1. Aqueous
Reactions

2. Key Concepts of Chapter
• Components of a solution.
• Electrolytes, Weak Electrolytes, Nonelectrolytes
• Precipitation Reactions and Solubility Guidelines
• Molecular, Complete Ionic, Net Ionic Equations
• A Few Concepts Related to Acids and Bases
• Redox Reactions and Determining Oxidation Numbers
• Activity Series of Metals
• Molarity / Solution Preparation / Dilution

3. Properties of water High melting and boiling point.
Expands upon freezing.
Dissolves a variety of substances easily.
Pure water is rare in nature.
These properties are mainly due to polarity, which will be discussed a bit later.
High Specific Heat
Cohesion and adhesion (meniscus)
Polarity: oxygen is (-) and H is (+)

4. Most solids—molecules are tightly packed, and therefore they contract upon freezing.
Ice—molecules take on an open hexagonal arrangement, and therefore it expands upon freezing. (More about this in Chapter 11)

5. Reactions in aqueous solutions
Solution in which water is the solvent (dissolving agent).
3 major types of chemical processes of aqueous solutions:
Precipitation reactions
Acid-base reactions
Redox reactions

6. Solution: Solvent: Solute: Homogenous mixture of 2 or more substances.
Dissolving medium, usually present in greater quantity.
Solute:
The other substance(s) in the solution.

7. Electrolytic Properties
Electrolyte:
A substance whose aqueous solution forms ions; conducts electricity.
Ionic compounds.
Nonelectrolyte:
Substance that does not form ions in an aqueous solution; poor conductor.
Molecular compounds.

9. Ionic compounds in water:
Dissociate into its component ions.

11. A closer Look

12. Unequal sharing of electrons leads to partial positive and negative charges in a water molecule. These charges attract the ions which causes dissociation of the ionic compound in water.

13. Equations showing ionization or dissociation
Do not get to wrapped up in the difference between the terms ionization and dissociation. Consider them to mean the same thing, the separation of a substances ions.
Equations showing ionization or dissociation

14. Molecular compounds in H2O
Molecular compounds – nonmetal + nonmetal
Structural integrity of molecule is usually maintained meaning no ions form (C12H22O11)
Exception:
Some molecular solutes interact with water to form ions. These would be electrolytes.
Examples: Acids  HCl, H2C3O Ammonia  NH3

16. Strong Electrolytes
Exists in solution completely or almost completely as ions.
All ionic compounds and a few molecular compounds.(Ex: Strong Acids)
The arrow in one directions means there is no tendency of the ions to recombine, they fully or nearly fully ionize.m

17. Weak Electrolytes
Molecular compounds that produce a small concentration of ions when dissolved in H2O.
Ex: Acetic acid (HC2H3O2) only slightly ionizes when dissolved in water.
HC2H3O2(aq) H+(aq)+ C2H3O2-(aq)
Arrows in both directions means that the Chemical equilibrium results in ions forming at the same rate that molecules recombine.
Weak electrolytes get double arrow, strong electrolytes get the single arrow
Strong or weak electrolytes is not dependent on how well it dissolves.
Ba(OH)2 is not very soluble but is a strong electrolyte
Acetic acid dissolves well, but is a weak electrolyte.
Dilute concentrations of acetic acids are better conductors than concentrated solutions of acetic acid. This is because adding water pushes the equilibrium to the right in the following equation: HC2H3O2 + H2O  C2H3O2- + H3O+ Acetic acid generally ionizes about 10%. Stability occurs when the substances in the equation are at equilibrium with one another. If the equilibrium is put under stress, the reaction will shift to one side to relieve the stress. This is what occurs when acetic acid is diluted with water. The addition of water causes a shift to the right creating more available ions. Therefore, adding water increases the conductivity.
Weak acids are better conductors if they are dilute, as you will see in lab. Explain.

18. Precipitation Reactions
Reactions that result in an insoluble product.
Insoluble:
Substance with solubility less than 0.01 mol/L
Water molecules cannot overcome the attraction between the ions.

