Chemistry A Molecular Approach, First Canadian Edition Nivaldo J. Tro Solution Manual.pdf (22)

Chapter 17. Gibbs Energy and Thermodynamics Teaching Tips

Suggestions and Examples 17.1Spontaneous and Nonspontaneous Processes

Misconceptions and Pitfalls A common misconception is that spontaneous processes occur at high rates.

This is often but certainly not always true, e.g., the combustion of octane at 25 oC in the absence of a spark or catalyst.

Mechanical potential energy and chemical potential energy typically increase through nonspontaneous processes. Ask the class for additional examples of nonspontaneous processes.  17.2Entropy and the Second Law of Thermodynamics

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Endothermic processes can occur spontaneously if the arrangement of particles changes from an orderly one to a more disorderly one. Even so, the spontaneity of these processes depends on temperature (e.g., the melting of an ice cube).

Entropy is a thermodynamic state function that increases with the number of energetically equivalent arrangements (microstates).

Spontaneity exothermic

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is not synonymous with . Endothermic processes also can be spontaneous. Exothermic processes can be nonspontaneous.

A beginner’s view of entropy relies on the analogy to disorder, but the better picture of entropy is energy equivalence. Spontaneous processes occur such that energy is evenly distributed, e.g., thermal equilibration.

There are many statements of the second law of thermodynamics. The one that is particularly relevant in this section is that for any spontaneous process, the entropy of the universe increases.

Another example to consider using is the probabilities associated with flipping multiple coins.

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The entropy of materials increases in transitioning from the solid phase to the liquid and then to the gas phase. There are more different ways to define the locations of the particles as a result of those phase changes, and there are more “places” for energy (vibration, rotation, translation).

Conceptual Connection 17.1 Entropy

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Chapter 17. Gibbs Energy and Thermodynamics Lecture Outline

Terms, Concepts, Relationships, Skills 17.3 Heat Transfer and Changes in the Entropy of the   o S = Ssys + Ssurr Surroundings Entropy change of the universe univ Temperature dependence of Ssurr energy dispersal per unit temperature o Entropy: o S = Hsys/T

Figures, Tables, and Solved Examples 

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Unnumbered figures: bar graphs of entropy changes for universe, surroundings, and system

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surr

Figure 17.5 Maxwell–Boltzmann Distribution of Kinetic Energy for Two Different Temperatures Example 17.2 Calculating Entropy Changes in the Surroundings

 Process occurring reversibly at constant temperature 17.4 Entropy for Phase Transitions  Changes  rS = rH/T

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17.5Gibbs Energy o G = TSuniv Gibbs energy (chemical G= H  TS potential) o o 

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Spontaneity be determined Spontaneous: G < 0with data for just the o Can system Nonspontaneous: G > 0 o Effect of T on spontaneity o 

Table 17.1 The Effect of rH, rS, T Figure 17.6 Gibbs Energy on Spontaneity Example 17.3 Computing Gibbs Spontaneity from rH and rS Energy Changes and Predicting

Four cases in Table 17.1

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276

Chapter 17. Gibbs Energy and Thermodynamics Teaching Tips

Suggestions and Examples

Misconceptions and Pitfalls

17.3 Heat Transfer and in by theeither Entropy of the Spontaneity canChanges be driven Ssys, Ssurr, or  both.

Surroundings Conceptual Connection 17.2 Entropy and Biological Systems

17.4Entropy Changes for Phase Transitions For constant pressure processes occurring reversibly the entropy can be calculated from the enthalpy of the system.   17.5 Gibbs Energy The greatchange advantage of system. using Gibbs energy is that it entropy for the  allows one to focus only on the enthalpy change and Since G = H  TS, the change in free energy can be composed of positive or negative enthalpy and terms, as well as the magnitude of T, determine entropy terms. The signs and magnitudes of those  Conceptual Connection 17.3 H, S, and G whether a reaction or process is spontaneous.

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A decrease in free energy  corresponds to a spontaneous process. watch the unit consistency. H is Students occasionally forget to often in kJ mol‐1, while S is often in  J.(molK)‐1.

For a reaction in which G > 0, the reverse reaction is spontaneous.

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Chapter 17. Gibbs Energy and Thermodynamics Lecture Outline

Terms, Concepts, Relationships, Skills Figures, Tables, and Solved Examples   rS° Entropy Changes in Chemical Reactions:Calculating ∆17.6  Figure Zero EntropySubstances at Values 17.7 (So) for Selected Table 17.2 Standard Molar Entropy o Unnumbered table: comparative So Standard states  o 298 K Gas: pressure of 1 bar o Liquid or solid: stable form at 1 bar (often  values for H2O(l) and H2O(g)  at 25 oC)  oS  Unnumbered Solution: concentration of 1 mol L‐1 Unnumbered table: table: comparative comparative So So Third law of thermodynamics values for noble gases = 0 for perfect crystal at 0 Kelvin o  Unnumbered table: comparative So Standard molar entropies values for diamond and graphite Tables  of values  Unnumbered table: comparative So  Phases values for Ar and NO gases  Molar mass   Allotropes values for CO and C2H4 gases  Molecular complexity table: comparative  Unnumbered Figure 17.8 “Places” for Energy inSo Dissolution o Gaseous NO SornpSo(products)nrS(reactants) Standard entropy change for a reaction table: comparative So  Unnumbered values for NO, NO2, and N2O4 gases  

values for KClO3(s) and KClO3(aq) Figure 17.9 Water Before and After Dissolution of AlCl3 Entropy Changes (∆rS°) Example 17.4 Calculating Standard

 Calculating ∆ rG° 17.7 Gibbs Energy Changes Chemical Reactions: Change in Gibbs freeinenergy from ∆rH° and ∆rS° o calculate ∆rH° from tabulated values o calculate ∆rS° from tabulated values o Go=HoTSo r

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r

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Example 17.6 Estimating the Standard Change in Gibbs Energy for a 25 Reaction Using Go=HoTSo at a Temperature Other than oC

r

Change in Gibbs free energy from ΔGof rG° =  np fG° (products) –  nr fG° (reactants)

r

only for reactions at 298 K o in Gibbs energy for stepwise reaction Change "Free" o energy o Energy available to do work Reversible reactions Irreversible reactions

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r

r

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Example 17.5 Calculating the Standard using Go=HoTSo r Free rEnergy for r a Reacting Change in

Energies of Formation (∆fG°) for Table 17.3 Standard Molar Gibbs Example 17.7 Calculating ∆rG° from Selected Substances at 298 K Standard Gibbs Energies of Formation Example 17.8 Calculating ∆rG° for Chemistry in Your Day: Making a a Nonspontaneous Process Spontaneous Stepwise Reaction Figure EnergyProcess Figure 17.10 17.11 Gibbs A Reversible 278

Chapter 17. Gibbs Energy and Thermodynamics Teaching Tips

Suggestions and Examples Calculating ∆ rS° 17.6 Entropy Changes in Chemical Reactions:

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Pitfalls Misconceptions The equationand for calculating ∆rS° is nearly identical to that for calculating ∆rH°.

The third law of thermodynamics defines the entropy of a perfect crystal at 0 Kelvin as zero, suggesting that for all other cases the standard molar entropy values are positive. These values are related to phase, molar mass, structure, and molecular complexity.

Standard entropy changes are determined from the standard entropies of the products and reactants; a larger table of these values (e.g., OH) are can be found innegative.

AppendixThis II. is because The standard entropies of aqueous ions aqueous standard entropiesmany are referenced to that of H+, defined to be 0.

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17.7 Gibbs Energy rG° Changes in Chemical Reactions:  Calculating ∆  The free energy change for a reaction can be calculated through several methods.

T r So is 298 K, the equationr Go=Ho r For reactions at temperatures other than

approximately valid. Molar enthalpy and entropy are functions of temperature, but we use the values tabulated at 298 K as a good approximation even for other  temperatures assuming a small variation over that temperature range. free energies of formation can be done only

The conditions for reversible reactions make them unrealistic; real reactions are irreversible and do not achieve the theoretical limit of available free energy.

Calculating the free energy change using for T = 298 K.

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279

Chapter 17. Gibbs Energy and Thermodynamics Lecture Outline

Terms, Concepts, Relationships, Skills rG° and ∆rG 17.8 Gibbs Energy Changes for Nonstandard States:The  RelationshipoBetween ∆ G=Go+RTln Q r r Nonstandard conditions o G o o G at standard conditions at equilibrium at other nonstandard conditions

Figures, Tables, and Solved Examples   

Nonstandard Conditions

rG°to the

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17.9 Gibbs Energy and Equilibrium: Relating ∆ EquilibriumoConstant G (K) Freeoenergy change at equilibrium Go=RTln K r = 0 at equilibrium  Temperature dependence of K o o

Unnumbered figure: photo of a wet floor and molecular models of H2O Example 17.9Gibbs Calculating Under Figure 17.12 Energy∆rG° versus Pressure for Water

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Ho1 So r  + r R T R o K  H 1  1    ln 2 = r R T2 T1 K1 

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ln K=

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Figure 17.13 Gibbs Energy and the Equilibrium Constant andConstant ∆rG° Example 17.10 The Equilibrium Example 17.11 Using the van’t Hoff Equation to Determine the Standard Enthalpy and Entropy of Reaction Example 17.12 Using the Two‐Point Form of the van’t Hoff Equation to Predict the Equilibrium Constant at a Given Temperature

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Chapter 17. Gibbs Energy and Thermodynamics Teaching Tips

Suggestions and Examples Relationship Between ∆ rG° and ∆rG 17.8 Gibbs Energy Changes for Nonstandard States:The

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Misconceptions and Pitfalls (G < 0), the free energy change increases Beginning with a spontaneous reaction until equilibrium is reached (G = 0).

For nonstandard conditions, the free energy quotient, Q. change can be calculated using the reaction The free energy change for a reaction at equilibrium is 0.

Conceptual Connection 17.4 Gibbs Energy Changes and Le Châ telier’s Principle 

rG°to the 17.9 Gibbs Energy and Equilibrium: Relating ∆ Equilibrium Constant (K) be determined by calculating ∆rG at the The equilibrium constant for any reaction can  temperature of interest.

This chapter has a relatively large number of equations, and many students try to memorize all of them rather than working with a few to derive the rest.

The natural logarithm of the equilibrium temperature. A plot of ln K vs. 1/T can be used constant is linearly related to the inverse of the y to find ∆rH° (from the slope) and ∆rS° (from the ‐intercept).

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281

Chapter 17. Gibbs Energy and Thermodynamics Additional Problem for Computing Gibbs Energy Changes and Predicting Spontaneity from rH and rS (Example 17.3)

Consider the combustion reaction for calcium: 2 Ca(s) + O2(g)  2 CaO(s) H = 1270 kJ, S = 365 J K-1

a) Calculate G at 25 °C and determine whether the reaction is spontaneous. b) Determine at what temperature (if any) the reaction is spontaneous.

a) Use G = H – TS to calculate  G from values of H and S (in the same units).

T = 273 + 25 = 298 K

G = H – TS = (1270 × 103 J) – 298 K (365 J K-1) = 1.16 × 106 J = 1.16 × 103 kJ = 1160 kJ

The reaction is spontaneous.

b) Since S is negative, G will become less negative with increasing temperature. Use G = H – TS to calculate the temperature at which G changes from positive to negative (i.e., G = 0).

G = H – TS 0 = (1270 × 103 J) – T (365 J K-1) T (365 J K-1) = 1270 × 103 J T = 1270 × 103 J / 365 J K-1 = 3479 K

The reaction will remain spontaneous until 3479 K. Above that temperature, G will be positive.

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Chapter 17. Gibbs Energy and Thermodynamics Additional Problem for Calculating the Standard Change in Gibbs Energy for a Reaction using rG° = rH° – TrS° (Example 17.5)

Consider the reaction: NO(g) + ½O2(g)  NO2(g)

Calculate rG° at 25 °C and determine whether the reaction is spontaneous.

Look up standard enthalpies and entropies fH°(kJ mol)-1

Begin by looking up the standard enthalpy of formation and standard entropy for each reactant and product.

-1 S° (in J molK)

NO(g)

91.3

210.8

O2(g)

0

205.2

NO2(g)

33.2

240.1

Calculate rH° Use Equation 6.15.

rH° =  np fH° (products) –  nr fH° (reactants) = 33.2 – [½(0) + 91.3] = 58.1 kJ

Calculate rS° rS° =  np S° (products) – nr S° (reactants)

Use Equation 17.16.

= 240.1 – [½(205.2) + 210.8] = 73.3 J K-1

Calculate rG° T = 273 + 25 = 298 K

Use Equation 17.17.

rG° = rH° – TrS° = 58.1 × 103 J – 298 K (73.3 JK-1) = 3.62 × 104 J = 3.62 × 101 kJ = 36.2 kJ The reaction is spontaneous at this temperature. .

