1. Chapter 4, part 2: Reactions in Aqueous Solution
Chemical Reactions
Chapter 4, part 2:
Reactions in Aqueous Solution

2. The dissolving process
When a solid is put into a liquid, solute-solute attractions compete with solute-solvent & solvent-solvent attractions

3. Solubility Let’s assume our solvent is water . . .
If solute-water attractions > solute-solute & water-water attractions, solute particles are pulled out one by one into the water:
The solute is SOLUBLE in water

4. Solubility
But if solute-water attractions < solute-solute & water-water attractions, solute particles remain together:
The solute is INSOLUBLE in water

5. Solution conductivity
Solution conductivity depends on type of solute particles

6. Solution conductivity
Ionic solutes (salts) made of cations & anions
Ions DISSOCIATE (separate) during dissolving
Molecular solutes made of molecules
each solute particle that moves into solution is identical

7. Solution conductivity
Solutions of ionic solutes contain independent mobile ions
Solution conducts electricity
Solute is an ELECTROLYTE
Solute also conducts when melted, but not when solid (ions can’t move)

8. Solution conductivity
Solutions of most molecular solutes contain independent neutral molecules
Solution does not conduct electricity
Solute is a NONELECTROLYTE
Solute also does not conduct when melted or solid

9. Solution conductivity
Some molecular compounds (acids & bases) react with water to produce ions, as if they dissociated
A few acids & bases do this very well, producing lots of ions = strong electrolytes, strong acids & bases
Most acids & bases do this weakly, producing a few ions = weak electrolytes, weak acids & bases

10. All other acids & bases are WEAK
Strong acids & bases
STRONG ACIDS
STRONG BASES
HCl
HNO3
LiOH
Mg(OH)2
HBr
H2SO4
NaOH
Ca(OH)2 HI
HClO4
KOH
Sr(OH)2
etc.
Ba(OH)2
All other acids & bases are WEAK

12. Predicting Electrolytes
All soluble salts and strong acids & bases are strong electrolytes
Weak acids & bases are weak electrolytes
All other molecular compounds are nonelectrolytes

13. Precipitation reactions

14. Precipitation reactions
spectator ions
precipitate

15. Predicting Precipitation Reactions
To predict whether a precipitate will form, you need to know which compounds are soluble (no ppt) and which are insoluble (ppt forms)
Memorize the guidelines in Table 4.1 on page 154 in your text!
Add this guideline: CrO42– acts like SO42–

16. Solubility guidelines: soluble
Compounds of these ions are generally soluble and do NOT form precipitates:
Alkali metals (group 1A, except Li1+)
Ammonium (NH41+)
Nitrates (NO31–) and acetates (C2H3O21–)
Chlorides, bromides, iodides, except Pb2+, Ag1+, Hg22+
Sulfates (SO42–) & chromates (CrO42–), except Sr2+, Ba2+, Pb2+, and Hg22+

17. Solubility guidelines: insoluble
Compounds of these ions are generally insoluble and DO form precipitates:
Hydroxides (OH1–) and Sufides (S2–)
Alkali metals (group 1A) and ammonium (NH41+) are soluble
Sulfides of group 2A metals are generally soluble
Hydroxides of Ca2+, Sr2+, and Ba2+ are soluble
Carbonates (CO32–) and phosphates (PO43–)

18. Net ionic equations AgNO3 (aq) + NaCl (aq) 
AgNO3 (aq) + NaCl (aq)  AgCl + NaNO3
AgNO3 (aq) + NaCl (aq)  AgCl (s) + NaNO3 (aq)
The (aq) substances are dissociated:
Ag1+ (aq) + NO31– (aq) + Na1+ (aq) + Cl1– (aq)  AgCl (s) + Na1+ (aq) + NO31– (aq)

