A.P. CHEMISTRY
SCHEDULE: CHAPTER 14
TERM 2 2015-16
ACTIVITIES
Determine New Groups. Be sure to read prior to class.
We will fly through review topics – properties of acids/bases; ionization equations, Ka, Arrhenius definition;
Bronsted-Lowry definition; conjugate acids and bases; acid/base strength; proton transfer reactions
Lecture, Chapter 14.1 pp.622-635.
Tues.
DATE
1-12
Wed.
1-13
Introduce Experiment: Acid Base Titration.
DUE: Problem Set 14.1. work time = 15 min.; QUIZ
Prob’s,Even #’s 16, 28, 30, 32ab, 34-40, 42ace, 44,48, 50,78-82 E.C. 144 pp. 673-679
Discuss 1, 2, 22
Make Base after Quiz - follow Acid Base lab instructions
Thurs.
1-14
HAND IN PRE-LAB BEFORE STARTING EXPERIMENT.
Complete Experiment: Acid Base Titration.
Note: If you don’t have time to finish lab, you can forgo work time for PS 14.2 in two days to finish in class, otherwise
you will need to come in before/after school to complete the lab.
Fri.
1-15
Post lab WS Due. Titration calculations DUE. Turn in lab notebooks.
In your notebook: Include a hand-written conclusion with your calculated GEM of your acid. (You won’t be able
to do a % error yet as you won’t have the theoretical value until after you turn in your notebook).
Experiment Quiz.
Start writing an error analysis. (Due day of PS 14.3 Quiz – in 5 days)
Tues.
1-19
Lecture Chapter 14.2 pp. 635-655.
Wed.
1-20
PS 14.2 a DUE: Problems Even #’s 56,58, 64, 66, 70& supplements pp. 673-679; E.C. #60
Finish PS 14.2 notes; Experiment Quiz (Note: You may also be able to do 72,74,76 from PS 14.2b)
Thurs
1-21
Problem Set. 14.2b DUE work time = 10 min. - Quiz
Problems Even #’s 62,72-76, 86b, 96 pp. 698-705; E.C. #134
Start Lecture Chapter 14.3 w/35 min left pp. 655-667; 683-196.
Fri.
1-22
DUE: Finish pg. 13-14 in notepacket (as homework)
Problem Set 14.3 Due – 30 min group time; Quiz
Even #’s 19, 100-106, 112, 114, 118, & supplements pp 673-679; E.C. 88, 138abc, 150
Know what ions are acids/bases (#100), know how to write rxns (#112) etc.
Clicker Practice Quiz
Mon.
1-25
Quiz- PS 14.3 (You will need to study for this quiz.)
Review Chapter 14 & Titration Experiment.
Magic Show Preparation – Make all solutions/prepare all chemicals for tomorrow
Tues.
1-26
TEST CHAPTER 14 & TITRATION EXPERIMENT
114 – see notes pg. 7 or book (review)
118- review from ch. 4 and 7 (also last
page of note packet – not done in
class)
NOTE: There will be a predicting reactions portion of the test. The only reactions included will be those presented in this unit - AcidBase reactions (all models), as well as acid anhydride & base anhydride reactions. See the predicting reactions section of your 3ring binder.
Examples and points explained on the back.
Directions for Predicting Reactions.
You will be asked to give the balanced net ionic equation for 3 chemical reactions.
NOTE: 1 pt. will be awarded for correctly written reactants
2 pt. will be awarded for correctly written products. (1 pt. will be deducted for including spectator ions.)
1 pt. for properly balancing the equation
1 pt. for properly answering a follow-up question
(5 pt. total for each reaction = 15 pt. total)
Example: A strip of magnesium is added to a solution of silver nitrate. (not from this unit)
Mg + 2 Ag+ --> Mg2+ + 2 Ag
What are the spectator ions? NO3- ions
Example: Solutions of Acetic acid and sodium hydroxide are mixed.
CH3COOH + OH- --> H2O + CH3COOWhat would you expect the resulting solution to be acidic, basic, or neutral? basic
REMEMBER:
* These are NET IONIC EQUATIONS:
• soluble ionic compounds & strong acids (HCl, HBr, HI, HNO3, HClO4, H2SO4)must be written as ions.
• Molecular compounds are written as molecules
* H2CO3 always decomposes into H2O +CO2
• A “solution of CO2 ” reacts as would H2CO3
ex. Solutions of carbon dioxide and sodium fluoride are mixed would look as follows:
CO2 + H2O + F- --> HF + HCO3- (note: forgetting charges results in no points)
Ch. 14 Essential Questions:
• How can you predict and write net ionic reactions for the following reaction types:
acid/base and precipitate?
