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Unformatted text preview: CHAPTER 8 ACIDS AND BASES
(TOPICS 8 AND 18) SUMMARY
Characteristic Properties of Acids corrosive, taste sour, and turn blue litmus red.
Aqueous solutions of acids have pH < 7. Acids react with hydroxides and oxides to form salt
and water; Acids react with “active”* metals to form hydrogen gas; Acids react with
carbonates to produce carbon dioxide; Similarly acids react with hydrogen carbonates.
Characteristic Properties of bases: Bases feel slippery, taste bitter, and turn red litmus blue.
Aqueous solutions of bases have pH > 7. React with acids to form salt (and water if a
Alkalis are bases that dissolve in water, e.g., NaOH. Substances such as ammonia, NH3,
soluble carbonates (e.g., Na2CO3) and hydrogen carbonates (e.g., NaHCO3) are called bases;
undergo base hydrolysis to form OH−; With ethylamine, ethyl ammonium ion is formed:
Neutralization is the reaction of an acid and a base to give salt and water; in the salt, cations
(e.g., Na+) come from the base and anions (e.g., Cl–) come from the acid:
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
Brønsted-Lowry and Lewis theories
• A Brønsted-Lowry acid is a proton donor.
• A Brønsted-Lowry base is a proton acceptor.
• A Brønsted-Lowry acid-base reaction is characterised by proton transfer.
• When an acid transfers a proton, it produces a conjugate base.
• When a base accepts a proton, a conjugate acid is produced.
• The 2 related species in a proton transfer reaction, known as a conjugate acid-base pair,
differ by a proton.
HCl (aq) + CO32– (aq) HCO3– (aq) + Cl– (aq)
• Lewis acid: a species that can accept an electron pair. Lewis acids are cations or neutral
molecules with available empty orbitals, e.g., H+, Cl+, NO2+, CH3+, AlCl3, BF3. These are
called electrophiles. • Lewis base: a species that can donate an electron pair. Lewis bases are anions or neutral
molecules with lone e– pair(s), e.g., OH–, Cl–, H2O:, :NH3. These are called nucleophiles
(or nucleophilic reagents). Electrolytes are substances that dissolve in water to give solutions containing ions.
Strong Electrolyte is a substance that dissolves and completely dissociates into ions (ionises
completely). This is shown by a single arrow.
Weak Electrolyte is a substance that dissolves and dissociates into ions partially (to a limited
A strong acid ionizes or dissociates essentially completely in water, and Ka, the acid
dissociation constant for strong acids is large e.g., HCl, HBr, HI, HNO3, H2SO4, HClO4.
"Strong" refers to complete dissociation, and “weak” refers to partial dissociation into ions.
Acid strength relates to the extent of dissociation, not concentration. Concentrated solution
contains larger amount of solute per dm3 volume solution; dilute contains lower amount of
solute per dm3 volume solution (in Chemistry, strong is not the same as concentrated; weak is
not the same as dilute).
A weak acid or a weak base ionizes only partially in water. HF (≈ 90% undissociated at 25°C)
and organic acids such as ethanoic acid are examples of weak acids. Ammonia and
aminoethane are examples of weak bases:
© IBID Press 2007 1 CHAPTER 8 ACIDS AND BASES
(TOPICS 8 AND 18) SUMMARY
Experimental Data to Distinguish Between Strong and Weak Acids and Bases
1. Since HCl is a strong acid, in solution the concentration of H+ = 0.10 mol dm−3 and thus its
pH is 1.0. Since CH3COOH (aq) is a weak acid, [H+] will be much less than 0.10 mol dm−3
with a higher pH (≈ 3). Thus, can be distinguished by differences in the pH value of equimolar
2. Can also be distinguished by measuring the conductivities of equimolar aqueous solutions.
Because the strong acid is completely ionized, there are more ions in solution and as a result it
will show greater conductivity, whereas the weak acid is only partially ionised, with fewer
ions and lesser conductivity.
3. Rate of reaction of acids with Mg or carbonate: The strong acid reacts faster to produce H2
or CO2 gas more quickly as there are more H+ ions present whereas the weak acid reacts
slowly as there are fewer H+ ions present.
