CHAPTER 8: ACID – BASE EQUILIBRIUM
4 U CHEMISTRY
P 526 – SEE KEY IDEAS
- examples f acids and bases.
8.1: The Nature of Acid – Base Equilibria:
- Arrhenius acids – H ions in water, bases – hydroxide ions in water
- acids – sour, turn blue litmus to red (pink), conduct electricity
- bases – bitter, soapy taste, slippery, conduct electricity, turn red litmus to blue
Bronsted – Lowry Theory:
- acids – donate proton, bases accept proton
- ex: HCl donates proton to water ( not reversible )
- ex: water donates proton to ammonia – reversible
- Amphoteric substances: ____________________________________________________________
- Advantage of B-L definition:
- 1. reactions can be identified as acid-base neutralization reactions without water
- 2. salts that form acidic or basic solutions when dissolved can be explained
Reversible Acid – Base Reactions
- many B-L acid-base reactions are reversible
- Conjugate acid – base pair – differ only by one H
- ex: H2O and H3O +
A Competition for Protons
- stronger c-acid succeeds in donating the most protons, has a weaker c-base
- stronger c-base succeeds in accepting the most protons, has a weaker c-acid
- the strong acid and strong base are always on the same side of the reaction arrow
- 99 % ionized = strong acid, ex: HCl
SUMMARY p 531
p 532 practice # 1, 2, 3.
The Autoionization of Water
- 2 molecules may collide and one molecule accepts a proton from the other – see p 532
- equilibrium of the reversible reaction – can write the constant
- Kw = [H +] [OH –1 ] = (1.0 x 10 –7 )2 = 1.0 x 10 –14
- Kw changes at different temp ( remember concept from ch 7 )
- so pH of pure water changes with diff temp
- but pure water always neutral since # protons = # hydroxide ions
Strong Acids
- ionize more than 99 % in water ex: HCl
- monoprotic acids – have only one ionizable H ex: HNO3
p 537 practice # 4, 5, 6.
Strong Bases:
- ionic substances, dissociate 100 % in water, but not all are very soluble
- ex: NaOH and other gp 1 hydroxides (soluble), Mg(OH)2 and other gp 2 hydroxides (slightly soluble)
- DON’T drink NaOH ! it will react with the tissues of your mouth ant throat and destroy them!!
- but people regularly eat Mg(OH)2 - as an antacid – it doesn’t dissolve on the way to the stomach, and
then only dissolves enough to react with stomach acid to neutralize it
- [OH –1]= concentration of strong base x # OH –1 groups per formula unit. Ignore OH –1 from water
p 540 practice # 8, 9, 10, 11.
Hydrogen Ion Concentration: pH = – log [H +]
- used to express conc that are smaller than 1 because neg exponents can sometimes be confusing
- SIG DIG’s count ONLY for the decimal numbers of a pH value
- pOH = – log[OH –1]
- pKw = – log Kw = 14 at SATP
- pOH + pH = 14 at SATP
Measuring pH
- indicators – weak acids or weak bases that change colour with the addition or removal of the ionizable
hydrogen ex: bromthymol blue ( BTT ) that we used in an investigation
- pH meters – change of electrical conductivity is converted into pH
pH of Strong Acids and Bases:
- conc of hydrogen ion = conc of the strong acid if the acid is monoprotic, H+ from water is negligible
- OMIT calculations of acids that are diprotic or higher.
- [OH –1] = concentration of strong base x # OH –1 groups per formula unit. Ignore OH –1 from water
SUMMARY p 549
p 546 practice # 12, 13, 14, 15ab, 16a.
p 549 practice # 17, 18, 19.
p 549 – 550 section 8.1 practice # 1, 2, 3, 4.
8.2: Weak Acids and Weak Bases
ex: _________________________________________________________________________________
- form equilibrium solutions between ionized and molecular forms
Weak Acids: _________________________________________________________________________
-ex: ________________________________________________________________________________
Weak Bases:
- react with water to ionize the water so OH –1 ions exist in the solution
- weak bases are proton acceptors – many are the neg ions of weak acids such as the acetate ion
- ex: NH3 + H2O  NH4+ + OH –1
Percent Ionization of Weak Acids: p = ( [H +] / [ HA ] ) x 100 %
p 554 practice # 1, 2
Ionization Constants for Weak Acids:
- for HA (aq)  H+ (aq) + A –1 (aq) Ka = [H +] [A –1 ]
[HA]
p 556 practice # 3, 4, 5.
