Download Topic 8 Acids and Bases File

Topic 8: Acids and Bases
1
Version 2/2014
Table of Contents
Topic 8: Acids and Bases ............................................................................ 1
Definitions ........................................................................................... 4
8.1 Introduction to Acids and Bases .............................................................. 6
8.1.1 Define acids and bases according to the Bronsted-Lowry and Lewis theories. .........................6
Reactions of Acids ..................................................................................................................................................9
8.1.2 Deduce whether or not a species could act as a Bronsted-Lowry and/or Lewis acid or base.12
8.1.3 Deduce the formulas of the conjugate acid (or base) of any Bronsted-Lowry base (or acid)13
8.1 Introduction to Acids and Bases Revision .............................................................................................. 15
8.2 Properties of Acids and Bases .............................................................. 16
8.2.1 Outline the characteristic properties of acids and bases in aqueous solutions...................... 16
Titration Calculations-1 ....................................................................................................................................... 17
Properties of Acids and Bases Revision Notes ............................................................................................ 18
8.3 Strong and Week Acids and Bases ......................................................... 19
Factors that Affect Acid Strength ................................................................................................................ 19
Dissociation of Strong and Weak Electrolytes ........................................................................................... 20
Factors that Affect Acid Strength ................................................................................................................ 21
Acids and Bases Practice .................................................................................................................................... 21
Topic 8.3 Strong and Weak Acids and Bases Revision ............................................................................. 23
8.4 The pH Scale ................................................................................. 25
8.4.1 Distinguish between aqueous solutions that are acidic, neutral or alkaline using the pH scale25
Acids, Bases and pH ............................................................................................................................................. 25
8.4.2 Identify which of two or more aqueous solutions is more acidic or basic, using the pH scale.
..................................................................................................... 26
pH Calculations ...................................................................................................................................................... 26
2
Version 2/2014
pH and pOH Calculations .................................................................................................................................... 27
pH (and Quatitative Chemistry) Review Problems ..................................................................................... 29
Extension: Acid Rain Research ......................................................................................................................... 29
Summary Sheet: Acids and Bases ................................................................................................................... 30
Topic 8 Acids and Bases Exam Questions ....................................................... 35
Topic 8 Acids and Bases Exam Question Markscheme ......................................... 38
3
Version 2/2014
Definitions
Learn these off by heart
Amphoteric: Can have the properties of both a base and an acid, depending on whether it is reacting
with a base or an acid.
Brønsted-Lowry: An acid is defined as a proton donator, while a base is a proton acceptor.
Buffer: A solution that resists changes in pH when small amounts of acid or alkali are added to it.
When a small amount of acid is added, the excess of H+ ions causes the equilibrium to shift to the left -> balances the difference. When a small
amount of alkali is added, the OH- ions react with the H+ ions to form water. The decrease in [H+] is compensated for by an equilibrium shift to the
right. Vice versa for alkali buffers. Buffer solutions are made by several means:
**strong base + excess weak acid; **strong acid + excess weak base; **weak acid + same acid’s salt; **weak base + same base’s salt.
Concentrated: High number of moles of solute per volume of solution.
Conductivity: The more a solution is dissociated into its ions, the greater its conductivity.
Conjugate: The species remaining after an acid has lost a proton (conjugate base) or a base has
gained one (conjugate acid). pKa + pKb = pKw
Diprotic: Where one mole of an acid produces two moles of hydrogen ions, e.g. H2SO4.
End point: The point at which the indicator changes colour most rapidly.
Equimolar: Containing moles at a ratio equal to the stoichiometric ratio.
Equivalence point: Where the acid and base are in equimolar quantities. Exactly enough to react with
each other.
Indicator: A weak acid or base in which the dissociated form is a different colour to the
undissociated form. The end point occurs when the pH is approximately equal to the pKin value.
Ideally, the end point corresponds to the equivalence point in a titration.
Lewis theory: An acid is defined as an electron pair acceptor (e.g. BF3) and a base is an electron
donator (e.g. NH3).
Monoprotic: Where one mole of the acid produces one mole of hydrogen ions, e.g. HCl.
pH: Power of hydrogen. – log[H+]
Salt hydrolysis: The process by which a salt is broken down by water.
Strong: An acid or a base that dissociates completely into its ions. Ka >> 1.
Some strong acids: hydrochloric, sulphuric, nitric (weaker than other two). Strong bases: hydroxides of alkali metals.
4
Version 2/2014
Water, ionic product of: The equilibrium constant for the dissociation of water into its ions, where
[H2O] is taken to be constant. Value of Kw increases as temperature is increased, as the dissociation
is an endothermic process.
Weak: An acid or base that only slightly dissociates into its ions. Ka << 1. Some weak acids: ethanoic,
carbonic. *Weak bases: ammonia, aminoethane.
5
Version 2/2014
8.1 Introduction to Acids and Bases
8.1.1 Define acids and bases according to the Bronsted-Lowry and Lewis theories.
Bronsted-Lowry Acids and Bases
Why?
Is water an acid, a base, neither, or both? There are three different ways to define a substance as an acid or
base. One definition is based on the ions found in a compound (Arrhenius), another is based on how a
compound behaves when added to water (Bronsted-Lowry), and a third is based on how a molecule reacts with
other molecules (Lewis). These definitions address different behaviours of compounds and explain how
seemingly different compounds can be classified as behaving like an acid or a base.
Success Criteria
•
•
•
•
Define an acid according to the Arrhenius definition and the Bronsted-Lowry definition.
Define a base according to the Arrhenius definition and the Bronsted-Lowry definition.
Identify acids and bases that illustrate the Arrhenius definition and Bronsted-Lowry definition.
Explain the acid-base properties of amphiprotic substances.
Information
Acid: A compound that yields hydronium ions, H 3O+ (aq) as positive ions in aqueous solution (Arrhenius
definition). A compound that donates a hydrogen ion (H +) to another species (Bronsted- Lowry definition).
Base: A compound that yields hydroxide ions, OH (aq)- ,as negative ions in aqueous solution (Arrhenius
definition). A compound that accepts a hydrogen ion, (H +), from another species (Bronsted-Lowry definition).
Neutral solution: Contains hydrogen ions and hydroxide ions in equal concentrations.
Note: In the context of acid-base chemistry, the hydrogen ion usually is referred to as a proton because an
atom of hydrogen contains one proton and one electron - when it loses the electron during ion formation all
that is left is the nucleus, which is one proton.
Model
1.
2.
3.
4.
NaOH(s) + H2O(l)  Na+(aq) + OH (aq)
HCl(aq) + H2O(l)  H3 O+(aq) + Cl- (aq)
NH3(g) + H 2O(l)  NH+4 (aq) + OH (aq)
H2CO3(g) + H2O(l)  H3O+(aq) + HCO3 (aq)
5. HCl(aq) + NH3(aq)  NH4+ (aq) + Cl- (aq)
Key Questions
1. In equation 1, is NaOH(s) an acid or a base? Explain.
6
Version 2/2014
2. In equation 2, is HCl(aq) an acid or a base? Explain.
3. In equation 3, is NH3(g) an acid or a base? Explain.
4. In equation 3, is H2O(l) and acid or a base? Explain.
5. In equation 4, is H2O(l) and acid or a base? Explain.
6. Is H2CO 3(g) in equation 4 an acid or a base? Explain.
7. Compare the behaviour of NH3 in equations 3 and 5. Identify any similarities and differences. Explain.
8. Identify the substances in the Model that behave as both an acid and a base?
9. Explain how this duplicity in behaviour can or cannot occur.
Exercises
1. In the reaction below identify which of the reactants is an acid and which is a base:
HC2H3O2(aq) + H2O(l)  C2H3O2-(aq) + H3O+
(aq)
2. Consider the atomic structure of the H + ion. Complete the table below indicating the correct number of each
subatomic particle.
Composition of the H
Subatomic Particle
+
ion (1 1H+)
Number of Subatomic
Particles
Protons
Electrons
Neutrons
3. In some textbooks, when explaining the Brønsted - Lowry definition, acids and bases are described as
proton donors and proton acceptors. Based on your response to Exercise 2, explain why these are correct
terms.
4. A definition of the prefix amphi is;
- "both or of both kinds." Define the term "amphiprotic" and based
on the insight you gained from examining the model, explain why the term is used to describe water.
