1. Unit 4 (Chapter 4): Aqueous Reactions & Solution Stoichiometry
Chemistry, The Central Science, 10th edition
Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten
Unit 4 (Chapter 4): Aqueous Reactions & Solution Stoichiometry
John D. Bookstaver
St. Charles Community College
St. Peters, MO
 2006, Prentice Hall, Inc.

2. Salts: Ionic Solids: (metal-nonmetal)
dissociate (dissolve) by separation into ions
Electrolytes:
ions in solution that conduct electricity

3. Non Weak Strong C11H22O11 CH3OH H2O NaOH HNO3 KCl CH3COOH HNO2 NH3
ions
SOME
ions
ALL
ions
only molecules
partially ionize HW
p.159 #33
completely dissociate

4. Electrolytes: Strong, Weak, or Non?
(ions conduct electricity)
Compound
metal-nonmetal
nonmetals (Covalent)
Ionic
Molecular
Acid
(H____)
Not Acid
STRONG
KBr
CaI2
FeCl3
NaOH
Ca(OH)2
(strong bases)
STRONG
(6)
WEAK
(& NH3)
NON
C11H22O11
C2H5OH
H2O
HCl, HBr, HI
HNO3
H2SO4
HClO4
CH3COOH
HNO2 HF
HW p #1,2,4,5,38

5. Precipitation Reactions
Double Replacement:
(precipitate)
2 KNO3(aq) + PbI2(s)
2 KI(aq) + Pb(NO3)2(aq) 
precipitate:
insoluble ionic compound
(as predicted by solubility rules)
Pb2+ I–

6. Solubility Rules ALWAYS Soluble Ions: * WS Solubility & NIE’s #1
Na+, K+, etc. group I (alkali metals)
NH ammonium
NO3– nitrate
HCO3– bicarbonate
C2H3O2– (CH3COO–) acetate (ethanoate)
ClO3–, ClO4– chlorate, perchlorate *
WS Solubility & NIE’s #1
Common Precipitates form with: examples
Ag+, Pb2+, Hg2+ (halogens) AgCl, PbI2
OH– (hydroxide) Cu(OH)2
CO32– (carbonate) CaCO3

7. HCl + H2O  H3O+ + Cl– NH3 + H2O  NH4+ + OH– + – + – O O Cl Cl
Acid: proton (H+) donor
Base: proton (H+) acceptor
NH3 + H2O  NH OH– H + – H N O N O H H H H H H H H

8. Strength of Acids and Bases
STRONG (complete dissociation)
HA(aq) H2O(l) H3O+(aq) A–(aq)
B(aq) H2O(l) BH+(aq) OH–(aq)
WEAK (partial dissociation)
HA(aq) H2O(l) H3O+(aq) A–(aq)
B(aq) H2O(l) BH+(aq) OH–(aq)
Why is HF considered a weak acid if it etches glass? H bond together and even form complexes like F- - H3O+

9. Strong Acids: Only 6 strong acids: Nitric (HNO3) Sulfuric (H2SO4)
Hydrochloric (HCl)
Hydrobromic (HBr)
Hydroiodic (HI)
Perchloric (HClO4)
Strong Acids:
HI + H2O  H3O+ + I–
proton (H+) donors

10. ase Strong Bases: The strong bases are soluble hydroxides (OH–) of…
Group 1 (Li,Na,K)
CBS (Ca, Ba, Sr)
Mg(OH)2 not as soluble
Strong Bases:
OH– + H3O+  H2O + H2O
proton (H+) acceptors
ase

11. Electrolytes: Strong, Weak, or Non?
QUIZ!!!
(at the bell)
Electrolytes: Strong, Weak, or Non?
Compound
metal-nonmetal
nonmetals (Covalent)
Ionic
Molecular
Acid
(H____)
Quiz – Electrolytes (to start class)
Not Acid
STRONG
STRONG
(6)
WEAK
(& NH3)
NON

12. Acid-Base Neutralization Reactions
strong acid (H+A–)
strong base (M+OH–)
ionic
compound
(M+A–)
water
H2O
(HOH)
ACID BASE SALT + WATER + – O Cl Na Na Cl H O H H H
HCl(aq) NaOH(aq) NaCl(aq) + H2O(l)
HW p.158 #40a

13. AgNO3(aq) + KCl(aq)  AgCl(s) + KNO3(aq)
Molecular Equation
The molecular equation lists the reactants and products in their molecular form.
AgNO3(aq) + KCl(aq)  AgCl(s) + KNO3(aq)

