This document provides an overview of chapter 4 from a chemistry textbook, which covers chemical reactions. It begins by defining electrolytes, non-electrolytes, and discussing the properties of aqueous solutions. It then covers the three main types of reactions that occur in aqueous solutions: precipitation reactions, acid-base reactions involving proton transfer, and redox reactions involving electron transfer. Specific examples of each reaction type are provided. Key concepts around oxidation, reduction, and oxidation numbers are also explained.

1. CHEMISTRY 161
Chapter 4

2. CHEMICAL REACTIONS
2 HgO(s) → 2Hg(l) + O2(g)
aq
1. properties of aqueous solutions
2. reactions in aqueous solutions
a) precipitation reactions
b) acid-base reactions (proton transfer)
c) redox reactions (electron transfer)

3. 1.PROPERTIES OF AQUEOUS SOLUTIONS
homogeneous mixture of two or more substances
solvent solute
substance in a large amount substance in a small amount
N2 gas phase O2
(air)
Ag solid phase Au
(alloys)
H2O liquid phase NaCl
(sea water)

4. EXP1
iodine in ethyl alcohol (C2H5OH)
does not conduct electricity
(molecular solid)
I2
EXP2
table salt in water (H2O)
does conduct electricity
(ionic solid)
Na+Cl-

5. AQUEOUS SOLUTION
solute
water (H2O)
solutes
solution conducts
electricity
solution does not
conduct electricity
EXP3
electrolytes non-electrolytes

6. electrolytes non-electrolytes
solution conducts
electricity
solution does not
conduct electricity

7. non-electrolyte weak electrolyte strong electrolyte
methanol
sugar
ethanol
water
ionic compounds
(NaCl, KF)
NaOH
HCl
H2SO4
CH3COOH
HCOOH
HF
EXP5
dark medium
bright

8. SOLUTION
concentration

9. SOLUTION
percentage concentration
% = g [solute] / g solvent X 100
12 g of sodium chloride are solved in 150 g of water.
Calculate the percentage concentration
8 %

10. SOLUTION
solubility of a solute
number of grams of solute that can dissolve
in 100 grams of solvent at a given temperature
36.0 g NaCl can be dissolve in 100 g of water at 293 K

11. GAS PHASE SOLUTION
Saturn
solvent
H2/He
solute
CH4, PH3

12. LIQUID SOLUTION
Europa
solvent
H2O
solute
MgSO4

13. SOLID SOLUTION
Triton
solvent
N2
solute
CH4

14. methanol
sugar
ethanol
water
ELECTROLYTES
ionic compounds
(NaCl, KF)
NaOH
HCl
H2SO4
CH3COOH
HCOOH
HF

15. migrating negative and positive charges
Kohlrausch NaCl

16. DISSOCIATION
‘breaking apart’
NaCl (s) → Na+ (aq) + Cl- (aq)
NaOH (s) → Na+ (aq) + OH- (aq)
HCl (g) → H+ (aq) + Cl- (aq)
Ca(NO3)2 (s) → Ca2+(aq) + 2 NO3
- (aq)
strong electrolytes are fully dissociated
EXP5
polyatomic ions do NOT dissociate

17. δ-
O
H H
δ+ δ+

18. SOLVATION
cations anions

19. SOLVATION
non-electrolyte

20. NaCl (s) → Na+ (aq) + Cl- (aq)
strong electrolytes are fully dissociated
←→
CH3COOH (aq) H+ (aq) + CH3COO- (aq)
weak electrolytes are not fully dissociated
reversible reaction
(chemical equilibrium)

21. CHEMICAL REACTIONS
1.properties of aqueous solutions
2. reactions in aqueous solutions
a) precipitation reactions
b) acid-base reactions (proton transfer)
c) redox reactions (electron transfer)

22. 2.1. PRECIPITATION REACTIONS
solution 1 solution 2 solution 1 + solution 2

23. 2.1. PRECIPITATION REACTIONS
formation of an insoluble product
(precipitate)
NaCl(aq) + AgNO3(aq)
AgCl(s) + NaNO3(aq)
EXP 6

24. insoluble compounds
1.M+ compounds (M = H, Li, Na, K, Rb, Cs, NH4)
2. A- compounds (A = NO3, HCO3, ClO3, Cl, Br, I)
(AgX, PbX)
23. SO2-
4
(Ag, Ca, Sr, Ba, Hg, Pb)
4. CO2-, PO3-, CrO3
4
4
2-, S2-
(Ag, Ca, Sr, Ba, Hg, Pb)

25. balanced molecular equation
NaCl(aq) + AgNO3(aq) → AgCl(s) + NaNO3(aq)
(table to determine which compound precipitates)