• How do chemical reactions reach a state of equilibrium?
• How can Le Chatelier’s principle be used to explain how various factors affect
chemical equilibria?
• How can you use equilibrium constants to calculate pH?
Common Misconceptions Acids and Bases
1.
2.
3.
4.
5.
6.
7.
8.
Solutions of strong acids always have lower pH than solutions of weak acids.
If there is OH in the formula then you have a base and if there is an H it has to be an acid.
Strength and concentration mean the same thing.
A neutral solution always has a pH of 7 regardless of the temperature.
The endpoint of all titrations is at pH 7.
Only pH of 1-14 exist.
Concentrated acids are pure substances.
The endpoint and equivalence point of a titration are the same thing.
Problem Set 14.1 –acid properties, pH, acid strength, Bronsted-Lowry model
1) Students will be able to define an acid and base in terms of both the Arrhenius and the Bronsted-Lowry definitions and provide
examples of each (including examples that follow the Bronsted-Lowry definition better than the Arrhenius definition).
2) Students will be able to predict relative concentrations of species (HA, H+, A-, HB, B-, OH-) in solution given Ka for various
solutions including:
a. An acid dissolved in water.
b. Two acids dissolved in water.
c. An acid and a weak (Bronsted-Lowry base) in water.
d. A polyprotic acid dissolved in water.
3) Students will be able to compare relative concentrations of H+ given the concentration of the dissolved acid as well as any
pertinent Ka values.
4) Students will be able to identify substances as acids/bases, or amphoteric.
5) Students will be able to compare the strength of an acid’s conjugate base to the strength of water as a base.
6) Students will be able to compare the strength of base’s conjugate acid to the strength of water as an acid.
7) Students will be able to identify (by name/formula) all six strong acids -.
8) Students will be able to recognize all types of acids including:
a.
mineral acids
ex. HF, HCl
b.
oxyacids
ex. HNO2, H2SO3
c.
organic acids
ex. CH3COOH, C6H5COOH
d.
conjugate bases of polyprotic acids ex. HS-, HCO3e.
ammonium salts
ex. NH4Cl (NH4+), CH3CH2NH3Cl (CH3CH2NH3+
f.
hydrated metal ions
ex. Cu(H2O)42+
9) Students will be able to explain the ionization of acid in water in terms of…
a. relative concentrations of species (H2O molecules, H+ions, and OH- ions) in solution.
b. Value of Kw.
c. H2O as an amphoteric substance - a conjugate base (of H3O+) and a conjugate acid (of OH--).
d. How concentrations of H+ and OH- would change (and thus pH) if the temperature of the solution would change.
e. Equal concentrations of H+ and OH- making water neutral.
10) Students will be able to interconvert between [H+], [OH-], pH, and pOH.
11) Students will be able to calculate the pKa of an acid (given Ka) and relate pKa to the strength of an acid.
12) Students will be able to predict the strength of an oxyacid based on its structure (i.e. number of oxygens or relative
electronegativity of the central atom).
13) Students will be able to list all properties of acids and bases discussed in class.
Problem Set 14.2a – Weak acid Ka and pH calculations
1) Students will be able to explain why the pH of a weak acid (i.e. 0.10 M HF) is not the same as the pH of a strong acid (i.e. 0.10
M HCl) by giving the pH of the strong acid and a quantitative estimate of the pH of the weak acid.
2) Students will be able to calculate the pH of a weak acid if given the value of Ka and the concentration of the weak acid.
3) Students will be able to calculate the Ka of a weak acid if given the pH and concentration of the acid.
4) Students will be able to calculate the pH of a polyprotic acid as well as the concentrations of all species in solution. (e.g.
Calculate the pH of 0.100 M H2SO3 and the concentrations of H2SO3, HSO3-, and SO32- in solution.)
Problem Set 14.2 b – Weak bases Kb and pH calculations
1) Students will be able to write the ionization of a weak base including:
a.
Molecules (amines)
ex. CH3NH2 + H2O  CH3NH3+ + OHb.
Anions (the conjugate base of a weak acid) ex. F- + H2O  HF + OH2) Students will be able to calculate Ka of the conjugate acid of an amine. (e.g. the Ka of CH 3NH3+ from Kb CH3NH2)
3) Students will be able to calculate the pH of a weak base if given the value of Kb and the concentration of the weak base.
4) Students will be able to calculate the Kb of a weak base if given the pH and concentration of the base.
Problem Set 14.3 – pH of salt solutions
1) Students will be able to define the meaning of a “salt” solution.
2) Students will be able to identify a salt solution as acidic/basic/neutral.