Applications
1. Ammonium chloride is one component of ordinary dry cell batteries. Ammonia gas can react with
hydrogen chloride gas to form the solid salt ammonium chloride. Write the balanced equation for this
reaction including the phases of each substance.
2. Label the acid and the base in the reactants of your equation in Application 1.
7
Version 2/2014
3. Your blood contains an acid-base buffer system. A buffer system is a chemical system that resists
changes in pH when small amounts of either acid or base are added to the system. It is important that
our blood pH does not change suddenly. A pH balance ensures that chemical reactions in the body take
place correctly. If the pH drops below 6.8 or rises above 7.8, death can occur. The buffer in blood is
the bicarbonate ion, HCO3- (aq). Two equations that illustrate bicarbonate's buffering action are shown in
these equations:
HCO3-(aq) + H+(aq)  H2CO3(aq)
HCO3-(aq) + OH-(aq)  CO32-(aq) + H2O(l)
Label the acid and the base in each of these equations.
4. Explain why bicarbonate ions are said to be amphiprotic.
5. When we exercise, CO2 builds up in our blood and the following reactions occur.
CO2 + H2O(l)  H2CO3(aq)
H2CO3(aq) H+(aq) + HCO3-(aq)
How does the buffer system in our blood respond to this reaction in order to keep the pH within the
acceptable range?
Got It?
Is water an acid, base, neither or both? Explain.
Naming Acids
1. Compounds that begin with H (it is really an H+1 ion) are probably acids
Example:
a. HF and HNO3 and H3PO4 are acids
b. CH4 is not an acid
2. There are different rules depending on what the H+1 ion is bonded to
3. Monoatomic ions vs. Polyatomic ions
Examples
a. HF  H is bonded to a monoatomic ion
b. HNO3  H is bonded to a polyatomic ion
c. H3PO4  H is bonded to a polyatomic ion
4. Rules for monoatomic:
a. Use “hydro” prefix
b. Use “ic” suffix
c. End with “acid”
8
Version 2/2014
Examples
d. HF  hydrofluoric acid
e. HCl  hydrochloric acid
f. HBr  Hydrobromic acid
5. Rules for polyatomic
a. Do NOT use prefixes
b. Use “ic” suffix for ions that end with “ate”
c. Use “ous” suffix for ions that end with “ite”
d. End with acid
Examples
e. HNO3  Nitric acid
f. HNO2 Nitrous acid
g. H2SO4  Sulfuric acid
h. H3PO4  Phosphoric acid
i. HC2H3O2  acetic acid
Monoatomic
Polyatomic
“ate”
Polyatomic
“ite”
Prefix
Root
Suffix
Example Name
ic
ic
Example
Formula
HCl
HNO3
Hydro
NONE
Ion name
Ion name
NONE
Ion name
ous
HNO2
Nitrous acid
Hydrochloric Acid
Nitric acid
Reactions of Acids
1. Synthesis of Acids
a) Period 3 non-metal oxides with water
P4O10 + 6 H2O  4 H3PO4
P4O6 + 6 H2O  4 H3PO3
SO3 + H2O  H2SO4
SO2 + H2O  H2SO3
Cl2O7 + H2O  2 HClO4
Cl2O + H2O  2 HClO
b) (HL) Period 3 non-metal chlorides (all produce HCl)
AlCl3 + 6H2O → H+(aq) + 3Cl-(aq) + [Al(H2O)5OH]2+
SiCl4 + 2H2O → SiO2 + 4HCl
PCl3 + H2O → H3PO3 + HCl or PCl5 + 4H2O → H3PO4 + 3HCl
S2Cl2 + H2O → complex, no need to write eqtn, but does produce HCl
9
Version 2/2014
Cl2 + H2O → HCl + HClO
c) Oxidation of organic alcohols
a. Complete Oxidation: primary alcohol + oxidizing agent carboxylic acid
b. complete oxidation occurs when reactants are heated under reflux (more on that in organic rxns
unit)
+
OH
K2Cr2O7
acid
O
OH
propanoic acid
propan-1-ol
2. Reactions of Acids
a) with metals
Reactions with metals produce salts and hydrogen gas
Zn(aq) + HCl(aq)  ZnCl2(aq) + H2(g)
b) with carbonates
Reactions with Carbonates and hydrogen carbonates produce carbon dioxide, salt, and water
CaCO3(aq) + 2 HCl(aq)  CaCl2(aq) + CO2(g) + H2O(l)
NaHCO3(aq) + HCl(aq)  NaCl(aq) + CO2(g) + H2O(l)
c) with bases
Acids and bases neutralize each other to form a salt
Example: HCl + NaOH  NaCl + H2O
Reactions of acids and bases are called neutralization reactions
d) Organic acids with alcohols
c. Carboxylic acid + alcohol  ester + water
d. Reaction conditions: acidic solution
e. The OH group on the carboxylic acid is replaced by the alcohols O-R group
f. Condensation reaction: produces water (also called dehydration synthesis and esterification)
g. Uses: flavouring agents, plasticizers, as solvents in perfume, polyesters
O
OH
+
O
O
+
O
propan-1-ol
acetic acid
OH
acid
OH
+
OH
H2O
ethyl acetate
acid
O
+
H2O
O
propan-1-ol
propionic acid
10
propyl propionate
Version 2/2014
Acid Base Definitions
Three definitions of
acids/bases:
Arrhenius
Bronsted Lowry
Lewis
Acid
Base
Donates protons (H+) in water
Donates (H+) in water
Accepts and electron pair
Donates (OH-) in water
Accepts (H+) from water
Donates electron pair
Table 1: equilibrium constants for some acid-base
reactions
HCl(g) + H2O(l)  H3O+(aq) + Cl-(aq)
NH3(aq) + H2O(l)  NH4+ (aq) + OH-(aq)
HCN(aq) + H2O(l)  H3O+(aq) + CN-(aq)
1.
2.
3.
4.
5.
6.
7.
Keq
2 x 104
3.3 x 10-7
1.1 x 10-11
(1)
(2)
(3)
Which chemical species are the Arrhenius acids in the forward reactions in table 1?
What chemical species are Arrhenius bases in the forward reactions in table 1?
What chemical species are Bronsted-Lowry Acids in the forward reactions in table 1?
What chemical species are Bronsted-Lowry bases in the forward reactions in table 1?
Is it possible for a substance to act as both an acid and a base?
Which do you think is the stronger acid, HCl or HCN? Explain your reasoning.
Consider reaction 1 from table 1 only
a. What species results from the loss of a proton from the Bronsted-Lowry acid in the forward
reaction?
b. Does your answer to a) act as an acid or base when the reverse of reaction 1 occurs?
c. What species results from the gain of a proton by the Bronsted-Lowry base in the forward
reaction?
d. Does you answer to c) act as an acid or base when the reverse of reaction 1 occurs?
e. Do the same relationships exist in reactions 2 and 3?
Conjugate Pairs:
Within the Bronsted-Lowry model, certain pairs of molecules are described as a conjugate acid-base pair.
The two species in a conjugate acid-base pair differ by a proton only. A base is said to have a conjugate acid
and an acid is said to have a conjugate base.
Table 2: examples of conjugate acid-base pairs
Acid
H2CO3
HCO3H2S
Base
HCO3CO32HS-1
Give the conjugate base for the following
1. NH4+
11
Version 2/2014
2. H2O
3. HSO44. H2SO4
5. OH6. CH3CH2NH3+
Give the conjugate acid for the following
1. NH3
2. H2O
3. CN-1
4. NO3-1
For the following reactions give the conjugate pairs and state which is the acid and which is the base.
1. H2SO4 + H2O  H3O+ + HSO4-1
2. H2S + NH3  NH4+ + HS-
8.1.2 Deduce whether or not a species could act as a Bronsted-Lowry and/or Lewis acid
or base.
Acid and Base Worksheet
1) Using your knowledge of the Brønsted-Lowry theory of acids and bases, write equations for the following
acid-base reactions and indicate each conjugate acid-base pair:
a) HNO3 + OH- 
b) CH3NH2 + H2O 
c) OH- + HPO4-2 
2) The compound NaOH is a base by all three of the theories we discussed in class. However, each of the
three theories describes what a base is in different terms. Use your knowledge of these three theories to
describe NaOH as an Arrhenius base, a Brønsted-Lowry base, and a Lewis base.