26. balanced ionic equation
1. NaCl(s) → Na+(aq) + Cl-(aq)
2. AgNO3(s) → Ag+(aq) + NO3
-(aq)
3. Na+(aq) + Cl-(aq) + Ag+(aq)+ NO3
-(aq) →
AgCl(s) + Na+ (aq) + NO3
-(aq)
spectator ions

27. Ba(NO3)2 (aq) + Na2SO4 (aq)
Ba(NO3)2(aq) + Na3PO4(aq)
Cs2CrO4(aq) + Pb(NO3)2(aq)
1. which compound falls out?
2. balanced molecular equation
3. balanced ionic equations
4. identify spectator ions

28. CHEMICAL REACTIONS
1.properties of aqueous solutions
2. reactions in aqueous solutions
a) precipitation reactions
b) acid-base reactions (proton transfer)
c) redox reactions (electron transfer)

29. ACIDS AND BASES
Arrhenius (1883)
ACIDS
HAc → H+ (aq) + Ac- (aq)
ionization
HCl (g) → H+ (aq) + Cl- (aq)
BASES
MOH → M+ (aq) + OH- (aq)
NaOH (s) → Na+ (aq) + OH- (aq)

30. IDENTIFICATION
Litmus Paper
acid
base
red
blue
Säure
Base
EXP7

31. ACIDS AND BASES
ACIDS and
BASES
NEUTRALIZE
EACH OTHER
HAc (aq) + MOH (aq) → MAc (aq) + H2O
HCl (aq) + NaOH (aq) → NaCl (aq) + H2O
acid + base salt + water

32. ACIDS AND BASES
Na+ ≈ 10-10 m
H+ ≈ 10-15 m

33. ACIDS AND BASES
HCl (g) → H+ (aq) + Cl- (aq)
H+(aq) + H2O H3O+(aq)
HCl (g) + H2O → H3O+ (aq) + Cl- (aq) one step
hydronium ion

34. (aq) (l) (aq) (aq)
acid base
hydronium ion

35. cation hydronium ion

36. PROPERTIES OF ACIDS
1. acids have a sour taste
vinegar – acetic acid
lemons – citric acid
2. acids react with some metals to form hydrogen
2 HCl(aq) + Mg(s) → MgCl2(aq) + H2(g)
3. acids react with carbonates to water and carbon dioxide
2 HCl(aq) + CaCO3(s) → CaCl2(aq) + [H2CO3]
H2CO3 → H2O(l) + CO2(g)
EXP8
EXP9
4. some acids are hygroscopic
H2SO4 (conc)

37. BASES
1. bases have a bitter taste
2. bases feel slippery
soap
3. aqueous bases and acids conduct electricity

38. EXAMPLES
KOH(aq) and HF(aq)
Mg(OH)2(aq) and HCl(aq)
Ba(OH)2(aq) and H2SO4(aq)
NaOH(aq) and H3PO4(aq)
(stepwise)

39. Bronsted (1932)
ACIDS
proton donors
HAc → H+ (aq) + Ac- (aq)
BASES
proton acceptor
B + H+ (aq) → BH+ (aq)

40. electrolyte
HCl(aq) + H2O(l) → H3O+(aq) + Cl-(aq)
HNO3(aq) + H2O(l) → H3O+(aq) + NO3
weak electrolyte
CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO-(aq)
NH3(aq) + H2O(l) NH4
+ + OH-strong
-(aq)
donor versus acceptor

42. CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO-(aq)
NH3(aq) + H2O(l) NH4
+(aq)+ OH-(aq)
H2O(l) + H2O(l) H3O+(aq) + OH-(aq)
water can be either an
acid or a base
AUTO DISSOCIATION

43. monoprotic acids
HF, HCl, HBr, HNO3, CH3COOH
diprotic acid
H2SO4 → H+(aq) + HSO4
-(aq)
HSO4
-(aq) H+(aq) + SO4
2-(aq)
triprotic acid
H3PO4 H+(aq) + H2PO4
-(aq)
-(aq) H+(aq) + HPO4
H2PO4
2-(aq)
HPO4
2-(aq) H+(aq) + PO4
3-(aq)
EXP10

44. CHEMICAL PROPOERTIES
1. Non-metal oxides react with water to form an acid
(acetic anhydrides)
g + ®
aq
g + ®
aq
g + ®
aq
SO ( ) H H SO ( ) sulfuric acid
3 2 O
2 4
N O ( ) H O
2HNO ( ) nitric acid
2 5 3
CO ( ) H O H CO ( ) carbonic acid
2 2 2 3
Cl2O7, SO2, Br2O5
22 + H2O

