ch.4 - the major classes of chemical...
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CHEMISTRY - SILBERBERG 7E
CH.4 - THE MAJOR CLASSES OF CHEMICAL REACTIONS
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CONCEPT: WRITING CHEMICAL REACTIONS
EXAMPLE: Predict whether a reaction occurs, and write the balanced molecular equation.
a. LiOH (aq) + MgSO4 (aq)
EXAMPLE: Predict whether a reaction occurs, and write the balanced molecular equation, the total and net ionic equations.
Molecular: Na2CO3 (aq) + HBr (aq)
Total Ionic:
Net Ionic:
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PRACTICE: Predict whether a reaction occurs, and write the balanced total and net ionic equations.
Total Ionic:
Net Ionic:
Molecular: Ag2SO4 (aq) + KCl (aq)
PRACTICE: Predict whether a reaction occurs, and write the balanced total and net ionic equations.
Total Ionic:
Net Ionic:
Molecular: MgBr2 (aq) + NaC2H3O2 (aq)
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CONCEPT: AQUEOUS SOLUTIONS
The ___________________________ of a compound represents the maximum amount of solute that dissolves in a solvent.
SOLUBILITY RULES
SOLUBLE IONIC COMPOUNDS
INSOLUBLE IONIC COMPOUNDS
1. Group 1A ions (Li+, Na+, K+, etc.) and ammonium ion (NH4+) are soluble.
1. (Hydroxides) OH- and (Sulfides) S2-, are insoluble
except when with Group 1A ions (Li+, Na+, K+, etc.),
ammonium ion (NH4+) and Ca2+, Sr2+, Ba2+.
2. (Nitrates) NO3- , (acetates) CH3COO- or C2H3O2-,
and most perchlorates (ClO4-) are soluble.
2. (Carbonates) CO32- and (Phosphates) PO43- are
insoluble except when with Group 1A ions
(Li+, Na+, K+, etc.), ammonium ion (NH4+).
3. Cl- , Br- , and I- are soluble, except when paired
with Ag+ , Pb2+ , Cu+ and Hg22+.
4. (Sulfates) SO42- are soluble, except those of Ca2+ ,
Sr2+ , Ba2+ , Ag+ , and Pb2+ .
When we classify a compound as soluble it means that the compound is _______________________, it is also known as
a(n) _______________________ because it conducts electricity.
NaNO3 (s) H2O
Na+ (aq) + NO3– (aq)
When we classify a compound as insoluble it means that the compound is a _______________________, it is also known
as a(n) _______________________ because it doesn’t conduct electricity.
CH3OH (l) H2O
CH3OH (aq)
BaSO4 (s) H2O
BaSO4 (aq)
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CONCEPT: ELECTROLYTES
Whenever we add a solute into a solvent three outcomes are possible:
• the solute will _________________ dissolve ( STRONG electrolytes).
• the solute will _________________ dissolve ( WEAK electrolytes).
• the solute will _________________ dissolve ( NON electrolytes).
Classification of Solutes in Aqueous Solution
STRONG ELECTROLYTES
WEAK ELECTROLYTES
NONELECTROLYTES
1. STRONG ACIDS: HCl, ______ , HI ,
HNO3 , _______ , _______ , _______ .
2. STRONG BASES:
Group 1A Metal with OH-, H-, O2- or
NH2-
Groups 2A Metal, Calcium or Lower, with
OH-, H-, O2- or NH2-
3) SOLUBLE IONIC COMPOUNDS:
1. WEAK ACIDS: HF, ____________ ,
________ , ________ , ________ .
2. WEAK BASES: Be(OH)2 , Mg(OH)2 ,
_________ , _________ .
1. MOLECULAR
COMPOUNDS:
______________
C6H12O6 (glucose)
C12H22O11 (sucrose)
______________
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PRACTICE: ELECTROLYTES
EXAMPLE: Each of the following reactions depicts a solute dissolving in water. Classify each solute as a strong electrolyte,
a weak electrolyte or a non-electrolyte.
a. PbSO4 (s) PbSO4 (aq)
b. HC2H3O2 (aq) H+ (aq) + C2H3O2– (aq)
c. CaS (s) Ca2+ (aq) + S2- (aq)
d. Hg (l) Hg (aq)
PRACTICE: Classify each of the following solutes as either a strong electrolyte, a weak electrolyte or a non-electrolyte.
a. Perbromic acid, HBrO4
b. Lithium chloride, LiCl
c. Formic Acid, HCO2H
d. Methylamine, CH3NH2
e. Zinc bromide, ZnBr2
f. Propanol, C3H8OH
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CONCEPT: OXIDATION-REDUCTION REACTIONS
Chemists use some important terminology to describe the movement of electrons.