3) When hydrogen chloride reacts with ammonia, ammonium chloride is formed. Write the equation for this
process, and indicate which of the reagents is the Lewis acid and which is the Lewis base.
4) Write an equation for the reaction of potassium metal with hydrochloric acid.
5) Borane (BH3) is a basic compound, but doesn’t conduct electricity when you dissolve it in water. Explain this,
based on the definitions of acids and bases that we discussed in class.
6) Write the names for the following acids and bases:
a) KOH ____________________________________
b) H2Se ____________________________________
c) C2H3O2H ____________________________________
d) Fe(OH)2 ____________________________________
e) HCN ____________________________________
7) Write the formulas for the following chemical compounds
a) ammonium sulfate ____________________________________
b) cobalt (III) nitride ____________________________________
c) aluminum carbonate ____________________________________
d) chlorine ____________________________________
12
Version 2/2014
8.1.3 Deduce the formulas of the conjugate acid (or base) of any Bronsted-Lowry base
(or acid)
Conjugate Pairs
For each of the following equilibrium systems, identify and label the conjugate acid-base pairs. Use A to label
acids, B to label bases, CA for the conjugate acid and CB for the conjugate base. Then connect the conjugate
pairs with a line
1. HCl + H2O 
H3O+ + Cl2. NH2- + H2O  NH3 + OH3. CN- + H2O 
HCN + OH4. HClO4 + CH3COOH  ClO4- + CH3COOH2+
5. HCN + H2O  H3O+ + CN6. HSO4- + HCl  H2SO4 + Cl7. SO42- + HNO3  HSO4- + NO38. NH4+ +
HSO4-  NH3 + H2SO4
9. HCl + Al(H2O)5(OH)2+  Cl- + Al(H2O)63+
10. H3O+ + OH-  2 H2O
For these write the equation first......
11. NH3 with CH3COOH
12. N2H5+ with CO3213. H3O+ with OH
14. HSO4- with HCOO
15. ammonia with hydrochloric acid
16. hydrogen carbonate ion with nitric acid
17. formic acid with cyanide ion
18. acetate ion with water
13
Version 2/2014
Conjugate Acid-Base Pairs Arranged by Strength


The stronger the acid, the weaker the conjugate base. The stronger the base, the weaker the conjugate acid.
Conjugate bases of diprotic acids are often atypical (see entries in italics for examples).
ACID
Strength of Acid
BASE
Name
perchloric acid
sulfuric acid
STRONG
ACIDS
Strong Acid
Moderately Strong Acid!
Neutral
Basic Molecule
NeutralMolecules
Strong Base!
Formula
–
HClO4
ClO4
H2SO4
HSO4–
–
Name
I
Br–
Cl–
NO3–
H2O
iodide ion
bromide ion
chloride ion
nitrate ion
water
hydrogensulfate ion
HSO4–
SO42–
sulfate ion
methylammonium ion
water
ammonia
hydrogen
methane
hydroxide ion
Neutral
Anions
Neutral
Neutral Anion!
HF
HNO2
HC2H3O2
H2CO3
H2S
NH4+
HCN
HCO3–
F
NO2–
C2H3O2–
HCO3–
HS–
NH3
CN–
CO32–
fluoride ion
nitrite ion
acetate ion
hydrogencarbonate ion
hydrogensulfide ion
ammonia
cyanide ion
carbonate ion
CH3NH3+
CH3NH2
methylamine
H 2O
NH3
H2
CH4
OH–
OH
NH2–
H–
CH3–
O 2–
14
–
Moderately Strong Acid!
hydrogensulfate ion
HI
HBr
HCl
HNO3
H3O+
–
Strength of Base
Neutral Anion
perchlorate ion
hydroiodic acid
hydrobromic acid
hydrochloric acid
nitric acid
hydronium ion
hydrofluoric acid
nitrous acid
WEAK ACIDS
acetic acid
carbonic acid
Acid strength INCREASES hydrosulfuric acid
as you go UP the column.
ammonium ion
hydrocyanic acid
hydrogencarbonate ion
Basic Anion!
Weak Acid
Formula
WEAK BASES
Base strength INCREASES
as you go DOWN the column.
hydroxide ion
amide ion
hydride ion
methide ion
oxide ion
Strong Base
STRONG
BASES
Version 2/2014
8.1 Introduction to Acids and Bases Revision
8.1.1 Define acids and bases according to the Brønsted–Lowry and Lewis theories.(1)
8.1.2 Deduce whether or not a species could act as a Brønsted–Lowry and/or a Lewis acid or base. (3)
8.1.3 Deduce the formula of the conjugate acid (or base) of any Brønsted–Lowry base (or acid). (3)
Focus (theory)
Acid Definition
Base Definition
Brønsted–Lowry
Lewis
Arrhenius
a. According to Brønsted–Lowry, label the acid and base for both the forward and reverse reactions:
HCl (g) + H2O (l) ⇌ H3O+(aq) + Cl-(aq)
b. What is a conjugate acid base pair?
c. What does it mean for a substance to be amphiprotic?
d. Show that water is amphiprotic by labelling the conjugate pairs for the reaction of water with:
i.
hydrochloric acid:
ii.
ammonia:
e. How do the titles of weak/strong and acid/base change in an equation regarding conjugate pairs?
f. What does the proticity of acids mean?
i. Give an example of each, and provide an equation for the deprotonation (1 at a time):
i.
Monoprotic:
ii.
Diprotic:
iii.
Triprotic:
g. Show how HSO3- can act as both an acid and a base:
h. What is a dative bond?
i. Show how NH3 can form a dative bond with
1. H+
2. BF3
ii.
Many transition metals allow for the formation of dative bonds through their empty 3d
orbitals, draw an example of [Fe(H2O)6]3+:
i. Provide a diagram for how acids react with different substances:
j. Show the difference between an alkali and a base:
k. How is it that all Brønsted–Lowry acids are Lewis acids, but the reverse cannot be said?
15
Version 2/2014
8.2 Properties of Acids and Bases
8.2.1 Outline the characteristic properties of acids and bases in aqueous
solutions.
Properties of Acids and Bases Review:
1. Acids react with metals to form ____________ and _____________
a. Mg + HCl 
b. Ca + HNO3 
c. Cu + H3PO4 
2. Acids react with carbonates to form ________, ________and _________
a. HCl + Na2CO3
b. HCl + KHCO3
c. HNO3 + NaHCO3
3. Acids react with bases to form ____________ and _____________
a. HCl + NaOH
b. H2SO4 + KOH
c. LiOH + H3PO4
4. Acids taste ______________
5. Bases taste ______________ and feel _________________
6. Acids and bases react with litmus paper.
a. Acids turn litmus paper (blue/red).
b. Bases turn litmus paper (blue/red).
7. If a substance has a pH < 7, it is a ________________
8. If a substance has a pH > 7, it is a ________________
9. If substance A has a pH of 3 and substance B has a pH of 4
a. Which is the stronger acid?
b. Which is the stronger base?
10. What volume of 0.10 M HCl is required to completely neutralized 10 mL of 2.0 M Mg(OH) 2?
11. Name the following acids
a. HCl
b. HBr
c. H2CO3
d. H3PO4
e. HC2H3O2
f. H2SO4
g. H2SO3
h. HNO3
i. HNO2
16
Version 2/2014
Write a dissociation reaction for
j. HCl
k. HNO3
l. H2SO4
m. NaOH
12. Give the equilibrium expression for 12 a and d
Titration Calculations-1
Acid-Bass titration involves the gradual addition of acids and bases to one another, generally for the
purpose of determining reacting volumes or concentrations. In a laboratory titration, 1M HCl was
gradually added to 50 mL of 1M NaOH. The pH was monitored throughout the analysis and the following
readings were obtained.
Plot these data on the grid above, and then use the graph to answer questions #1 – 3.
1. What volume of 1M HCl was required to reach the equivalence point? ____________
2. How would the graph differ if the concentration of titrating HCl were increased to 2 M?
3. List two indicators that could be used to recognize the equivalence point of this titration.
Describe the colour change that each would undergo. (Look at your results from the pH Indicator
Activity).