Water Dissociation: H2O (l)
at 25°C. H+ (aq) + OH− (aq); Kw = [H+] [OH−] = 1.0 x 10−14 mol2 dm−6 Ka x Kb = Kw for any combination of a weak acid and its own conjugate base (or weak base and
its own conjugate acid).
pH is a measure of acidity; pOH is a measure of basicity. One way to measure H+
concentration is to measure the pH of a solution. The relationship between pH of a solution
and hydrogen ion concentration (or [H+]) is:
pH = log10
= − log10 [ H + ] . ∴[ H ] = 1.0 x 10
pOH = − log 10 [OH − ] = log 10 1
[OH − ] Salt Hydrolysis
A salt, an ionic compound, contains cations (e.g., Na+) from a base and anions (e.g., Cl−) from
an acid. Equal amounts of monoprotic acids react with equal amounts of monohydroxy bases,
but the aqueous solutions formed will have different pH values depending on the strength of
the acid and the base. Thus, in aqueous solutions, properties of salts, which are completely
dissociated into ions, are those of their ions.
• • Strong acids (HCl, HBr, HI, HNO3, H2SO4) produce neutral anions (Cl−, Br−, I−, NO3−,
Stronger the acid, weaker the conjugate base; weaker the acid, stronger the conjugate base.
Strong bases (NaOH, KOH…) produce neutral cations (Na+, K+, …).
Stronger the base, weaker the conjugate acid; weaker the base, stronger the conjugate acid.
Neutral ions, called spectator ions, do not affect acid-base behaviour of salts.
Salt formation (note: Cation of salt comes from the base, anion from the acid):
o Strong acid + Strong base → Salt will be neutral; pH = 7.0
o Weak acid + Strong base → Salt will be basic; pH > 7
o Strong acid + Weak base → Salt will be acidic; pH < 7
o For weak acid + weak base → Salt will be approximately neutral (pH ≈7) depending
on the relative strength of the weak acid and the weak base.
An amphoteric substance can behave either as an acid or as a base: e.g., Al(OH)3.
As a base, it reacts with H+: Al(OH)3 (aq) + 3 H+ (aq) → Al3+ (aq) + 3 H2O (l) (note:
cations come from base)
As an acid, it reacts with OH−: Al(OH)3 (aq) + OH− (aq) → Al(OH)4− (aq) (note: anions
come from acid) © IBID Press 2007 2 CHAPTER 8 ACIDS AND BASES
(TOPICS 8 AND 18) SUMMARY
To determine acid-base property of a salt solution:
• Identify the cation and anion present.
• Deduce whether each is acidic, basic, or neutral. (1) If both are neutral, salt is neutral;
(2) if one is neutral and the other acidic, the salt is acidic; (3) if one is neutral and the
other basic, the salt is basic.
• If one ion is acidic and the other basic, the salt may be approximately neutral
depending on the relative strengths of the acid and the base.
A buffer is a solution that resists change in pH when a small amount of a strong acid or a
strong base is added to it. A buffer consists of a weak acid and its conjugate base or a weak
base and its conjugate acid in equilibrium with each other. A buffer is therefore made up of a
salt of a weak acid plus the weak acid e.g. sodium ethanoate + ethanoic acid: Na+CH3COO– +
CH3COOH is an example of an acidic buffer. Or a salt of a weak base plus the weak base e.g.,
NH4+Cl– + NH3 (a basic buffer). It is necessary to have both components.
(i) If now OH– ions from a strong base are added to the buffer solution, these react with the
acid of buffer; effectively, the strong base OH– has been replaced by the weaker base, and its
effect on pH minimized.
(ii) If H+ ions from a strong acid are added to the buffer, these will react with the base of
buffer; ehe strong acid H+ has been replaced by the weaker acid and its effect on pH
minimized. Thus effect of adding any strong base, OH– or strong acid, H+ to a buffer is
When [acid] = [conjugate base], an optimum buffer is present and has equal ability to
minimize the effect of added [H+] or [OH–]. Also, when [acid] = [conjugate base], then [H+] =
Ka or pH = pKa.
