1. Chemical Reactions in Aqueous Solutions
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Honors Unit 4:
Chemical Reactions in Aqueous Solutions
Dr. Mihelcic Honors Chemistry

2. Solute Concentrations, Molarity
Solution: homogeneous mixture of two or more substances
Solute: the substance being dissolved
Solvent: the substance doing the dissolving

3. Concentration of a solution:
11/6/2017
Concentration of a solution:
the quantity of a solute in a given quantity of solution (or solvent)
A concentrated solution contains a relatively large amount of solute vs. the solvent.
A dilute solution contains a relatively small concentration of solute vs. the solvent.
“Concentrated” and “dilute” aren’t very quantitative .
Dr. Mihelcic Honors Chemistry

4. Molarity Molarity (M), or molar concentration –
the amount of solute, in moles per liter of solution
(Formula) Molarity = moles of solute
liters of solution
A solution that is 0.35 M sucrose contains moles of sucrose in each liter of solution.
Keep in mind that molarity signifies moles of solute per liter of solution, not liters of solvent.

5. Add more water to reach the 1.000 liter mark.
Volumetric Flasks
Weigh mol (1.580 g) KMnO4.
Dissolve in water. How much water? Doesn’t matter, as long as we don’t go over a liter.
Add more water to reach the liter mark.

6. Molarity Calculations
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Molarity Calculations
Example #1: What is the molarity of a solution prepared by dissolving 5.0 mol of NaCl in enough water to make 2 L of solution?
Dr. Mihelcic Honors Chemistry

7. Molarity Calculations
11/6/2017
Molarity Calculations
Example #2: What is the molarity of a solution prepared by dissolving g of AgNO3 in enough water to make 1500 mL of solution?
Dr. Mihelcic Honors Chemistry

8. Molarity Calculations
11/6/2017
Molarity Calculations
Example #3: How many grams of sodium carbonate are needed to prepare L of a M solution?
Dr. Mihelcic Honors Chemistry

9. Arrhenius’s Theory of Electrolytic Dissociation
Why do some solutions conduct electricity?
Arrhenius’s theory:
Certain substances dissociate into cations and anions when dissolved in water.
These ions allow electricity to flow!

10. Review: Unit 2 Ionic Compounds
***Electrolytes dissociate to produce ions.
The more the electrolyte dissociates,
the more ions it produces.

12. Types of Electrolytes A strong electrolyte dissociates completely.
A strong electrolyte is present in solution almost exclusively as ions.
Strong electrolyte solutions are good conductors.
Ex: most ionic compounds

13. Types of Electrolytes A weak electrolyte dissociates partially.
Weak electrolyte solutions are poor conductors.
Different weak electrolytes dissociate to different extents.
Ex: weak acids and bases

14. Nonelectrolytes A nonelectrolyte does not dissociate.
A nonelectrolyte is present in solution almost exclusively as molecules.
Nonelectrolyte solutions DO NOT conduct electricity.
Ex: sugar, ethanol

15. Is it a strong electrolyte, a weak electrolyte, or a nonelectrolyte?
Strong electrolytes include:
Strong acids & Strong bases
Most water-soluble ionic compounds
Weak electrolytes include:
Weak acids and weak bases
A few ionic compounds
Nonelectrolytes include:
Most molecular compounds
Most organic compounds (most are molecular)
How do we tell whether an acid (or base) is weak? Memorize the strong acids/bases!

16. Dissolution Equations (Strong Electrolytes)
Write cations and anions as individual particles with the appropriate amount of each:
NaCl (aq) 
MgSO4(aq) 
Na2SO4(aq) 

17. Calculating Ion Concentrations in Solution
Multiply molarity (M) x COEFFICIENT of each ion to find the ion concentration!
EXAMPLE:
Na2SO4  2Na SO42-
.010 M M M

18. Example
Write the dissociation equation and then calculate the concentrations of all ions present in the solution:
a M solution of NaCl
b. 1.5 M solution of AlF3
c. 0.5 M solution of Al2(CO3)3

19. Types of Reactions (Review)
Synthesis (combination)
A + B  AB
Decomposition
AB  A + B
Single Replacement
A + BX  AX + B
Double Replacement
AX + BY  AY + BX
Combustion

20. Predicting Products (Synthesis)
Synthesis Reactions (A + B  AB)
Examples:
Silver + oxygen 
Calcium + chlorine 
Sodium + fluorine 

21. Predicting Products (Decomposition)
Decomposition Reactions (AB  A + B)
Examples:
Water 
Potassium chloride 
Iron (II) oxide 

22. Chemical Reactions in Water
Single Replacement Reactions
Metal replaces a Metal in solution
Nonmetal replaces a Nonmetal in solution
But a reaction doesn’t always happen!