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Chapter 17. Gibbs Energy and Thermodynamics Additional Problem for Calculating rG° from Standard Gibbs Energies of Formation (Example 17.7)

Consider the reaction: CH3Cl(g) + Cl2(g)  CH2Cl2(l) + HCl(g)

Use the standard free energies of formation to determine rG° for this reaction at 25°C.

Look up standard free energies of formation Begin by looking up standard free energies of formation for each reactant and product. Remember that the standard free energy of formation of a pure element is zero.

fG° (kJ mol-1) CH3Cl(g)

60.2 0

Cl2(g) CH2Cl2(g)

63.2

HCl(g)

95.3

Calculate rG° Use Equation 17.18.

rG° =  np fG°(products)   nr fG° (reactants) = [63.2 + (95.3)] – [60.2 + (0)] = 98.3 kJ mol-1

The reaction is spontaneous.

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284

Chapter 17. Gibbs Energy and Thermodynamics Additional Problem for The Equilibrium Constant and rG° (Example 17.10)

Use the tabulated free energies of formation to calculate the equilibrium constant for the reaction at 298 K: H2(g) + Cl2(g)

2HCl(g)

Look up standard free energies of formation G f° (kJ mol) -1

Begin by looking up the standard free energies of formation for each reactant and product.

H2(g)

0

Cl2(g)

0

HCl(g)

95.3

Calculate rG° Use Equation 17.18.

rG° =  np fG° (products) –  nr fG° (reactants) = [2(95.3)] – [0 + 0] = 190.6 kJ

Calculate K Calculate K from rG° by solving equation 17.20.

rG° = RT ln K

ln K  

-Go r RT

190.6103 Jmol1

8.314 J(molK)1298 K 

K  e76.9  2.501033

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285

Chapter 18. Electrochemistry Student Objectives

18.1Pulling the Plug on the Power Grid   Know that electrical power comes from a power grid, a distribution system for electricity.

Know that fuel cell technology can replace the power source for an automobile or a home.

Know that the oxidation and reduction reactions for the combination of hydrogen and oxygen to make water can power a fuel cell.

Define oxidation and reduction.

18.2Balancing Oxidation–Reduction Equations   Balance aqueous redox reactions in acidic solution by using the half‐reaction method.

Balance aqueous redox reactions in basic solution by using the half‐reaction method.

18.3Voltaic (or Galvanic) Cells: Generating Electricity from Spontaneous Chemical Reactions        

Understand that the components in redox reactions be separated two compartments or cells Define electrochemical cell, voltaic (galvanic) cell,can electrolytic cell,into half‐cell, and electrode. and that the energy released by these reactions can be used to do electrical work.

Define potential difference and know its unit.

Understand the rate of flowing electrons as current, and know the unit for current.

Define electromotive force, cell potential, and standard cell potential.

Know and understand that electrochemical reactions are a combination of half‐cell reactions, oxidation occurring at the anode, and reduction occurring at the cathode.

Know and understand the components of an electrochemical cell, including the half‐cells, electrodes, and salt bridge.

Understand the direction of electron flow and the potential of a voltaic cell.

Use cell notation to represent electrochemical cells.

18.4Standard Electrode Potentials   Know that the standard hydrogen electrode has an electrochemical potential of 0.00 V. Understand standard potentials and use them to generate potentials for electrochemical cells.  Calculate standard potentials for electrochemical cells from standard electrode potentials of the half‐  reactions.

Calculate standard potentials and predict the direction of spontaneity reactions.

Use standard to predict metals will dissolve inenergy, acid.for redox Calculate and potentials interconvert betweenwhether the standard change in Gibbs G°, the standard potential 18.5Cell Potential, Gibbs Energy, and the Equilibrium Constant for an electrochemical cell, E°, and the equilibrium constant, K.

18.6Cell Potential and Concentration  Define concentration cell and understand how a concentration cell works.  Use the Nernst equation to calculate the cell potential under nonstandard conditions.

Know and understand nerve cells as a biological example of a concentration cell.

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286

Chapter 18. Electrochemistry

18.7Batteries: Using Chemistry to Generate Electricity  

Know some characteristics of dry‐cell, lead–acid storage, nickel–cadmium, nickel–metal hydride, and lithium–ion batteries.

Know difference between batteries and fuel Know the andprimary understand how the hydrogen–oxygen fuel cellcells. and the fuel‐cell breathalyzer work.

18.8Electrolysis: Driving Nonspontaneous Chemical Reactions with Electricity   Know and understand how an electrolytic cell works.

Predict the products of some common electrolysis reactions.

Calculate and interconvert between current, time, and amount for electrolysis reactions.

18.9Corrosion: Undesirable Redox Reactions   Know that corrosion generally refers to the oxidation reactions of metals.

Know the steps by which iron rusts.

Know some strategies for preventing corrosion.

Section Summaries

 Outline Lecture  Terms, Relationships, Skills Figures,Concepts, Tables, and Solved Examples  Teaching Tips  Suggestions andand Examples Misconceptions Pitfalls

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287

Chapter 18. Electrochemistry Lecture Outline

Figures, Tables, and Solved Examples

Terms, Concepts, Relationships, Skills 18.1Pulling the Plug on the Power Grid o Fuel cells o Definition o Uses Reaction

 Intro figure: illustration of an MP3 player powered by a fuel cell  Unnumbered figure: photo of a fuel cell capable of providing power for a house Unnumbered figure: hydrogen fuel cell bus

–Reduction Equations  18.2 Balancing o Oxidation Redox reactions o Oxidation: loss of electrons, increase in oxidation number  Reduction: gain of electrons, decrease o number Balancing redox reactions o in oxidation o Oxidation numbers o Separation of half‐reactions o Mass balance o Charge balance Electron balance Recombination of half‐reactions

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Example 18.1 Half‐Reaction Method of Balancing Aqueous Redox Equations in Acidic Solution Example 18.2 Balancing Redox Reactions Occurring in Basic Solution

 18.3 Voltaic (or Galvanic) Cells: Generating Electricity o from Spontaneous PropertiesChemical of voltaicReactions cells  Electrical current: ampere (A) o 1 A = 1 C s‐1   Potential difference: 1V = 1 J C‐1 volt (V) o Components of voltaic cells o Half‐cells Electrodes   Anode: oxidation Cathode: reduction o  Inert electrodes Salt bridge Cell notation

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Figure 18.1 A Spontaneous Oxidation– Reduction Reaction Figure 18.2 A Voltaic Cell Figure 18.3 An Analogy for Electrical Current Figure 18.4 Inert Platinum Electrode

288

Chapter 18. Electrochemistry Teaching Tips

Suggestions and Examples 18.1Pulling the Plug on the Power Grid

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Misconceptions and Pitfalls The combustion of hydrogen with oxygen occurs rapidly and often explosively. The hydrogen–oxygen fuel cell controls the kinetics but harnesses the same amount of energy.

This example of a fuel cell is based on the reaction of hydrogen with oxygen. This reaction can be demonstrated as an explosion; the same reaction in a fuel cell releases the same amount of energy but in a controlled way.

Batteries are ubiquitous devices and should motivate interest in this subject. Ask for applications of batteries in everyone’s daily lives.

–Reduction Equations 18.2 Balancing Oxidation  Identifying the individual half‐reactions is  generally straightforward. 

 Half‐reactions that are not balanced initially with respect to mass (e.g., oxygen on one side but not the other side) are initially confusing to some students.

Particularly in redox reactions, both mass and charge must be balanced.

There are other slightly different workflows for balancing redox reactions.

Working through a few examples in class helps students get the hang of balancing these reactions. 

18.3 Voltaic (or Galvanic) Cells: Generating Electricity  from Spontaneous Chemical Reactions The voltage or potential between the two half‐cells drives the current. 

An oxidation or reduction cannot occur in the absence of its counterpart. It can appear that they are independent since each half‐  cell reaction appears to occur in an individual compartment.

The presence of inert electrodes in cell notation is initially confusing to some students.

Demonstrations of redox reactions and voltaic cells are simple and quite satisfying to students.

Electrochemical cell notation summarizes the critical components of an electrochemical cell in a compact form.

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289

Chapter 18. Electrochemistry Lecture Outline

Terms, Concepts, Relationships, Skills 18.4Standard Electrode Potentials o Standard hydrogen electrode o o  Reaction Half‐cell potential Standard electrode Eo Eo  Eocellpotentials cathode anode Measured versus reference like SHE o  Table  of values Sign and value o Relative oxidizing strength o Predicting spontaneous Metals dissolving in aciddirection

Figures, Tables, and Solved Examples  

Figure 18.5 An Analogy for Electrode  Potential Figure 18.6 The Standard Hydrogen  Electrode 

Figure 18.7 Measuring Standard Electrode Potential

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Figure 18.8 Electrochemical Cell with SHE as the Anode

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 18.5 CellGibbs Potential, Gibbs Energy, and the Equilibrium energy change, rGo o Faraday’s constant, F = 96,485 C/mol e Constant o  rGo= nFEo  0.0592 V RT o E=o constant ln K= log K Equilibrium n nF

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Table 18.1 Standard Electrode Potentials at 25 oC Example 18.3 Calculating Standard Potentials for Electrochemical Cells from Standard Electrode Potentials—1 Example 18.4 Calculating Standard Potentials for Electrochemical Cells from Standard Electrode Potentials—2 Figure 18.9 Mn/Ni2+ Electrochemical Example 18.5 Predicting Spontaneous Cell Redox Reactions and Sketching Electrochemical Cells Figure 18.10 Electrochemical Cell Zn with HCl(aq) Composed of Mg2+/Mg and Fe2+/Fe Unnumbered figure: photo of reaction of Electrodes Figure 18.10 Electrochemical Cell Composed of Mg2+/Mg and Cr3+/Cr Unnumbered relating Gibbs Example 18.6figure: Relating rGo and Eocell energy to equilibrium constant Electrodes

Example 18.7 Relating Eocell and K Unnumbered figure: relating Gibbs energy to equilibrium constant with formulas

290

Chapter 18. Electrochemistry Teaching Tips

Suggestions and Examples 18.4Standard Electrode Potentials 

The electrode potential of SHE has been set at zero. Other potentials are measured

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relatively.

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Misconceptions and Pitfalls

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The standard potential for an electrochemical cell determines the direction of spontaneity.

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Conceptual Connection 18.1 Selective Oxidation Conceptual Connection 18.2 Metals Dissolving in Acids

Half‐reactions with negative electrode potentials can still be spontaneous depending on the potential of the other half‐ reaction.

Stronger oxidizing agents have higher electrode potentials; stronger reducing strong itselfelectrode is reduced. agents oxidizing have moreagent negative potentials. Students sometimes forget that a When balancing redox reactions, one must often multiply a half‐reaction by a factor such as 2 or 3 to balance the electrons. Students are tempted to multiply the standard half‐cell potential by that same factor.

 18.5 Cell Potential, Gibbs Energy, and the Equilibrium

Students often try to memorize all these equations and their variations, but they all can be derived easily using relationships the students already know from thermodynamics.

Constant The standard cell potential, the standard  Gibbs energy change, and the equilibrium 

constant can be used to predict the spontaneity of a reaction under standard conditions.

Conceptual Connection 18.3 Periodic Trends and the Direction of Spontaneity for Redox Reactions

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291

Chapter 18. Electrochemistry Lecture Outline

Terms, Concepts, Relationships, Skills 18.6Cell Potential and Concentration o Reaction quotient, Q Nonstandard conditions o o 

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Figures, Tables, and Solved Examples   

Concentration cells o Identical electrodes, different concentrations o Concentration equilibration Human o nerve cells o Ion channels in membranes Sodium–potassium pump Membrane potential

Figure 18.12 PotentialEcell and Under Example 18.8Cell Calculating Concentration

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Nonstandard Conditions Example 18.9Cu/Cu2+ Determining the Ksp of Cell a Figure 18.12 Concentration Slightly Soluble Salt Using a Concentration



Cell

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Chemistry and Medicine: Concentration Cells in Human Nerve Cells

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Figure 18.14 Concentration Changes in Nerve Cells Figure 18.15 Potential Changes Across the Nerve Cell Membrane

  18.7 Batteries: o Using Chemistry to Generate Electricity Battery types  Dry‐cell Standard: graphite and zinc  electrodes; moist MnO2 and

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NH4Cl

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Alkaline: graphite and zinc  electrodes; moist MnO2 and Lead–acid storage  Rechargeable o KOH  Six cells for total of 12 V  rechargeable Other  Nickel–cadmium Nickel–metal hydride Lithium–ion o Fuelocell o Definition Example: hydrogen–oxygen cell Example: fuel‐cell breathalyzer o

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.

Figure 18.16 Dry‐Cell Batteries Figure 18.17 Lead–Acid Storage Battery Unnumbered figure: photo of rechargeable batteries and charger Table 18.2 Energy Density and Overcharge Tolerance of Several Rechargeable Batteries Figure 18.18 Lithium–Ion Battery Figure 18.19 Hydrogen–Oxygen Fuel Cell Chemistry in Mesoporous Your Day: The Lithium Battery Figure 18.20 Carbon

292

Chapter 18. Electrochemistry Teaching Tips

Misconceptions and Pitfalls the equation for G from thermodynamics.