19. Net ionic equations
Na1+ and NO31– are spectator ions
Ag1+ (aq) + NO31– (aq) + Na1+ (aq) + Cl1– (aq)  AgCl (s) + Na1+ (aq) + NO31– (aq)
The net ionic equation omits spectators: Ag1+ (aq) + Cl1– (aq)  AgCl (s)

20. Is that really a spectator?
An ion is a spectator if and only if it is in exactly the same form in the products and reactants:
Na2CO3 (aq) + BaCl2 (aq)  BaCO3 (s) + 2 NaCl (aq)
CO32– (aq) + Ba2+ (aq)  BaCO3 (s)
Na2CO3 (s) + 2 HCl (aq)  2 NaCl (aq) + H2O + CO2 (g)
Na2CO3 (s) + 2 H1+ (aq)  2 Na1+ (aq) + H2O + CO2 (g)
Only Cl1– is a spectator
Na1+ is not a spectator because it was (s), then (aq)

21. Examples
Indicate whether a ppt forms and if so, complete the reaction as a balanced net ionic equation:
AlCl3 (aq) + KOH (aq) 
K2SO4 (aq) + FeBr3 (aq) 
CaI2 (aq) + Pb(NO3)2 (aq) 
Na3PO4 (aq) + AlCl3 (aq) 
Al2(SO4)3 (aq) + BaCl2 (aq) 
(NH4)2CO3 (aq) + Pb(NO3)2 (aq) 

22. Acids Acids produce H3O1+ in aqueous solution:
Strong acids are molecular compounds that react completely with water to produce H3O1+:
HCl (g) + H2O  H3O1+ (aq) + Cl1– (aq)
For convenience, we often show acids as simply dissociating to produce H1+: HCl (g)  H1+ (aq) + Cl1– (aq)
There are only 6 strong acids (memorize them, pg 161)

23. Acids Acids produce H3O1+ in aqueous solution:
Weak acids are molecular compounds that react incompletely with water:
HC2H3O2 (aq) + H2O  H3O1+ (aq) + C2H3O21– (aq)
HC2H3O2 (aq)  H1+ (aq) + C2H3O21– (aq)
All acids that are not strong are weak

24. Acids H1+ is a proton The reaction with water is proton transfer:
HCl (g) + H2O  H3O1+ (aq) + Cl1– (aq)
HC2H3O2 (aq) + H2O  H3O1+ (aq) + C2H3O21– (aq)
Acids with one H1+ to transfer are monoprotic acids
HCl HNO HC2H3O2
Acids with more than one H1+ to transfer are polyprotic acids H2SO H3PO H2C3H2O4

25. Acids Polyprotic acids produce H3O1+ in steps:
H2SO4 (aq) + H2O  H3O1+ (aq) + HSO41– (aq)
HSO41– (aq) + H2O  H3O1+ (aq) + SO42– (aq)
For H2SO4 the first step is strong and the second weak
H2C2O4 (aq) + H2O  H3O1+ (aq) + HC2O41– (aq)
HC2O41– (aq) + H2O  H3O1+ (aq) + C2O42– (aq)
For all other polyprotic acids, the first step is weak and the second step is weaker

26. Bases Bases produce OH1– in aqueous solution:
Strong bases are ionic hydroxide compounds that are completely dissociated in water: NaOH (s)  Na1+ (aq) + OH1– (aq)
The strong bases are the hydroxides of group 1A and 2A metals (memorize them)
Weak bases are molecular compounds that react incompletely with water to produce OH1–: NH3 (aq) + H2O (l)  NH41+ (aq) + OH1– (aq)
Amines (–NH2 ) are weak bases

27. Neutralization Acids and bases neutralize each other
The H1+ from the acid is transferred to the base: HCl (aq) + NaOH (aq)  HOH (l) + NaCl (aq) H1+ (aq) + OH1– (aq)  HOH (l)
The base is not always an OH1– compound: HCl (aq) + NH3 (aq)  NH41+ (aq) + Cl1– (aq) H1+ (aq) + NH3 (aq)  NH41+ (aq)