23. Predicting Products – The Activity Series
A more active METAL has a higher oxidation potential
Greater tendency to form cations
A more active NONMETAL has a higher reduction potential
Greater tendency to form anions

24. Predicting Products (Single Replacement Reaction)
Must use the activity series!!!
Aluminum and lead (II) chloride 
Lead and calcium chloride 
Chlorine and iron (III) bromide 
Iodine and lithium fluoride 

25. Chemical Reactions in Water
DOUBLE REPLACEMENT REACTIONS
Pb(NO3) 2(aq) + 2 KI(aq) ----> PbI2(s) KNO3 (aq)
The cations change places.

26. Precipitation Reactions
Precipitate: insoluble product
Solid formed as a result of a reaction
In this unit, solid formed from a double replacement reaction!
Example: PbI2 is the precipitate above
Pb(NO3) 2(aq) + 2 KI(aq) ---> PbI2(s) KNO3 (aq)

27. Precipitation Reactions
The “driving force” is the formation of an insoluble compound — a precipitate.
Pb(NO3)2 (aq) KI (aq)
-----> 2 KNO3(aq) PbI2(s)

28. Solubility Rules See handout. (Or look at pg
Solubility Rules See handout!! (Or look at pg. 13/14 in your reference book!)

29. Use of Solubility Tables— See Reference Booklet (pg. 13/14)
s = soluble -- dissolves in water, is (aq)
i = ss = insoluble or slightly soluble -- forms a precipitate, is (s)

30. Solubility Rules Example #1: Are the following compounds soluble or insoluble? NaNO3 Ba(OH)2 NaSO4 CaSO4

31. Example #2
Using the solubility rules, predict the products for the following reactions and indicate if they are soluble or insoluble.
Cu(NO3)2 + (NH4)2SO4 
FeCl3 + AgNO3 

32. Net Ionic Equations
Net Ionic Equations– equation showing only those compounds and ions that undergo a chemical change in a reaction (when dissolved in aqueous solution)
Charges & atoms must still balance!
Spectator Ions – ions that are present in aqueous solution but are not involved in a chemical reaction

33. Sample Problem Ba(NO3)2 + Na2CO3 
Predict the products of the following reactions. Then indicate which compounds would be soluble and which would be insoluble. Finally, write the full ionic equation and the net ionic equation.
Ba(NO3) Na2CO3 

34. Sample Problems (continued)
CuCl (NH4)2SO4 
MgCl AgNO3 

35. END OF
MATERIAL
FOR
PART I !!!!

36. Acids and Bases

37. Properties of Acids Taste sour
React with indicators: turn blue litmus red turn phenolphthalein colorless
React with certain metals to form H2
Corrosive to metals and skin (strong acids)
Neutralize bases
to form water and
salts

38. Properties of Bases Taste bitter Feel slippery React with indicators:
turn red litmus blue
turn phenolphthalein pink
Corrosive to skin (strong bases)
Neutralize acids to form
water and salts

39. Arrhenius Definitions of Acids and Bases
Acids produce H+ in water solution
Bases produce OH- in water solution
Examples: HCl (aq)  H+ + Cl-
NaOH (aq)  Na+ + OH-

40. Arrhenius definitions are limited!!!
Not all bases contain OH-
H+ does not exist by itself in aqueous systems

41. Definition: Hydronium Ion
In aqueous solution, H+ does NOT exist!
Note: In problems, [H+] = [H3O+]
H+ + H2O  H3O+
(hydronium ion)

42. Strong Acids/Bases HNO3  H+ + NO3- Strong Acids: THERE ARE ONLY SIX!
HCl (hydrochloric) HNO3 (nitric)
HBr (hydrobromic) HClO4 (perchloric)
HI (hydroiodic) H2SO4 (sulfuric)
Strong acids are 100% ionized in solution!
(Use single arrow; no equilibrium established)
HNO3  H+ + NO3-

43. Strong Bases Sr(OH)2  Sr2+ + 2 OH - Strong Bases:
Group I metals + OH –
(Some Group II metals + OH –)
Some examples: NaOH, LiOH, Sr(OH)2
Strong bases are 100% ionized in solution!
(Use single arrow; no equilibrium established)
Sr(OH)2  Sr OH -

44. Common Strong Acids and Strong Bases- memorize!!!
MEMORIZE the strong acids and know how to recognize the strong bases!