Nernst equation is can easily using  The An electrochemical cell be derived established

Suggestions and Examples 18.6Cell Potential and Concentration



 

The potential of a cell can be determined from the standard potential and the conditions, i.e., the cell Q expression. concentration ratios at nonstandard Concentration cells continue to operate until

using two identical electrodes; if the concentrations are different, the cell potential will not be zero.

the concentrations become equal. This is the same concept as thermal equilibration and heat flow.

A nerve cell is an excellent example of a Conceptual Connection 18.4 Relating Q, K, concentration , and E°cellcell with biological relevance.

E cell

18.7Batteries: Using Chemistry to Generate Electricity



Batteries are necessary for keeping electronics operating, so there may be considerable interest. A type of battery comes in a specific voltage; the size determines the amount of current available. Larger potentials represent multiple cells connected in series.

Fuel cells also generate a specific voltage.

Some devices are capable of high throughput of the chemical reactants and can generate considerable power.

.

293

Chapter 18. Electrochemistry Lecture Outline

Terms, Concepts, Relationships, Skills

Figures, Tables, and Solved Examples   Figure 18.21 Electrolysis of Water  Figure 18.22 Silver Plating  Figure 18.23 Voltaic versus Electrolytic Cells 18.24 Electrolysis ofof Molten NaCl  Figure Unnumbered figure: photo battery‐

18.8 Electrolysis: Driving Nonspontaneous Chemical o Reactivity Reactions with Electricity o Spontaneous: voltaic cell  Nonspontaneous: electrolytic cell  Design o Examples  oo Electrolysis of water Silver plating Predicting products o o Pure molten salts o Mixtures of cations and anions  Aqueous solutions Aqueous NaCl and overvoltage Stoichiometry

 powered electrolysis of water  Figure 18.25 Electrolysis of Aqueous NaI Figure 18.26 Electrolysis of Aqueous NaCl:  The Effect of Overpotential 

Example 18.10 Predicting the Products of Electrolysis Reactions Figure 18.27 Electrolytic Cell for Copper Plating

Example 18.11 Stoichiometry of Electrolysis     

18.9Corrosion: Undesirable Redox Reactions o Corrosion definition o Oxidation of metals o Reduction of oxygen  o Examples of rust Preventing corrosion o o Paint and covering Sacrificial electrodes Galvanizing



.

Unnumbered figure: photo of rusted ship Unnumbered figure: photo of aluminum pipe Unnumbered figure: scratch in paint Figure 18.28 Corrosion of Iron: Rusting Unnumbered figure: photo of galvanized nails Unnumbered figure: illustration of sacrificial anode

294

Chapter 18. Electrochemistry Teaching Tips

Suggestions and Examples

Misconceptions and Pitfalls

18.8 Electrolysis: Driving Nonspontaneous Chemical

Students sometimes have trouble predicting the products of electrolysis reactions in aqueous solution because they neglect the possibility of decomposing the water itself.

Electrolytic cells use electrical input from an Reactions with Electricity external source to drive a nonspontaneous electrochemical reaction. The applied potential determines what kinds of reactions can occur; the current and time determine the moles of material converted.

18.9Corrosion: Undesirable Redox Reactions





Corrosion typically occurs when metals are converted to their metal oxides. Rust, a metal oxide of iron, does not have the desired structural properties that iron and steel have.

Corrosion is prevented by physically blocking the oxidation reaction or by providing a sacrificial electrode.

Conceptual Connection 18.5 Sacrificial Anodes

.

295

Chapter 18. Electrochemistry Use tabulated half-cell potentials to calculate Additional Problem for Calculating Standard Potentials for standard Electrochemical the standard cell potential for the reaction in an Cells from Standard Electrode Potentials electrochemical cell at 25 °C: (Examples 18.3 and 18.4)

Zn2+(aq) + H2O2(aq)  Zn(s) + O2(g) + 2 H+(aq)

Separate into half-cells Begin by separating the reaction into oxidation and reduction half-reactions. [In this case, the zinc cations form zinc metal, a reduction reaction.]

Oxidation: H2O2(aq) + 2 e  O2(g) + 2 H+(aq)

Reduction: Zn2+(aq)  Zn(s) + 2 e-

Assign Potentials Look up the standard half-cell potentials for each half-reaction. For the oxidation,

Ox: 2 e + H2O2(aq) 

remember to change the sign of Eored to obtain Red: Zn2+(aq) Eoox. Add the half-cell reactions together.

O2(g) + 2 H+(aq)

 Zn(s) + 2 e

_________________________________________

Zn2+(aq) + H2O2(aq)  Zn(s) + O2(g) + 2 H+(aq)

Eoox = Eored = 0.70 V Eored = 0.76 V

Solve Add the half-cell potentials together to obtain the overall standard potential.

red Eocl = Eoo + Eo el x

= 0.70 V + (0.76 V) = 1.46 V

.

296

Chapter 18. Electrochemistry Additional Problem for Relating o Ecell and K (Example 18.7)

Use the tabulated half-cell potentials to calculate K for the oxidation of nickel by chlorine:

Clg 2() + Ni(s)  2 Cl-(aq) + Ni2+ (aq)

2 + Ni()s 2 Cl(aq) + Ni2+(aq) Given Cl(g)

Sort You are given the redox reaction and asked to find K.

Find K

Conceptual Plan

Strategize Use the tabulated values of the half-cell potentials

red Eoox, Eo

to calculate Eocell. Then use Equation 18.7 to



Eocell

red Eo + Eo = Eo cell ox

calculate K from Eocell.

Eocell  o Ecell =

K 0.0592 V log K n

Solve Break the reaction up into oxidation and reduction half-reactions and find the standard half-cell

Solution

potentials for each. Sum the values to find Eocell.

Red: Cl2(g) + 2 e  2 Cl-(aq) _________________________________

Ox:

Ni(s)  2 e + Ni2+(aq)

Cl2(g) + Ni(s)  2 Cl-(aq) + Ni2+(aq)

Eo ox =Eored = +0.23 V Eored = 1.36 V

red Eo + Eo = Eo cell ox

= +0.23 V + 1.36 = +1.59 V .

297

Chapter 18. Electrochemistry

Compute K from Eocell. The value of n (the number of 0.0592 V Eo cell = log K n moles of electrons) corresponds to the number of 0.0592 V electrons that are cancelled in the half-reactions.

1.59 V =

2 2 (1.59) log K = 0.0592 K = 1053.72

log K

= 5.251053

The reaction is spontaneous and strongly favours the products. Chlorine is a powerful oxidizing agent.

Check

.

298

Chapter 18. Electrochemistry Additional Problem for Calculating Ecell Under Nonstandard Conditions (Example 18.8)

An electrochemical cell is based on two halfreactions:

Oxidation: Fe(s)  Fe2+(aq, 0.010 mol L-1) + 2 e

Reduction: Br() 2l + 2  e 2 Br-(

aq, 1.0 mol L-1)

Compute the cell potential.

Sort

Given [Fe2+] = 0.010 mol L-1; [Br-] = 1.0 mol L-1

You are given the half-reactions for a redox reaction and the concentrations of the aqueous reactants and products. You are asked to find the cell potential.

Strategize

Find Ecell

Conceptual Plan Use the tabulated values of the half-cell potentials

red  Eoox, Eo

cell Eo

to calculate Eocell. Then use Equation 18.10 to calculate Ecell.

2+ [Br]-  Ecell cell, [Fe], Eo

ell

Ec

o = Ecell 

Solve

Solution

Write the oxidation and reduction half-reactions and sum them. Find the standard half-cell potentials of

Ox:

0.0592 V log Q n

Fe(s)  2 e + Fe2+(aq)

Red: Br2(l) + 2 e  2 Br(aq)_____________________________________

each and sum them to find Eocell.

- + Fe(aq)2+ Br2(l) + Fe(s)  2 Br(aq)

o ox = Ered = +0.45 V Eo

Eored = 1.09 V

.

299

Chapter 18. Electrochemistry Eo cell =Eoox + Eored = +0.45 V + 1.09 = +1.54 V

Compute Ecell from Eocell. The value of n is the

0.0592 V log Q n 2 0.0592 V 2+ = Eo cell  log [Fe ] [Br] n 0.0592 = 1.54 V  log (0.010)(1.0) 2 = 1.54 V  0.0296 (2) = 1.60 V

Ecell = Eocell 

number of moles of electrons. Determine Q based on the overall balanced equation and the given concentrations of reactants and products.

The reaction is spontaneous. Bromine is a powerful oxidizing agent.

Check

.

300

Chapter 18. Electrochemistry Additional Problem for Stoichiometry of Electrolysis (Example 18.11)

Aluminum can be plated out of a solution containing Al3+ according to the half-reaction: Al3+(aq) + 3 e  Al(s) What mass of aluminum (in grams) will be plated by the flow of 50 A of current 10 hours?

Given 3 mol e = 1 mol Al

Sort You are given the half-reaction for the plating of aluminum, which shows the stoichiometry between the moles of aluminum and the moles of electrons. You are also given the current and time. You are asked to find the mass of aluminum that will be deposited in that time.

50 A 10 hours

Find g Al

Strategize You need to find the amount of aluminum, which is related stoichiometrically to the number of electrons that have flowed through the cell.

Begin with time in hours and convert to seconds.

Given the current and time, find the number of coulombs. Then use Faraday’s constant to calculate the number of moles of electrons and the stoichiometry of the reaction to find the number of moles of aluminum. Finally, use the molar mass of aluminum to convert to the mass of aluminum.

Conceptual Plan

Solve

Solution

Follow the conceptual plan, cancelling units to arrive at the mass of aluminum.

= 10 hr 

hr  s  C  mol e-  mol Al  g Al 60 min60 s 50 C s 1 hr 1 min

1 mol e 96,485 C

1 mol Al 3 mol e

26.98 g Al 1 mol Al

3600 s 50 C 1 mol e 1 mol Al 26.98 g Al     1 mol Al hr s 96,485 C 3 mol e

= 168 g Al

Check

The units (g Al) are correct. The magnitude of the answer seems to make sense. (This amount is about 14 drink cans.) It takes a considerable amount of electricity to make the aluminum. In fact, aluminum is made electrolytically and is energy-intensive and expensive for this reason. This is why there is incentive to recycle aluminum cans.

.

301

Chapter 19. Radioactivity and Nuclear Chemistry Student Objectives

19.1Medical Isotopes 

Know that radioactivity, the natural emission of subatomic particles from atoms, can be used in Define radioactivity. medical applications like radiology, the practice of diagnostic imaging.

19.2The Discovery of Radioactivity 

Know that radioactive elements emit energy that is capable of penetrating some types of matter, including tissues of the body.

Know that radioactive substances can expose photographic plates, a phenomenon that was discovered by Becquerel and further studied by Marie and Pierre Curie.

Define nuclide and recall the conventions for writing nuclear symbols.  19.3 Types of Radioactivity   Define ionizing power and penetrating power.  Know that an alpha particle is the same as a helium‐4 nucleus.  Know that a beta particle is the same as an electron and is emitted when a neutron changes to a  proton.   

Know beta decay, the is mass number of photon the radioactive nucleus staysinthe same.

Know that that in a gamma particle a high‐energy and is often emitted conjunction with the other types of radiation.

Know that a positron is the antiparticle of an electron and that it is emitted when a proton changes to a neutron.

Know that electron capture occurs when an electron from an inner orbital is combined with a nuclear proton to form a nuclear neutron.

Write nuclear symbols for all the subatomic particles, and write balanced reactions for all the types of radioactive decay.

19.4The Valley of Stability: Predicting the Type of Radioactivity  Know that atomic nuclei are held together by the strong force.

N/Z ratio.  and Predict the relative stability and type of radioactive decay of isotopes based on their atomic number  Know stability ofdecay an isotope on itsthe absolute number neutrons.

Know that that the a radioactive seriesdepends characterizes steps by whichofanprotons isotopeand transforms into other smaller and lighter atoms.

19.5Detecting Radioactivity 

Know that radioactivity exposure in humans is often measured with film‐badge dosimeters, photographic film in a light‐proof case that responds to other forms of radiation besides light.



Know that the Geiger‐Mü ller counter is a device that can count particles emitted by radioactive nuclei.

Know that a scintillation counter is another device that detects low levels of radiation.

.

302

Chapter 19. Radioactivity and Nuclear Chemistry

19.6The Kinetics of Radioactive Decay and Radiometric Dating 

Know that natural radioactivity is a part of our environment, unrelated to human activity, and understand that it is a consequence of the instability of all nuclei beyond atomic number 83.



Know and understand that all radioactive nuclei decay via first‐order kinetics and that the half‐life of a radioactive substance is independent of the amount of the substance.

   

Know that radon is a colourless, odourless gas that is emitted from the ground via the natural radioactive decay of uranium.

Know that the natural, radioactive isotope carbon‐14 decays with a known half‐life and that its concentration in dead organic material can be used to measure the age of the material through radiocarbon dating.