4. Calculate the concentration of a monoprotic (single Hydrogen) acid solution, given that 140.0 mL
of the acid solution is neutralized by 55.0 mL of a 0.750 M solution of NaOH.0.295M
5. Calculate the concentration of a monoprotic basic solution, given that 99.3 mL of the basic
solution is neutralized by 83.0 mL of a 0.0325 M solution of HCl.0.027M
17
Version 2/2014
6. Calculate the number of milliliters of 0.540 M HNO3 that is required to neutralize 750 mL of
0.283M KOH.393.1 mL
7. By titration it is found that 12.4 mL of H2SO4 is required to neutralize 19.8 mL of 0.0100
MCa(OH)2. What is the molarity of the H2SO4? 0.016M
8. What volume of 0.12M Ba(OH)2 is needed to neutralize 12.2 mL of 0.25M HCl? 25.4 mL
9. A 0.055 g sample of Al(OH)3 is reacted with 0.200 M HCl. How many mL’s of the acid is needed
to neutralize the Al(OH)3 ?n3.0 mL
10.
How many grams of of the acid KHP (mwt. - 204.23 g/mol) are needed to neutralize 21.2
mL of 0.15M NaOH? 0.649 g
Properties of Acids and Bases Revision Notes
1. 8.2.1 Outline the characteristic properties of acids and bases in aqueous solution. (2)
Properties of Acids
Properties of Bases
a. What is an electrolyte?
b. What happens to a base in water?
What happens?
Acid with Metal
Relevant Equations:
In general:
Acid with Metal
Carbonates
Acid with Bases
Acid with Metal
Oxides
Acid with Metal
Hydroxides
18
Version 2/2014
c. In general, provide a diagram (web) for the generic reactions acids are able to undergo:
d. Use the example of hydrogen chloride (s) dissolved in water and the organic material methyl benzene to describe
why water is important to the function of acids:
Test
Solution of HCl in water
Solution of HCl in C6H5CH3
Universal Indicator Paper
Addition of CaCO3
Conductivity
Enthalpy of Solution
e. What is the difference between “hard” and “soft” water?
f. What is the difference between a “normal salt” and an “acid salt?”
8.3 Strong and Week Acids and Bases
Factors that Affect Acid Strength
Polarity of bond
3.
Molecules containing H will only donate a proton if the bond holding the H is
polarized with the H atom slightly positive.
Strength of Bonds
4.
5.
Very strong bonds are less easily dissociated than weak bonds.
The weaker the bonds, the stronger the acid.
6.
As the stability of the conjugate base increases, so does the strength of the acid.
Stability of the Conjugate Base
Binary Acids
a)
Smaller atoms bonded to H ions can contain stronger bonds because there is less shielding.
 What is the trend of bond strength down a family in the periodic table?
2. ______________ Acidity down the table.
3.
E.g. The hydrogen halide HF is a weak acid while the other hydrogen halides
are strong acids. This is because the HF bond is much stronger than the others.
a.
Bond polarity is the main factor determining strength of binary acids across a period of the
table.
4.
Acidity increases as the electronegativity of the element bonded to the H atom increases.
 What is the trend of electronegativity across a row of the periodic table?
This means that there is an __ of acidity across a period.
E.g. Acidity of CH4 < NH3 < H2O < HF
Oxyacids
5.
Many compounds contain OH groups. Whether this causes the molecule to be acidic or
basic depends on the bond between the OH and the other element.
6.
Elements with low electronegativities ( metals) tend to behave as bases.
19
Version 2/2014
7.
Elements with higher electronegativities ( non-metals) tend to behave as neutral or
acidic. This is because the bond between the element and the O is non-polar and instead of
dissociating the OH, it is just the H that is lost.
A. As the electronegativity of the element increases, so does the acidity.
8.
Electrons drift from the H atom towards the electronegative atom, which promotes
ionization.
9.
The conjugate base will be stabilized by the electronegative atom.
10.
E.g. HClO > HBrO > HIO
a. Oxyacids may contain additional O atoms attached to the central atom.
11.
These additional electronegative O atoms increase the effect of electron drift from the
H atoms.
3. How does an additional O atom affect the acidity of an oxyacid?
12.
E.g. HClO< HClO2 < HClO3 < HClO4 the acidity increases from hypochlorous to perchloric
acid.
Carboxylic Acids
13.
The largest category of organic acids
14.
The additional O atom attached in the functional group increases the electronegativity
of the anion portion and the stability of the conjugate base ( partly because of resonance structures)
compared to alcohols.
15.
As the electronegativity of atoms in the acid increases, the strength of the
acid_____________
.
16.
E.g. CF3COOH > CH3COOH trifluoric acid is stronger than acetic acid.
Dissociation of Strong and Weak Electrolytes
Instructions: Identify each of the substances below as either strong or weak electrolytes. In each
case, write a balanced equation that represents how the substance dissociates in water.
Background: General Rules
Strong electrolytes = soluble salts, strong acids and strong bases. Six common strong acids encountered
most often are HCl, HBr, HI, HNO3, HClO4, and H2SO4. The common strong bases include the Group 1
hydroxides, and the heavier group 2 hydroxides, Ca, Sr, and Ba.
Weak electrolytes are most acids (other than those mentioned above) and most bases (other than those
mentioned above).
1.
2.
3.
4.
5.
6.
7.
8.
KCl
HNO3
NaOH
Ba(OH)2
NH3
HF
(NH4)2CO3
Na3PO4
20
Version 2/2014
9. Na2S2O3
10. Mg(OH)2
11. H2C2O4
12. Fe(NO3)3
13. Al(OH)3
14. NH4Cl
15. H2SO4
16. NaHCO3
17. KNO2
18. K2CrO4
19. Pb(C2H3O2)4
20. HClO
Factors that Affect Acid Strength
1. Explain the following observations:
a) HNO3 is a stronger acid than HNO2
b) H2S is a stronger acid than H2O
c) H3PO4 is a stronger acid than H3AsO4
d) HCl is a stronger acid than H2S
e) Benzoic acid (C6H5COOH) is a stronger acid than phenol (C6H5OH)
f) H2SO4 is stronger than H2SeO4
g) CCl3COOH is stronger than CH3COOH
2. Select the stronger base in each of the following pairs, explain your choice:
a) BrO-1 or ClO-1
b) BrO-1 or BrO2-1
c) PO43- or AsO4-1
3. Select the stronger acid in each of the following pairs, explain your choice:
a) HBr or HF
b) PH3 or H2S
c) H2SeO3 or H2SeO4
Acids and Bases Practice
1. Which of the following is/are formed when a metal oxide reacts with a dilute acid?
I. A metal salt
II. Water
III. Hydrogen gas
A. I only
B. I and II only
C. II and III only
D. I, II and III
21
Version 2/2014
2. The pH of a solution is 2. If its pH is increased to 6, how many times greater is the [H+] of the
original solution?
A. 3
B. 4
C. 1000
D. 10 000
3. The equation for the reaction between nitric acid and sulfuric acid is shown below.
H2SO4 + HNO3 H2NO3+ + HSO4–
Which species are acting as acids in this reaction according to the Brønsted-Lowry theory?
A. H2SO4 and HNO3
B. H2SO4 and H2NO3+
C. HNO3 and H2NO3+
D. H2NO3+ and HSO4–
4. Four aqueous solutions, I, II, III and IV, are listed below.
I. 0.100 mol dm–3 HCl
II. 0.010 mol dm–3 HCl
III. 0.100 mol dm–3 NaOH
IV. 0.010 mol dm–3 NaOH
What is the correct order of increasing pH of these solutions?
A. I, II, III, IV
B. I, II, IV, III
C. II, I, III, IV
D. II, I, IV, III
5. Which is a conjugate acid-base pair in the following reaction?
HNO3 + H2SO4 H2NO3+ + HSO4–
A. HNO3 and H2SO4
B. HNO3 and H2NO3+
C. HNO3 and HSO4–
D. H2NO3+ and HSO4–
6. Which equation represents an acid-base reaction according to the Lewis theory but not the
Brønsted-Lowry theory?
A. NH3 + HCl NH4Cl
B. 2H2O H3O+ + OH–
C. NaOH + HCl NaCl + H2O
D. CrCl3 + 6NH3 [Cr(NH3)6]3+ + 3Cl–
22
Version 2/2014
7. Which one of the following species can act as both a Brønsted-Lowry acid and base in aqueous
solution?
