Chapter Four: TYPES OF CHEMICAL REACTIONS AND SOLUTION STOICHIOMETRY

Chapter Four:

TYPES OF CHEMICAL REACTIONS AND

SOLUTION STOICHIOMETRY

Copyright © Houghton Mifflin Company. All rights reserved.Chapter 4 | Slide 2

Water, the Common Solvent

• One of the most important substances on earth

• Can dissolve many different substances

• A polar molecule

4.1


Dissolution of a Solid in a Liquid

4.1


Nature of Aqueous Solutions

• Solute – substance being dissolved

• Solvent – liquid water

• Electrolyte – substance that when dissolved in water produces a solution that can conduct electricity

4.2


Electrolytes

• Strong Electrolytes – conduct current very efficiently (bulb shines brightly)

• Weak Electrolytes – conduct only a small current (bulb glows dimly)

• Nonelectrolytes – no current flows (bulb remains unlit)

4.2


Electrolyte Behavior

4.2


Chemical Reactions of Solutions

• We must know:– Nature of the reaction– Amounts of chemicals present in the solutions

4.3


Molarity

• Molarity (M) = moles of solute per volume of solution in liters:

4.3

M

M

molaritymoles of soluteliters of solution

HClmoles of HClliters of solution

362


Concept Check

Which of the following solutions containsthe greatest number of ions?

a) 400.0 mL of 0.10 M NaCl.

b) 300.0 mL of 0.10 M CaCl2.

c) 200.0 mL of 0.10 M FeCl3.

d) 800.0 mL of 0.10 M sucrose.

4.3


Let’s Think About It

• Draw molecular level pictures showing each solution. Think about relative numbers of ions.

• How many moles of each ion are in each solution?

4.3


Notice

The solution with the greatest number of ions is not necessarily the one in which:

– the volume of the solution is the largest.– the formula unit has the greatest number of

ions.

4.3


Dilution

• The process of adding water to a stock solution to achieve the molarity desired for a particular solution.

• Dilution with water does not alter the numbers of moles of solute present.

• Moles of solute before dilution = moles of solute after dilution

M1V1 = M2V2

4.3


Types of Chemical Reactions

• Precipitation Reactions

• Acid-Base Reactions

• Oxidation-Reduction Reactions

4.4


Precipitation Reaction

• A double displacement reaction in which a solid forms and separates from the solution.– When ionic compounds dissolve in water, the

resulting solution contains the separated ions– Precipitate – the solid that forms

4.5


The Reaction of K2CrO4(aq) and Ba(NO3)2(aq)

4.5

• Ba2+(aq) + CrO42–(aq) → BaCrO4(s)

solid BaCrO4 formed


Precipitation of Silver Chloride

4.5


Precipitates

• Soluble – solid dissolves in solution; (aq) is used in reaction

• Insoluble – solid does not dissolve in solution; (s) is used in reaction

• Insoluble and slightly soluble are often used interchangeably

4.5


Simple Rules for Solubility

1. Most nitrate (NO3) salts are soluble.

2. Most alkali (group 1A) salts and NH4+ are soluble.

3. Most Cl, Br, and I salts are soluble (except Ag+, Pb2+, Hg2

2+).

4. Most sulfate salts are soluble (except BaSO4, PbSO4, Hg2SO4, CaSO4).

5. Most OH salts are only slightly soluble (NaOH, KOH are soluble, Ba(OH)2, Ca(OH)2 are marginally soluble).

6. Most S2, CO32, CrO4

2, PO43 salts are only slightly

soluble.

4.5


Describing Reactions in Solution

• Formula Equation (Molecular Equation)– Reactants and products as compounds

AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq)

• Complete Ionic Equation– All strong electrolytes shown as ions

Ag+(aq) + NO3(aq) + Na+(aq) + Cl(aq)

AgCl(s) + Na+(aq) + NO3(aq)

4.6


Describing Reactions in Solution

• Net Ionic Equation– Show only components that actually react

Ag+(aq) + Cl(aq) AgCl(s)

– Na+ and NO3 are spectator ions

4.6


Concept Check

You have two separate beakers with aqueous solutions, one with 4 “units” of potassium sulfate and one with 3 “units” of barium nitrate.

a) Draw molecular level diagrams of both solutions.

4.6/4.7


Concept Check


b) Draw a molecular level diagram of the mixture of the two solutions before a reaction has taken place.