19. KI (aq) + Pb(NO3)2 (aq)  PbI2(s) + KNO3 (aq)
Precipitate

20. To do so, you must memorize how specific polyatomic ions act in water.
You must be able to determine whether a substance is soluble in water by simple examination of the chemical formula.
To do so, you must memorize how specific polyatomic ions act in water.
Not as hard as it sounds. We will focus mainly on 10 anions. This will give you the tools to predict the solubility of many compounds.

21. Solubility of Ionic Compounds
All acetates and nitrates are soluble in water.
All ionic compounds of alkali metals and ammonium are soluble.
Table on next slide is on page 118 of your textbook.
Cl, Br, I, soluble accept Hg, Ag, Pb.
Sulfate: soluble accept Hg, Pb, Sr, Ba
Insoluble:
CO3 , PO4(other than alkali metal and NH4)
OH accept Ca, Sr, Ba

23. Soluble Exceptions Ag+, Hg22+, Pb2+ None Cl- Acetate ion C2H3O2- Br-
Sr2+, Ba2+, Hg22+, Pb2+
Cl-
Acetate ion C2H3O2-
Br- I-
NO3-
SO42-

24. Insoluble Exceptions NH4+ and Group 1 metals
Group 1 metals and Ba2+, Sr2+, Ca2+
NH4+ and Group 1 metals and Ba2+, Sr2+, Ca2+
CO32-
OH-
PO43-
S2-

25. Equation Types Molecular Complete Ionic Net ionic equation 3
Molecular equations:
Shows complete chemical formulas of reactants and products. Does not show ionic character.
Complete ionic equation:
Show all strong electrolytes as ions.
Net ionic equation:
Show only those ions involved in the reaction.

26. Ionic Equations
Those ions that appear on both sides of a complete ionic equation are known as Spectator Ions.
Net ionic equations do not include spectator ions.
Spectator Ions do not play a role in the reaction.

27. Exchange Reactions
Metathesis reactions
Double displacement
Double replacement

28. Writing Net Ionic Equations
Write a balanced molecular equation.
Rewrite the equation showing ions of strong electrolytes only.
Identify and cancel all spectator ions.

29. All this acid rain is killing my complexion!
Acid-Base Reactions
All this acid rain is killing my complexion!
Acids:
Ionize in H2O, causes increase in H+ ions.
H+ ions are bare protons.
Acids are proton donors.
Acids and bases are very common electrolytes

30. Monoprotic Acids: (HCl, HNO3)
Acids that can only yield one H+ per molecule upon ionization.
HCl  H+ + Cl-
Triprotic Acid- H3PO4

31. Ionization occurs in 2 steps.
Diprotic Acids: (H2SO4)
Ionization occurs in 2 steps.
Only the first ionization is complete.

32. Is HF a weak or strong acid?
weak acid
Although it is a weak acid, this acid is extremely reactive because of the F- ion.
Must be kept in special polypropylene container because it eats through glass.
Used to etch glass.
Has caused major accidents in lab.

33. Bases
Substances that increase the OH- when added to water. (NaOH)
NH3 is a base. In water it accepts an H+ ion from HOH, leaving an OH- in solution.
NH3 is a weak electrolyte
About 1% ionizes to form NH4+/OH-

34. Strong acids and bases
Acids and bases that ionize completely in solution are strong acids and bases.
Those that only ionize partially are weak acids and bases.
You must memorize these.
HF is a weak acid because it only partially ionizes.
It is very reactive with many substances, this is due to not just the H+ c concentration, but also the F- ion, which is extremely reactive.
See list of strong acids and bases on page 115.
Most acids are weak.

35. Strong Acids Hydrochloric Acid – Hydrobromic Acid – HCl
Hydroiodic Acid –
Nitric Acid –
Sulfuric Acid –
Chloric Acid –
Perchloric Acid –
HCl
HBr HI
HNO3
H2SO4
HClO3
HClO4

36. Strong Bases All group 1 Metal Hydroxides
(LiOH, NaOH, KOH, RbOH, CsOH)
Heavy Group 2 Metal Hydroxides
Ca(OH)2, Sr(OH)2, Ba(OH)2

37. Once you memorize the strong acids and bases, you will have enough information to determine if a substance is a strong or weak electrolyte.