Calculate and predict age ofofuranium‐238 a sample tobased onwithin its igneous concentration of becarbon‐14.

Know and understand thatthe the ratio lead‐206 rocks can used to measure the age of the rocks.

19.7TheCalculate Discoveryand of Fission: Atomic predictThe the age Bomb of a and rockNuclear basedPower on its ratio of uranium‐238 to lead‐206.  Define nuclear fission.  Understand how the fission of an isotope like uranium‐235 can lead to a self‐amplifying chain  reaction.

Know that the first nuclear weapons, developed by the United States in the 1940s, were fission weapons.

Know and understand that nuclear reactors take the energy from a small, controlled fission reaction and convert the energy to heat to produce electrical power through conventional methods, e.g., steam‐driven turbines.

19.8Converting Mass to Energy: Mass Defect and Nuclear Binding Energy  Define mass defect and nuclear binding energy.  Know and understand why nuclear reactions release far more energy than chemical reactions.

Calculate the mass defect and nuclear binding energy for a radioactive isotope.

Define nuclear fusion.

19.9Nuclear Fusion: The Power of the Sun  Know that attempts at human‐made nuclear fusion are still at experimental stages and are being pursued because fusion does not involve radioactivity or radioactive byproducts.

Define transmutation.  Nuclear Transmutation and Transuranium Elements 19.10  Know and understand that linear accelerators and cyclotrons are devices used to transmutate elements for experimentation and to produce medically important isotopes for diagnosis and treatment of disease.

.

303

Chapter 19. Radioactivity and Nuclear Chemistry

 The Effects of Radiation on Life 19.11 

Know that radioactivity in the body can do damage and is classified into three groups: acute radiation damage, increased cancer risk, and genetic effects.



Know that exposure to radioactivity can come from several sources including natural radiation, diagnostic medical procedures, and consumer products.

Know and understand that the dose and time extent of a radiation exposure give rise to a variety of adverse effects, including death.  Radioactivity 19.12 in Medicine and Other some Applications Define radiotracer and recognize examples.   Know that radioactive isotopes are used for radiotherapy in medicine.

Know that radioactivity is used in a variety of applications, including the sterilization of implanted medical devices, the destruction of bacteria and parasites that spoil foods, and the sterilization of insects that carry disease.

.

304

Chapter 19. Radioactivity and Nuclear Chemistry Section Summaries

 Outline Lecture  Terms, Relationships, Skills Figures,Concepts, Tables, and Solved Examples  Teaching Tips  Suggestions andand Examples Misconceptions Pitfalls

.

305

Chapter 19. Radioactivity and Nuclear Chemistry Lecture Outline

Terms, Concepts, Relationships, Skills 19.1 Medical Isotopes  Radioactivity Medical isotopes

 Tables, and Solved Examples Figures,  Intro figure: illustration of decaying atom, scan of infected appendix Unnumbered figure: Chalk River Laboratories

 19.2 The Discovery of Radioactivity  Phosphorescence: glow‐in‐the‐dark o Radioactivity history o o Antoine Henri‐Becquerel (1896) Marie Curie Pierre Sklodowska Curie



Unnumbered figure: photograph of phosphorescence Figure 19.1 The Discovery of Radioactivity Unnumbered figure: photo of Marie Curie

 19.3 Types of Radioactivity o 2 protons + 2 neutrons: 42He nucleus Alpha decay o o  High ionizing power, low penetrating power Electron: 01e Beta decay o o Conversion of neutron to proton and electron Medium ionizing power and penetrating  power o Photon:  Gamma o ray emission o o No charge or mass o Accompanies other decays  Low ionizing power, high penetrating power Positron emission 



Unnumbered figure: photo of glow‐in‐  Example 19.1 Writing Nuclear the‐dark watch Figure 19.2 Equations for AlphaAlpha Decay Decay Figure 19.3 Beta Decay   Figure 19.4 Positron Emission Example 19.2 Writing Nuclear Equations for Beta Decay, Positron Emission, and Electron Capture Table 19.1 Modes of Radioactive Decay 

Positron: 01e

o



Antiparticle of electron Conversion of proton to neutron and o positron Electron capture Inner‐orbital electron + proton  neutron .

306

Chapter 19. Radioactivity and Nuclear Chemistry Teaching Tips

Suggestions and Examples 19.1Medical Isotopes

Misconceptions and Pitfalls  X‐rays and gamma rays are different only by their source. X‐rays come from high‐energy electronic transitions, while gamma rays come from nuclear transitions.

Many students likely have personal experience with some type of radiological procedure. X‐rays are not a form of radioactivity but are high‐energy electromagnetic photons similar to gamma rays. Ask  the students to suggest other radiological techniques.

Various radiological techniques collect information by penetrating tissues. Only a few involve radioactivity or nuclear  chemistry.

19.2 The Discovery of Radioactivity Becquerel, the Curies, and other early experimenters did not understand the power of the radioactivity that they observed by exposing photographic film. It is likely that Marie Curie shortened her life via the  biological damage she sustained by conducting Radium Girls: Women and Industrial Health Reform by Claudia Clark recounts women who painted radium experiments without adequate protection. and other radioactive isotopes on watch and clock faces. They often licked their paint brushes to maintain a "sharp" tip, thereby ingesting the radioactive material.

19.3 Types of Radioactivity 



The different types of radioactivity have different properties with respect to their detection and to their biological hazards.

 Radioactivity for the most part cannot be seen, but it can stimulate phosphorescence. Film and the Geiger‐Mü ller counter are commonly used.

 Only certain atoms and their isotopes are naturally radioactive, and each decays by a specific pathway.

Conceptual Connection 19.1 Alpha and Beta Decay

.

307

Chapter 19. Radioactivity and Nuclear Chemistry Lecture Outline

Figures, Tables, and Solved Examples  Figure 19.5 Stable and Unstable Nuclei Example 19.3 Predicting the Type of  Radioactive Decay

Terms, Concepts, Relationships, Skills 19.4 The Valley of Stability: Predicting the Type of   Ratio of neutrons to protons: N/Z Radioactivity Strong force o o  Up to Z = 83 o Valley stability /Z tooofhigh: beta decay N/Z too low: positron emission or N  o electron capture Magic numbers Absolute numbers of protons, neutrons

 Table 19.2 Number of Stable Nuclides with Even and Odd Numbers of Nucleons Figure 19.6 The Uranium‐238 Radioactive Decay Series

 

19.5Detecting Radioactivity   Film‐badge dosimeter Geiger‐Mü ller counter Scintillation counter

Figure 19.7 Film‐Badge Dosimeter Figure 19.8 Geiger‐Mü ller Counter

 19.6 The Kinetics of Radioactive Decay and o Radiometric Radioactive oDating nuclei: first‐order decay to number of nuclei 0.693 o Rate proportional Half‐life: t1/2=depends on isotope Rate constant k 



     

Radon o o o Colourless, odourless gas Naturally Nt=ktoccurs from uranium in soil ln Integrated rate law N0

 

 o Applications 14 C: t1/2 = 5730 yr Radiocarbon dating  o

Table 19.3 Selected Nuclides and Their Half‐ Lives Figure 19.9 The Decay of Radon‐220 Chemistry in Your Day: Uranium Isotopes and the CANDU Reactor Unnumbered figure: CANDU nuclear power plant in Qinshan Example 19.4 Radioactive Decay Kinetics Unnumbered figure: photo of Dead Sea Scrolls Unnumbered figure: photo of bristlecone pine tree Example 19.5 Radiocarbon Dating Example 19.6 Uranium/Lead Dating to Estimate the Age of a Rock

Dating method for organic  238 U: t1/2 <= 50,000 4.5 × 109 materials yr yr Uranium/lead dating

.

308

Chapter 19. Radioactivity and Nuclear Chemistry Teaching Tips

Misconceptions and Pitfalls regardless of N/Z, are radioactive. atomic number 83,  All Thenuclei stableabove N/Z changes as atomic number

Suggestions and Examples 19.4 The Valley of Stability: Predicting the Type of  Protons are in close proximity (~1015 m) in the nucleus and repel each other. The strong Radioactivity force or nuclear force holds the nucleus together. The number of neutrons is important to the stability of isotopes, as demonstrated by the neutron‐to‐proton ratio.  Stable nuclides have specific patterns of

increases. A 1:1 ratio is stable for low numbers of protons but not for relatively high numbers.

numbers of neutrons and protons referred to as magic numbers. Even–even is most common, while odd–odd is rare.

 19.5Detecting Radioactivity 

Radioactivity detection is simple and still makes use of a historic system: the exposure



of film.

Radioactivity is invisible.

The Geiger‐Mü ller counter is a portable device for counting decay events.

The scintillation counter is used to measure radioactivity in low levels in liquid samples.

19.6 The Kinetics of Radioactive Decay and 

 Radiocarbon dating can be used to determine the age of a material that contained a biological source of carbon, i.e., one that exchanged carbon with the environment. As long as carbon is being exchanged, the relative amount of carbon‐14 remains constant.

Radiometric Dating The first‐order rate law, integrated rate law, and half‐life equation can be used for all radioactive decays. The rate constants for  these decays vary considerably from fractions of a second to billions of years. 

 

Radon is a radioactive gas that must be trapped to be analyzed. It is typically absorbed onto carbon or charcoal, a material with a large surface area.

Radiocarbon dating is a powerful technique but is limited to times less than 50,000 years; very short times also are not as accurately measured.

Uranium/lead dating is used for very long times and volcanic rock samples.

Conceptual Connection 19.2 Half‐Life

.

309

Chapter 19. Radioactivity and Nuclear Chemistry Lecture Outline

Terms, Concepts, Relationships, Skills 19.7 The Discovery of Fission: The Atomic Bomb and  o 235 Nuclear Power Uranium U   Can undergo fission and create a self‐sustaining chain reaction  Only 1% of naturally occurring o uranium History o o Enrico Fermi Lise Meitner, Fritz Strassmann, OttoJ.Hahn Manhattan Project: Albert Einstein, R.  Oppenheimer o Applications o Atomic weapons: Hiroshima, Nagasaki Nuclear reactor: electricity generation 19.8 Converting Mass to Energy: Mass Defect and Nuclear    

Einstein's equation: E = mc2 Mass defect Nuclear binding energy

Figures, Tables, and Solved Examples  

Unnumbered figure: photo of Lise  Meitner Unnumbered figure: illustration of  uranium‐235 fission Figure 19.10 A Self‐Amplifying Chain Reaction   

Unnumbered figure: on July 16, 1945, in the New Mexico desert, the world’s first atomic bomb Table 19.4 Selected Countries that Use Nuclear Power Figure 19.11 Afigure: Nuclear Reactor Unnumbered photo of Chernobyl reactor

 

Example 19.7 Mass Defect and Nuclear Binding Energy Figure 19.12 Nuclear Binding Energy per Nucleon

o Binding o energy curve Fusion vs. fission Predicting reactions

 

 19.9 Nuclear Fusion: The Power of the Sun Higho temperature requirement of fusion Examples o o Hydrogen bombs Energy source of stars Production of electricity

Figure 19.13 A Nuclear Fusion Reaction Figure 19.14 Tokamak Fusion Reactor

  Nuclear Transmutation and Transuranium Elements 19.10  Transmutation o Facilities o Stanford Linear Accelerator Fermilab

.

  

Unnumbered figure: photo of Joliot‐ Curies Figure 19.15 The Linear Accelerator Unnumbered figure: photos of SLAC and Fermi lab Figure 19.16 The Cyclotron 310

Chapter 19. Radioactivity and Nuclear Chemistry Teaching Tips

Suggestions and Examples 19.7 The Discovery of Fission: The Atomic Bomb and Nuclear  Power Nuclear fission is not a natural phenomenon; it requires energetic neutrons bombarding target atoms.

Furthermore, the target atom is chosen so that an  increase in mass is sufficient to trigger the

Misconceptions and Pitfalls  Uranium‐238 does not undergo fission under normal circumstances, but it can be made to undergo fission in thermonuclear devices.

fragmentation. 

Controlled nuclear fission provides useful energy in nuclear reactors, while uncontrolled fission is used for nuclear weapons.

Nuclear reactors and the production of nuclear weapons create nuclear waste, a potential environmental problem. More than 80 percent of France’s electricity comes from nuclear power, and the French maintain a reactor capable of accommodating the waste.

19.8 Converting Mass to Energy: Mass Defect and Nuclear   Energy Binding Unlike chemical reactions, mass is not conserved in nuclear processes. 

The large energies from small mass changes make nuclear reactions far more powerful than chemical reactions.

The binding energy per nucleon vs. mass number enables prediction of atoms susceptible to fusion or 19.9 Nuclear Fusion: The Power of the Sun fission.   

 The mass of a nucleus is not the simple sum of the masses of the protons and neutrons. Some of  that mass is converted to energy, referred to as the binding energy.

Keeping careful track of units is key for calculations in this section.

Energy changes can be calculated per atom or per mole of atoms. It is useful to remind students that a  joule is a derived SI unit.