45. CHEMICAL PROPERTIES
2. Soluble metal oxides react with water to form a base
(base anhydrides)
s aq
s aq
+ H2O
+ H2O
CaO( ) H O Ca(OH) ( ) calcium hydroxide
2 2
Na O( ) H O 2NaOH( ) sodium hydroxide
MgO, Al2O3
2 2
®
+ ®

46. NAMING ACIDS AND BASES
binary acids
prefix hydro- the suffix –ic
to the stem of the nonmetal name followed by the word acid
g aq hydro ic acid
g aq hydro ic acid
HCl( ) hydrogen chloride HCl( ) chlor
H S( ) hydrogen sulfide H S( ) sulfur
2 2

47. NAMING ACIDS AND BASES
oxo acids acids
contain hydrogen, oxygen, plus another element
main group 5
HNO3 nitric acid
HNO2 nitrous acid
H3PO4 phosphoric acid
H3PO3 phosphorous acid

48. main group 6
H2SO4 sulfuric acid
H2SO3 sulfurous acid
main group 7
HClO4 perchloric acid
HClO3 chloric acid
HClO2 chlorous acid
HClO hypochlorous acid

49. Acids in the Solar System
Venus
H2SO4(g)
Europa
H2SO4(s)

50. Acids in the Interstellar Medium

51. Orion
NH3, H2O, H2S
CH3COOH
HCOOH
HF, HCl

52. CHEMICAL REACTIONS
1.properties of aqueous solutions
2. reactions in aqueous solutions
a) precipitation reactions
b) acid-base reactions (proton transfer)
c) redox reactions (electron transfer)

53. KEY CONCEPTS
1. oxidation
loss of electrons
2. reduction acceptance of electrons
NUMBER OF ELECTRONS MUST BE CONSERVED

54. 1. oxidation
EXAMPLE
2. reduction
Na+Cl-
Na ® Na+ + e
Cl2 + 2 e ® 2 Cl-
!!!balance electrons!!!
CaO, Al2O3

55. substance that lost the electrons
reduction agent
substance that gained the electrons
oxidizing agent
oxidizing agent is reduced
reducing agent is oxidized
2 Na + Cl2 ® 2 Na+Cl-

56. EXAMPLE 1
solid state reaction of potassium with sulfur
to form potassium sulfide
EXAMPLE 2
solid state reaction of iron with oxygen
to form iron(III)oxide

57. OXIDATION NUMBER
ionic compounds ↔ molecular compounds
NaCl HF, H2
Na+Cl- ?
electrons are fully transferred covalent bond
charges an atom would have if electrons are
transferred completely

58. EXAMPLE 1
HF H+ + F-molecular
compound ionic compound
H+ oxidation state +1
F- oxidation state -1

59. H2O
2 H+ + O2-
EXAMPLE 2
molecular compound ionic compound
H+ oxidation state +1
O2- oxidation state -2

60. H2
H+ + H-EXAMPLE
3
molecular compound ionic compound
OXIDATION NUMBER OF FREE ELEMENTS IS ZERO

61. RULE 1
OXIDATION NUMBER OF FREE ELEMENTS IS ZERO
H2, O2, F2, Cl2, K, Ca, P4, S8

62. RULE 2
monoatomic ions
oxidation number equals the charge of the ion
group I M+
group II M2+
group III M3+ (Tl: also +1)
group VII (w/ metal) X-

63. RULE 3
oxidation number of hydrogen
+1 in most compounds
(H2O, HF, HCl, NH3)
-1 binary compounds with metals (hydrides)
(LiH, NaH, CaH2, AlH3)

64. RULE 4
oxidation number of oxygen
-2 in most compounds
(H2O, MgO, Al2O3)
-1 in peroxide ion (O2
2-) (H2O2, K2O2, CaO2)
-1/2 in superoxide ion (O2
-) (LiO2)

65. RULE 5
oxidation numbers of halogens
F: -1 (KF)
Cl, Br, I: -1 (halides) (NaCl, KBr)
Cl, Br, I: positive oxidation numbers if combined
with oxygen (ClO4
-)

66. RULE 6
charges of polyatomic molecules must be integers
(NO3
-, SO4
2-)
oxidation numbers do not have to be integers
-1/2 in superoxide ion (O2
-)

68. MENUE
1.oxidation states of group I – III metals
2.oxidation state of hydrogen (+1, -1)
3. oxidation states of oxygen (-2, -1, -1/2, +1)
4.oxidation state of halogens
5.remaining atoms

69. oxidizing agents
?????
OCl- Cl-
EXP10

70. reducing agent
2 Na + 2 H2O ® H2 + 2 NaOH
EXP11/12

71. NO+
NO NO 2
3- SO4
NO3
NO -
2
NO- -
PO4
2-
SO3
SO2
KO2
K2O
BrO-KClO
4

72. REVISION
1.redox reactions
2. oxidation versus reduction
3. oxidation numbers versus charges
4. calculation of oxidation numbers

