ch.4 - reactions in solution: aqueous chemistry in...

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CHEMISTRY - GILBERT 5E

CH.4 - REACTIONS IN SOLUTION: AQUEOUS CHEMISTRY IN NATURE

CONCEPT: MOLARITY

Molarity (M) can serve as the connection between the interconversion of ____________ to ____________ and vice versa.

For example, a 5.8 M NaCl solution really means __________________________ per __________________________.

Molarity =MolesSolute)(LitersSolution)(

A typical mixture consists of a smaller amount of one substance, the ________________, dissolved in a larger amount of

another substance, the __________________. Together they form a __________________.



PRACTICE: MOLARITY

EXAMPLE 1: 2.64 grams of an unknown compound was dissolved in water to yield

150 mL of solution. The concentration of the solution was 0.075 M. What was the

molecular weight of the substance?

EXAMPLE 2: A solution is prepared by dissolving 0.1408 mol calcium nitrate, Ca(NO3)2, in enough water to make 100.0 mL

of stock solution. If 20.0 mL of this solution is then mix with an additional 90 mL of deionized water, calculate the

concentration of the calcium nitrate solution.

PRACTICE 1: What is the molarity of calcium ions of a 650 mL solution containing 42.7 g of calcium phosphate?

PRACTICE 2: A solution with a final volume of 750.0 mL was prepared by dissolving 30.00 mL of benzene (C6H6, density =

0.8787 gmL

) in dichloromethane. Calculate the molarity of benzene in the solution.

M =MolesSolute)(LitersSolution)(



CONCEPT: NORMALITY

Another measurement for concentration usually encountered is normality (N), which represents the number of equivalents per liter of solution.

N =equivalents of soluteLiters of solution

equivalent (eq) = n ×moles

• An equivalent is the mass of a compound that can either donate or accept an __________ or __________.

An equivalent (eq) and n are both determined by the compound being used (acid or base) and if the compound is undergoing a redox reaction.

Acids

For an acid the number for n is based on the number of _________________________ present.

EXAMPLE 1: Determine the number of equivalents for each of the acids given.

a) 2.5 moles CH3COOH b) 133.4 g H3PO4

Bases

For a base the number for n is based on the number of _________________________ present.

EXAMPLE 2: Determine the number of equivalents for the following base.

a) 50.0 mL of 0.165 M Ca(OH)2

Redox Reactions

For a redox reaction the number for n is based on the number of ________________________ transferred.

EXAMPLE 3: Based on the given redox reaction determine the value for n.

MnO4– (aq) + H+ (aq) Mn2+ (aq) + H2O (l)



PRACTICE: NORMALITY EXAMPLE 1: What is the normality of a solution made by dissolving 325.1 g HNO3 in enough water to create a 750.0 mL solution?

EXAMPLE 2: Determine the equivalent weight of the following compounds.

a) Al(OH)3 b) H2CO3

EXAMPLE 3: What volume, in mL, of 50.0 g H2SO4 is needed to create a 0.300 N H2SO4 solution?

EXAMPLE 4: If a concentrated 3.25 M H3PO4 solution possesses a density of 1.350 g/mL, what is its normality?



CONCEPT: MOLARITY & CHEMICAL REACTIONS

Whenever we are provided given information in a reaction we use ___________________ to find any unknown information.

• In aqueous reactions, this given information is typically in units of __________________ or __________________ .

Entities means ______________________ , ______________________ or ______________________.

Volume of Given Moles of Given Moles of Unknown Volume of Unknown

Entities of Unknown

Grams of Unknown

Use this chart when given a chemical equation with the known quantity in either ________ or ________ of a compound or

element and asked to find the unknown quantity of another compound or element.

EXAMPLE: How many grams of sodium metal are needed to react with 38.74 mL of 0.275 M NaOH?

2 Na (s) + 2 H2O (l) H2 (g) + 2 NaOH (aq)



PRACTICE: MOLARITY & CHEMICAL REACTIONS

PRACTICE 1: How many milliliters of 0.325 M HCl are needed to react with 16.2 g

of magnesium metal?

2 HCl (aq) + Mg (s) MgCl2 + H2 (g)

PRACTICE 2: What is the molarity of a hydrobromic acid solution if it takes 34.12 mL of HBr to completely neutralize 82.56

mL of 0.156 M Ca(OH)2?

2 HBr (aq) + Ca(OH)2 (aq) CaBr2 (aq) + 2 H2O (l)

PRACTICE 3 (CHALLENGE): 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 30.0 mL of 0.100 M K2Cr2O7 to titrate a 25 mL Fe2+ solution, what is the molar concentration of Fe2+?

M =MolesSolute)(LitersSolution)(



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: 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: 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: BALANCING REDUCTION-OXIDATION REACTIONS

When balancing a redox reaction we balance them in terms of the number of electrons transferred between reactants. Balancing A Redox Reaction in Acidic Reactions: STEP 1: Write the equation into 2 half-reactions.

STEP 2: Balance elements that are not oxygen or hydrogen.

STEP 3: Balance Oxygens by adding _______.

STEP 4: Balance Hydrogens by adding _______.

STEP 5: Balance overall charge by adding e– to the more _______ side. Both reactions must have an equal number of e –.

STEP 6: Combine the half-reactions and cross out reaction intermediates.

Balancing A Redox Reaction in Basic Reactions: Follow Steps 1-6 from above.

STEP 7: Balance remaining H+ by adding an equal amount _______ ions to both sides of the chemical reaction.

EXAMPLE: Balance the following reaction in an acidic solution.

O2– + F2 O2 + F–



PRACTICE: BALANCING REDUCTION-OXIDATION REACTIONS

EXAMPLE 1: Balance the following reaction in an acidic solution. NO2– NO3– + NO

EXAMPLE 2: Balance the following reaction in a basic solution.

Cr2O72– + Hg Hg2+ + Cr3+



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 - reactions in solution: aqueous chemistry in...

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