38. Soluble Ionic? YES NO Acid? YES NO NH3 ? StrongAcid? NO YES NO YES
Electrolyte
Acid?
YES NO
NH3 ?
StrongAcid? NO
YES NO
YES
Strong
Electrolyte
Weak
Electrolyte
Nonelectrolyte

39. Another Flowchart (Choose the one you like for practice)
Acid/base? No
Yes
Ionic or Molecular
Compound?
Strong or weak
acid/base?
Ionic
(metal + nonmetal)
Molecular
(nonmetals)
Strong
Weak
Soluble?
nonelectrolyte
Strong Electrolyte
Weak Electrolyte
Yes
Strong Electrolyte

40. Example problems: KF Na3PO4 NH3 CH3CH2OH HCl NO2 HC2H3O2 CH4 NH4Cl
CH3Cl
strong electrolyte
strong electrolyte
weak electrolyte
nonelectrolyte
strong electrolyte
nonelectrolyte
weak electrolyte
nonelectrolyte
strong electrolyte
nonelectrolyte

41. Handout: Flowchart
Homework:
Memorize Strong Acids and Bases
4.27 – 4.28
worksheet

42. Some Properties of Acids and Bases
Acid Properties
Sour taste
Turn blue litmus red
pH < 7
Base properties
Bitter taste
Turns red litmus blue
pH >7
slippery
Ask students to help you determine properties of acids and bases by having them test different types at their seats. Have students get goggles at beginning of class. For larger class have a few students get goggles. Have watch glasses ready and white paper. Show phenolphthalein, universal indicator, pH paper, red and blue litmus paper.

43. Acid + Base Neutralization
Products of a neutralization reaction have none of the properties of an acid or a base.
An acid reacts with a metal hydroxide to form a salt plus water.
A salt is defined as an ionic compound whose cation comes from a base and anion comes from an acid.

44. Neutralization Reactions
Acid + Base (Metal Hydroxide)  Salt + Water
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
H+ + Cl- + Na+ + OH-  Na+ + Cl- + H2O(l)
H+ + OH-  H2O(l)
Neutralization reactions between metal hydroxides and acids are metathesis reactions
The ions exchange partners.

45. Potassium Hydroxide + Sulfuric Acid
Write the net ionic equation for the following reaction. It might help to first write the molecular equation, and then the complete ionic equation, followed by the net ionic equation.
Ionic equation:
Net Ionic equation:
Potassium Hydroxide + Sulfuric Acid
Products K2SO4(aq) + HOH(l)

46. Neutralization Reaction of Weak Acid
*Remember, only strong electrolytes are written as ions.*
Acetic Acid + Sodium Hydroxide 
HC2H3O2(aq) + NaOH(aq)  NaC2H3O2(aq) + H2O(l)
Weak acid strong base soluble salt water
HC2H3O2 + Na+ + OH-  Na+ + C2H3O2- + H2O(l)
HC2H3O2(aq) + OH-(aq)  C2H3O2- (aq)+ H2O(l)

47. Acid/Base Rx’s with gas formation
Other bases besides OH- react with H+ to form molecular compounds.Two common bases are CO3-2 and S-2.
Carbonates and bicarbonates react with acid to form CO2.
Name the ions: Carbonate and sulfide

48. unstable Overall Net Hydrochloric acid + Sodium Sulfide 
2HCl (aq) + Na2S(aq)  H2S(g) + 2NaCl(aq)
2H+ (aq) + S2-(aq)  H2S(g)
Hydrochloric acid + Sodium Hydrogen Carbonate 
HCl (aq) + NaHCO3(aq)  NaCl(aq) + H2CO3(aq)
unstable
H2CO3(aq)  H2O(l) + CO2(g)
Overall
HCl (aq) + NaHCO3(aq)  NaCl(aq) + H2O(l) + CO2(g)
Net
H+ (aq) + HCO3-(aq)  H2O(l) + CO2(g)

49. Acid Spills and Antacids
NaHCO3 and Na2CO are used to clean up acid spills in the lab.
See page 127 for a list of antacids used in over the counter medications.

50. Oxidation-Reduction
Reactions

51. Reactions in which electrons are transferred between substances
Oxidation-Reduction
(Redox)
Reactions in which electrons are transferred between substances

52. Use of Oxidation numbers in determining redox reactions is basically a bookkeeping method for keeping track of electrons
You must be able to identify an oxidation-reduction reaction. But first, we must learn the rules for assigning oxidation #’s to different species.