A. CH3COOH
B. NO3–
C. H2PO4–
D. OH–
8. Define the terms strong acid and weak acid. Using hydrochloric and ethanoic acid as examples, write
equations to show the dissociation of each acid in aqueous solution.
(Total 4 marks)
9. (i) Calcium carbonate is added to separate solutions of hydrochloric acid and ethanoic acid of the
same concentration. State one similarity and one difference in the observations you could make. (2)
(ii) Write an equation for the reaction between hydrochloric acid and calcium carbonate. (2)
(iii) Determine the volume of 1.50 mol dm–3 hydrochloric acid that would react with exactly 1.25 g of
calcium carbonate.
(iv) Calculate the volume of carbon dioxide, measured at 273 K and 1.01×105 Pa, which would be produced
when 1.25 g of calcium carbonate reacts completely with the hydrochloric acid. (2)
(Total 9 marks)
10. The pH values of solutions of three organic acids of the same concentration were measured.
acid X pH = 5
acid Y pH = 2
acid Z pH = 3
(i) Identify which solution is the least acidic. (1)
(ii) Deduce how the [H+] values compare in solutions of acids Y and Z. (2)
(iii) Arrange the solutions of the three acids in decreasing order of electrical conductivity, starting with
the greatest conductivity, giving a reason for your choice. (2)
(Total 5 marks)
11. The equilibrium reached when ethanoic acid is added to water can be represented by the following
equation:
CH3COOH(l) + H2O(l) CH3COO–(aq)+H3O+(aq)
Define the terms Brønsted-Lowry acid and Lewis base, and identify two examples of each of these
species in the equation.
(Total 4 marks)
12. Put on another sheet if needed: Identify one example of a strong acid and one example of a weak
acid. Outline three different methods to distinguish between equimolar solutions of these acids in the
laboratory. State how the results would differ for each acid. (Total 5 marks)
Topic 8.3 Strong and Weak Acids and Bases Revision
1. 8.3.1 Distinguish between strong and weak acids and bases in terms of the extent of
dissociation, reaction with water and electrical conductivity. (2)
2. 8.3.2 State whether a given acid or base is strong or weak. (1)
23
Version 2/2014
a. Students should consider hydrochloric acid, nitric acid and sulfuric acid as examples of
strong acids, and carboxylic acids and carbonic acid (aqueous carbon dioxide) as weak acids.
b. Students should consider all group 1 hydroxides and barium hydroxide as strong bases, and
ammonia and amines as weak bases.
3. 8.3.3 Distinguish between strong and weak acids and bases, and determine the relative strengths of
acids and bases, using experimental data. (2)
a. Complete the following table with a list of common strong and weak acids and bases needed for IB:
Strong Acids
Weak Acids
Strong Bases
Weak Bases
b. When a strong acid dissolves, show what happens with equations and a simple bar diagram:
c. Why is an equilibrium symbol used for weak acids and bases while a yields symbol is used for strong
ones?
d. Give the general equations for the strong acids of hydrochloric, nitric, sulfuric, and perchloric acid in
aqueous solutions with water:
e. When a weak acid dissolves, show what happens with equations and a simple bar diagram:
f. Give the general equations for the weak acids of ethanoic, carbonic, and phosphoric acid in aqueous
solutions with water:
i. Stream water has a naturally acidic pH around 5.7, explain why the presence of carbon dioxide in in the
air causes this:
ii. Why does carbonic acid not actually exist?
g. List five methods to approach distinguishing between a strong and weak acid experimentally
To experimentally determine difference in strong and weak acids
h. Does an acid lose its strength when it’s diluted in solution?
i. For equimolar (same concentration) solutions of hydrochloric and ethanoic acid, complete the following
table with comparable properties for the strong and weak acid respectively:
24
Version 2/2014
0.1 mol dm-3 HCl (aq)
0.1 mol dm-3 CH3COOH (aq)
[H+(aq)]
pH
Conductivity
Reaction Rate with Mg
Reaction Rate with CaCO3
j. Explain how the effect of strong and weak bases are similar and yet different to that of like bases
when dissolved in solution by providing the proper generic equations and explanations:
k. Give the general equations for the strong bases of sodium hydroxide, potassium hydroxide, and
barium hydroxide in aqueous solutions with water:
l. Give the general equations for the weak bases of ammonia and ethylamine in aqueous solutions with
water:
m. Calcium hydroxide is considered a strong base, provide an equation for the addition of Ca(OH)2(s) to
aqueous solution and explain the limitations of effectiveness of this compound:
8.4 The pH Scale
8.4.1 Distinguish between aqueous solutions that are acidic, neutral or
alkaline using the pH scale
Acids, Bases and pH
1. Use Eq (1) and the log10 button on your calculator to determine the pH of solutions with the specified
hydrogen ion concentrations [H+]:
(a)
(b)
(c)
(d)
(e)
(f)
[H+]
0.10 M
0.0010 M
10-7 M
5.0 x 10-10 M
6.0 M
1.0 M
pH
Acidic, basic
or neutral?
(a) 1.00(b) 3.00(c) 7(d) 9.30(e) - 0.78
(f) 0.00
Solutions (a), (b), (e), and (f) are acidic; solution (c) is neutral; solution (d) is basic.
2. Use Eq (2) and the +/- and 10x buttons on your calculator to determine the [H+] of solutions with the
following pH values:
(a)
(b)
(c)
(d)
(e)
pH
2.00
4.00
0.30
12.80
4.500
+
[H ]
25
Version 2/2014
Acidic, basic or
neutral?
(a) 0.010 M(b) 1.0 x 10-4 M(c) 0.50 M(d) 1.6 x 10-13 M(e) 3.16 x 10-5 M
Solutions (a), (b), (c), and (e) are acidic; solution (d) is basic.
3. Complete the table for each aqueous solution at 25°C.
State whether the solutions are acidic or basic.
[H3O+]
[OH ]
pH
5
2.0 x 10
6.25
5.6 x 10
pOH
acidic or basic
2
9.20
8.7 x 10
10
8.4.2 Identify which of two or more aqueous solutions is more acidic or basic,
using the pH scale.
1.




Rank the following solutions in order of increasing acidity (from lowest to highest):
Solution A with pH = 6.50
Solution C with [OH–] = 6.0x10–10M
Solution B with [H3O+] = 3.5x10–5M
Solution D with pH = 5.85
2. What is an ion?
3. A solution with a pH of 6.5 is said to be _______________. A solution with a pH 10 is said to be
_____________.
4. Which is more acidic vinegar or tomato juice?
5. Which is more basic toothpaste or human blood?
6. What subatomic particle (proton, neutron, or electron) do atoms gain or lose in order to become
charged?
7. Hydroxide ions have a negative charge, why?
8. Hydrogen ions have a positive charge, why?
9. If you add lemon juice to pure water what would happen to the pH? What is in the lemon juice
10. that makes the solution this way?
pH Calculations
Find the pH of the following acidic solutions:
(Note: All the acids in this exercise are STRONG ACIDS)
26
Version 2/2014
1. A 0.001 moldm-3 solution of HCl (hydrochloric acid).
2. A 0.09 moldm-3 solution of HBr (hydrobromic acid).
3. A 1.34 x 10-4 moldm-3 solution of hydrochloric acid.
4. A 2.234 x 10-6 moldm-3 solution of HI (hydroiodic acid).
5. A 7.98 x 10-2 moldm-3 solution of HNO3 (nitric acid).
6. 12 dm3 of a solution containing 1 mole of hydrochloric acid.
7. 735 dm3 of a solution containing 0.34 moles of nitric acid.
8. 1098 dm3 of a solution containing 8.543 moles of hydrobromic acid.
9. 660 dm3 of a solution containing .0074 moles of hydrochloric acid.
10. 120 cm3 of a solution containing 0.005 grams of hydrochloric acid.
11. 1.2 dm3 of a solution containing 5.0 x 10-4 grams of hydrobromic acid.
12. 2.3 dm3 of a solution containing 4.5 grams of nitric acid.
13. 792 cm3 of a solution containing 0.344 grams of hydrochloric acid.
14. 100 cm3 of a solution containing 1.00 grams of nitric acid.