28. Neutralization and net ionic equations
It is important to recognize strong acids and bases when writing net ionic equations
HCl (aq) + NaOH (aq)  H2O (l) + NaCl (aq) H1+ (aq) + OH1– (aq)  H2O (l)
HC2H3O2 (aq) + NaOH (aq)  H2O (l) + NaC2H3O2 (aq)
HC2H3O2 (aq) + OH1– (aq)  H2O (l) + C2H3O21– (aq)
The weak acid is not significantly ionized,
so acetate is not a spectator (even if it is aq)

29. Examples Write the molecular and net ionic equations for HNO3 + NaOH 
HF + KOH 
HC2H3O2 + NH3 
H2SO4 + Ba(OH)2 
H2C2O4 + NaOH 
HCHO2 + Ca(OH)2 

30. Gas forming reactions Some neutralizations produce a gas:
CO32– + 2 H1+  H2CO3
H2CO3 is unstable and decomposes immediately
H2CO3  CO2 (g) + H2O
The overall reaction is
CO32– + 2 H1+  CO2 (g) + H2O

31. Gas forming reactions Memorize these gas-formers (pg 166):
SO32– H1+  SO2 (g) + H2O
HSO31– + H1+  SO2 (g) + H2O
CO32– + 2 H1+  CO2 (g) + H2O
HCO31– + H1+  CO2 (g) + H2O
S2– + 2 H1+  H2S (g)
NH OH1–  NH3 (g) + H2O
burning sulfur smell
rotten egg smell
ammonia

32. Redox
Redox (oxidation-reduction) reactions are those in which electrons are transferred
The loss of electrons is oxidation
The gain of electrons is reduction
LEO says GER

33. Oxidation states
The oxidation state (O.S.) or oxidation number is a convenient but artificial way to describe the electron environment around an atom
It is related to the number of electrons gained, lost, or apparently used in forming compounds
Oxidation states are assigned using the rules on page 169 of your text (memorize these in order)

34. Assigning oxidation states
1. The O.S. of each atom in an element is zero
2. The O.S. of a monoatomic ion is equal to its charge
3. The total of the O.S. of all atoms in any species (formula unit, molecule or ion) equals the charge on that species
4. In compounds, metals always have a positive O.S.
• Group 1A metals are always O.S. +1 and Group 2A metals are always O.S. +2
5. For nonmetals in compounds,
• the O.S. of fluorine is –1.
• the O.S. of hydrogen is +1.
• the O.S. of oxygen is –2.
6. In binary compounds, the O.S. of a Group 7A element is –1, Group 6A element –2, and Group 5A element –3.

35. Examples What is the oxidation state of each element in
S Cr2O72– Cl2O KO2
S2O32– Hg2Cl KMnO H2CO

36. Redox In a redox reaction, atoms change (O.S.)
The element oxidized loses electrons and its O.S. becomes more positive
The element reduced gains electrons and its O.S. becomes more negative

37. Examples
Which of these reactions is a redox reaction? Identify the species oxidized and reduced.
HCl (aq) + NaOH (aq)  H2O + NaCl (aq)
2 Pb(NO3)2 (s)  2 PbO (s) NO2 (g) + O2 (g)
NH4Cl (s) + NaOH (aq)  NH3 (g) + H2O + NaCl (aq)
Identify the species oxidized and reduced:
5 VO2+ (aq) + MnO41– (aq) + H2O  5 VO21+ (aq) + Mn2+ (aq) H1+ (aq)

38. Agents of Oxidation and Reduction
An agent makes something happen
An oxidizing agent makes oxidation happen by being reduced
A reducing agent makes reduction happen by being oxidized
The “agent” is the entire species in which the oxidized or reduced atom appears