On dilution of a buffer, the ratio of [acid] to [conjugate base] does not change as the two
concentrations are affected equally. Thus the pH of a buffer does not change on dilution, but
the pH of a 0.10 mol dm–3 HCl or CH3COOH does change on dilution, e.g., 0.10 mol dm-3
HCl: pH = 1.0; 0.050 mol dm–3 HCl: pH = 1.3.
The pH of a buffer depends on:
(a) Ka of the weak acid, and (b) The relative concentrations of acid and conjugate base.
Experimental Determination of Ka (or Kb) value for a weak acid (or a weak base):
There are two ways in which to create the condition [acid] = [conj. base]:
(1) Take equal concentration of the weak acid and conjugate base and mix these to produce
(2) By the half-neutralisation method: e.g., dissolve a known sample of weak acid HX in a
fixed volume of solution. Neutralize half the amount (in moles) of HX with a strong base. At
half-neutralization, pH = pKa for an acidic buffer, and pOH = pKb for a basic buffer
Acid – Base Titrations
An indicator signals the endpoint of a titration by changing colour. End point of titration is the
experimentally determined point when the same amount (in moles) of acid and base have been
added. Equivalence point is the theoretical value when quantities of substances specified in a
chemical equation have reacted together.
If an unknown base is present, to this is added a known amount of excess acid. The excess
acid can be determined by titration with a standard base. Thus the amount of acid required to
neutralize the unknown base and therefore, the amount of unknown base can be calculated −
this procedure is called back titration.
© IBID Press 2007 3 CHAPTER 8 ACIDS AND BASES
(TOPICS 8 AND 18) SUMMARY
Indicators signal the endpoint of a titration by changing colour. Indicators used in acid-base
titrations are usually organic dyes and are mostly weak acids (or weak bases) where the acid
and the conjugate base of the indicator are different colours. Consider phenolphthalein, HIn, as
an example a weak organic acid, Ka = 1.0 x 10−9, in which HIn is colourless and its conjugate
base, In− is pink:
HIn (aq) + − H (aq) + In (aq); [ H + ][ In − ]
= 1.0 x 10−9 at 25oC
[ HIn] An indicator changes colour over ≈ 2 pH unit range. Phenolphthalein, is colourless below pH
≈ 8, pink above 10 and end point colour change occurs at pH = 9. The end point for any
indicator depends on its Ka value − the pH at the end point equals the pKa of the indicator.
In carrying out an acid-base titration, select an indicator that changes colour at or near the
equivalence point of the reaction at which equal quantities of the acid and base have been
added i.e. pKa of (weak acid) indicator should be close to pH of solution at equivalence point.
Acid-base Titration Summary:
(a) Strong acid + strong base: solution at equivalence point, pH = 7.0; large pH change; can
use phenolphthalein (pH range 8.3 – 10); or methyl red: (4.2 – 6.3); or methyl orange (pH
range: 3.1 – 4.4).
(b) Weak acid + strong base: solution at equivalence point, pH > 7; use phenolphthalein (pH
range 8 – 10)
(c) Strong acid + weak base: solution at equivalence point, pH < 7; use methyl orange (pH
range: 3.1 – 4.4)
(d) Weak acid + weak base: solution at equivalence point, pH about 7; gradual change in pH;
no suitable color indicator; record change in pH using a pH meter and obtain point of
inflection for equivalence point. Methods of Determining Ka of a Weak Acid or a Weak Base
1. Determine pH of a solution of known concentration of the weak acid. Knowing pH,
calculate [H+] = x.
Ka = H+ (aq) + X– (aq) 0.10 — — 0.10 – x x x Assume x is negligible;
i.e., x << 0.10 2
[H + ][X − ]
; if x = 1.0 x 10–pH, then Ka can be calculated.
0.10 − x 0.10 2. Can obtain Ka or Kb from titration curve by the half-neutralization method:
(N.B. Shading indicates Ch 18 (AHL) material.) © IBID Press 2007 4 ...
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This note was uploaded on 09/23/2010 for the course CS 001 taught by Professor Jix during the Spring '10 term at Riverside Community College.
- Spring '10