1. Reactions in Aqueous Solution
Chapter 4

2. A solution is a homogenous mixture of 2 or more substances
The solute is(are) the substance(s) present in the smaller amount(s)
The solvent is the substance present in the larger amount
Solution
Solvent
Solute
Soft drink (l)
H2O
Sugar, CO2
Air (g) N2
O2, Ar, CH4
Soft Solder (s) Pb Sn
4.1

3. An electrolyte is a substance that, when dissolved in water, results in a solution that can conduct electricity.
A nonelectrolyte is a substance that, when dissolved, results in a solution that does not conduct electricity.
nonelectrolyte
weak electrolyte
strong electrolyte
4.1

4. Ionic compounds that dissolve in water
always produce solutions that are strong
electrolytes.
Molecular compounds that dissolve in water may produce nonelectrolytes, weak electrolytes (weak acids or bases), or strong electrolytes (strong acids).

5. Conduct electricity in solution?
Cations (+) and Anions (-)
Strong Electrolyte – 100% dissociation
NaCl (s) Na+ (aq) + Cl- (aq)
H2O
Weak Electrolyte – not completely dissociated
CH3COOH CH3COO- (aq) + H+ (aq)
Reversible reaction - Chemical equilibrium
reversible

6. Nonelectrolyte does not conduct electricity?
No cations (+) and anions (-) in solution
C6H12O6 (s) C6H12O6 (aq)
H2O
Strong Electrolyte
Weak Electrolyte
Nonelectrolyte
HCl
CH3COOH
(NH2)2CO
HNO3 HF
CH3OH
HClO4
HNO2
C2H5OH
NaOH
H2O
C12H22O11
All Ionic Compounds
4.1

7. Water’s power as an ionizing solvent results from the distribution of its electrons and its overall shape.
A polar molecule is one that has a + end and a - end. Water is one of the most polar molecules.

8. Hydration is the process in which an ion is surrounded by water molecules arranged in a specific manner. d+ d-
H2O

10. The Solubility of Ionic Compounds in Water
The solubility of ionic compounds in water depends upon the relative
strengths of the electrical forces between ions in the ionic compound
and the attractive forces between the ions and water molecules in the
solvent. There is a tremendous range in the solubility of ionic
compounds in water! The solubility of so called “insoluble” compounds
may be several orders of magnitude less than ones that are called
“soluble” in water, for example:
Solubility of NaCl in water at 20oC = 365 g/L
Solubility of MgCl2 in water at 20oC = g/L
Solubility of AlCl3 in water at 20oC = 699 g/L
Solubility of PbCl2 in water at 20oC = 9.9 g/L
Solubility of AgCl in water at 20oC = g/L
Solubility of CuCl in water at 20oC = g/L

11. Problems from Chapter 4
Pages
1-4, 6-12, 14-18, 20-23, 25-36, 40,
42, 46, 48, 50, 54, 56, 58, 60-62,
64-66, 72, 74, 80, 81, 92, 94

12. Precipitation Reactions
Precipitate – insoluble solid that separates from solution
PbI2
precipitate
Pb(NO3)2 (aq) + 2NaI (aq) PbI2 (s) + 2NaNO3 (aq)
molecular equation
Pb2+ + 2NO3- + 2Na+ + 2I PbI2 (s) + 2Na+ + 2NO3-
ionic equation
Pb2+ + 2I PbI2 (s)
net ionic equation
Na+ and NO3- are spectator ions
4.2

13. Writing Net Ionic Equations
Write the balanced molecular equation.
Determine precipitate from solubility rules
Write the ionic equation showing the soluble strong electrolytes as ions.
Cancel the spectator ions on both sides of the ionic equation
Write the net ionic equation for the reaction of silver
nitrate solution with sodium chloride solution.
AgNO3 (aq) + NaCl (aq) AgCl (s) + NaNO3 (aq)
Ag+ + NO3- + Na+ + Cl AgCl (s) + Na+ + NO3-
Ag+ + Cl AgCl (s)
4.2

14. Molecular Equation: reactants and products are
written as if they were undissociated compounds.
Total Ionic Equation: all soluble ionic substances
(strong electrolytes) are dissociated into ions.
Net Ionic Equation: Eliminate spectator ions.
(These are ions which are unchanged in the
chemical reaction.)

15. Solubility Rules for Common Ionic Compounds In water at 250C
Soluble Compounds
Exceptions
Compounds containing alkali metal ions and NH4+
NO3-, HCO3-, ClO3-
Cl-, Br-, I-
Halides of Ag+, Hg22+, Pb2+
SO42-
Sulfates of Ag+, Ca2+, Sr2+, Ba2+, Hg2+, Pb2+
Insoluble Compounds
CO32-, PO43-, CrO42-, S2-
OH-
Compounds containing alkali metal ions and Ba2+
4.2

17. These precipitation reactions are one type of a class of reactions sometimes called double displacement or metathesis reactions. The neutralization reactions we will study next are also in the class of reactions.