39. Examples
Identify the elements oxidized & reduced and the oxidizing & reducing agents in
2 NO2 (g) H2 (g)  2 NH3 (g) H2O (g)
5 H2O2 (aq) MnO41– (aq) + 6 H1+ (aq)  8 H2O Mn2+ (aq) O2 (g)
S2O32– (aq) Cl2 (aq) H2O  2 HSO41– (aq) H1+ (aq) Cl1– (aq)
6 Fe2+ (aq) H1+ (aq) + Cr2O72– (aq)  6 Fe3+ (aq) Cr3+ (aq) H2O
S2O32– (aq) H1+ (aq)  S (s) + SO2 (g) + H2O

40. Reactions You now know how to write & balance 4 types of reactions
combustion
precipitation
acid-base neutralization
gas-forming
and how to recognize redox reactions

41. Reaction quizzes Reaction quizzes (RQ) will replace NQ
I give you the names of the reactants
You write the formulas of the reactant and the formulas of the products
Cross out spectator ions
Write the balanced net ionic equation
You need not include state symbols such as (aq) or (g)

42. Example Solutions of silver nitrate and potassium chloride are mixed
AgNO3 (aq) + KCl (aq) 
AgNO3 (aq) + KCl (aq)  AgCl (s) + KNO3 (aq)
cross out K1+ and NO31– as spectators
Ag Cl1–  AgCl
check: it is already balanced

43. Example
Hydrochloric acid solution is added to solid sodium hydrogen carbonate
HCl (aq) + NaHCO3 (s) 
HCl (aq) + NaHCO3 (s)  H2O + CO2 (g) + NaCl (aq)
cross out Cl1– as a spectator
do NOT cross out Na1+ because NaHCO3 is solid
H1+ + NaHCO3  H2O + CO2 + Na1+
check: it is already balanced

44. Example Ethoxyethane burns in air C2H5OC2H5 + O2 
C2H5OC2H5 + O2  CO2 + H2O
there are no spectators in combustion
balance
C2H5OC2H O2  4 CO H2O

45. Example Solutions of nitrous acid and sodium hydroxide are mixed
HNO2 (aq) + NaOH (aq) 
HNO2 (aq) + NaOH (aq)  H2O + NaNO2 (aq)
cross out Na1+ as a spectator
do NOT cross out NO21– because HNO2 is a weak acid and is not significantly dissociated in solution!
HNO2 + OH1–  H2O + NO21–
check: it is balanced

46. Titrations
The text includes examples of acid-base titration and redox titration
These are just stoichiometry problems, based on a particular type of reaction
Do examples 5-9 and 5-10 to practice these
They will be on the test, and you may have to write and/or balance your own equations

47. Exercise 26
Both NH3 (aq) and HC2H3O2 (aq) conduct weakly, but when these solutions are mixed, the resulting solution conducts electricity very well. Explain.

48. Exercise 32
Assuming volumes are additive, what is [NO31–] in a solution obtained by mixing 275 mL M KNO3, 328 mL M Mg(NO3)2, and 784 mL H2O?

49. Exercise 34
Predict whether a reaction occurs, and if so, write a net ionic equation:
a. AgNO3 (aq) + CuCl2 (aq) 
b. Na2S (aq) + FeCl2 (aq) 
c. Na2CO3 (aq) + AgNO3 (aq) 

50. Exercise 39
Write net ionic equations to represent these neutralizations of stomach acid (HCl):
a. sodium bicarbonate
b. calcium carbonate
c. magnesium hydroxide

51. Exercise 42
Give net ionic equations for the preparation of each salt by acid-base neutralization: (NH4)2HPO4, NH4NO3, (NH4)2SO4

52. Exercise 47 Balance these redox reactions in acid solution:
a. MnO41– + I1–  Mn I2 (s)
b. BrO31– + N2H4  Br1– + N2
c. VO43– + Fe2+ 
U  UO NO (g)

