H+ will be <10-7 mol dm-3
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4. • pH scale runs from 0 to 14
• A change in one unite in the pH
– Corresponds to a tenfold change H+
• 0.1 mol dm-3 solution of a strong
monoprotic acid will have pH of 1
• 0.001 mol dm-3 solution of the same
acid will have a pH of 3
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5. Determination of pH
• Using pH meter “universal” indicator
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6. Question
• pH value of a solution rises from 4 to 7.
What is the change in concentration of
hydrogen ions in it?
– Increase by factor 3
– Increase by factor 1000
– Decrease by factor 3
– Decrease by factor 1000
• Decrease by factor 1000 #

7. Question
• 10cm3 of an HCl solution with a pH
value of 2 was mixed with 90cm3 of
water. What will be the pH of the
resulting solution?
• Ratio 1:10
– pH 3
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8. THEORIES OF ACIDS AND
BASES AND SALT
HYDROLYSIS
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9. The Ionic Theory
• Acid: oxide of a non-metal combined with
water
• Substance which produces hydrogen
ions, H+ (aq)
• HCl  H+(aq) + Cl-(aq)
• Base: substance that can neutralize acid
• Alkali is a base that is soluble in water
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10. Bronsted-Lowry Acid and Bases
• Bronsted-Lowry Acid can donate a proton
• Bronsted-Lowry Base can accept a proton
– Must contain a non-bonding pair of electrons
• Conjugate base of an acid is the species remaining
after the acid has lost a proton
• Conjugate acid is the species formed after the base
has accepted a proton
• HCl + H2O ↔ H3O+ + Cl-
– Water is conjugate base of H3O+ and Cl- is conjugate
base of HCl #

11. Bronsted-Lowry Acid and
Bases
• Water can act both as acid and base 
amphiprotic
• Strong acids have weak conjugate base
• Weak acids have strong conjugate base
– ACID: Ethanol (C2H5OH) weak
– CONJUGATE BASE: Ethoxide ion (C2H5O-) 
strong
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12. Lewis Acids and Bases
• Bases can donate a pair of electrons
• Acids can accept a pair of electrons
– Covalent bond is formed
• Many Lewis Acids don’t contain
hydrogen
– BF3: Lewis acid, only six electrons so
leaves room for two more
– Pair up with NH3
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13. PROPERTIES OF ACIDS AND
BASES
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14. Properties of Acids and Bases
• Neutralization reaction with bases
– With hydroxides to form salt and water
• CH3COOH + NaOH  NaCH3COO + H2O
– With metal oxides to form a salt and water
• H2SO4 + CuO  CuSO4 + H2O
– With ammonia to form salt
• HCl + NH3  NH4Cl
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15. Properties of Acids and Bases
• With reactive metals to form salt and hydrogen
– 2HCl + Mg  MgCl2 + H2
• With carbonates to form salt, CO2, water
– 2HCl + CaCO3  CaCl2 + CO2 + H2O
• With hydrogencarbonates to form salt, CO2, water
– HCl + NaHCO3  NaCl + CO2 + H2O
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16. Strong Acids and Bases
• Strong acid and base is completely dissociated
(ionized) into its ions in aqueous solutions
• Strong acids
– HCl (Hydrochloric acid)
– HNO3 (Nitric Acid)
– H2SO4 (Sulfuric Acid)
• Strong bases
– NaOH (Sodium Hydroxide)
– KOH (Potassium Hydroxide)
– Ba(OH)2 (Barium Hydroxide)
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17. Strong Acids and Bases
• Monoprotic acid
– One mole of HCl produces one mole of
Hydrogen ions
• Diprotic acid
– One mole of H2SO4 produces two moles of
Hydrogen ions
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18. Weak Acids and Bases
• Weak acids and bases is slightly
dissociated into their ions in aqueous
solution
• Weak acids
– CH3COOH (Ethanoic Acid)
– H2CO3(Carbonic acid)
• Weak bases
– NH3 (Ammonia)
– C2H5NH2 (Aminoethane)
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19. The Differences
• Can be seen in the reaction with water:
– Strong Acid
• HCl + H2O  H3O+ + Cl- (strong)
– Reaction goes to completion
– Weak Acid
• CH3COOH + H2O ↔ CH3COO- + H3O+ (weak)
– Equilibrium lies on left
 A solution of HCl consist only of hydrogen ions and
chloride ions in water
 Solution of Ethanoic Acid contains mainly
undissociated ethanoic acid with only very few
hydrogen and ethanoate ions
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20. The Differences
– Strong Base
• KOH  K+ + OH (strong)
– Weak Base
• NH3 + H2O ↔ NH4+ + OH- (weak)
– Equilibrium lies on the left
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21. Distinguish Between Strong & Weak
Acid & Bases
• Strong acids has lower pH
• Strong base has higher pH
• Strong: more positively charged
• Strong: larger atomic radius
• Strong: higher electronegativity
• Strong: have much higher conductivity
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22. Question
• Which of the following statements about
aqueous solutions of most weak acids is
correct?
a. They react with carbonates to produce
CO2
b. They conduct electricity better than strong
acid
• A. Because strong acid conducts
electricity better #

23. Question
• Ammonia (weak base) + water = alkaline
Why?
• H2O + NH3  NH4OH
Dissociation of ammonia to produce hydroxide
This will have smaller pH because hydroxide is
stronger
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24. Question
• What mass of ethanoic acid would you use
to prepare 0.5dm3 of a 0.5mol dm-3 ethanoic
acid solution? (Mt ethanoic acid = 60.0)
• 0.5 x 0.5 x 60 = 15 grams
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25. Question
• Which will be the same for separate
1mol dm-3 solutions of a strong acid
and a weak acid?
– Electrical conductivity
– Concentration of H+ ions
• Neither
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26. BUFFER SOLUTION
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27. Buffer Solution
• Resist change in pH
– When small amounts of acid or base are
added to it
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28. Buffer Solution
• Acidic buffer solution
– Made by mixing a strong base and excess of
weak acid so the solution contains salt and
unreacted weak acid
– Example:
• NaOH + CH3COOH  NaCH3COOH + H2O + CH3COOH
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29. Buffer Solution
• Basic buffer solution
– Made from a weak base with the salt of the
base with a strong acid
– Example:
• NH4Cl  NH4+ + Cl-
• NH3 + H2O  NH4+ + OH-
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30. Question
• Will these solutions form a buffer solution?
– 1.0dm3 solution containing 0.1mol NH3 and
0.2mol HCl
• No
– After reaction: 0.1 mol NH4Cl + 0.1 mol HCl
– Weak acid and strong acid  strongly acidic solution
– Not a weak acid-conjugate base combination
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31. Question
• Will these solutions form a buffer
solution?
– 1.0dm3 solution containing 0.2mol NH3
and 0.1mol HCl
• Yes
– After reaction: 0.1mol NH3 + 0.1mol NH4Cl
– Weak acid and conjugate base
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32. THANK YOU 
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