Fusion has been accomplished in thermonuclear weapons in which a fission reaction released the energy and high temperatures required to initiate the fusion reaction.

Fusion involves the lightest isotopes, typically hydrogen, deuterium, and tritium.

Fusion provides more energy per mass of fuel and does not produce radioactive byproducts.

Extremely high temperatures are required for fusion becauseTransmutation of the electrostatic repulsion between the 19.10 Nuclear and Transuranium Elements  positively charged nuclei.



Transmutations create one element from another in devices like linear accelerators and cyclotrons. They are an important source of knowledge for specific isotopes with medical uses.

 Cyclotrons are not necessarily huge devices. Some are nearly small enough to fit on a tabletop.

Conceptual Connection 19.3 Nuclear Transformations .

311

Chapter 19. Radioactivity and Nuclear Chemistry Lecture Outline

Figures, Tables, and Solved Examples  Table Weighting Factors Table 19.5 19.6 Radiation Effects of Instantaneous

Terms, Concepts, Relationships, Skills  The Effects of Radiation on Life 19.11 Types o o o Acute radiation damage Increased cancer risk  o Genetic defects Measuring exposure o o Becquerel (Bq) o Gray (Gy) o  Sievert (Sv) Radiation weighting factor (WR) Sources o o Natural radiation Diagnostic medical procedures  Consumer products Exposure effects

Radiation Exposure

    

 Radioactivity in Medicine and Other Applications 19.12 o Diagnosis in medicine o o  Radiotracers Positron emission tomography (PET)  Radiotherapy in medicine o Gamma radiation Other o applications o Kill bacteria on medical devices Kill bacteria and parasites in food Sterilize insects

.

Figure 19.17 A Bone Scan Table 19.7 Common Radiotracers Figure 19.18 A PET Scan Figure 19.19 Radiotherapy for Cancer Figure 19.20 Irradiation of Food

312

Chapter 19. Radioactivity and Nuclear Chemistry Teaching Tips

Misconceptions and Pitfalls

Suggestions and Examples  The Effects of Radiation on Life 19.11



Exposure to radiation can cause acute radiation damage (rare exposure), increased cancer risk (more common), and genetic defects.

Conceptual Connection 19.4 Radiation Exposure

 19.12 Radioactivity in Medicine and Other 

There is common misunderstanding about food irradiation. The food is treated with gamma rays and not with the radioactive materials themselves.

Applications There are substantial beneficial uses of radioactivity: diagnosis and radiotherapy in medicine, sterilization of foods, control of insects.

.

313

Chapter 19. Radioactivity and Nuclear Chemistry Additional Problem for Radioactive Decay Kinetics (Example 19.4)

Fluorine-18 is a radioactive isotope that emits positrons and is used in PET scans (positron emission tomography). It has a halflife of 110 minutes. If a patient is given a 1.5 ng sample of 18F, how much of the sample in molecules persists after a day?

Sort You are given the initial mass of F-18 in a sample and asked to find the mass after 1 day.

Given mass(F-18, initial) = 1.5 ng t1/2 = 110 min t = 1 day

Find number of molecules(F-18, final)

Strategize Use the integrated rate law (Equation 19.3) to solve this problem. First determine the rate constant k using the expression for the half-life (Equation 19.1).

Conceptual Plan

The integrated rate law gives the ratio of Nt/N0, the ratio of numbers of atoms. Use the mass, the molar mass, and Avogadro’s number to find the initial number of atoms.

massinit



t1/2

0.693

t1 =

k

2



N0 g F-18 1 mol F-18  6.0221023 atoms F-18 18 g F-18 1 mol F-18

N0, k, t N ln N0

Solve Follow the conceptual plan.

k



Nt

t =  kt

Solution 0.693 k

t1 = 2

110min = k=

0.693

k 0.693 = 6.30103 min1 110 min

1 g F-18 mol F-18 6.0221023 atoms F-18 1.5 ng F-18 1109 ng F-18 118 g F-18  1 mol F-18 = 5.021013 atoms F-18

.

314

Chapter 19. Radioactivity and Nuclear Chemistry 24 hr 60 min = 1440 min 1 day  1 day  1 hr

Solve the integrated rate law by substituting the values.

N ln t = kt N0 6.30103 Nt = ( ln )(1440 min) min 5.021013 = 9.07 ln Nt 5.02  1013 Nt = e9.07 5.02  1013 4

  atoms Nt = 1 .1510(5.0210) = 5.78109 atoms

Check

13

The units (atoms) are correct. The magnitude of the answer makes sense since it is much smaller than the starting value.

.

315

Chapter 19. Radioactivity and Nuclear Chemistry Additional Problem for Radiocarbon Dating (Example 19.5)

A papyrus scroll was buried with King Tutankhamen 3330 years ago. The half-life of carbon-14 is 5730 years.

The papyrus scroll sample had an initial decay of 15.3 disintegrations per minute per gram carbon. What is the current decay rate of the sample? (The decay rate is proportional to the amount of C-14 present.)

Given rate0 = 15.3 dis / (ming C)

Sort You are given the initial rate of decay for the papyrus sample and its current age. You are asked to find the current decay rate.

t1/2 = 5730 years t = 3330 years

Find ratet

Strategize Use the expression for the half-life (Equation 19.1) to find the rate constant.

Conceptual Plan 

1/2

t

k 0.693

t1 =

k

2

Use the value of the rate constant, initial rate, and age of the sample in the integrated rate law (Equation 19.4).

rateinit, k, t  rate ln rate t =

ratet kt

0

Solve

Solution

Follow the conceptual plan to solve the problem.

0.693 k

t1 = 2

5730 yr = k=

ln Use the rate constant, time, and initial rate in the integrated rate expression.

0.693 k 0.693 = 1.21104 yr1 5730 yr

rate

t  kt rate0

1.21104  ln  ln 15.3  yr  rate

1

rate 1  0.402 15.3 rate1  e0.402  0.669 15.3 ratet  0.669(15.3) 10.2

dis ming C

dis ming C

.

316

Chapter 19. Radioactivity and Nuclear Chemistry Check

The units of the answer are correct since they are the same as the initial rate. The magnitude makes sense since the age of the sample is less than one half-life.

Greater than half of the original sample still exists (i.e., the current rate of decay is greater than half the original decay).

.

317

Chapter 20. Organic Chemistry I: Structures Student Objectives

20.1Fragrances and Odours 

Know molecules called odourants bind with olfactory receptors in the nose, leading to the Definethat organic chemistry. sensation of smell.

20.2Carbon: Why It Is Unique  

Know thatorigin carbon tendency form four covalent bonds, including the possibility of double Know the ofhas the the terms organictoand inorganic. and triple bonds.

Know that carbon has the tendency to catenate or bond to itself to form chain, branched, and ring structures.

20.3Hydrocarbons 

Know that hydrocarbons contain only carbon and hydrogen and comprise alkanes, alkenes, alkynes, and aromatic compounds.



Know the generic formulas, the structures, and some examples of alkanes, alkenes, alkynes, and aromatics.

Draw and interpret drawings of hydrocarbon structures using structural formulas, condensed structural formulas, carbon skeleton formulas, ball‐and‐stick models, and space‐filling models.

20.4Functional Groups 

Know that functional groups are specific atoms bonded to each other that can be inserted onto a hydrocarbon.

Know and identify the common functional groups: alcohol, ether, aldehyde, ketone, carboxylic acid, ester, and amine.

20.5Constitutional Isomerism Know how to write constitutional isomers for hydrocarbons and for molecules containing functional groups.

20.6Stereoisomerism I: Conformational Isomerism   Understand that isomerism is due to rotation around a single bond.

Describe alkane conformations.

Identify ring conformations of isomer, cycloalkanes.

Define stereoisomer, optical enantiomer, and chiral.

20.7Stereoisomerism II: Configurational Isomerism   Identify chiral molecules.  Describe the experiment demonstrating plane‐polarized light and how an optically active sample can  Assign rotate the of polarization.

E/Zplane configurations to alkenes.  Establish the difference between geometric isomers and stereoisomers.

Know how to assign absolute configurations to chiral carbons.

.

318

Chapter 20. Organic Chemistry I: Structures

20.8Structure Determination  Use the index of hydrogen deficiency to establish the number of pi bonds or rings in a molecule.

Know the different spectroscopic methods of structure determination: IR, NMR. Section Summaries

 Outline Lecture  Terms, Relationships, Skills Figures,Concepts, Tables, and Solved Examples  Teaching Tips  Suggestions andand Examples Misconceptions Pitfalls

.

319

Chapter 20. Organic Chemistry I: Structures Lecture Outline

Terms, Concepts, Relationships, Skills 20.1Fragrances and Odours   Odour and smell Odourants o Organic o chemistry Chemistry of carbon Foods, drugs, petroleum

 Figures, Tables, and Solved Examples Intro figure: illustration of cologne bottle, organic molecules Unnumbered figure: photo of skunk, molecular models of thiols



products, pesticides

  20.2 Carbon: Why It Is Unique o Bonding o o Formation of four bonds Double and triple bonds Catenation

Unnumbered figures: Lewis structures, molecular models of methane, ethane, ethene, ethyne Unnumbered figure: Lewis structures of propane, isobutane, cyclohexane The Nature offigure: Science: Vitalism and the Wö Perceived Unnumbered photo of Freidrich hler Difference Between Organic and Inorganic

  

  20.3 Hydrocarbons o Hydrocarbons: only C and H o o Alkanes Alkenes o o Alkynes  Aromatics Representations o Structural formula o Condensed structural o formula o Carbon skeleton formula Ball‐and‐stick model Space‐filling model

Unnumbered figures: structural formula, condensed structural formula, carbon skeleton formula, ball‐and‐ stick model, space‐filling model of butane, isobutane Unnumbered figure: different representations of butane Unnumbered figures: structural formula, condensed structural formula, carbon skeleton formula, ball‐and‐ stick model, space‐filling model of propene, propyne Example 20.1 Writing Structural Formulas for Hydrocarbons Table 20.1 Alkanes, Alkenes, Alkynes Unnumbered figure: orbital overlap illustrating sigma bonds in methane Unnumbered figure: Lewis structure and valence bond model of ethane Unnumbered figure: acetylene torch Unnumbered figure: Lewis structure and valence Unnumbered figure: p orbital overlap in 1,3‐ bond model of ethyne butadiene Unnumbered figure: Lewis structure and ball‐and‐ stick model of benzene Figure 20.1 Polycyclic Aromatic Compounds

 

        

.

320

Chapter 20. Organic Chemistry I: Structures Teaching Tips

Misconceptions and Pitfalls

Suggestions and Examples 20.1Fragrances and Odours sensitivity; the isopropyl and t‐butyl thiols in The sense smell is fascinating. natural gasofcan be smelled at partsOne perfactor billionis concentration.

  20.2 Carbon: Why It Is Unique  background for the term organic chemistry.

The history of chemistry gives and the The tendencies of carbon to catenate form four bonds lead to large numbers of organic compounds, ~20 million characterized.

Other elements (e.g., silicon) catenate, but not to the extent of carbon.

 20.3Hydrocarbons  Hydrocarbons provide a good way to illustrate the idea of structural isomerism. 

Free rotation about single bonds can sometimes mislead students into thinking that slightly different representations of the same molecule are different molecules altogether. Making molecular models and  rotating about the bonds helps alleviate this confusion.

Reminding students of the octet rule will help them draw better structures for organic compounds.

Encourage students to make their own models of hydrocarbons. Toothpicks, coffee stirrers, paper, tape, marshmallows, and ingenuity are all one needs for models.

Browser plug‐ins for Chime, Rasmol, JMol, and other related visualization programs are free. Websites have libraries of files of molecules.

.

321

Chapter 20. Organic Chemistry I: Structures Lecture Outline

Terms, Concepts, Relationships, Skills 20.4Functional Groups  Specific o collections or groups of atoms Examples o o Alcohol o Ether o Aldehyde o Ketone acid o Carboxylic Amine Ester

Figures, Tables, and Solved Examples   Table 20.2 Properties of Methane and Methanol  Unnumbered figure: carbon tetrachloride, zero  dipole moment  Unnumbered figure: Teflon tape  Table 20.3 Bonds in Organic Halides  Unnumbered figure: quaternary ammonium salt  Unnumbered figure: fresh fish  Figure 20.2 The Structure of Aniline Unnumbered figure: Lewis structure and space‐  filling model of ethanol  Unnumbered figure: Petro‐Canada gas price  Unnumbered figure: primary, secondary, and  tertiary alcohols Table 20.4 Solubilities of Alcoholsmodel of Unnumbered figure: space‐filling  dimethyl ether

 

Figure 20.3 Charge Density Plots of the Carbonyl Group Unnumbered figure: Lewis structure and space‐ filling model of methanal Unnumbered figure: bananas (pentyl ethanoate) Table 20.4 Acyl Functional Groups

  20.5 Constitutional Isomerism Constitutional isomers in different classes of compounds

Unnumbered figure of propan‐1‐ol and propan‐2‐ ol

.