14. Ionic Equation AgNO3(aq) + KCl(aq)  AgCl(s) + KNO3(aq)
In the ionic equation all strong electrolytes (strong acids, strong bases, and soluble ionic salts) are dissociated into their ions.
This more accurately reflects the species that are found in the reaction mixture.
AgNO3(aq) + KCl(aq)  AgCl(s) + KNO3(aq)
Ag+(aq) + NO3–(aq) + K+(aq) + Cl–(aq) 
AgCl(s) + K+(aq) + NO3–(aq)

15. Net Ionic Equation Ag+(aq) + NO3–(aq) + K+(aq) + Cl–(aq) 
Cross out Spectator Ions that do not change (same state & same charge) from the left side of the equation to the right.
The only species left are those things that react (change) during the course of the reaction.
Ag+(aq) + NO3–(aq) + K+(aq) + Cl–(aq) 
AgCl(s) + K+(aq) + NO3–(aq)
NIE: Ag+(aq) + Cl–(aq)  AgCl(s)

16. Balanced Net Ionic Equations
comp – diss – cross – net – bal
Write a complete molecular equation.
Dissociate all strong electrolytes (aq) .
Cross out spectators (same charge & state)
Write the net ionic equation with the species that remain and balance it.
(solubility rules)

17. Balanced Net Ionic Equations
comp – diss – cross – net – bal + 2– 2+ – + –
(NH4)2SO4 + Ba(NO3)2 →
NaOH + MgBr2 →
BaSO4 + NH4NO3
Ba SO42– → BaSO4 + – 2+ – + –
NaBr + Mg(OH)2
Mg OH– → Mg(OH)2(s)
HW p.158 #21

18. Neutralization Reactions
When a Strong Acid reacts with a Strong Base, the net ionic equation is…
H+ + OH–  H2O
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
H+ + Cl– + Na+ + OH–  Na+ + Cl– + H2O
Salt is (metal/nonmetal) is dissolved (aq)

19. Neutralization Reactions
When a Weak reacts with a Strong, the net ionic equation is…
HX + OH–  X– + H2O
HF(aq) + KOH(aq)  KF(aq) + H2O(l)
HF + Na+ + OH–  Na+ + F– + H2O
Salt is (metal/nonmetal) is dissolved (aq)
HW p.159 #40 (finish)

20. Balanced Net Ionic Equations
comp – diss – cross – net – bal + 2– 2+ – + –
(NH4)2SO4 + Ba(NO3)2 →
BaSO4 + NH4NO3
Ba SO42– → BaSO4(s) + – + –
HF(aq) + KOH(aq)  KF(aq) + H2O(l)
HF + OH–  F– + H2O
WS Solubility & NIE’s #2

21. Gas-Forming Reactions
H2 Demo
(M0)
(H+)
(M+)
(gas)
Single Rep: Metal + Acid Metal Ion + H2
Ex: Zn(s) + H2SO4(aq) ZnSO4(aq) + H2(g)
NIE: Zn(s) + 2 H+(aq) Zn2+(aq) + H2(g) + 2– 2+ 2–
(gas)
H2O(l) + CO2(g)
CO2 Demo
(H+)
(CO32–)
Double Rep: Acid + Carbonate Salt + Ex: HCl(aq) + CaCO3(s) CaCl2(aq) + H2O(l) + CO2(g) NIE: 2 H+(aq) + CaCO3(s) Ca2+(aq) + H2O(l) + CO2(g)
H2CO3(aq)
(or Bicarbonate)
(HCO3–)
(decomposes
immediately)
H2 Demo – small bit of Zn in large test tube, then add but of 6 M H2SO4, hold burning splint for test.
CO2 - Demo
HW p. 159 #43
CH3COOH + NaHCO3 
CH3COONa +
H2O + CO2

22. Oxidation-Reduction Reactions (REDOX)
video clip
(One
cannot
occur
without
the other)
LEO says GER
lose e-, charge up….gain e-, charge down (Video – OxRed I)

23. Oxidation Numbers Is it a redox reaction? To find out…
assign oxidation numbers* (or oxidation states) to each element in a reaction.
check if any oxidation states changed
(↓ reduced , ↑ oxidized)
*oxidation numbers
of elements describe electrons that would be lost or gained IF the compound was 100% ionic.
of ions show electrons transferred IN an ionic compound
*charges

24. Assigning Oxidation Numbers
All elements are 0. (all compounds are 0)
Monatomic ion is its charge. (Ex. Na+ ion)
Most nonmetals tend to be negative, but some are positive in certain compounds or ions. (Ex. SO3)
O is −2 always, but in peroxide ion is −1 (O2–2).
H is +1 with nonmetals, −1 with a metals.
F is always −1.
other halogens are −1, but can be positive, like in oxyanions.
Ex. ClO3– or NO3– or SO42–