15. 8.7 dm3 of a solution containing 1.1 grams of nitric acid.
16. 1.5 dm3 of a solution containing 5.6 grams of hydroiodic acid.
17. 10.7 dm3 of a solution containing 0.01 grams of hydrochloric acid.
18. 8,000 dm3 of a solution containing 6.7 grams of nitric acid and 4.5 grams of hydrochloric acid.
19. 150,000 L of a solution containing 45 grams of nitric acid and 998 grams of hydrobromic acid.
20. 50 L of a solution containing 0.09 grams of HCl, 0.9 grams of HBr, 9.0 grams of HI, and 90.0
grams of HNO3.
21. What is the pH of a solution that contains 25 grams of hydrochloric acid (HCl) dissolved in 1.5
liters of water?
22. What is the pH of a solution that contains 1.32 grams of nitric acid (HNO3) dissolved in 750 mL
of water?
23. What is the pH of a solution that contains 1.2 moles of nitric acid (HNO3) and 1.7 moles of
hydrochloric acid (HCl) dissolved in 1000 liters of water?
24. If a solution has a [H+] concentration of 4.5 x 10-7 M, is this an acidic or basic solution? Explain.
25. An acidic solution has a pH of 4. If I dilute 10 mL of this solution to a final volume of 1000 mL,
what is the pH of the resulting solution?
pH and pOH Calculations
1. Determine the pH of the following solutions:
a) A 4.5 x 10-3 moldm-3 HBr solution.
b) A 3.67 x 10-5 moldm-3 KOH solution.
c) c) A solution made by diluting 25 cm3 of 6.0 moldm-3 HCl until the final
volume of the
3
solution is 1.75 dm .
d) 5 dm3 of an aqueous solution that contains 1.0 grams of HBr and 1.0 grams of nitric acid.
2. What are the pOHs for the solutions in problem 1?
a) ___________________________
b) ___________________________
27
Version 2/2014
c) ___________________________
d) ___________________________
3. Explain why even a basic solution contains some H+ ions.
4. Explain why even an acidic solution contains some OH- ions.
5. More challenging: What is the pH of a 1.5 x 10-10 HBr solution?
6. Does your answer from problem #5 make sense? Explain.
7. Determine the pH of a 0.0034 M HNO3 solution.
8. Determine the pOH of a 0.0034 M HNO3 solution.
9. Determine the pH of a 4.3 x 10-4 M NaOH solution.
10. If a solution is created by adding water to 2.3 x 10-4 moles of NaOH and 4.5 x 10-6 moles of HBr
until the final volume is 1 L, what is the pH of this solution?
11. Determine the pH of a 4.5 x 10-11 M NaOH solution.
12. Why would we say that a solution with a H+ concentration of 1.00 x 10-7 M is said to be neutral. If
it contains acid, shouldn’t it be acidic?
Answers
1.
a)
2.4
b)
9.56
c)
1.1 (the diluted solution is 0.086 M HCl)
d)
2.2 (the solution has an overall acid concentration of 0.0056 M)
2.
a)
11.6
b)
4.44
c)
12.9
d)
11.8
3. All aqueous solutions contain H+ ions from the autoionization of water, H2O  H+ + OH-.
4. The same answer from #3 applies here.
5. More challenging: What is the pH of a 1.5 x 10-10 HBr solution?
If you do the calculation to find –log (1.5 x 10-10), you get an answer of 9.8. However, this is
intuitively incorrect – after all, how can a solution that’s made of nothing but pure water (pH = 7)
with an acid added be basic overall? The answer, it can’t. The actual pH of the solution is just
about 7, with the main acid source being the H+ formed from the autoionization of water.
6. See answer from #5.
7. pH = -log[H+] = -log(0.0034) = 2.47
8. pH = -log[H+] = -log(0.0034) = 2.47pOH = 14 – pH = 14 – 2.47 = 11.53
9. pOH = -log[OH-] = -log(4.3 x 10-4) = 3.37
pH = 14 – pOH = 14 – 3.37 = 10.63
10. To solve:
 Both acid and base are present. Since they neutralize each other, you must first figure out
how much acid or base is left over after it neutralizes. Since the amount of base is larger
than the amount of acid, there will be more base than acid. The amount of base is 2.3 x 10 -4 –
4.5 x 10-6 = 2.26 x 10-4 moles.
 Since there is one L of solution, the molarity of the base is 2.26 x 10-4 M.
 To find pOH, take the –log of 2.26 x 10-4, which is 3.65.
 To find pH, subtract 3.65 from 14. The pH of this solution is 10.35.
28
Version 2/2014
11. Although there is some NaOH present in the solution, the pH isn’t found by taking the –log of
anything. The reason for this is that the concentration of base is much, much smaller than the
concentration of acid which is naturally found in neutral water. As a result, this base doesn’t really have
any effect on the pH, so the pH of the solution is 7.00.
12. It isn’t acidic because while there is some acid in the solution, there is an equal quantity of base. In
neutral solutions, the H+ and OH- concentrations are identical, because water breaks up to form them.
As a result, the solution is neither acidic nor basic.
pH (and Quatitative Chemistry) Review Problems
What is the molarity of a solution that has 450 grams of sodium chloride in 800 cm3 of water?
What is the molarity of a solution that contains 100 grams of iron (II) nitrate in 2.4 dm3 of
water?
What is the pH of a solution that contains 2.4 x 10-5 moles of hydrobromic acid in 0.5 dm3 of
water?
What is the pH of a solution that contains 25 moles of nitric acid dissolved in 5000 liters of
water?
What is the pH of a solution that contains 0.009 grams of hydrochloric acid in 100 cm3 of water?
What is an acid/base indicator used for?
Define “titration”:
In a few steps, describe how you would titrate a base of unknown concentration with an acid with
concentration 1 moldm-3.
I did a titration where it took 50 cm3 of 0.1 moldm-3 hydrochloric acid to neutralize 500 cm3 of a
base with unknown concentration. Using this titration information, what was the concentration of
the base?
I did a titration where it took 25 cm3 of 5 moldm-3 NaOH to neutralize 1000 cm3 of an acid with
unknown concentration. Using this information, what was the concentration of the acid?
1)
2)
3)
4)
5)
6)
7)
8)
9)
10)
Extension: Acid Rain Research
1.
2.
3.
4.
5.
6.
What is acid rain?
What are three problems caused by acid rain listed in the article?
Write the reaction for limestone (CaCO3) and nitric acid (use state symbols).
Write the reaction for iron and nitric acid.
With acid in the air, what effect do you think this has on human health?
Given that sulfur and nitrogen oxides are brown colored, what other problems occur because of
these gases?
You should now be able to use the questions above to answer the overall question: What are the
problems caused by acid rain?
7. What is the generic formula for nitrogen oxides? Sulfur oxides?
29
Version 2/2014
8. What is a natural source of nitrogen and sulfur oxides?
9. What are the anthropogenic (man-made) sources of nitrogen oxides and sulfur oxides?
10. Write the reaction for the following with water:
a. carbon dioxide
b. sulfur trioxide
c. sulfur dioxide
d. nitrogen dioxide (hint: products are nitric and nitrous acid)
11. All of the above are acid deposition reactions. Which one is considered natural?
12. Acids can exist as both gases and as aqueous solutions. When oxides of nonmentals (e.g. NOx)
react with water in the form of rain, snow or fog to form an acid this is ________ (wet/dry)
deposition. When acidic gases or particles land on a surface due to gravity, this is ________
(wet/dry) deposition.
You should now be able to use the questions above to answer the overall question: What causes acid
rain?
13.
14.
15.
16.
17.
18.
What is a scrubber?
Write the reaction for the scrubber.
What are two drawbacks to this method of removing pollutants?
Describe selective catalytic reduction.
What is a drawback to this method of removing pollutants?
Identify 2 more methods of reducing the presence of SOx and NOx in the air.
You should now be able to use the questions above to answer the overall question: What are some
ways of counteracting and preventing acid rain?
Summary Sheet: Acids and Bases
1. Properties
a. Acids
i.
ii.
iii.
iv.