Percent Ionization and Concentration:
- the more dilute the conc, the greater the % ionization
- this is an application of Le Chatelier’s Principle
NB: - dilute solutions have a small # mol in 1 L of solution
- a strong acid may be in a concentrated or a dilute solution
- a weak acid may be in a concentrated or a dilute solution
Ionization Constants for Weak Bases
- for B (aq) + H2O (l)  HB+ (aq) + OH –1(aq)
Kb = [HB+ ] [OH –1 ]
[B]
Organic Bases:
- organic molecules often contain amine groups – based on ammonia with one or more of the H’s from
the ammonia substituted with a carbon chain
- the N has an unshared e- pair and so can form a coordinate covalent bond with a proton
- so: amines act as weak bases in water
- common organic weak bases: _________________________________________________________
The Relationship Between Ka and Kb
- Ka x Kb = Kw
- pKa + pKb = pKw = 14
p 563 practice # 6
The pH of Weak Acid Solutions:
- use the equation for Ka
- the HUNDRED rule: shortcut for VERY weak acids: from the ICE table – since the % ionization is very
low, [HA] at equ’m is assumed to equal the original [HA] - the HUNDRED rule if [HA] initial > 100
Ka
- the 5 % rule: to check if the short cut hundred rule was justified: if
x
<
5%
[HA] initial
SUMMARY p 568
p 568 practice # 7, 8.
p 570 # 9, 10.
The pH of Weak Base Solutions:
SUMMARY p 574
Polyprotic acids:
- have more than 1 ionizable hydrogen
- Ka for each H is greater than the Ka for the one after
- ex> H3PO4 : Ka1 > ka2 > ka3
- most often: [H ] is due to the first ionizable H only – the only type of calculation we will do with these
acids
p 579 Section 8.2: practice # 1, 2, 3, 4, 5, 6, 7, 9, 10b, 11, 12, 13, 14, 15, 16, 17, 18, 21a, 22abc, 23a.
8.3 Acid – Base Properties of Salt Solutions
Salts that form Neutral Solutions:
- salts with cations from strong bases: sodium and potassium ions ( from NaOH and KOH ):
- salts with anions from strong acids: chloride and bromide ions from HCl and HBr
- ex: NaCL and KBr
Salts that Form Acidic Solutions:
- contain the cations of weak bases ex: NH4+ (aq) salts
- contain highly charged metal ions ex: Fe3+ (aq) – cause water to hydrolyze
Salts That Form Basic Solutions:
- contain anions of weak acids ex: carbonate or hydrogen carbonate ions from carbonic acid – cause
water to hydrolyze
Salts that Act as Acids and Bases:
- contain cation of a weak base and anion of a weak acid ex: NH4CN
- check the Ka and Kb values – the largest will have the most effect
- Kb for CN-1 (aq) > Ka for NH4+ (aq) so the solution will be basic
SUMMARY p 585
p 588 practice # 1, 2, 3, 4, 5.
Hydrolysis of Amphoteric Ions
ex: HCO3-1 (aq) will be basic, and HSO4-1 (aq) will be acidic
- check the Ka and Kb values
- the larger value determines the dominant nature of the ion and so you can predict if the solution will be
acidic or basic
p 589 practice # 8
Hydrolysis of Metal Oxides and Non-Metal Oxides
- metal oxides react with water to form hydroxides – basic solutions
- non-metal oxides react with water to form acidic solutions
- acid rain from CO2, SO2 and NO2 all dissolve minerals that aren’t very soluble in water – results in
caves from limestone and stalactites and stalagmites
- tooth enamel is insoluble in water but has an OH group and so reacts with acids – replacing OH group
with F ion ( fluoride tooth pastes ) reduces solubility even more and prevents reaction with acids and so
prevents tooth decay
p 591 practice # 9, 10
SUMMARY p 592.
Lewis Model of Acids and Bases:
includes B-L acids and bases and more substances which neutralize when they react
- Lewis acid – electron pair acceptor – have an incomplete valence shell
- Lewis base – electron pair donor – have an unshared electron pair in the valence shell
- ex: H+ (aq) + H2O (l)  H2O+ (aq) ex: HCl (g) + NH3 (g)  NH4Cl (s)
p 594 practice # 12
p 594 8.3 practice # 1, 2, 4, 5, 6, 7.
8.4: Acid – Base Titrations
titration: ___________________________________________________________________________
titrant: _____________________________________________________________________________
sample: ____________________________________________________________________________
equivalence point: ____________________________________________________________________
endpoint: ___________________________________________________________________________
-
-
technique: use plain water to wash a clinging drop from the end of the burette into the reaction flask –
because the level of titrant in the burette has already accounted for that drop, and the water will not
affect the # mol of H+ or OH-1 and so will not alter the endpoint
common indicators – phenolphthalein and bromthymol blue – both change colours near pH 7
Titrating a Strong Acid with a Strong Base:
- use stoichiometry to calculate the # mol reacted, and then calculate the pH or pOH of the reaction
solution
- pH graphed against # mL of NaOH solution added to a sample acid shows almost vertical line when
one more drop of NaOH is added at the equivalence point
p 599 practice: # 1, 2, 3.