322

Chapter 20. Organic Chemistry I: Structures Teaching Tips

Suggestions andGroups Examples 20.4 Functional 

Misconceptions and Pitfalls Students often have a hard time differentiating between certain functional groups such as aldehydes and ketones.

Emphasize the importance of functional groups in giving organic compounds special physical and chemical properties.

20.5Constitutional Isomerism



Emphasize that structural isomers have the same molecular formula but a different carbon atom skeleton



Emphasize the possibility of structural isomerism among molecules with different functional groups Conceptual Connection 20.1 Constitutional Isomers

.

323

Chapter 20. Organic Chemistry I: Structures Lecture Outline

Terms, Concepts, Relationships, Skills 20.6 Stereoisomerism I: Conformational  Isomerism Conformational isomerism—rotation o about o a single bond projection formulas o Newman Eclipsed and staggered o o

  

conformations

Gauche conformation Steric strain Ring conformations Angle strain 20.7 Stereoisomerism II: Configurational Isomerism  Cis‐trans  isomerism in alkenes o Chirality o o Enantiomers o Optical activity Chemical behaviour in chiral enviroment Absolute configurations 

Figures, Tables, and Solved Examples Unnumbered figure: three‐dimensional drawing of carbon atom Figure 20.5 Energy versus Rotation Angle for Ethane Figure 20.6 Energies of Butane Conformations Unnumbered figure: chair conformation of cyclohexane

      

   

20.8 Structure Determination   Index o of hydrogen deficiency Spectroscopic methods of structure determination o  IR spectroscopy NMR spectroscopy  Chemical shift Proton Carbon NMR NMR

.

of cis and trans dichloroethane Unnumbered figure: space‐filling models Table 20.6 Physical Properties of cis‐ and 1,2‐Dicloroethene trans‐ Example 20.2 Assigning the Configurations of Alkenes Figure 20.7: Mirror Images Figure 20.8 Rotation of Plane‐Polarized Light Unnumbered figure: sodium rubidium d‐ tartrate Unnumbered figure: illustration of tetrahedral molecule and its mirror image Unnumbered figures: illustration of enantiomers in chiral environments Figure 20.9 Absolute Configurations of Example 20.3 Assigning R,S Butan‐2‐ol Enantiomers Configurations Chemistry and Medicine: Anesthetics and Alcohol

     

Example 20.4 Index of Hydrogen Deficiency Table 20.7 Infrared Absorptions of Selected Bonds Figure 20.10 Infrared Absorption Spectrum of Butan‐2‐one Unnumbered figure: tetramethylsilane Figure 20.11 13C‐NMR Spectrum of Cyclohexanone Example 20.5 Determining the Structure of an Unknown Organic Compound 324

Chapter 20. Organic Chemistry I: Structures Teaching Tips

Suggestions and Examples 20.6Stereoisomerism I: Conformational Isomerism 

Misconceptions and Pitfalls Students often have a hard time understanding the difference between the conformers obtained as a result of a carbon– carbon single bond rotation and a constitutional isomer.

Allow students to visualize the single bond rotation and the fact that no single bonds break during this rotation



In the chair conformation of cyclohexane, introduce the concept of the axial up and axial down position, for better visualization.

It is also useful in the case of the equatorial

positions.

20.7 Stereoisomerism II: Configurational Isomerism  Conceptual Connection 20.2 Staggered Conformations in Rings 

 Emphasize thatonassigning absolute track chirality paper and that R and S configuration to chiral carbon is a way to configurations to not relate directly to d and l configurations.

Emphasize the difference between configuration and confirmation from a three‐dimensional perspective.

Conceptual Connection 20.3 Chirality

  20.8 Structure Determination Emphasize the importance of the index of hydrogen deficiency in determining possible structures for organic compounds.

.

325

Chapter 20. Organic Chemistry I: Structures Additional Problem for Naming Alkanes

Name the alkane:

Longest Chain Count the longest continuous chain of carbons. There are six carbons. The path straight across only includes five carbons in a row.

Substituents or Groups This compound has two substituents, both of which are named methyl.

Numbering Try numbering from the end at right. That would give “3,5-dimethyl.”

Numbering from the left gives smaller numbers: “2,4-dimethyl” General Form 2,4-dimethylhexane

Substituent number(s) + substituent name(s) + base name

This molecule can also be drawn as

.

326

Chapter 20. Organic Chemistry I: Structures

Name the alkene:

Additional Problem for Naming Alkenes and Alkynes

Longest Chain Count the longest continuous chain of carbons that contains the double bond. There are six carbons. The path straight across only includes five carbons in a row.

Counting from the left, the double bond is between the third and fourth; counting from the right gives the same outcome.

Substituents or Groups This compound has two substituents, both of which are named methyl.

Numbering Try numbering from the end at right. That would give “3,5-dimethyl.”

Numbering from the left gives smaller numbers: “2,4dimethyl.”

General Form Substituent number(s) + substituent name(s) + base name

2,4-dimethyl-3-hexene

.

327

Chapter 21. Organic Chemistry II: Reactions Student Objectives

21.1Discovering New Drugs Understand ideathe of multistep Know how tothe write expressionsynthesis. for Ka for organic acids.

21.2Organic Acids and Bases  Know how to calculate pKa for organic acids.  Know how to write the conjugate bases of organic acids.  Know the different factors affecting acidity: inductive effect, hybridization, resonance.  Familiarize yourself with mechanisms in organic chemistry from an acid–base point of view and the curved arrows involved in movement of electrons in these mechanisms.  Understand acids and bases and their use as chemical reagents.

21.3Oxidation and Reduction 

Know the rules for assigning oxidation states to atoms in organic molecules based on electronegativity of atoms connected to the atom to be determined.

Understand organic redox reactions and the transformation of functional groups using redox reactions.

21.4Nucleophilic Substitution Reactions at Saturated Carbon   Know the definition of nucleophilic and electrophilic species.  Know the meaning of the leaving group and substrate.  Know and understand the steps and reaction energy diagram of the SN1 mechanism.

Know and understand the steps and reaction energy diagram of the SN2 mechanism.  Know the factors affecting nucleophilic reactions: nucleophile type, the leaving group type, and substrate structure.

Understand carbocation stability.  21.5 Elimination Reactions   Know the different types of elimination reactions: dehydration and dehydrohalogenation.  Understand the concept of 1,2 elimination.  Understand Zaitsev’s rule for product formation in elimination reactions.

Know and understand the steps and reaction energy diagram of the E1 mechanism.

Know and understand the steps and reaction energy diagram of the E2 mechanism.

21.6Electrophilic Additions to Alkenes  Understand mechanism of hydrohalogenation an example of electrophilic addition to alkenes.

Understand the other addition reactions and be able toaspredict their product.

21.7Nucleophilic Additions to Aldehydes and Ketones  

Understand nucleophilic addition of alcohols to aldehydes and ketones, as well as the formation of acetals and ketals.

Understand mechanism reagents of acetal and formation.

Understand the organomettalic and ketal the Grignard reaction. .

328

Chapter 21. Organic Chemistry II: Reactions

21.8Nucleophilic Substitutions to Acyl Compounds   

Understand the general mechanism of nucleophilic substitution of acyl compounds and know how to apply it in the case of Fischer esterification and hydrolysis of esters.

Know the steps of the esterification mechanism.

Understand the relative reactivity within the carboxylic acid derivative series.

Understand acyl group interconversion.

21.9Electropholic Aromatic Substitution   Understand the preference of electrophilic substitution versus addition for aromatic compounds.

Know the steps of the general mechanism of electrophilic aromatic substitution.

Know electrophilic aromatic reactions and their products.

Definecommon monomers and polymers.  Polymerization 21.10  Know the difference between step growth polymerization and addition polymerization.

Section Summaries

 Outline Lecture  Terms, Relationships, Skills Figures,Concepts, Tables, and Solved Examples  Teaching Tips  Suggestions andand Examples Misconceptions Pitfalls

.

329

Chapter 21. Organic Chemistry II: Reactions Lecture Outline

Terms, Concepts, Relationships, Skills 21.1Discovering New Drugs

 Figures, Tables, and Solved Examples Intro figure: montelukast sodium (Singulair) chemical structure and pills, child with an inhaling mask

Multistep synthesis of organic compounds and their importance

21.2 Organic Acids and Bases    Relative acidity of organic acids Inductive o effects and acidity Resonance effects on acidity  Acidic hydrogens on o hydrocarbons Mechanisms in organic chemistry o Curved arrows Acids and bases as reagents

 

Table 21.1 pKa Values for Selected Acids Unnumbered figure: inductive effect in hydrogen



Unnumbered figure: ethyne and its pKa value cyanide Unnumbered figure: acetic acid and trifluoroacetic acid and their inductive effects Figure 21.1 Electrostatic Potential Surfaces Unnumbered figure: delocalization in phenoxide ion and benzene molecule Unnumbered figure: carbanion and pentan‐2,4‐dione Example 21.1 Mechanisms for Acid–Base Reactions Example 21.2 Selecting an Appropriate Acid–Base

      

Unnumbered figure: selected Brønsted–Lowry acids Reagent and bases in order of decreasing basicity and acidity 21.3

Oxidation and Reduction   Hydrogenation as redox reaction  Interconversion of functional groups through redox reactions

 

Example 21.3 Oxidation States in Organic Molecules



Table 21.2 Reagents for Selected Oxidations and Reductions Example 21.4 Choosing Reagents for Oxidation and Reductions Chemistry in Your Day: Hydrogen and the Oil Sands



Choosing reagents for redox reactions

.

330

Chapter 21. Organic Chemistry II: Reactions Teaching Tips

Misconceptions and Pitfalls

Suggestions and Examples 21.1 Discovering New Drugs Identifying certain functional groups in common drugs might help students understand the significance of drug discoveries. 

21.2 Organic Acids and Bases  

In writing acid–base reactions, students tend to have the arrow indicate that the acidic hydrogen “departs” instead of the arrow indicating that the acidic hydrogen is being “picked up” by the base.

Relate the strength of organic acids to the magnitude of their equilibrium constant.

Use curved arrows, electron flow, to write both organic and inorganic acid–base reactions.

21.3 Oxidation and Reduction

 



Establish a connection between the oxidation state of carbon and the different functional groups.

Review the rules of assigning oxidation states for inorganic ions and elements and emphasize the difference of assigning oxidation states to atoms in organic molecules.

Conceptual Connection 21.1 Oxidation and Reduction

.

331

Chapter 21. Organic Chemistry II: Reactions Lecture Outline

Terms, Concepts, Relationships, Skills 21.4 Nucleophilic Substitution Reactions at  Saturated Carbon Identification of substrate,  nucleophile, electrophile, and leaving  

group SN1 mechanism SN2 mechanism o Factors substitutions o affecting nucleophilic Carbocation stability Steric hindrance in substrate

 Figures, Tables, and Solved Examples  Unnumbered figure: partial charges in a polar C—X  bond Figure 21.2 Energy Profile for SN1 mechanism  Unnumbered figure: the two steps of the SN1 mechanism  Unnumbered figure: SN1 mechanism on 2‐chloro‐2‐ methylpropane   Unnumbered figure: SN1mechanism on a chiral carbon  Unnumbered figure: SN2 mechanism Unnumbered figure: SN2 mechanism on   bromomethane Figure 21.3 Energy Profile for SN2 Mechanism  Unnumbered figure: wind inverting an umbrella  Table 21.3 Selected Nucleophiles and Their Relative Effectiveness Unnumbered figure: increasing carbocation stability  Unnumbered figure: steric hindrance in bromoethane, 2‐bromopropane, 2‐bromo‐2‐

21.5 Elimination Reactions 

  

o 1,2 eliminations o Dehydration and dehydrohalogenation

methylpropane  Example 21.5 Predicting Nucleophilic Substitutions   Unnumbered figure: position 1,2 for elimination Unnumbered figure: stepwise E1 mechanism Unnumbered figure: E2 mechanism

Zaitsev’s rule E1 mechanism E2 mechanism o Substitution vs. elimination Strong bases favour elimination

.

332

Chapter 21. Organic Chemistry II: Reactions Teaching Tips

Suggestions and Examples 21.4 Nucleophilic Substitution Reactions at Saturated  Carbon Emphasize that through nucleophilic  substitution a variety of different functional undergoing these reactions has to be sp3 groups can the be introduced incarbon molecules.

Emphasize fact that the atom  hybridized. 

Misconceptions and Pitfalls Students have a hard time differentiating between nucleophiles and bases.

Emphasize that inversion of configuration can happen only if the carbon atom is chiral.

Conceptual Connection 21.2 Leaving Groups and Base Strength 21.5 Elimination Reactions

 









Emphasize the formation of a double bond, compounds. and therefore the formation of a new class of The strength of different bases has to be discussed in detail for students to appreciate

Base strength can be a challenging concept to students.

Acid and conjugate base pairs should be reviewed.

the full extent of the reaction.

Size of molecule and base strength is also an important concept for differentiation of products.