53. Rules for oxidation numbers
Atoms in elemental form are 0.
Monatomic ion; charge of the ion is its oxidation number.
Nonmetals; usually negative numbers.
a.) oxygen = -2 unless a peroxide = -1
b.) Hydrogen +1 with nonmetals, -1 with metals
c.) Halogens (-1) unless bonded to oxygen (+)
in a polyatomic ion (Ex: ClO3-; Cl = +5)
4) Sum of oxidation numbers must = 0
Most electronegative (furthest to right and up) element gets a negative charge.
Rule 2: Group I A always +1; Group IIA always +2; Al is always +3
Rule 3:
a.) Oxygen is –2 (accept peroxides are –1)
b.) Hydrogen is -1 bonded to metal, +1 when bonded to nonmetals.
c.) F is always –1. Other halogens are –1 in binary compounds. When they are combined with oxygen, they are positive.
Rule 4:
a.) Neutral compound sum equals zero.
b.) Polyatomic ions: sum of the oxidation numbers equals the charge of the ion.

54. Atoms in elemental form are 0. Examples Ag Pb Cl2 O2
Oxidation # = 0 for 7 diatomic elements and for all
other elements when by themselves.

55. Monatomic ion--
charge of the ion is its oxidation number.
Examples
AgCl Ag = Cl = -1
PbI Pb = I = -1
Fe2O Fe = O = -2

56. 3) Nonmetals; usually negative numbers.
a.) oxygen = -2 unless a peroxide = -1
b.) Hydrogen +1 with nonmetals, -1 with metals
c.) Halogens (-1) unless bonded to oxygen (+)
in a polyatomic ion (Ex: ClO3-; Cl = +5)
Examples
PbO oxygen = Na2O2 oxygen = -1
H2S hydrogen = NaH hydrogen = -1
KI iodine = -1 KIO iodine =

57. Determine Oxidation # of element red element in each of the following:
MnO2 +4
KMnO4 +7
BrO2- +3
BrO3- +5
Br2
HClO4 +7
H2SO4 +6
PO33- +3
CaH2 -1
SO42- +6
Na2S -2
Mg(NO3)2 +5

58. If one reactant gains electrons another must lose electrons.
Again, oxidation reduction reactions occur when there is a transfer of electrons from one species to another in a reaction.
If one reactant gains electrons another must lose electrons.
Reduction is always accompanied by oxidation.

59. Oxidation-Reduction Reactions
An atom that becomes more positively charged is oxidized.
This is due to loss of e-.
The gain of electrons by an atom is called reduction.
Precipitations Reactions: Cations and anions come together to form insoluble ionic compounds
Neutralization Reactions: H and OH come together to from water.
Third type of reaction can result in a precipitate.
By transferring electrons between them. This is an oxidation reduction reaction.

60. Two mnemonics for remembering which substance is undergoing oxidation and which is undergoing reduction? -
OIL -- RIG
Oxidation Involves Loss -- Reduction Involves Gain
“Leo the lion says Ger”
Loss of electrons oxidation -- Gain of electrons reduction

61. Many metals react with O2 in the air to form metal oxides.
Metals lose electrons to oxygen.
2 Fe O2  FeO
As Fe is oxidized (loses e-),
oxygen is reduced (gains e-).
Reduction is gain

62. 2 Fe O2  FeO +2 -2
oxidation
reduction

63. Oxidation of metals by acids and salts
Reaction of a metal with either an acid or metal salt follows general form of:
A + BX AX +B
Single displacement reaction
Combustion reactions are a form of redox reaction.
Elemental oxygen is converted to compounds of oxygen.

64. CuSO4(aq) + Zn(s)  ZnSO4(aq) + Cu(s) +2+6-2 0 +2+6-2 0
CuSO4(aq) + Zn(s) 
ZnSO4(aq) + Cu(s)
Reduced
Oxidized
What are the products?
What are the charges on each species?
What is oxidized and what is reduced?