Titrating a Weak Acid with a Strong Base:
- use stoichiometry to calculate the # mol consumed
- use equilibrium reaction to calculate the new conc and so the pH at any point in the titration
- equivalence point will be above pH 7 because of the action of the anoin of the weak acid
- use an indicator which will change colour in the range of pH 8 – 9 (phenolphthalein)
p 607 practice # 4, 5
Titrating a Weak Base with a Strong Acid
- same steps as above
- equivalence point at pH below 7 due to the action of the weak base
- use an indicator that changes colour in the acid range – bromocresol green or methyl red
p 608 practice # 6
- OMIT the rest of the section – read for interest if you wish.
p 613 8.4 practice # 1ac, 2, 4, 6, 8, 9.
8.5: Buffers
- buffer: solution which changes in pH very little with the addition of a small quantity of strong acid or
strong base
- usually weak acid with added anion of the acid, or weak base with added cation of the weak base
- ex: acetic acid + sodium acetate
- ex: aqueous ammonia + ammonium chloride
Explaining Buffers:
- for the reaction: HA  H+ (aq) + A-1 (aq)
1.0 M
1.0 M
- if a small amount of base is added, the acid molecules react and [ H+ ] remains virtually unchanged so
the pH remains virtually unchanged
- if a small amount of acid is added, the anion accepts H+ and becomes molecular acid, the [H ] remains
virtually unchanged
- if almost all the molecular acid is used up, or in almost all the anion is used up, then the pH will change
dramatically with the addition of one drop more of strong acid or base
The Capacity of a Buffer
calculations with buffers
Buffers in Action:
- cells need constant pH for enzymes to work
- blood needs to be at pH 7.4
- blood at pH 7.2 is deadly enough acid in a glass of orange juice to kill if the blood were not buffered
- important buffering systems in living organisms: H2PO4-1 (aq) with HPO4-2 (aq)
and H2CO3 (aq) with HCO3-1 (aq)
p 620 preactic # 1, 2, 3.
p 620 8.5 practice # 1, 3, 4, 5, 6, 7, 8, 9.
8.6: The Science of Acid Deposition
1. Source of the acid? ________________________________________________________________
2. Affects of the acid? _________________________________________________________________
3. Technology to reduce the acid? _______________________________________________________
4. Social costs of the acid? _____________________________________________________________
p 624 practice # 2, 3ab.
Careers in Chemistry – read!
Problem Set: part 1: p 632 – 633
you must do for marks # 1, 2, 4, 5, 6, 7, 8, 9, 10a, 11, 20a,
plus: p 641 # 27, 28, 29, 30, 32, 34, 38, 40,
you should do, for practice, # 3, 12, plus p 641 # 31, 33, 35, 36, 37, 39, 41,
Problem Set: part 2 p 633 you must do: 15, 18, plus
you should be able to do # 43, 45,
p 641
# 41, 42, 44, 46c, 51, 52.
Introductory Lab: informal. May be done over 2 - 3 days while reviewing concepts
Purpose: - to remind students about concepts from grade 11
- to practice stoichiometry calculations
Safety: concentrated acids and bases used. Goggles must be worn.
Procedure:
1. mass 1.0 g Cu metal
2. add it to 7 mL of 6.0 M nitric acid in a beaker with a glass plate covering it
3. allow reaction to finish in the fume hood
4. Place beaker in an ice water bath.
5. add 7 mL of COLD 6.0 M sodium hydroxide
6. Swirl to mix.
7. Warm the beaker and mixture on a hot plate until it is all black.
8. add 10 mL of water, swirl to mix.
9. let settle and decant water
10. repeat steps 8 & 9 again ( washing out excess nitrate and sodium ions )
11. add 5 mL of 6.0 M sulfuric acid
12. add 10 mL water
13. add up to 1.5 g powdered zinc metal, a small amount at a time,
14. swirl to mix after each addition
15. stop adding zinc when no more brown ppt is seen
16. add 6.0 M sulfuric acid, drop by drop, until no more bubbling is seen.
17. mass a piece of filter paper.
18. filter the mixture and let dry
19. mass the filter paper plus residue to determine the amount of Cu metal obtained.
20. Calculate the % yield.
Reactions:
Cu (s) + 3 HNO3 (aq) 
Cu(NO3)2 (aq)
Cu(NO3)2 (aq)
NO2 (g) + H2O (l)
brown gas
(poisonous)
+ H+ (aq)
single displacement
+ decomposition

Cu(OH)2 (s)
+ 2 NaNO3 (aq) double displacement
blue, jelly-like
( all nitrates are soluble, most hydroxides are insoluble )
( almost all Cu ion solutions are blue or blue-green )
Cu(OH)2 (s)
+ 2 NaOH (aq)
+

CuO (s)
+
CuO (s) + H2SO4 (aq) 
( most sulfates are soluble )
H2O (l) decomposition
CuSO4 (aq)
CuSO4 (aq) + Zn (s)  ZnSO4 (aq)
( metal activity series )
( theoretical yield and % yield )
+
2 H2O (l) acid – base neutralization
double displacement
+ Cu (s) double displacement
Zn (s) + H2SO4 (aq)  ZnSO4 (aq) + H2 (g) single displacement
(to isolate the Cu metal and attempt to purify it )
( reasons for % yield < or > 100 % )