ch.4 - the major classes of chemical...

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CHEMISTRY - SILBERBERG 7E

CH.4 - THE MAJOR CLASSES OF CHEMICAL REACTIONS

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:



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)



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)



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)

______________



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



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)



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



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



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+)



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)



16. How many milligrams of NaCN are required to prepare 712 mL of 0.250 M NaCN?



17. What volume (in µL) of 0.100 M HBr contains 0.170 moles of HBr?



18. How many moles of Ca2+ ions are in 0.100 L of a 0.450 M solution of Ca3(PO4)2?



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



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?



23. If 880 mL of water is added to 125.0 mL of a 0.770 M HBrO4 solution what is the resulting molarity?



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?



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-?



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?



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?



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.



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)




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)




ch.4 - the major classes of chemical...

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