65. Magnesium + Hydrochloric Acid Aluminum + Cobalt(II) Nitrate
For the following reactants:
Write the reaction that occurs.
Identify what is being oxidized and reduced.
Magnesium + Hydrochloric Acid
Aluminum + Cobalt(II) Nitrate

66. Mg(s) + HCl (aq)  MgCl2 (aq) + H2(g)+1 +2
reduction
oxidation

67. Al(s) + Co(NO3)2 (aq)  Al(NO3)3(aq) + Co(s)+2 +3
reduction
oxidation

68. Types of Redox Reactions
Combination (synthesis)
Decomposition
Displacement
hydrogen, metal, halogen
Disproportionation (When an element is simultaneously oxidized and reduced).
Ex: H2O2  H2O + O2
Synthesis: S (s)+ O2(g)  _SO2(g)
Decomposition: 2KClO3(s)  2KCl(s) + 3O2(g)
Displacement:
hydrogen: 2Na(s) + 2H2O  2NaOH (aq) + H2(g)
Metal: Zn(s) + CuSO4(aq)  ZnSO4(aq) + Cu(s)
Halogen: Cl2(g) + 2KBr(aq)  KCl(aq) + Br2(l)
Disproportionation: an element in one oxidation state is simultaneously oxidized and reduced.
2H2O2(aq)  2H2O(l) + O2(g)
Cl2(g) + 2OH-  ClO-(aq) + Cl- + H2O(l)
Hypochlorite- Cl =+1 & Cl= -1

69. Activity Series
List of metals in order of decreasing ease of oxidation.
Alkali and alkaline earth metals are at the top. (active metals)
Gold, Silver, Platinum, and palladium are considered to be (noble metals) because they resist oxidation.

70. Using activity series to predict reactions
Activity series can be used to predict reactions between metals and metal salts or acids.
Any metal listed on the series can be oxidized by the ions of elements below it on the list.
This is used to show that the metal is oxidized by another metal.
HCl(aq) + Zn(s) > ZnCl2 + H2
HCl(aq) + Cu(s) > No reaction
HNO3(aq) + Cu(s) > Cu(NO3)2(aq) +2H2O(l) + 2NO2(g)

72. Using activity series of metals, which metals from the list below can be oxidized by H+?Ni Al Cu Pb Ag Mg Au
Demo (If time): Cu + AgNO3=  Cu(NO3)2 + Ag
Copper is oxidized, Ag+ is reduced.

73. Chemical Reaction Types
This should be a review before handing out All reaction types worksheet. Must edit all reaction type worksheet. It has to many weak acids reacting with strong bases, which is confusing for students. Make sure most have reactions (see note in folder on handout).
Chemical Reaction Types
Decomposition
Synthesis
Single Replacement
Precipitation
Neutralization
Combustion
Types of Redox Reactions
Types of Double Replacement Reactions

74. What are the 7 Diatomic Elements
H2 – hydrogen
N2 – nitrogen
O2 – oxygen
F2 - fluorine
Cl2 – chlorine
Br2 – bromine
I2 – Iodine

75. Synthesis A + B  AB Examples H2 (g) + O2 (g)  H2O (g)
Mg (s) + O2 (g)  MgO (s)
Na (s) + Cl2 (g)  NaCl (s)

76. Decomposition AB  A + B Examples NaCl (s)  Na (s) + Cl2 (g)
KClO3 (s)  KCl (s) + O2 (g)
elec

77. Single Replacement A + BC  AC + B Examples
Na (s) + HOH (l)  NaOH (aq) + H2 (g)
sodium replaces hydrogen in water
Cl2 (g) + NaBr (aq)  Br2 (l) + NaCl (aq)
chlorine replaces bromine in sodium bromide

78. Double Replacement (Metathesis)
AB + CD  AD + BC
Examples
AgNO3 (aq) + NaCl (aq)  AgCl (s) + NaNO3 (aq)
HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)
NH4Cl (aq) + NaOH (aq)  NH3 (g) + H2O (l) + NaCl (aq)
Blue color for the products represents the driving force
which allows the chemical reaction to occur.