45. Weak Acids and Bases Weak acids are not fully ionized in water.
Weak organic acids contain the –COOH group
One of the best known is acetic acid = CH3COOH

46. Weak Acids and Bases Weak bases are not fully ionized in water.
Use double arrow; equilibrium established
One of the best known weak bases is ammonia
NH3(aq) + H2O(liq)  NH4+(aq) + OH-(aq)

47. Acid/Base Theories
The most general theory for determining acids and bases in aq. soln. is the BRØNSTED - LOWRY theory
DEFINITIONS:
ACIDS DONATE H+ IONS
BASES ACCEPT H+ IONS

48. Bronsted-Lowry Theory
The Brønsted definition means NH3 is a BASE in water — and water is an ACID (in this case)

49. Conjugate Pairs NH3 / NH4+ is a conjugate pair
Related by the gain or loss of H+
Every acid has a conjugate base, formed when H+ is removed from the acid.
Every base has a conjugate acid, formed when H+ is added to the base.

50. Conjugate Pairs

51. Example #1: Write the Bronsted Lowry equation for the weak acid HNO2 and the weak base NH3, identifying the conjugate acid-base pairs in each equilibrium.

52. Example #2: Write the Bronsted Lowry equation for the weak acid HSO4- and the weak base pyridine (C5H5N), identifying the conjugate acid-base pairs in each equilibrium.

53. Amphiprotic or Amphoteric Substances
Some substances can function as both an ACID OR a BASE, depending on what they are reacting with.
These can donate OR accept H+
They are called amphiprotic or amphoteric
H2O + H2O  H3O OH-

54. Naming Acids Acids fall into 2 categories based on their formulas
Binary acids - H+ w/ 1 other element (HCl, H2S, etc.)
H+ with a polyatomic ion
These two categories each have their own set of rules for naming the acid

55. Naming Acids – Binary Acids
“Hydro__________ic acid”
**In the blank, put the root of the non-metal element’s name
Ex.) HCl = hydrochloric acid [chlorine  “chlor”]
HI = hydroiodic acid [iodine  “iod”]

56. Naming Acids – Acids w/ Polyatomic Ions
Change the ending of the polyatomic ion as shown below:
“-ate” ending  “-ic” ending
“-ite” ending  “-ous” ending
Add the word “acid” after the name of the polyatomic ion with the changed ending
I ate ic, and now I have ite-ous!

57. Naming Acids – Acids w/ Polyatomic Ions
Examples:
H2SO4 – Sulfate ion  Sulfuric acid
HNO3 –
HNO2 –
I ate ic, and now I have ite-ous!

58. Naming Acids – Acids w/ Polyatomic Ions
Examples:
H2SO4 – Sulfate ion  Sulfuric acid
HNO3 – Nitrate ion  Nitric acid
HNO2 – Nitrite ion  Nitrous acid

59. Practice!! Name the following acids:
H2CO3 HCl
HBr HF
H3PO4 H2SO3
HClO2 CH3COOH

60. Practice!! Name the following acids:
H2CO3 HCl
carbonic acid hydrochloric acid
HBr HF
hydrobromic acid hydrofluoric acid
H3PO4 H2SO3
phosphoric acid sulfurous acid
HClO2 CH3COOH
chlorous acid acetic acid

61. The pH Scale pH = - log [H3O+]
A common way to express acidity and basicity is with pH (the “power of hydrogen”)
pH = - log [H3O+]
***pH calculations look at the pH, pOH, concentration of H+ ion in solution [H+], and concentration of OH- in solution [OH-]

62. ***pH scale ranges from 0-14
Size of pH
***pH scale ranges from 0-14
Basic solution pH > 7
Neutral pH = 7
Acidic solution pH < 7
Big number = Basic

63. Neutral, Acid and Basic Solution
Neutral solution:
[H+] = [OH-] = 1.0 x 10-7 M
Acid solution:
[H+] > [OH-]; [H+] > 1.0 x 10-7 M
Basic solution:
[H+] < [OH-]; [OH-] > 1.0 x 10-7 M

64. The pH Scale

65. pH of Common Substances

66. pH Relationships pH = -log[H+] or pH = -log[H3O+] pOH = -log[OH-]
pH + pOH = 14
[H3O+][OH-] = 1 x = Kw

67. Water Dissociation Constant a.k.a. Auto-ionization Constant of Water
For any sample of water molecules:
H2O (liq) + H2O (liq)  H3O+ (aq) + OH- (aq)
Kw = [H3O+] [OH-] = x (at 25 oC)

68. But how would you find the [H+] or [OH-] from the pH or pOH??
pH Relationships
pH = -log[H+] pOH = -log[OH-]
But how would you find the [H+] or [OH-] from the pH or pOH??