53. Exercise 49 Balance these redox reactions in basic solution:
a. CN1– + MnO41–  MnO2 (s) + CNO1–
b. [Fe(CN)6]3– + N2H4  [Fe(CN)6]4– + N2
c. Fe(OH)2 (s) + O2 (g) (OH)3 (s)
C2H5OH + MnO41–  C2H3O21– + MnO2 (s)

54. Exercise 52 Write a balanced equation for each redox reaction:
a. Oxidation of nitrite ion to nitrate ion by permanganate ion in acid solution, producing Mn2+
b. Reduction of H2S with SO2. Both are converted to elemental sulfur and water.
c. Reaction of sodium metal with hydroiodic acid.
d. Reduction of VO2+ to V3+ by zinc metal, in acidic solution.

55. Exercise 53
Identify the oxidizing and reducing agents in these reactions:
a. 5 SO32– MnO41– H+  SO42– Mn H2O
b. 2 NO H2  2 NH H2O
c. 2 [Fe(CN)6]4– + H2O H+  [Fe(CN)6]3– H2O

56. Exercise 60
a. How many mL of concentrated HCl (d = 1.19 g/mL, 38% HCl by mass) must be diluted to 20.0 L to prepare 0.10 M HCl?
b mL of the 0.1 M HCl requires mL of M NaOH for titration. What is the molarity of the HCl?
c. Why could you not prepare M HCl simply by diluting concentrated HCl?

57. Exercise 64
5.00 mL battery acid (H2SO4) requires Ml M NaOH for complete neutralization. Is the concentration of the battery acid within the desired range (4.8 – 5.3 M)?

58. Exercise 67
Iron ore weighing g is dissolved in HCl to produce Fe2+, which is then titrated with mL M K2Cr2O7. What is the % Fe by mass in the ore? Fe H+ + Cr2O72–  Fe Cr H2O

59. Half-reactions
A half reaction describes either the oxidation or the reduction part of a redox reaction
It includes the number of electrons lost or gained
The overall reaction is the sum of two half-reactions, one oxidation and one reduction

60. Examples 5-5A & 5-5B
Represent reaction 4.2 with half-reactions and the overall net reaction:
2 Al (s) HCl (aq)  2 AlCl3 (aq) H2 (g)
Represent the reaction of chlorine gas with aqueous sodium bromide to produce liquid bromine and aqueous sodium chloride with half-reactions and a net equation

61. Balancing Redox Equations
Redox equations require electrical balance as well as material (atom) balance
Write balanced half reactions for oxidation & reduction
First balance atoms, adding H1+ and H2O as needed
When atoms are balanced, add e– to balance charge.
Multiply each half-reaction by a factor such that the number of electrons lost and gained are the same
Combine the half reactions and cancel species that are identical on both sides

62. Examples 5-6A & 5-6B Balance these equations in acid solution
Fe2+ (aq) + MnO41– (aq)  Fe3+ (aq) + Mn2+ (aq)
UO2+ (aq) + Cr2O72– (aq)  UO22+ (aq) + Cr3+ (aq)

63. Redox in basic solution
To balance a redox equation in basic solution,
Balance it as if it is in acid solution
After combining the half-reactions, add equal moles of OH1– to both sides such that all H1+ is neutralized to form H2O
Simplify

64. Examples 5-7A & 5-7B Balance these equations in basic solution
S (s) + OCl1– (aq)  SO32– (aq) + Cl1– (aq)
MnO41– (aq) + SO32– (aq)  MnO2 (s) + SO42– (aq)

65. Disproportionation
In disproportionation reactions, the same element is both oxidized and reduced
2 H2O2 (aq)  2 H2O + O2 (g)
S2O32– (aq)  S (s) + SO2 (g)

66. Common Redox agents Common oxidizing agents include
MnO41– (permanganate) in acid solution
Cr2O72– (dichromate) in acid solution
O2 or O3
Common reducing agents include
S2O32– (thiosulfate) H2