79. Combustion CO2 + H2O Examples CH4 (g) + O2 (g)  CO2 (g) + H2O (l)
hydrocarbon + oxygen  carbon dioxide + water
Examples
CH4 (g) + O2 (g)  CO2 (g) + H2O (l)
C3H8 (g) + O2 (g)  CO2 (g) + H2O (l)
CH3OH (g) + O2 (g)  CO2 (g) + H2O (l)
CO2 + H2O

80. Neutralization Reactions
Strong Acid + Strong Base  Salt + Water
Example
HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)
H+ + Cl- + Na+ + OH-  Na+ + Cl- + H2O(l)
H+ + OH-  H2O

81. Solution Preparation

82. Concentration (Molarity)
Concentration — the amount of solute per unit of solution.
Molarity (M) — expresses concentration:

83. Must have the correct units.
Calculate the molarity of a solution that contains 20.0g copper(I) chloride and has a total volume of mL.
Must have the correct units.

84. Given: 20.0 grams CuCl; 300.0 ml solution
Need: Moles CuCl; L of solution 1
mol CuCl
20.0 g CuCl
0.202 mol CuCl =
99.0
g CuCl 1 L
300.0 mL L =
1000 mL

85. 0.673 M CuCl

86. Must calculate # of moles, and then convert it into grams.
How many grams of NaCl are needed to make 2.5 L of 0.20 molar solution?
Given: 0.20 M solution; 2.5 L solution
Need: grams of NaCl
(L)
(L)
Add 29g NaCl to a volumetric flask and add distilled water to a total volume of 2.5L
Must calculate # of moles, and then convert it into grams.
mol = ML

87. Mol = 0.20 mol NaCl x 2.5 L = 0.50 mol NaCl
Mol = ML
Mol = 0.20 mol NaCl x L = 0.50 mol NaCl L
55.84 g NaCl
0.50 mol NaCl
28 g NaCl = 1
mol NaCl

88. Dilution of Stock Solutions
Chemicals are purchased in concentrated form. They need to be diluted for most lab use.
Formula for dilution:
Mi Vi = Mf Vf
Mi is the stock molarity/ vi is the volume of stock need to make a desired volume(Vf) and molarity (Mf).

89. Must rearrange equation above to solve for initial volume.
How much stock (12 M) HCl (aq) is
required to make mL of 3M HCl (aq)?
Mi Vi = Mf Vf
Mi = 12 M
Vf = mL
Mf = 3 M
Vi = ? mL
Must rearrange equation above to solve for initial volume.

90. 50. mL
Measure 150 mL of water in a beaker. Slowly add 50.0 mL of 12 M HCl for a final volume of mL.

91. moles before dilution = moles after dilution
Two things to note:
Always add concentrated acid to water, and not the reverse to avoid unwanted splashing due to the heat generated.
2) When diluting a solution, the amount of solute doesn’t change, only the final volume. WA AW
moles before dilution = moles after dilution

92. If diluting a solution other than acids, start with initial volume of concentrated solution, and then dilute with distilled water until you have the desired volume.
You try one!

93. We want to prepare 500. mL of 1. 00 M acetic acid from a 17
We want to prepare 500. mL of 1.00 M acetic acid from a 17.5 M stock solution of acetic acid. What volume of the stock solution is required?
Mi Vi = Mf Vf
Mi = 17.5 M
Vf = 500. mL
Mf = 1.00 M
Vi = ? mL

94. 28.6 mL
Pour mL of distilled water into a beaker. Slowly pour the 28.6 ml of acid into the water and swirl. Fill the container with distilled water to 500. mL.

95. There may be times when you must consider the concentration of ions in a solution.
(You must consider the subscripts for this)
MgCl2  Mg Cl-
In a solution of 0.25 M MgCl2 you have:
M of Mg2+ = 0.25 M
M of Cl- = 2 x 0.25 M = 0.50 M
What is the concentration of each ion in the following?
0.15 Na3P
M of Na+ = 0.45 M
M of P3- = 0.15 M

96. Titrations Determining the concentration of an unknown solution.
Use a 2nd solution of known concentration (standard solution) that undergoes a reaction with the unknown solution.
Use the ratios in the balanced equation along with the M = mol/L equation to determine molarity of unknown.

97. The point at which the two solutions are stoichiometrically equal is known as the equivalence point.
The reaction is complete and no excess reactant is present.
How do we know when this occurs during the reaction?

98. In acid base reactions dyes known as indicators are used.
Phenolphthalein is colorless in acid solution, and pink in basic solution.
End point is reached when a drop of the base remains pink. There is no acid for this drop to react with and the solution is now basic.