Sour
Turn litmus red
pH less than 7
Reactions with Carbonates and hydrogen carbonates produce carbon dioxide, salt, and
water
1. CaCO3(aq) + 2 HCl(aq)  CaCl2(aq) + CO2(g) + H2O(l)
2. NaHCO3(aq) + HCl(aq)  NaCl(aq) + CO2(g) + H2O(l)
v. Reactions with metals produce salts and hydrogen gas
1. Zn(aq) + HCl(aq)  ZnCl2(aq) + H2(g)
vi. Corrode metals
vii. Dissolves carbonate rocks
30
Version 2/2014
viii.
b. Bases
i.
ii.
iii.
iv.
Conducts electricity
Bitter
Turn litmus blue
pH greater than 7
Soluble alkali metals make strong bases and hydrogen gas
1. 2Na + 2H2O  2NaOH + H2
v. Slippery
c. Acids and bases neutralize each other to form a salt
i. Example: HCl + NaOH  NaCl + H2O
ii. Reactions of acids and bases are called neutralization reactions
2. Strong vs. weak
a. Strength depends on amount of dissociation/ionization
i. HA
+ H2O 
H3O+
+
A(acid)
(base)
(conjugate acid) (conjugate base)
b. Strong = complete dissociation/ionization
i. Example, when HCl is dissolved in water, all of it becomes H+ (actually, H3O+)+ Clii. Ka >> 1
iii. Strong acids to know: Nitric, sulfuric, hydrochloric
iv. Strong bases to know: All group 1 hydroxides, barium hydroxide
c. Weak = incomplete dissociation/ionization
i. Ka << 1
ii. Acetic acid (vinegar) is a weak acid
iii. CH3COOH, small amounts dissociate into H+ and CH3COO-)
iv. Use the Keq to determine strength of weak acids
Acid
Keq
Acetic Acid
1.76 x 10-5
Phosphoric Acid
7.5x10-3
v.
Keq 
[ H ][ A]
HA
vi. Keq is also called Ka or acid dissociation constant since it is the equilibrium expression
for the dissociation of an acid
vii. Keq for the dissociation of a base = Kb The larger Keq, the more dissociation
viii. More dissociation = stronger acid
ix. Therefore, phosphoric acid is stronger than acetic acid
x. Weak acids to know: ethanoic (acetic) CH3COOH, carbonic
xi. Weak bases to know: ammonia (NH3), ethylamine (C2H5NH2)
d. More dissociation = more conductivity
i. Strong acids and bases are more conductive then weak acids and bases
ii. Conductivity increases with the number of ions in solution
iii. Conductivity is measured in Siemens. The higher the value, the more conductive, the
stronger the acid/base.
iv.
31
Version 2/2014
3. Unless there are ions present a substance cannot be an acid
a. Acids are not acids until they are dissolved in water
b. Acids must be dissolved in water in order to create the H3O+ ion through dissociation.
c. A substance that is a solid but will form an acid when it dissolves is called an acid anhydride
d. (Bronsted Lowry Acids and Bases can be gases, however)
4. Classification
a. Arrhenius Acid
i. Acids donate protons
ii. HA  H+ + Aiii. HA +H2O  H3O+ + Aiv. Presence of H+ (really H3O+, hydronium ion) ions in solution creates acidic solution
v. Definition we’ve been using
vi. Examples:
1. Sulfuric: H2SO4
2. Hydrochloric: HCl
3. Nitric: HNO3
4. Perchloric: HClO4
5. Phosphoric: H3PO4
6. Acetic: CH3COOH
b. Arrhenius Base
i. Bases donate hydroxide ion
ii. BOH  B+ + OHiii. Presence of hydroxide (OH-) ions in solution creates basic solution
iv. Definition we’ve been using
v. Examples
1. Sodium Hydroxide: NaOH
2. Potassium Hydroxide: KOH
3. Calcium Hydroxide: Ca(OH)2
4.
c. Bronsted-Lowery
i. Acids donate protons (same as Arrhenius)
ii. Bases accept protons (different from Arrhenius)
1. Example: Ammonia
2. NH3 + H2O  NH4+ + OH3. Ammonia takes a proton from a water molecule which creates a hydroxide ion
4. Presence of OH- creates basic solution
d. Lewis Acids and Bases – an even broader definition of acids and bases
i. Acids accept an electron pair
ii. Bases donates an electron pair (forms a dative covalent bond!)
1. Example: Ammonia
2. NH3 + H2O  NH4+ + OH32
Version 2/2014
3. NH3 has a lone pair (draw the lewis structure to verify for yourself). It
donates the electron pair to a hydrogen ion from water. Water is acting as a
Lewis acid because it is accepting the electron pair from ammonia.
4. Example: Metal ions
5. Fe2+ + H2O  Fe(H2O)6]2+
6. Metal ions accept a lone pair from the oxygen on the water molecule. Up to 6
water molecules can do this for various metal ions. These are dative covalent
bonds, also called coordinate covalent bonds.
7. Anything with a lone pair can act as a Lewis Base.
8. All Lewis Bases are Bronsted Lowery bases…the definitions are consistent with
each other. However, not all Lewis bases are Bbronsted-Lowry bases (metal ions
that for coordination compounds – see iron example above)
5. pH
e. Monoprotic vs Polyprotic
i. Acids that have one proton to lose are monoprotic
1. Example: HCl, HNO3
ii. Acids that have more than one proton to lose are polyprotic
1. Diprotic: H2SO4
2. Triprotic: H3SO4
f. Oxyacids = acids involving a polyatomic oxyanion (anion containing oxygen)
i. Ex: H2SO4
ii. Non example: HCl
g. Amphoteric compounds
i. Compounds that can act as both acid and base
1. Example: Water
2. H2O + CO32-  OH- + HCO3-1 (water acts as proton donor)
3. H2O + H+  H3O+ (water acts as proton acceptor)
a.
b.
c.
d.
e.
f.
presence of H3O+ tells the strength of the acid
This is measured by pH
pH = -log [H3O+]
pH < 7 = acid
pH > 7 = base
Why? It comes from the dissociation of water
33
Version 2/2014
Keq  [ H 3 O  ][OH  ]  1x10 14
log[ H 3 O  ]  log[ OH  ]  log(1x10 14 )
log[ H 3 O  ]  log[ OH  ]  14
 log[ H 3 O  ]  log[ OH  ]  14
pH  pOH  14
In a neutral solution there is just as much H3O+ as OH-, therefore the pH must be 7 when the
solution is neutral
Keq  [ H 3 O  ][OH  ]  1x10 14
If solution is neutral, [ H 3 O  ]  [OH  ]
[ H 3 O  ] 2  1x10 14
[ H 3 O  ]  1x10 14  1x10 7
 log(1x10 7 )  7
How does Kw change with temperature? Since the dissociation/ionization of water is
endothermic, as temperature is increased equilibrium shifts to the right (towards the products)
and the dissociation of water increases. Therefore, Kw increases with temperature.
Therefore, at temperatures other than 25 C, the Kw is not equal to 1x10-14 and the pH will not be
7.
Example: Calculate the pH of a water at 40 C if the Kw at 40 C is 2.5x10-14
[H+][OH-] = 2.5x10-14, [H+]=[OH-] = x
x2 = 2.5x10-14
x = [H+] = square root (2.5x10-14) = 1.58 x 10-7 M
pH = -log (1.58 x 10-7) = 6.8
Calculations involving pH
1. Find pH given [H3O+]
a. If the concentration of [H3O+] is 1.3 x 10-7
b. pH = - log (1.3 x 10-7) = 6.9
2. Find pOH given pH
a. If pH = 6.9
b. pOH = 14 – 6.9 = 7.1
3. Find [H3O+] given pH
34
Version 2/2014
a. If pH = 5
b. [H3O+] = 10^(-pH) = 1 x 10-5 M
4. Find [OH-] given [H3O+]
a. If [H3O+] = 1 x 10-5 M
b. 1 x 10 -14 M = [H3O+][OH-]  [OH-] =(1 x 10 -14 M)/ (1 x 10-5 M) = 1 x 10-9 M
5. When pH changes by 1, this indicates a 10 fold change in hydrogen ion concentration
a. pH is a log scale, so a change of 1 means the change in [H+] =101= 10
b. pH change of 2  [H+] =102= 100
Neutralization of Acids and Bases
1. Reaction with acid and base can make a neutral solution
2. H+ ions of acid cancel out OH- ions of base
3. This property can be used to determine concentrations
If it takes 10 L of HCl to neutralize 20 L of 4 M NaOH, what is the concentration of HCl?