3) Students will be able to write an equation showing the acidic/basic properties of a salt solution.
4) Students will be able define an acid and base in terms of the Lewis definition by doing the following:
a. providing an example that demonstrates a substance that acts only as a lewis acid
b. providing an example that demonstrates a substance that acts as both a Lewis acid as well as a BronstedLowry/Arrhenius acid.
c. explaining an example that demonstrates a substance that acts as both a Lewis base as well as a BronstedLowry/Arrhenius base.
5) Students will be able to identify the ligand in a Lewis acid/base reaction and identify the coordination number.
6) Students will be able to predict the products of the following reactions (by writing net ionic equations):
a.
Base anhydride (metal oxide) + water  base
ex. MgO + H2O  Mg2+ + 2 OHb.
Acid anhydride (nonmetal oxide) + water  acid
ex. CO2 + H2O  H2CO3
c.
Acid anhydride (nonmetal oxide) + Base anhydride (metal oxide)  water + salt
ex. MgO + CO2  MgCO3
Problem Set 14.2 Supplements
1) Write the ionization for the following acids.
a. ZnCl2 (Zn(H2O)42+)
b. CH3NH3Br (CH3NH3+)
c. C6H5COOH
d. phosphoric acid
e. ammonium ion
2) Write the formula for the hydrated nickel (II) ion in water.
3) Which of the following ions would be hydrated in solution?
a. Cu2+
b. Mg2+
c. Na+
d. Ba2++
e. Co2+
f. K+
4) Draw the lewis dot structure for each of the following, AND label as acids or bases.
a. C2H5NH2
b. CH3NH3+
c. (CH3)2NH
d. (C2H5)2NH2+
5) Calculate the pKa for each of the following.
a. H2S
Ka = 1.0 x 10-7
b. HF
Ka = 6.7 x 10-4
c. HS
Ka = 1.2 x 10-15
+
d. NH4
Ka = 5.7 x 10-10
6) A. Which of the acids in #5 is the strongest? B.The weakest?
7) Which acid is stronger – HClO3 or HClO2? Explain in terms of acid structure.
Problem Set 14.3 Supplements
1) Identify the Lewis acid and the Lewis base in each of the following reactions.
a. B(OH)3 (aq) + H2O (l)  B(OH)4- (aq) + H+ (aq)
b. 2 NH3 (aq) Ag+ (aq)  Ag(NH3)2+
c. BF3 (g) + F- (aq) BF4- (aq)
2) Identify the Lewis acid and the Lewis base in each of the following reactions.
a. 6 H2O (l) + Fe3+ (aq)  Fe(H2O)63+ (aq)
b. CN- (aq) + H2O (l) +  HCN (aq) + OH- (aq)
c. HgI2 (s) + 2 I- (aq)  HgI4- (aq)
3) Write the formula for the species formed in the following lewis acid/base reactions. (Note: all products are complex ions.)
a. Nickel (II) ions in solution
b. Silver ions and ammonia
c. Iron (III) ions and the cyanide ion
d. Aluminum ions and the hydroxide ion
4) Identify the ligand for each set of reactants.
a.
b.
c.
d.
5) Would you expect Fe3+ or Fe2+ to be the stronger Lewis acid? Explain.
6) Cu2+ (aq) + 4 H2O (l)  Cu(H2O)42+ (aq )
a.
Identify the acid and the base in the above reaction.
b.
Is the acid in the equation above a(n) Arrhenius acid/Bronsted-Lowry acid/Lewis acid? (Circle all that apply.)
7) Calculate the value for the equilibrium constant for each of the following aqueous reactions.
a. NH3 + H3O+  NH4+ + H2O
b. NO2- + H3O+  HNO2 + H2O
c. NH4+ + OH-  NH3 + H2O
d. HNO2 + OH-  H2O + NO2_
8) Determine the oxidation number for the nonmetal other than oxygen.
a. NO2-
____
h.
PO33-
____
b. NO3-
____
i.
CO32--
____
c. SO32-
____
j.
N2O4
____
d. SO42-
____
k.
N2O5
____
e. ClO4-
____
l.
P2O5
____
f. ClO3-
____
m. SO2
____
g. PO43-
____
n.
____
SO3
9) Write balanced net ionic equations for each of the following.
a.
Solid lithium oxide is placed in distilled.
b.
Solid calcium oxide is placed in distilled water.
c.
Carbon dioxide gas is bubbled through distilled water.
d.
Sulfur dioxide gas is bubbled through distilled water.
e.
Sulfur trioxide gas is bubbled through distilled water.
f.
Potassium oxide and sulfur dioxide are heated strongly.
g.
Lithium oxide and sulfur trioxide are heated strongly.