1. CHEMISTRY
PROPERTIES OF SOLUTION
__RaNNY ROLINda RUSMaN__

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

3. 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-

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

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

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

7. concentration

8. 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%

9. 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

10. Saturn
solvent
H2/He
solute
CH4, PH3

11. Europa
solvent
H 2O
solute
MgSO4

12. Triton
solvent
N2
solute
CH4

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

14. migrating negative and positive charges
Kohlrausch
NaCl

15. 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
polyatomic ions do NOT dissociate
EXP5

16. δO
δ+
H
H
δ+

17. SOLVATION
cations
anions

18. SOLVATION
non-electrolyte

19. 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)

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

21. 2.1. PRECIPITATION REACTIONS
solution 1
solution 2
solution 1 + solution 2

22. 2.1. PRECIPITATION REACTIONS
formation of an insoluble product
(precipitate)
NaCl(aq) + AgNO3(aq)
AgCl(s) + NaNO3(aq)
EXP 6

23. insoluble compounds
1.M+ compounds (M = H, Li, Na, K, Rb, Cs, NH4)
2. A- compounds (A = NO3, HCO3, ClO3, Cl, Br, I)
(AgX, PbX2)
3. SO42(Ag, Ca, Sr, Ba, Hg, Pb)
4. CO32-, PO43-, CrO42-, S2-

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

25. 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

26. 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

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

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

29. Litmus Paper
acid
red
Säure
base
blue
Base
EXP7

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

31. Na+
H+
≈ 10-10 m
≈ 10-15 m

32. 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

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

34. cation
hydronium ion

35. 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)
EXP8
3. acids react with carbonates to water and carbon dioxide
2 HCl(aq) + CaCO3(s) → CaCl2(aq) + [H2CO3]
H2CO3 → H2O(l) + CO2(g)
4. some acids are hygroscopic
H2SO4 (conc)
EXP9

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

37. 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)

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

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

41. 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

42. monoprotic acids
HF, HCl, HBr, HNO3, CH3COOH
diprotic acid
H2SO4 → H+(aq) + HSO4-(aq)
HSO4-(aq)
H+(aq) + SO42-(aq)
triprotic acid
H3PO4
H+(aq) + H2PO4-(aq)
H2PO4-(aq)
H+(aq) + HPO42-(aq)
HPO42-(aq)
H+(aq) + PO43-(aq)
EXP10

43. CHEMICAL PROPOERTIES
1. Non-metal oxides react with water to form an acid
(acetic anhydrides)
SO3 (g ) + H 2O → H 2SO 4 (aq) sulfuric acid
N 2 O5 (g ) + H 2O → 2HNO3 (aq) nitric acid
+ H+H
CO 2 (g )2O 2O → H 2CO 3 (aq) carbonic acid
+ H 2O
+ H 2O
Cl2O7, SO2, Br2O5

44. CHEMICAL PROPERTIES
2. Soluble metal oxides react with water to form a base
(base anhydrides)
CaO(s ) + H 22O→ Ca(OH) 2 (aq) calcium hydroxide
+ HO
Na 2 O(s ) + H 2O → 2NaOH(aq)
+ H 2O
MgO, Al2O3
sodium hydroxide

45. NAMING ACIDS AND BASES
binary acids
prefix hydrothe suffix –ic
to the stem of the nonmetal name followed by the word acid
HCl(g ) hydrogen chloride
HCl(aq) hydrochloric acid
H 2S(g ) hydrogen sulfide
H 2S(aq) hydrosulfuric acid

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

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

48. Europa
Venus
H2SO4(s)
H2SO4(g)

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

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

52. 1. oxidation
loss of electrons
2. reduction
acceptance of electrons
NUMBER OF ELECTRONS MUST BE CONSERVED

53. Na+Cl-
1. oxidation
Na → Na+ + e
2. reduction
Cl2 + 2 e → 2 Cl-
!!!balance electrons!!!
CaO, Al2O3

54. 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-

55. 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

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

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

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

59. H2
molecular compound
H+ + Hionic compound
OXIDATION NUMBER OF FREE ELEMENTS IS ZERO

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

61. 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-

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

63. RULE 4
oxidation number of oxygen
-2 in most compounds
(H2O, MgO, Al2O3)
-1 in peroxide ion (O22-) (H2O2, K2O2, CaO2)
-1/2 in superoxide ion (O2-) (LiO2)

64. 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-)

65. RULE 6
charges of polyatomic molecules must be integers
(NO3-, SO42-)
oxidation numbers do not have to be integers
-1/2 in superoxide ion (O2-)

67. 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

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

69. reducing agent
2 Na + 2 H2O → H2 + 2 NaOH
EXP11/12

70. K2O
PO
34
NO+
SO42-
KO2
SO3
NO3NO2
NO
KClO4
BrO-
SO2
NO2-
NO-

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

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

73. 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

74. 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

75. 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

76. 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

77. 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)

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

79. 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

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

82. 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)

83. 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)

84. 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)

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

86. molar concentration
Molarity
(M)
molarity (M) =
moles of solute
= liters of solution

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

88. 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

89. 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

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

91. 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
0.0120 mol BaSO 4

92. 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?
Vdil X M dil
=
Vconcd X M concd

93. LIMITING REACTANT
EXP14
C2H4 + H2O
C2H5OH

94. limiting reactant
excess reactant

95. 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