Conceptual Connection 21.3 Can a Strong Acid React Like a Base?

.

333

Chapter 21. Organic Chemistry II: Reactions Lecture Outline

Terms, Concepts, Relationships, Skills  Electrophilic Additions to Alkenes 21.6.  Hydrohalogenation mechanism Other addition reactions

 Figures, Tables, and Solved Example Unnumbered figure: stepwise mechanism of hydrohalogenation Table 21.4 Selected Additions to Alkenes



21.7 Nucleophilic Additions to Aldehydes and Ketones

  

Unnumbered figure: partial charges in the carbon–magnesium covalent bond Unnumbered figure: mechanism of the Grignard addition to a carbonyl group Example 21.6 Grignard Reactions

 o Addition of alcohols Hemiacetal and hemiketals The Grignard reaction



21.8 Nucleophilic Substitutions of Acyl Compounds

  

Unnumbered figure: general mechanism of nucleophilic substitution of an acyl group Unnumbered figure: equilibrium between esterification and hydrolysis Unnumbered figure: mechanism of Fischer esterification Unnumbered figure: reactivity ranking of carboxylic acid derivatives Table 21.5 Reagents for Common Interconversions of Acyl Groups Example 21.7 Acyl Group Interconversions

 Fischer esterification Synthesis of different acyl compounds by interconversions

   

.

334

Chapter 21. Organic Chemistry II: Reactions Teaching Tips

Misconceptions and Pitfalls

Suggestions and Examples  Electrophilic Additions to Alkenes 21.6. 

Emphasize the formation of two new sigma bonds in the molecule as a result of the addition reaction.

 

Emphasize the formation of new functional groups as a result of the reaction.

Conceptual Connection 21.4 H—X as an Electrophile 21.7 Nucleophilic Additions to Aldehydes and Conceptual Connection 21.5 KetonesMechanism of Hydration 



 Clarify for students the ring formation of monosaccharides.

Emphasize the difference between electrophilic and nucleophilic additions.



Emphasize the formation of new classes of compounds: acetals and



hemiacetals.

Explain the formation of the cyclic form of glucose.

Emphasize the importance of the 21.8 Nucleophilic Substitutions of Acyl chain Grignard reaction as a carbon Compounds  extension reaction.



 Students have a hard time distinguishing between compounds undergoing nucleophilic substitution and nucleophilic addition.

Emphasize the difference in reactivity of the these types of carbonyl compound (substitution vs. addition).

Conceptual Connection 21.6 The Reaction Between an Acyl Chloride and NH3

.

335

Chapter 21. Organic Chemistry II: Reactions Lecture Outline

Terms, Concepts, Relationships, Skills 21.9 Electrophilic Aromatic Substitutions Nitration of benzene

 Figures, Tables, and Solved Example Unnumbered figure: general mechanism of electrophilic substitution Unnumbered figure: nitration mechanism Table 21.6 Common Electrophilic Aromatic Substitution Reactions

 

21.10 Polymerization

  

Monomers vs. polymers Step‐growth polymers Addition polymers

   

Figure 21.4 Poly(ethylene trephthalate) Chemistry in Your Day: Kevlar Figure 21.6 Poly(vinyl chloride) Unnumbered figure: ethylene and polyethylene, chloroethane and poly(vinyl chloride)

.

336

Chapter 21. Organic Chemistry II: Reactions Teaching Tips

Suggestions and Examples 21.9 Electrophilic Aromatic Substitutions



Misconceptions and Pitfalls

Emphasize why aromatic compounds undergo electrophilic substitution rather than electrophilic addition.

Conceptual Connection 21.7 Resonance Structures 21.10 Polymerization  Emphasize the importance of polymers in everyday life.

.

337

Chapter 22. Biochemistry Student Objectives

22.1Diabetes and the Synthesis of Human Insulin 

Know that diabetes is a disease caused by the absence or insufficiency of insulin, a protein synthesized in the pancreas that helps regulate cellular and blood glucose levels.



Know that genetic engineering has allowed modified bacteria to produce human insulin, allowing it to be harvested for human use.

Know that biochemistry is the chemistry of living organisms, often involving macromolecules such as lipids, carbohydrates, proteins, and nucleic acids.

22.2Lipids   Give a general description of lipids with respect to structure and function.  Know that fatty acids can be saturated (all C–C bonds) or unsaturated (at least one C=C bond). and understand understand that howfats theare tri‐alcohol glycerol combines with fatty acids to form triglycerides.  Know Know and solids and oils are liquids room temperature, a consequence Know that natural geometric isomers of unsaturated fatty acidsatare typically cis, that trans forms are of the structures of the fatty acids in triglycerides. 

introduced when hydrogenation is used to control potential rancidity, and that trans fatty acids are

 

not considered healthy by nutrition experts.



Know that carbohydrates are polyhydroxy aldehydes and ketones that are nutritional sugars and starches and that serve as energy sources for the body.

Give a general description of phospholipids with respect to structure and function.

Draw and recognize representations of phospholipids and their role in lipid bilayers.

Recognize the structure of cholesterol and know that it serves as a structural component of cell membranes and as a precursor to other steroids in the body.

22.3Carbohydrates



Recognize the general structures of monosaccharides like and glucose and fructose.

Know and understand how monosaccharides like glucose fructose form rings that can combine to make more complex structures such as disaccharides (e.g., sucrose) or polysaccharides (e.g., cellulose, starch, glycogen).

22.4Proteins and Amino Acids 

Know and understand that proteins are macromolecules that serve as structural components, enzymes, and hormones and have additional roles in transport, storage, protection, and mobility.



Know that there are 20 common amino acids and that 10 essential ones must be consumed since the body cannot produce them from other molecules.

Know the general structure of amino acids and write reactions for the formation of peptide bonds.

22.5Protein Structure   Know that proteins are characterized as fibrous or globular.

Know that protein of fouroflevels: primary, secondary, tertiary, and‐pleated quaternary.  Recognize the twostructure commonconsists components secondary structure, ‐helices and sheets, and Know that the primary structure of a protein gives the sequence of amino acids.   know that hydrogen bonding enables the formation of both.

Understand the different intermolecular interactions that maintain tertiary structure.

Know that multimeric proteins exhibit quaternary structure. .

338

Chapter 22. Biochemistry

22.6Nucleic Acids: Blueprints for Proteins  Understand the basic structure of nucleic acids: sugars, bases, and phosphate links.  Know and understand how the nucleotides are covalently combined to make a single strand and how Define and understand codon, gene, and chromosome. two complementary strands form a double helix through hydrogen bonding.

22.7DNA Replication, the Double Helix, and Protein Synthesis  Know process offrom DNAan replication.

Know and howunderstand proteins arethe synthesized RNA copy of DNA and understand how the codons direct the ribosome to attach particular amino acids to each other to make the specific protein sequence.

Section Summaries

 Outline Lecture  Terms, Relationships, Skills Figures,Concepts, Tables, and Solved Examples  Teaching Tips  Suggestions andand Examples Misconceptions Pitfalls

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339

Chapter 22. Biochemistry Lecture Outline

Figures, Tables, and Solved Examples

Terms, Concepts, Relationships, Skills 22.1Diabetes and the Synthesis of Human Insulin o Diabetes o  Insulin Biotechnology: genetic engineering Biochemistry

 Intro figure: illustration of spiral timeline, structures of biomolecules Unnumbered figure: Frederick Banting

 22.2Lipids o Fatty o acid structure Even number of carbons, C4 to C20 Chains   



    

Saturated  cis Unsaturated  trans



o Glycerides o  Triglycerides Phospholipids  Phosphatidylcholine Lipid bilayers o Cholesterol o Membrane structure Biosynthetic precursor to testosterone, estradiol

    

Unnumbered figures: Lewis structures, space‐filling models of myristic acid, oleic acid Table 22.1 Fatty Acids Figure 22.1 The Effect of Unsaturation Figure 22.2 The Formation of Tristearin Unnumbered figure: photo of lard package Unnumbered figure: Lewis structure, space‐ filling model of tristearin, triolein Chemistry and Medicine: Dietary Fat: The Figure 22.3 Cis and Trans Fats Good, the Bad, and the Ugly Figure 22.4 Phosphatidylcholine Figure 22.5 Schematic for Phospholipid or Glycolipid Figure 22.6 Lipid Bilayer

cholesterol, testosterone, ‐estradiol Unnumbered figure: Lewis structure of 

 General structure: (CH2O)n 22.3 Carbohydrates o Monosaccharides o Carbonyl: aldehyde or ketone o o  OH groups Cyclic form Disaccharides o  Glycosidic linkage o Sucrose (from glucose and fructose) Polysaccharides o o Cellulose Starch Glycogen

       

.

Unnumbered figure: Lewis structure, space‐ filling model of glucose, fructose Figure 22.7 Fischer Projection of Carbon Atom with Four Different Substituents Example 22.1 Absolute Configurations of Carbohydrates Figure 22.8 Intramolecular Reaction of Glucose to Form a Ring Unnumbered figure: Lewis structures of ring forms of glucose, fructose, galactose Figure 22.9 Formation of a Glycosidic Linkage Figure 22.10 Hydrolysis of Disaccharides Unnumbered figures: Lewis structures of cellulose, starch Figure 22.11 Hydrolysis of Polysaccharides 340

Chapter 22. Biochemistry Teaching Tips

Suggestions and Examples 22.1Diabetes and the Synthesis of Human Insulin



Misconceptions and Pitfalls Not all missing or damaged proteins can be replaced. Insulin is an example of one that can be.

There are several examples of essential proteins that are missing and can be replaced.

Diabetes affects about 11 percent of the population. Other examples include human growth hormone and thyroid hormone.

Biotechnology or genetic engineering is commonly used to make important proteins 22.2 Lipids  and biomedical molecules.





 The structural properties of fats and oils come primarily from the chain length and degree of unsaturation of the fatty acids.

Fatty acids and glycerol are the building blocks of triglycerides, the components of nutritional fats and oils. The space‐filling models give a good sense of the shapes.

Melting points give a sense of how well the molecules can pack together.

Trans fats and the trans fatty acids from which they are made appear often in the nutrition literature and in the media. Trans fats do occur in nature but in small quantities.

The interactions of phospholipids with water make the membranes containing phospholipids even more stable. 

22.3Carbohydrates



Monosaccharides occur primarily in the cyclic form in water.

When compared with the lipids, carbohydrate structures are more complex, including issues such as the nature of the monosaccharide, ring structures, and the glycosidic linkage between any two monosaccharides in more complex molecules.

The microscopic and macroscopic properties of the more complex sugars arise from the nature of the monosaccharides and the glycosidic linkages between them.

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Chapter 22. Biochemistry Lecture Outline

Figures, Tables, and Solved Examples  Table 22.2 Protein Functions Unnumbered figure: Lewis structures of  amino acids  Unnumbered Unnumbered figure: figure: alanine illustration of different  types of amino acid side chains  Unnumbered figure: zwitterion  Table 22.3 Common Amino Acids  Unnumbered Example 22.2 figure: Peptidepeptide Bonds bond

Terms, Concepts, Relationships, Skills 22.4Proteins and Amino Acids o Classes o o Enzymes Hormones Structure, transport, storage,  o contraction, protection Amino o acids o General structure 20 common amino acids  10 essential amino acids Cannot be biosynthesized  22.5 Proteino Structure Structural categories  o Fibrous (linear)  Example: collagen  Globular Example: hemoglobin Levels o o Secondary: ‐helices, ‐pleated sheets o Primary: amino acid sequence o Tertiary: intramolecular interactions Quaternary: subunits

        

Figure 22.12 Fibrous and Globular Proteins Figure 22.13 Levels of Protein Structure Chemistry and Medicine: The Essential Amino Acids Table 22.4 Essential Amino Acids Figure 22.14 Primary Structure of Egg‐White Lysozyme Unnumbered photo Structure of normal and Figure 22.15 figure: The ‐Helix sickled 22.16 red blood Figure Thecells ‐Pleated Sheet Structure Figure 22.17 Interactions Within Proteins

 

22.6Nucleic Acids: Blueprints for Proteins o Basic o nucleotide structure o Phosphate Sugar: ribose or deoxyribose  o Base: purine or pyrimidine Genetic o code o Nucleotide o Codon Gene Chromosome

    

.

Figure 22.18 DNA Structure Unnumbered figure: Lewis structures of DNA and RNA sugars Unnumbered figure: DNA and RNA bases Figure 22.19 Base Pairing in DNA Figure 22.20 Short Strand of DNA Figure 22.21 Genetic Structure Figure 22.22 Chromosomes

342

Chapter 22. Biochemistry Teaching Tips

Suggestions and Examples  22.4 Proteins and Amino Acids





Misconceptions and Pitfalls  The amino onesessential that cannot be acids are biosynthesized by humans and therefore must be consumed (eaten).

The combination of amino acids to form polypeptides is an example of a condensation reaction. The acid portion of one amino acid combines with the amine portion of another, yielding a peptide bond and a water molecule.