73. TYPES OF REDOX REACTIONS
1.combination reactions
A + B → C
2. decomposition reactions
C → A + B
3. displacement reactions
A + BC → AC + B
4. disproportionation reactions

74. 1.combination reactions
A + B → C
two or more compounds combine to form a single product
S8(s) + O2(g) → SO2(g)
1. oxidation numbers
2. balancing charges

75. MENUE
1.oxidation states of group I – III metals
2.oxidation state of hydrogen (+1, -1)
3. oxidation states of oxygen (-2, -1, -1/2, +1)
4.oxidation state of halogens
5.remaining atoms

76. 2. decomposition reactions
C → A + B
breakdown of one compound into two or more compounds
HgO(s) → Hg(l) + O2(g)
KClO3(s) → KCl(s) + O2(g)
1. oxidation numbers
2. balancing charges

77. 3. displacement reactions
A + BC → AC + B
an ion or atom in a compound is replaced by an ion or atom
of another element
3.1. Hydrogen displacement
3.2. Metal displacement
3.3. Halogen displacement

78. 3.1. Hydrogen displacement
group I and some group II metals (Ca, Sr, Ba)
react with water to form hydrogen
Na(s) + H2O(l) → NaOH + H2(g)
less reactive metals form hydrogen and the oxide in
water (group III, transition metals)
Al(s) + H2O(l) → Al2O3(s) + H2(g)

79. 3.1. Hydrogen displacement
even less reactive metals form hydrogen in acids
Zn(s) + HCl(aq) → ZnCl2(aq) + H2(g)
EXP12

80. activity series of metals
Li K Ba Ca Na Mg Al Zn Cr Fe Cd Co Ni Sn Pb H Cu Hg Ag Pt Au
displace H from water
displace H from steam
displace H from acids

81. Li K Ba Ca Na Mg Al Zn Cr Fe Cd Co Ni Sn Pb H Cu Hg Ag Pt Au
likes to donate electrons does not like so much to donate electrons
EXP13

83. 3.2. Metal displacement
V2O5(s) + 5 Ca(s) → 2 V(s) + 5 CaO(s)
TiCl4(g) + 2 Mg (l) → Ti(s) + 2 MgCl2(l)

84. 3.3. Halogen displacement
F2 > Cl2 > Br2 > I2
reactivity (‘likes’ electrons)
0 +1 -1 +1 -1 0
Cl2(g) + 2 KBr(aq) → 2 KCl(aq) + Br2(l)
Br2(g) + 2 KI(aq) → 2 KBr(aq) + I2(s)

85. 4. disproportionation reactions
an element in one oxidation state is oxidized and reduced
at the same time
H2O2(aq) → 2 H2O(l) + O2(g)
Cl2(g) + 2 OH-(aq) → ClO-(aq) + Cl-(aq) + H2O(l)

86. SUMMARY
1.combination reactions
A + B → C
2. decomposition reactions
C → A + B
3. displacement reactions
A + BC → AC + B
4. disproportionation reactions

87. STOCHIOMETRY
(CONCENTRATION)
molar concentration
Molarity
(M)
molarity (M) == moles of solute
liters of solution

88. How many grams of AgNO3 are needed to prepare
250 mL of 0.0125 M AgNO3 solution?
3 0.531 g AgNO

89. How many mL of 0.124 M NaOH are required
to react completely with 15.4 mL of 0.108 M H2SO4?
2 NaOH + H2SO4 Na2SO4 + 2H2O
26.8 mL NaOH

90. How many mL of 0.124 M NaOH are required
to react completely with 20.1 mL of 0.2 M HCl?
NaOH + HCl NaCl + H2O

91. How many grams of iron(II)sulfide have to react with hydrochloric acid
to generate 12 g of hydrogen sulfide?

92. How many moles of BaSO4 will form if 20.0 mL of
0.600 M BaCl2 is mixed with 30.0 mL of 0.500 M MgSO4?
BaCl2 + MgSO4 BaSO4 + MgCl2
This is a limiting reagent problem
4 0.0120 mol BaSO

93. How many ml of a 1.5 M HCl will be used to neutralize
a 0.2 M Ba(OH)2 solution?
How many ml of a 1.5 M HCl will be used to prepare
500 ml of a 0.1 M HCl?
dil dil V X M = Vconcd X Mconcd

94. LIMITING REACTANT
C2H4 + H2O C2H5OH
EXP14

95. limiting reactant excess reactant

96. How many grams of NO can form when 30.0 g
NH3 and 40.0 g O2 react according to
4 NH3 + 5 O2 4 NO + 6 H2O