1. write a balanced neutralization reaction
a. HCl + NaOH  NaCl + H2O
2. Find the number of moles of NaOH used
a. 4 M = x /20 L  x = 80 moles NaOH
3. Find the moles of HCl
a. According the reaction, there is a 1:1 ratio of HCl to NaOH
b. 80mol NaOH 
1 mol HCl
 80mol HCl
1 mol NaOH
4. Find molarity
a. Molarity = 80 mole/ 10 L = 8 M
Topic 8 Acids and Bases Exam Questions
1.
For equal volumes of 1.0 mol dm–3 solutions of hydrochloric acid, HCl(aq), and methanoic acid,
HCOOH(aq), which statements are correct?
I.
HCl dissociates more than HCOOH
II.
HCl is a better electrical conductor than HCOOH
III. HCl will neutralize more NaOH than HCOOH
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
2.
When equal volumes of four 0.1 mol dm–3 solutions are arranged in order of increasing pH (lowest pH first),
what is the correct order?
A.
CH3COOH < HNO3 < CH3CH2NH2 < KOH
B.
HNO3 < CH3COOH < CH3CH2NH2 < KOH
35
Version 2/2014
C.
D.
CH3CH2NH2 < HNO3 < CH3COOH < KOH
KOH < CH3CH2NH2 < CH3COOH < HNO3
(Total 1 mark)
3.
What is the approximate pH of a 0.01 mol dm–3 ammonia solution?
A.
2
B.
More than 2 but less than 7
C.
More than 7 but less than 12
D.
12
(Total 1 mark)
4.
Which are definitions of an acid according to the Brønsted-Lowry and Lewis theories?
Brønsted-Lowry theory
Lewis theory
A.
proton donor
electron pair acceptor
B.
proton acceptor
electron pair acceptor
C.
proton acceptor
electron pair donor
D.
proton donor
electron pair donor
(Total 1 mark)
5.
According to the Brønsted-Lowry theory, how does each species act in the equilibrium below?
CH3COOH2+ + HSO4–
CH3COOH + H2SO4
CH3COOH
H2SO4
CH3COOH2+
HSO4–
A.
acid
base
base
acid
B.
acid
base
acid
base
C.
base
acid
base
acid
D.
base
acid
acid
base
(Total 1 mark)
6.
Which list contains only strong acids?
A.
CH3COOH, H2CO3, H3PO4
B.
HCl, HNO3, H2CO3
C.
CH3COOH, HNO3, H2SO4
D.
HCl, HNO3, H2SO4
(Total 1 mark)
7.
Which pH value is that of an aqueous solution of carbon dioxide?
A.
2.1
B.
5.6
C.
9.8
D.
12.2
(Total 1 mark)
8.
What is the formula of the conjugate base of the hydrogenphosphate ion, HPO42–?
36
Version 2/2014
A.
B.
C.
D.
H2PO4–
H3PO4
HPO4–
PO43–
(Total 1 mark)
9.
An example of a strong acid solution is perchloric acid, HClO4, in water. Which statement is correct for this
solution?
A.
HClO4 is completely dissociated in the solution.
B.
HClO4 exists mainly as molecules in the solution.
C.
The solution reacts only with strong bases.
D.
The solution has a pH value greater than 7.
(Total 1 mark)
10.
(i)
Define a Lewis acid and state an example that is not a Brønsted-Lowry acid.
(2)
(ii)
Draw structural formulas to represent the reaction between the Lewis acid named in (i) and a Lewis
base and identify the nature of the bond formed in the product.
(4)
(Total 6 marks)
11.
The equations of two acid-base reactions are given below.
Reaction A
NH3(aq) + H2O(l)

NH 4 (aq) + OH–(aq)
The reaction mixture in A consists mainly of reactants because the equilibrium lies to the left.
Reaction B
NH2–(aq) + H2O(l)
NH 3 (aq) + OH–(aq)
The reaction mixture in B consists mainly of products because the equilibrium lies to the right.
(i)
For each of the reactions A and B, deduce whether water is acting as an acid or a base and explain your
answer.
(2)
(ii)
In reaction B, identify the stronger base, NH2– or OH– and explain your answer.
(2)
(iii)
In reactions A and B, identify the stronger acid, NH4+ or NH3 (underlined) and explain your answer.
(2)
(Total 6 marks)
12.
(a)
Describe two different experimental methods to distinguish between aqueous solutions of a strong base
and a weak base.
(5)
(b)
Two acidic solutions, X and Y, of equal concentrations have pH values of 2 and 6 respectively.
37
Version 2/2014
(i)
Calculate the hydrogen ion concentrations in the two solutions and identify the stronger acid.
(2)
(ii)
Determine the ratio of the hydrogen ion concentrations in the two solutions X and Y.
(1)
(Total 8 marks)
13.
(i)
Define a Brønsted-Lowry acid.
(1)
(ii)
Deduce the two acids and their conjugate bases in the following reaction:
OH–(aq) + NH4+(aq)
H2O(l) + NH3(aq)
(2)
(iii)
Explain why the following reaction can also be described as an acid-base reaction.
F–(g) + BF3(g)
BF4–(s)
(2)
(Total 5 marks)
14.
Ethanoic acid, CH3COOH, is a weak acid.
(i)
Define the term weak acid and state the equation for the reaction of ethanoic acid with water.
(2)
(ii)
Vinegar, which contains ethanoic acid, can be used to clean deposits of calcium carbonate from the
elements of electric kettles. State the equation for the reaction of ethanoic acid with calcium carbonate.
(2)
(Total 4 marks)
Topic 8 Acids and Bases Exam Question Markscheme
1.
A
2.
B
3.
C
4.
A
5.
D
6.
D
7.
B
8.
D
9.
A
[1]
[1]
[1]
[1]
[1]
[1]
[1]
[1]
[1]
38
Version 2/2014
10.
(i)
(ii)
(Lewis acid) electron pair acceptor;
appropriate example (such as AlCl3, BF3 etc.);
2
structural formula of Lewis acid (e.g. BF3, AlCl3, Transition element etc.);
structural formula of Lewis base (e.g. NH3, H2O etc.);
structural formula of product (e.g. F3BNH3 etc.);
dative covalent (bond)/coordinate (bond);
Penalize missing structural formulas once.
4
[6]
11.
(i)
(ii)
(iii)
acid in both reactions;
because it loses a proton/hydrogen ion/H+ / proton/hydrogen ion/H+ donor;
Second mark can be scored if they do not identify it as an acid in
both reactions.
2
NH2–;
more readily accepts a proton / equilibrium lies to the right /
takes H+ from H2O;
If OH– chosen award [0]
2
NH4+;
donates a proton more readily than NH3 / equilibrium lies to the left;
If NH3 chosen award [0]
2
[6]
12.
(a)
solutions of the same concentration;
pH meter;
strong base has a higher pH / weak base has lower pH;
indicator paper/U.I solution;
strong base has a higher pH/more purple / weak base has lower pH/blue not
purple / OWTTE;
measuring conductivity (with conductivity meter);
strong base has a higher conductivity / weak base has lower conductivity;
comparing heat of neutralization with acid;
strong base releases more heat / weak base releases less heat;
Award [4 max] for two correct methods with expected results.
(b)
(i)
(ii)
5
X;
[X] = 10–2 (mol dm–3) and [Y] = 10–6 (mol dm–3);
2
10 000/104 :1;
Ratio should be in form above.
1
[8]
13.
(i)
donates a proton / H+ ion;
(ii)
(acid)
1
(conjugate base)
39
Version 2/2014
(iii)
H2O
OH–;
NH4+
NH3;
[1 max] if all four acids and bases given but not clearly paired.
2
Lewis acid accepts an electron pair / Lewis base donates an electron pair;
F– is the base / BF3 is the acid;
2
[5]
14.
(i)
(ii)
partially dissociated or ionized;
CH3COOH + H2O
CH3COO– + H3O+ / CH3COOH
required for mark.
CH3COO– + H+;
2CH3COOH + CaCO3 → Ca(CH3COO)2 + CO2 + H2O
Award [1] for correct reactants and products and [1] for balancing.
2
2
[4]
40
Version 2/2014