70. Example #1:
A sample of tap water has a [H+] = 2.8 x 10-6 M. What is the [OH-]?

71. Example #2:
Calculate the pH, pOH and [OH-] of an acid solution whose [H+] is 1.8 x 10-4 M.

72. Example #3:
If the pH of Coke is 3.12, what is the concentration of the hydrogen ion in solution? Also, is this solution acidic or basic?

73. pH of Strong Acids
Strong Acids dissociate completely in aqueous solution:
E.g. HCl  H Cl-
For a strong acid, [H+] can be calculated from the molarity of the acid.
2.0 M M M

74. Example #4:
Calculate the [H+], pH, pOH, and [OH-] of a 0.15 M solution of the strong acid, HNO3.
Strategy:
Write the disassociation equation and calculate [H+].
Calculate pH, pOH, and [OH-] like you would for a regular pH problem.

75. Example #4:
Calculate the [H+], pH, pOH, and [OH-] of a 0.15 M solution of the strong acid, HNO3.

76. Strong Bases Strong Bases dissociate completely in aqueous solution:
E.g. Sr(OH)2  Sr OH-
For a strong base, [OH-] can be calculated from the molarity of the base.
0.5 M M M

77. Example #5: Calculate the pH, pOH, [H+], and [OH-] of a solution made by dissolving mol of Ba(OH)2 - a strong base - in 5.00 L of water.
Strategy:
1. Calculate the molarity of Ba(OH)2.
Write disassociation equation; calculate molarity of OH-
Calculate pOH, pH, and [H+] as you would for a regular pH problem.

78. Example #5: Calculate the pH, pOH, [H+], and [OH-] of a solution made by dissolving mol of Ba(OH)2 - a strong base - in 5.00 L of water.

79. Neutralization Reactions & Titrations
Definitions:
Neutralization reaction
Titration
Buret
Titrant
Indicator
Endpoint
Equivalence point

80. Neutralization Reactions
Neutralization is the (usually complete) reaction of an ACID + BASE
The products of a neutralization rxn are a “salt” (ionic compound) + water
It is a double replacement reaction.
Example: HCl + NaOH  H2O + NaCl

81. Definitions
A titration is a carefully controlled neutralization reaction.
Titrations are used to determine the concentration of an unknown acid or base.
A buret is the piece of glassware used in a titration.

82. Definitions
The titrant is the substance of known concentration used to determine the unknown concentration of the other substance.
An indicator--substance that changes color at a certain pH—is added to tell us when the neutralization is complete.
Example: Phenolphthalein undergoes a color change between pH 8 and 10
clear in acid
Light pink in neutral
Dark pink in base

83. Definitions
The equivalence point is the point in the titration where the neutralization is complete:
[H3O+] = [OH-]
The endpoint is the point where the indicator changes color.
***If the indicator is chosen correctly, these two points are identical!

84. 1 2 3 4

85. Strong Acid/Base Titration
Chemical reaction:
ACID + BASE  SALT + WATER
HCl + NaOH  NaCl + H2O
A double replacement reaction!

86. Strong Acid/Base Titration: Net Ionic Equation
HCl + NaOH  H2O + NaCl
In the reaction above: HCl, NaOH, and NaCl all are strong electrolytes and dissociate completely.
The actual reaction occurs between ions.
H+ + Cl– + Na+ + OH–  H2O (l) + Na+ + Cl–
H+ + OH–  H2O (l)
Na+ and Cl– are spectator ions.
A net ionic equation shows the species actually involved in the reaction.