25. Oxidation Numbers
The sum of the ox. #’s in a neutral compound is 0.ex: NaCl = 0
The sum of the ox. #’s in a polyatomic ion is the charge on the ion. Ex: PO4=-3
Calculate the oxidation number of each:
Sulfur in… SO3
Chromium in… K2Cr2O7
Nitrogen in… NH4+
Cobalt in… [CoCl6]3–

26. Classifying REDOX Reactions
All rxns (but…NOT double replacement)
Synthesis
A + B → AB 2 → 1
( → +/–)
Decomposition
AB → A + B 1 → 2
(+/– → )
Single Replacement
AB + C → A + CB
(+/– → /–)
Combustion
CxHy + O2 → CO2 + H2O
(–/ → +/– /–)
Zn in acid releases H2 gas 26

28. Single Replacement (REDOX)
silver ions
oxidize
copper metal
What is reduced?
Cu(s) Ag+(aq)  Cu2+(aq) Ag(s) X
Cu2+(aq) Ag(s)  Cu(s) Ag+(aq)

29. Activity Series of Metals
increasing ease of oxidation
Cannot displace H+ from acid to make H2(g)

30. Writing REDOX Reactions
Write an overall equation for the synthesis of calcium sulfide from its elements. Then write the RED and OX half-equations to identify the REDOX process. +2 –2
Ca + S  CaS
OX:
Ca0  Ca+2
+ 2 e–
RED:
2 e– +
S0  S–2
Ca + S  CaS

31. Writing REDOX Reactions
Write an overall equation for the decomposition of aluminum oxide into its elements. Then write the RED and OX half-equations to identify the REDOX process. +3 –2
Al2O3  Al + O2
3 ( )
OX: 2
O–2  O20
+ 4 e– 6
O–2  3 O20
+ 12 e–
4 ( )
RED:
3 e– +
Al+3  Al0
12 e– +
4 Al+3  4 Al0
2 Al2O3  4 Al O2

32. Writing REDOX Reactions
Write the net ionic equation for the reaction of solid zinc in a solution of hydrochloric acid.
comp – diss – cross – net – bal +1 –1 +2 –1
Mg(s) + HCl(aq) 
MgCl2(aq) + H2(g)
Mg + H+  Mg H2 2
Classify the reaction in two ways.
Single-Replacement and Redox

33. Mg + 2 H+  Mg2+ + H2(g) ox red What is red & what is ox? WS 5c #1-2
Title:
Reaction of magnesium with acid.
Caption:
The bubbles are due to hydrogen gas.

34. + Solutions: homogeneous mixtures. ( _____ throughout) same
solvent is present in greatest abundance.
solute dissolved in/by solvent +
same

35. Molarity units: mol/L or mol∙L–1 moles of solute (mol) Molarity (M) =
Molarity (M) is a measure of the concentration of a solution.
moles of solute (mol)
liters of solution (L)
Molarity (M) =
units: mol/L or mol∙L–1
What is the molarity of a solution with 29.2 g of sodium chloride in 250. mL of water?
mol/L like g/mol
1 mol NaCl
58.44 g NaCl
2.00 M
NaCl
0.500 mol NaCl
29.2 g NaCl x = =
0.250 L 35

36. Preparing a Solution 1-mass solute 2-add solvent, swirl to dissolve
WS #1-2 Conc. Calc’s
1-mass solute
2-add solvent, swirl to dissolve
3-add solvent to mark
HW p.160 #60, 67 36

37. WS Concentration & Dilutions#1
1 mol NaHCO3
84.01 g NaHCO3
1 L NaHCO3
5.00 g NaHCO3 x x =
0.100 mol NaHCO3
0.595 L
NaHCO3 #2
1.20 mol CuSO4
1 L CuSO4
g CuSO4
0.275 L CuSO4 x x =
1 mol CuSO4
52.7 g
CuSO4

38. Dilution M1V1 = M2V2 1-calc M1V1=M2V2 2-pipet V1 from concentrated
3-fill to mark w DI 38

39. Dilution
WS #3-4 Dilutions
M1V1 = M2V2
What volume of water must be added to prepare 2.0 L of 3.0 M CuSO4 from a stock solution of concentration 8.0 M ? 39

40. WS Concentration & Dilutions
M1V1 = M2V2
M1V1 = M2V2 #3
(12.0 M)V1 = (1.25 M)(500 mL)
V1 = 52.1 mL ( L) #4
M1V1 = M2V2
(2.50 M)V1 = (0.200 M)(250 mL)
V1 = 20.0 mL ( L)

41. Solution Stoichiometry
Rxn: A(aq) B(aq)  C D
molar mass A
molarity A (M)
g A
mol A
L of A
g A
1 mol A
mol A
1 L
mol-to-mol ratio
HW p. 161 #81
g B
1 mol B
mol B
1 L
molar mass B
molarity B (M)
g B
mol B
L of B 41