18. Acids
Have a sour taste. Vinegar owes its taste to acetic acid. Citrus
fruits contain citric acid.
React with certain metals to produce hydrogen gas.
React with carbonates and bicarbonates to produce carbon
dioxide gas
Bases
Have a bitter taste.
Feel slippery. Many soaps contain bases.
4.3

19. Arrhenius acid is a substance that produces H+ (H3O+) in water
Arrhenius base is a substance that produces OH- in water
4.3

20. A Brønsted acid is a proton donor A Brønsted base is a proton acceptor
A Brønsted acid must contain at least one
ionizable proton!
4.3

21. Monoprotic acids Diprotic acids Triprotic acids HCl H+ + Cl-
Strong electrolyte, strong acid
HNO H+ + NO3-
Strong electrolyte, strong acid
CH3COOH H+ + CH3COO-
Weak electrolyte, weak acid
Diprotic acids
H2SO H+ + HSO4-
Strong electrolyte, strong acid
HSO H+ + SO42-
Weak electrolyte, weak acid
Triprotic acids
H3PO H+ + H2PO4-
Weak electrolyte, weak acid
H2PO H+ + HPO42-
Weak electrolyte, weak acid
HPO H+ + PO43-
Weak electrolyte, weak acid
4.3

22. Identify each as a Bronsted acid or base or both:
a. HI
b. CH3COO-
c. H2PO4-

23. EH Homework due Friday
Unit 5
Sec 1 Solubility Rules 90
Sec 2 Metathesis Rxns 80
Sec 3 Ionic/Net ionic Rxns 90
Sec 4 Balancing Eq. 90
Sec 5 Predicting Products omit
Sec 6 Activity Series
Unit 10
Sec 1 Species of Redox Rxns 90
Sec 2 Determining Oxidation No. 90

24. EH Homework due Sunday Unit 6
Sec 1 Molar Concentrations 90
Sec 2 Titration Problems 90
Sec 3 Volumetric Analysis 90
Sec 4 Molarity of Ions
Sec 5 Redox Titrations
Test on Chapters 3 and 4

25. Neutralization Reaction
acid + base salt + water
HCl (aq) + NaOH (aq) NaCl (aq) + H2O
H+ + Cl- + Na+ + OH Na+ + Cl- + H2O
H+ + OH H2O
4.3

26. Writing Balanced Equations for Neutralization Reactions
Problem: Write balanced chemical equations (molecular,
total ionic, and net ionic) for the following
chemical reactions:
a) barium hydroxide(aq) and hydroiodic acid(aq)
b) lithium hydroxide(aq) and nitric acid(aq)
c) sodium hydroxide(aq) and sulfuric acid(aq)

27. Oxidation-Reduction (redox) reactions involve a transfer of electrons
Oxidation-Reduction (redox) reactions involve a transfer of electrons. oxidation: loss of electrons reduction: gain of electrons reducing agent: substance oxidized oxidizing agent: substance reduced

28. Oxidation-Reduction Reactions
(electron transfer reactions)
2Mg (s) + O2 (g) MgO (s)
2Mg Mg2+ + 4e-
Oxidation half-reaction (lose e-)
O2 + 4e O2-
Reduction half-reaction (gain e-)
2Mg + O2 + 4e Mg2+ + 2O2- + 4e-
2Mg + O MgO
4.4

29. Oxidation number: the charge an atom would have if the shared electrons were transferred completely to the atom with the greater attraction for the electrons. An increase in oxidation number indicates oxidation. A decrease in oxidation number indicates reduction.

30. Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2
Oxidation number
The charge the atom would have in a molecule (or an
ionic compound) if electrons were completely transferred.
Free elements (uncombined state) have an oxidation number of zero.
Na, Be, K, Pb, H2, O2, P4 = 0
In monatomic ions, the oxidation number is equal to the charge on the ion.
Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2
The oxidation number of oxygen is usually –2. In H2O2 and O22- it is –1.
4.4

31. The oxidation number of hydrogen is +1 except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1.
Group 1 metals are +1, group 2 metals are +2 and fluorine is always –1.
6. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion.
Oxidation numbers of all the elements in HCO3-, HNO3, Na3PO4, IF7, KIO3, K2Cr2O7,
SO42- ?
4.4

32. Fig. 4.10

33. Types of Oxidation-Reduction Reactions
Combination Reaction
A + B C +4 -2
S + O SO2
Oxidized? Reduced? Oxidizing agent? Reducing agent?
Decomposition Reaction
C A + B +1 +5 -2 +1 -1
2KClO KCl + 3O2
4.4

34. Types of Oxidation-Reduction Reactions
(Single) Displacement Reaction
A + BC AC + B +1 +2
Sr + 2H2O Sr(OH)2 + H2
Hydrogen Displacement +4 +2
TiCl4 + 2Mg Ti + 2MgCl2
Metal Displacement -1 -1
Cl2 + 2KBr KCl + Br2
Halogen Displacement
Oxidized? Reduced? Oxidizing agent? Reducing agent?
4.4