• In ______________ reactions we have the movement of electrons from one reactant to another.
L
E
O
G
E
R
Agent Agent
Rules for Assigning an Oxidation Number (O.N.)
A. General Rules
1. For an atom in its elemental form (Na, O2, S8, etc.): O.N. = 0
2. For an ion the O.N. equals the charge: Na+ , Ca2+ , NO3 –
B. Specific Rules
1. Group 1A: O.N. = +1
2. Group 2A: O.N. = +2
3. For hydrogen: O.N. = +1 with nonmetals
O.N. = -1 with metals and boron
4. For Fluorine: O.N. = -1
5. For oxygen: O.N. = -1 in peroxides (X2O2 , X = Group 1(A) element)
O.N. = − 12
in superoxides (XO2 , X = Group 1(A) element)
O.N. = - 2 in all other compounds
6. Group 7A O.N. = -1 (except when connected to O)
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CONCEPT: OXIDATION-REDUCTION REACTIONS (PRACTICE)
EXAMPLE: In the following reaction identify the oxidizing agent and the reducing agent:
a. 2 C6H6 (l) + 15 O2 (g) 12 CO2 (g) + 6 H2O (g)
PRACTICE: What is the oxidation number of each underlined element?
a. P4 b. BO33-
c. AsO42- d. HSO4
–
PRACTICE: In the following reaction identify the oxidizing agent and the reducing agent:
a. Cr2O72- + 6 Fe2+ + 14 H+ 2 Cr3+ + 6 Fe3+ + 7 H2O
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CONCEPT: BASIC REDOX CONCEPTS OXIDATION-REDUCTION (REDOX) reactions deal with the transfer of electrons from one reactant to another.
Lose
Electrons
Oxidation
Gain
Electrons
Reduction
Reducing Agent (Reductant)
} }} Element
becomes
more positive
Oxidation
Number
Increases } Element
becomes
more negative
Oxidation
Number
Decreases
Oxidizing Agent (Oxidant)
Li (s) Li+ (aq) + e – Cl2 (g) + 2 e – 2 Cl – (aq)
Li (s) + Cl2 (g) Li+ (aq) + 2 Cl – aq)
Electrical Charge
The units for electrical charge are measured in ________________ (C).
} Charge of 1 electron Faraday Constant
}(1.602×10−19C) ⋅ (6.022×1023mol−1) = 9.647×104C
1mole e−charge mole
e –
q = n ⋅ F Faraday Constant
Electrical Current
The units for electrical current are in __________ (A).
Electrical Voltage
The relationship between work and voltage can be expressed as:
The relationship between Gibbs Free Energy and electric potential can be expressed as:
Ohm's Law
The units for resistance are in __________ (Ω).
Power
Power represents work done per unit of time. The units for power are in __________ (W).
w = E ⋅ qWork Voltage Charge
ΔG = − n ⋅ F ⋅ E GibbsFree Energy
mole e –
Faraday Constant
Voltage
I = ER
Voltage
ResistanceCurrent
P = E ⋅ IPower Voltage Current
Current
Charge
TimeI = q
t
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CONCEPT: SINGLE REPLACEMENT REACTIONS In a single replacement or displacement reaction one element displaces another element from a compound. More reactive or active metals displace less reactive metals or hydrogen from compounds.
Act
ivit
y In
crea
ses
Metals
Lithium (Li) > Potassium (K) > Barium (Ba) > Strontium (Sr) > Calcium (Ca) > Sodium (Na)
Activity
Metals in this category can displace hydrogen from liquid water, steam and acids:
Metals
Magnesium (Mg) > Aluminum (Al) > Zinc (Zn2+) > Chromium (Cr2+, Cr3+) > Iron (Fe2+, Fe3+)
Activity
Metals in this category can displace hydrogen from steam and acids:
Metals
Cadmium (Cd2+) > Cobalt (Co2+, Co3+) > Nickel (Ni2+) > Tin (Sn2+, Sn4+) > Lead (Pb2+, Pb4+)
Activity
Metals in this category can displace hydrogen from acids:
Hydrogen and Metals
Hydrogen (H) > Antimony (Sb3+) > Arsenic (As3+, As5+) > Bismuth (Bi3+) > Copper (Cu+, Cu2+) > Mercury (Hg22+, Hg2+) > Silver (Ag+) > Palladium (Pd3+) > Platinum (Pt2+, Pt3+) > Gold (Au+, Au3+)
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PRACTICE: SINGLE REPLACEMENT REACTIONS EXAMPLE 1: Based on your understanding of activities determine if a reaction occurs and if so provide the products formed.