4.6/4.7


Concept Check


c) Draw a molecular level diagram of the product and solution formed after the reaction has taken place.

4.6/4.7


Solution Stoichiometry Problems

• Identify the species present in the combined solution.

• Write the balanced net ionic equation.

• Calculate the moles of reactants.

• Determine limiting reactant.

• Calculate the moles of product(s).

• Convert to grams or other units.

4.7


Acid-Base Reactions (Brønsted-Lowry)

• Acid – proton donor

• Base – proton acceptor

• For a strong acid and base reaction:H+(aq) + OH–(aq) H2O(l)

4.8


Neutralization of a Strong Acid by a Strong Base

4.8


Key Titration Terms

• Titrant – solution of known concentration used in titration

• Analyte – substance being analyzed• Equivalence point – enough titrant added to

react exactly with the analyte• Endpoint – the indicator changes color so

you can tell the equivalence point has been reached

4.8


Performing Calculations for Acid-Base Reactions

1. List initial species and predict reaction.

2. Write balanced net ionic reaction.

3. Calculate moles of reactants.

4. Determine limiting reactant.

5. Calculate moles of required reactant or product.

6. Convert to grams or volume, as required.

4.8


Oxidation-Reduction Reactions (Redox Reactions)

• Reactions in which one or more electrons are transferred.

4.9


Reaction of Sodium and Chlorine

4.9


Rules for Assigning Oxidation States

1. Oxidation state of an atom in an element = 0

2. Oxidation state of monatomic element = charge

3. Oxygen = 2 in covalent compounds (except in peroxides where it = 1)

4. Hydrogen = +1 in covalent compounds

5. Fluorine = 1 in compounds

6. Sum of oxidation states = 0 in compounds

7. Sum of oxidation states = charge of the ion

4.9


Exercise

Find the oxidation states for each of the elements in each of the following compounds:

• K2Cr2O7

• CO32-

• MnO2

• PCl5• SF4

4.9


Redox Characteristics

• Transfer of Electrons

• Transfer may occur to form ions

• Oxidation – increase in oxidation state (loss of electrons); reducing agent

• Reduction – decrease in oxidation state (gain of electrons); oxidizing agent

4.9


Concept Check

Which of the following are oxidation-reduction reactions? Identify the oxidizing agent and the reducing agent.

a) Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g)

b) Cr2O72-(aq) + 2OH-(aq) 2CrO4

2-(aq) + H2O(l)

c) 2CuCl(aq) CuCl2(aq) + Cu(s)

4.9


Balancing Oxidation-Reduction Reactions

• Cr2O72-(aq) + SO3

2-(aq) Cr3+(aq) + SO42-(aq)

• How can we balance this equation?

• First Steps:– Separate into half-reactions– Balance elements except H and O

4.10


Method of Half Reactions

• Cr2O72-(aq) 2Cr3+(aq)

• SO32-(aq) SO4

2-(aq)

• How many electrons are involved in each half reaction?

4.10


Method of Half Reactions (continued)

• 6e- + Cr2O72-(aq) 2Cr3+(aq)

• SO32-(aq) + SO4

2-(aq) + 2e-

• How can we balance the oxygen atoms?

4.10



• 6e- + Cr2O72-(aq) Cr3+(aq) + 7H2O

• H2O +SO32-(aq) + SO4

2-(aq) + 2e-

• How can we balance the hydrogen atoms?

4.10



• This reaction occurs in an acidic solution.

• 14H+ + 6e- + Cr2O72- 2Cr3+ + 7H2O

• H2O +SO32- SO4

2- + 2e- + 2H+

• How can we balance the electrons?

4.10



• 14H+ + 6e- + Cr2O72- 2Cr3+ + 7H2O

• 3[H2O +SO32- SO4

2- + 2e- + 2H+]

• Final Balanced Equation:Cr2O7

2- + 3SO32- + 8H+ 2Cr3+ + 3SO4

2- + 4H2O

4.10


Exercise

Balance the following oxidation-reduction reaction that occurs in acidic solution.

Br-(aq) + MnO4-(aq) Br2(l)+ Mn2+(aq)

4.10


Half-Reaction Method – Balancing in Base

1. Balance as in acid.

2. Add OH that equals H+ ions (both sides!)

3. Form water by combining H+, OH. Cancel any water molecules that appear on both sides.

4. Check elements and charges for balance.

4.10

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Chapter Four: TYPES OF CHEMICAL REACTIONS AND SOLUTION STOICHIOMETRY