35. The Activity Series for Metals
(Single) Displacement Reaction
M + BC MC + B
Reaction will occur only
if M is above B.
Ca + 2H2O Ca(OH)2 + H2
Pb + 2H2O Pb(OH)2 + H2
4.4

36. Zn (s) + CuSO4 (aq) ZnSO4 (aq) + Cu (s)
Zn Zn2+ + 2e-
Zn is oxidized
Zn is the reducing agent
Cu2+ + 2e Cu
Cu2+ is reduced
Cu2+ is the oxidizing agent
Nickel wire reacts with silver nitrate to form silver metal.
What is the oxidizing agent in the reaction?
Ni (s) + 2AgNO3 (aq) Ni(NO3)2 (aq) + 2Ag (s)
Ni Ni2+ + 2e-
Ni is oxidized Ni is the reducing agent.
Ag+ + 1e Ag
Ag+ is reduced
Ag+ is the oxidizing agent
4.4

37. 4.4

38. For displacement reactions involving halogens the reactivity is
F2> Cl2>Br2>I2
Cl2(aq) + NaBr(aq)  ?
Cl2(aq) + NaF(aq)  ?

39. Types of Oxidation-Reduction Reactions
Disproportionation Reaction
Element is simultaneously oxidized and reduced. +1 -1
Cl2 + 2OH ClO- + Cl- + H2O
Chlorine Chemistry
4.4

40. Classify the following reactions.
Ca2+ + CO CaCO3(s)
Precipitation
NH3 + H NH4+
Acid-Base
Zn + 2HCl ZnCl2 + H2
Redox (H2 Displacement)
Ca + F CaF2
Redox (Combination)
4.4

41. Solution Stoichiometry
The concentration of a solution is the amount of solute present in a given quantity of solvent or solution.
M = molarity =
moles of solute
liters of solution
What mass of KI is required to make 500. mL of
a 2.80 M KI solution?
M KI
M KI
volume KI
moles KI
grams KI
1 L
1000 mL x
2.80 mol KI
1 L soln x
166 g KI
1 mol KI x
500. mL
= 232 g KI
4.5

42. 4.5

43. 500 mL of solution contains 0. 730 moles of glucose
500 mL of solution contains moles of glucose. What is the molarity?
What volume of 2.00 M CuSO4 will contain 10.0 g of CuSO4?

44. Dilution is the procedure for preparing a less concentrated solution from a more concentrated solution.
Dilution
Add Solvent
Moles of solute
before dilution (i)
after dilution (f) =
MiVi
MfVf =
4.5

45. How would you prepare 60.0 mL of 0.20 M
HNO3 from a stock solution of 4.00 M HNO3?
MiVi = MfVf
Mi = 4.00
Mf = 0.20
Vf = 0.06 L
Vi = ? L
Vi =
MfVf Mi =
0.20 M x 60.0 mL
4.00 M
= 3.0 mL
3.0 mL of acid
to 60 mL of solution
4.5

46. Gravimetric Analysis Dissolve unknown substance in water
React unknown with known substance to form a precipitate
Filter and dry precipitate
Weigh precipitate
Use chemical formula and mass of precipitate to determine amount of unknown ion
4.6

47. Practice Exercise page 113
A sample of g of an unknown compound containing Br- ion is dissolved in water and treated with an excess of AgNO3 solution. If the mass of AgBr ppt that forms is g, what is the percent by mass of Br in the unknown sample?

48. Volumetric Analysis - Titrations
In a titration a solution of accurately known concentration is added gradually added to another solution of unknown concentration until the chemical reaction between the two solutions is complete.
Equivalence point – the point at which the reaction is complete
Indicator – substance that changes color at (or near) the
equivalence point
Slowly add base
to unknown acid
UNTIL
the indicator
changes color
4.7

49. What volume of a 1.420 M NaOH solution is
Required to titrate mL of a 4.50 M H2SO4
solution?
WRITE THE CHEMICAL EQUATION!
H2SO4 + 2NaOH H2O + Na2SO4 M
acid rx
coef. M
base
volume acid
moles acid
moles base
volume base
4.50 mol H2SO4
1000 mL soln x
2 mol NaOH
1 mol H2SO4 x
1000 ml soln
1.420 mol NaOH x
25.00 mL
= 158 mL
4.7

50. (What is net ionic equation?)
A standard solution is one of known concentration. NaOH solutions are frequently standardized with potassium hydrogen phthalate (KHP).
NaOH(aq) + KHC8H4O2(aq)  KNaC8H4O2(aq) + H2O
(What is net ionic equation?)
In a titration mL of NaOH solution are required to neutralize g of KHP. What is the concentration of the NaOH solution?