Ba (s) + H2O (g)
EXAMPLE 2: Based on your understanding of activities determine if a reaction occurs and if so provide the products formed.
Zn (s) + NiCl2 (aq)
EXAMPLE 3: If the activity of halogens is stated as: Fluorine > Chlorine > Bromine > Iodine, determine if a reaction occurs and if so provide the products formed.
Cl2 (g) + AlBr3 (aq)
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16. How many milligrams of NaCN are required to prepare 712 mL of 0.250 M NaCN?
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17. What volume (in µL) of 0.100 M HBr contains 0.170 moles of HBr?
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18. How many moles of Ca2+ ions are in 0.100 L of a 0.450 M solution of Ca3(PO4)2?
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19. How many chloride ions are present in 65.5 mL of 0.210 M AlCl3 solution? a) 4.02 × 1023 chloride ions
b) 5.79 × 1024 chloride ions
c) 2.48 × 1022 chloride ions
d) 8.28 × 1021 chloride ions
e) 1.21 × 1022 chloride ions
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22. To what final volume would 100 mL of 5.0 M KCl have to be diluted in order to make a solution that is 0.54 M KCl?
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23. If 880 mL of water is added to 125.0 mL of a 0.770 M HBrO4 solution what is the resulting molarity?
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26. Consider the following balanced redox equation:
H2O + 2 MnO4 – + 3 SO32- 2 MnO2 + 3 SO42- + 2 OH – How many grams of MnO2 (MW: 86.94 g/mol) are produced when 32.0 mL of 0.615 M MnO4- (MW: 118.90 g/mol) reacts with excess water and sulfite?
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27. Iron (III) can be oxidized by an acidic K2Cr2O7 solution according to the net ionic equation:
Cr2O72- + 6 Fe2+ + 14 H+ 2 Cr3+ + 6 Fe3+ + 7 H2O
If it takes 35.0 mL of 0.250 M FeCl2 to titrate 50 mL of a solution containing Cr2O72-, what
is the molar concentration of Cr2O72-?
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28. Vinegar is a solution of acetic acid, CH3COOH, dissolved in water. A 5.54 g sample of
vinegar was neutralized by 30.10 mL of 0.100 M NaOH. What is the percent by weight of
acetic acid in the vinegar?
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29. What is the molar mass of a 0.350 g sample of a monoprotic acid if it requires 50.0 mL of 0.440 M Ca(OH)2 to completely neutralize it?
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30. Give the complete ionic equation for the reaction (if any) that occurs when aqueous solutions of sodium sulfide and copper (II) nitrate are mixed. a) Na+ (aq) + SO42-(aq) + Cu+(aq) + NO3-(aq) CuS(s) + Na+(aq) + NO3-(aq)
b) Na+ (aq) + S-(aq) + Cu+(aq) + NO3-(aq) CuS(s) + NaNO3(aq)
c) 2 Na+(aq) + S2-(aq) + Cu2+(aq) + 2 NO3-(aq) Cu2+(aq) + S2-(aq) + 2 NaNO3(s)
d) 2 Na+(aq) + S2-(aq) + Cu2+(aq) + 2 NO3-(aq) CuS(s) + 2 Na+(aq) + 2 NO3-(aq)
e) No reaction occurs.
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31. Give the net ionic equation for the reaction (if any) that occurs when aqueous solutions of H2SO4 and KOH are mixed. a) H+(aq) + OH-(aq) H2O(l)
b) 2 K+(aq) + SO42-(aq) K2SO4(s)
c) H+(aq) + OH-(aq) + 2 K+(aq) + SO42-(aq) H2O(l) + K2SO4(s)
d) H22+(aq) + OH-(aq) H2(OH)2(l)
e) No reaction occurs.
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32. Give the net ionic equation for the reaction (if any) that occurs when aqueous solutions of Na2CO3 and HCl are mixed. a) 2 H+(aq) + CO32-(aq) H2CO3(s)
b) 2 Na+(aq) + CO32-(aq) + 2 H+(aq) + 2 Cl-(aq) H2CO3(s) + 2 NaCl(aq)
c) 2 H+(aq) + CO32-(aq) H2O(l) + CO2(g)
d) 2 Na+(aq) + CO32-(aq) + 2 H+(aq) + 2 Cl-(aq) H2CO3(s) + 2 Na+(aq) + 2 Cl-(aq)
e) No reaction occurs.
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