87. Strong Acid/Base Titration
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Strong Acid/Base Titration
Short-cut Equation:
(# H+)MAVA = (# OH-) MBVB
Where #H+ and #OH- are obtained from the formula
Volumes can be in mL or L, but must be same units on both sides of equation!
Dr. Mihelcic Honors Chemistry

88. Strong Acid/Base Titration
At the equivalence point (end point):
Total mols of H+ from the acid
Total mols of OH- from the base =
*Remember:
(conc)(vol in L) = moles

89. Example #1: HCl is titrated with NaOH.
How many milliliters of M HCl are required to titrate mL of M NaOH?

90. Example #2: Ba(OH)2 is titrated with HCl.
How many milliliters of M Ba(OH)2 are required to titrate mL of M HCl?

91. END OF MATERIAL FOR PART 2 TEST!!

92. Reactions Involving Oxidation and Reduction
Oxidation: Loss of electrons
Reduction: Gain of electrons
Both oxidation and reduction must occur simultaneously.
If something loses electrons, something else has to gain them!
Historical: “oxidation” used to mean “combines with oxygen”; the modern definition is much more general.

93. Oxidation Numbers
An oxidation number is the charge on an ion, or a hypothetical charge assigned to an atom in a molecule or polyatomic ion.
Example:
In NaCl, the ox. number of Na is +1, that of Cl is –1 (the actual charges).
In CO2 (a molecular compound, no ions) the oxidation number of oxygen is –2, because oxygen as an ion would be expected to have a – 2 charge.
Why does the carbon in CO2 have an ox. number of +4?

94. Identifying Oxidation–Reduction (Redox) Reactions
In a redox reaction, the oxidation number of a species changes during the reaction.
Oxidation occurs when the oxidation number increases (species loses electrons).
Reduction occurs when the oxidation number decreases (species gains electrons).
If any species is oxidized or reduced in a reaction, that reaction is a redox reaction!!

95. A Redox Reaction: Mg + Cu2+  Mg2+ + Cu
Electrons are transferred from Mg metal to Cu2+ ions and …
… the products are Cu metal and Mg2+ ions.

96. Oxidizing and Reducing Agents
An oxidizing agent causes another substance to be oxidized.
The oxidizing agent is reduced.
A reducing agent causes another substance to be reduced.
The reducing agent is oxidized.
Mg + Cu2+  Mg2+ + Cu
What is the oxidizing agent?
What is the reducing agent?

97. Oxidation–Reduction Equations
Redox equations must be balanced according to both mass and electric charge.
For now, our main goals will be to:
Identify oxidation–reduction reactions.
Balance certain simple redox equations.
Recognize, in all cases, whether a redox equation is properly balanced.

98. Rules for Assigning Oxidation Numbers
Neutral Species
For an isolated atom or molecule the oxidation numbers is ex: Cl2, Fe
Monatomic Ions (Groups 1, 2, 17)
Group 1 elements all have an oxidation number of +1
Group 2 elements all have an oxidation number of +2.
Fluorine always has an oxidation number or -1.
Oxygen
In most compounds, oxygen has an oxidation number of –2.
Hydrogen
+1 when bonded to a nonmetal, -1 when bonded to a metal

99. Assigning Oxidation Numbers
The sum of the oxidation numbers of all atoms (or ions) in a neutral compound = 0.
The sum of the oxidation numbers of all atoms in a polyatomic ion = charge on the polyatomic ion.
ex: CaCl2

100. Example #1: Assign oxidation numbers to the elements in the following species:
CaC2O4 Cr2O72-
N2O N2O4
ClO ClO41-

101. Identifying Redox Reactions
Look for changes in oxidation numbers!!!
Al (s) + Fe2O3(s)  2 Fe (l) + Al2O3
What gets oxidized?
What gets reduced?

102. A) Fe2+(aq) + Cr2O72-(aq) + H+(aq)  Fe3+(aq) + Cr3+(aq) + H2O(l)
Example #2: Identify the element reduced, the element oxidized, the reducing agent and the oxidizing agent:
A) Fe2+(aq) + Cr2O72-(aq) + H+(aq) 
Fe3+(aq) + Cr3+(aq) + H2O(l)
B) 3 Cl2 (g) Cr(OH)3 (aq) + 10 OH- 
2CrO4-(aq) + 6Cl-(aq) + 8H2O(l)

103. Applications of Redox Reactions
In industry: to produce iron, steel, other metals, and consumer goods.
In foods and nutrition: redox reactions “burn” the foods we eat; antioxidants react with undesirable free radicals.

104. Applications of Redox Reactions
Everyday life: to clean (bleach) our clothes, sanitize our swimming pools (“chlorine”), and to whiten teeth (peroxide).