The 20 common amino acids can be placed in a few categories according to their R group: alkyl groups, alcohols and thiols, carboxylic acids, amines and amides, and aromatic rings.

Conceptual Connection 22.1 Peptides 

22.5 Protein Structure



Protein structure involves intramolecular interaction between amino acid components that extends far beyond the mere order of the

Protein structural categories arise from different functions.

The fibrous examples like collagen serve as structural elements in muscles for example. Globular proteins are often more water soluble and serve as enzymes or to

amino acids in a protein.

support other functions besides structure. 

The levels of protein structure arise from the size, complex structure, and functions of many proteins. The intramolecular interactions often arise from the side chains of the amino acids.

Sickle‐cell anemia is due to a protein defect in which a single amino acid is replaced with one that is less functionalized. The ability of the red blood cell to carry oxygen is reduced. 

 22.6 Nucleic Acids: Blueprints for Proteins



The structures of RNA and DNA are relatively simple considering their importance. For example, only four bases are used to code for a host of proteins. The simple structure, though, enables duplication and repair.

Conceptual Connection 22.2 The Genetic Code

.

Base‐pairing in DNA occurs through specific hydrogen bonding interactions. Base‐ pair mismatches do not have those proper hydrogen bonding interactions.

343

Chapter 22. Biochemistry Lecture Outline

Terms, Concepts, Relationships, Skills

Figures, Tables, and Solved Examples   Figure 22.23 Watson and Crick  Unnumbered figure: complementary strains  Figure 22.24 DNA Double Helix  Figure 22.25 DNA Replication Figure 22.26 Protein Synthesis Chemistry and Medicine: The Human

22.7 DNA Replication, the Double Helix, and Protein o Replication Synthesis o o DNA double helix structure  Complementary single strands  o Base‐pair match Protein o synthesis o RNA copies of DNA Codons Ribosomes

Genome Project

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344

Chapter 22. Biochemistry Teaching Tips

Suggestions and Examples  22.7 DNA Replication, the Double Helix, and Protein Synthesis The Human Genome Project has provided the identity and order of bases for all genetic material in humans.

Now the effort focuses on identifying which chromosome  each gene resides in and the mechanisms by which the

Misconceptions and Pitfalls

massive amounts of information are organized.

Protein synthesis makes use of three amino‐acid codons.

As each codon is read by the ribosome, an amino acid is added to the polypeptide.

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345

Chapter 22. Biochemistry Additional Problem for Carbohydrates and Optical Isomerism Examine the structure of deoxyribose, part of the backbone in DNA. Does deoxyribose exhibit optical isomerism? If so, which carbon atoms are chiral?

Chiral Centres Examine each carbon to see which have four different atoms or groups attached.

Carbons 3 and 4 are chiral centres. Carbon 1 is the carbonyl C=O; carbon 2 and carbon 5 have two H attached apiece.

With two chiral centres, deoxyribose exhibits optical isomerism.

Optical Isomerism

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346

Chapter 22. Biochemistry Additional Problem for Peptide Bonds (Example 22.2)

Show the tripeptide that results from linking the amino acids glycine, valine, and serine, in that order. Label the N-terminus and C-terminus ends of the tripeptide.

Peptide Bonds A peptide or amide bond forms when the hydroxyl –OH from the carboxylic acid of one amino acid is removed, together with the hydrogen from the amino group of the adjacent amino acid.

Water is removed each time a peptide bond is formed. They appear in the boxes.

Tripeptide The resulting tripeptide is:

H2N – glycine – valine – serine – CO2H

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347

Chapter 23. Chemistry of the Nonmetals Student Objectives

23.1Insulated Nanowires 

Know that carbon forms nanotubes and that boron nitride, a material that is isoelectronic with carbon, also forms nanotubes.

Know that carbon nanotubes are electrical conductors, while boron nitride nanotubes are electrical insulators.

23.2The Main‐Group Elements: Bonding and Properties Recall the periodic trends for radius, the general properties of metals and nonmetals, and the major divisions of the periodic chart.

23.3Silicates: The Most Abundant Matter in Earth’s Crust  Define silicate.  Know that silicon and oxygen together constitute the majority of the Earth’s crust.

Know that quartz and glass have the empirical formula SiO2 but consists of tetrahedral SiO4   

Know that aluminosilicates form with Al atoms in Si sites, giving the formula unit AlO2. structural units.

Know some general characteristics of the different silicate structures.

Predict the structure of a given silicate.

23.4Boron and Its Remarkable Structures   Know that the different allotropes of elemental boron, B12, are based on an icosahedral structure.

Know that boron reacts with halogens to form trihalides.  Define and describe the different borane structures: closo, nido, and arachno.

Know that boron reacts with oxygen to form network covalent solids.

23.5Carbon, Carbides, and Carbonates  common crystallineform allotropes of carbon: graphite and diamond.  Describe Know thatthe coal is an amorphous of carbon that forms under pressure from the decomposition of    

ancient organic material and consists of three main types: lignite, bituminous, and anthracite.

Know some general characteristics of coke, charcoal, activated carbon, soot, and carbon black.

Know that carbon also forms fullerenes and nanotubes.

Know that carbon forms ionic carbides, covalent carbides, and metallic carbides when combined with different elements.

Know with oxygen to form carbon and carbon monoxide.

Know that that carbon carbon combines dioxide and water combine to form thedioxide carbonates: carbonic acid, bicarbonate, and carbonate.

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Chapter 23. Chemistry of the Nonmetals

23.6Nitrogen and Phosphorus: Essential Elements for Life  Recall that elemental nitrogen is a diatomic gas that comprises 78 percent of Earth’s atmosphere.  Know some general characteristics of the main allotropes of phosphorus: white phosphorus, red  phosphorus, and black phosphorus.  Know the common nitrogen hydrides: ammonia, hydrazine, and hydrogen azide.

Know the common nitrogen oxides: nitrogen oxide, nitrogen dioxide, dinitrogen trioxide, nitrous  oxide, and dinitrogen tetraoxide.   Know the oxyanions and oxyacids of nitrogen: nitrate, nitric acid, and nitrite.

Know that phosphorus forms phosphine (PH3) and the phosphine halides PX3 and PX5.

Know reacts with to form and Know that that phosphorus phosphoric acid and the oxygen phosphates areP4O6 among theP4O10. most important and useful phosphorus compounds.

23.7Oxygen   Recall that most elemental oxygen is a diatomic gas that comprises 21 percent of Earth’s atmosphere.

Know the three oxide types: oxide, peroxide, and superoxide.

Know some general characteristics of another common allotrope of oxygen, ozone (O3).

23.8Sulfur: A Dangerous but Useful Element     

Know that elemental sulfur is found in a few natural deposits believed to have been formed by anaerobic bacteria and is extracted by the Frasch process.

Know some general characteristics of dihydrogen sulfide.

Assign oxidation states to sulfur in redox reactions.

Know that sulfur dioxide forms naturally from volcanic activity and forms as a pollutant from coal and oil combustion as well as metal extraction.

Know the most abundantly produced chemical in the world.

Know that somesulfuric generalacid usesisof sulfuric acid.

23.9Halogens: Reactive Elements with High Electronegativity   Determine the oxidation states of halogens in redox reactions.  Know that halogens are not found in their elemental forms because of their high electronegativities.

Know hydrofluoric acid,and andhalogen chlorine.

Know some some general general characteristics characteristics of of fluorine, interhalogen compounds oxides.

Section Summaries

 Outline Lecture  Terms, Relationships, Skills Figures,Concepts, Tables, and Solved Examples  Teaching Tips  Suggestions andand Examples Misconceptions Pitfalls

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349

Chapter 23. Chemistry of the Nonmetals Lecture Outline



Terms, Concepts, Relationships, Skills 23.1Insulated Nanowires o Carbon nanotubes o o  Structure Electrical conductivity Boron o nitride nanotubes Structure Electrical insulation

 

Crust  

   



o Quartz and glass o o Formula unit: SiO2 Structure and shape: SiO4 tetrahedra

o



        

Aluminosilicates AlO2 in place of SiO2 44 o  Silicate units, chains, sheets o o Orthosilicate: SiO 2 76 Pyrosilicate: SiO unit: SiO  Repeating Pyroxene  Formula unit: SiO32 264 o Amphibole  Repeating and formula unit: 4 116   SiO Phyllosilicate  Formula unit: SiO 4104 Repeating unit: SiO 252

Intro figure: illustration of carbon nanotubes Table 23.1 Properties of BN and C Figure 23.1 Boron Nitride Nanotube

 Unnumbered figure: periodic tables p‐ block elements

23.2The Main‐Group Elements: Bonding and Properties   Atomic radius Electronegativity Metals, nonmetals, metalloids

23.3 Silicates: The Most Abundant Matter in Earth’s

Figures, Tables, and Solved Examples

Figure 23.2 Major Elements in Earth’s Crust Figure 23.3 SiO4 Tetrahedron Figure 23.4 Structure of Quartz Example 23.1 Determining the Composition of an Aluminosilicate Figure 23.5 Pyrosilicate Structure Figure 23.6 Pyroxene Structure Unnumbered figure: photo of asbestos fibres Unnumbered figure: photo of mica sheets Figure 23.7 Amphibole Structure Figure 23.8 Phyllosilicate Structure Table 23.2 Types of Silicate Structures Example 23.2 Composition and Charge Balance of Silicates Examples 23.3 and 23.4 Predicting Silicate Structures

Predicting structures

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350

Chapter 23. Chemistry of the Nonmetals Teaching Tips

Misconceptions and Pitfalls

Suggestions and Examples 23.1Insulated Nanowires Nanotubes are modern examples of the descriptive chemistry of carbon, boron, and nitrogen. The microscopic structures are nearly the same, but the macroscopic properties (e.g., electrical conductivity) are different.

23.2 The Main‐Group Elements: Bonding and Review the periodic properties of the p‐block elements: electron configurations, Properties electronegativity, radius, electron affinity, and metallic character.

 23.3 Silicates: The Most Abundant Matter in Earth’s Crust 

Silicates, like SiO2, do not exist as discrete molecules but instead are network covalent solids.

The silicates form a variety of structures, though most of them are based on a tetrahedral shape. This recognizable green pyramid is shown in all of the ball‐and‐stick models.

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351

Chapter 23. Chemistry of the Nonmetals Lecture Outline Terms, Concepts, Relationships, Skills

Figures, Tables, and Solved Examples  Figure 23.9 B12 Icosahedron  Figure 23.11 closo‐Borane Structures Figure Structure  Figure 23.10 23.12 B2O3 nido‐ and arachno‐Boranes

23.4 Boron and Its Remarkable Structures   Elemental boron: B12 icosahedron   Boron halides, BX3 o Boron oxygen compounds Trigonal BO3 in hexagonal shapes o closo Boron o hydrides nido ‐boranes o arachno ‐boranes ‐boranes   

23.5Carbon, Carbides, and Carbonates o Crystalline carbon forms o  Coal o Graphite: planar, hexagonal o Diamond: tetrahedral o Lignite Bituminous  Anthracite  





      

Other types: coke, charcoal, activated carbon, soot, carbon black o Specialized structures: fullerenes, nanotubes Carbides o Ionic o Covalent o Metallic Carbon o oxides o Carbon dioxide Carbon monoxide Carbonates o o Carbonic acid o Washing soda Baking soda Baking powder

.

Figure 23.13 Graphite Structure Figure 23.14 Diamond Structure Table 22.3 Approximate Composition of the Main Types of Coal Unnumbered figure: photo of tires Figure 23.15 C60 and a Geodesic Dome Figure 23.16 Carbon Nanotubes Figure 23.17 The Nanocar Figure 23.18 Calcium Carbide Structure Unnumbered figure: micrograph of steel Unnumbered figure: photo of Alka‐Seltzer tablets in water

352

Chapter 23. Chemistry of the Nonmetals Teaching Tips

Misconceptions and Pitfalls

Suggestions and Examples 23.4Boron and Its Remarkable Structures Boron predominantly forms trigonal structures for boron halide and boron oxide compounds but cage structures for hydride compounds.

 23.5Carbon, Carbides, and Carbonates





The crystalline allotropes of carbon are well known; other forms of carbon (e.g., coal, charcoal) are widely used. The newer forms, fullerenes and nanotubes, have been discovered and developed in the last quarter century.





The carbides, particularly the covalent and metallic carbides, are important commercially because of their hardness.

Carbide tips are added to saw blades made of

Graphite has a high melting point because of the strong bonds within the sheets, but it is slippery because of the weak interactions between graphite sheets.

Large amounts of carbon are sequestered in ocean water, in shells of marine organisms, and in rocks.

Fullerenes, particularly C60 and C70, are naturally occurring.

steel to make them sharper and cut longer.   

The combustion products of carbon, CO2 and CO, are well known and highly relevant to daily life.

Students are always fascinated by the fullerenes and carbon nanotubes.

Conceptual Connection 23.1 Phase Changes and Pressure Conceptual Connection 23.2 Carbonate Solubility

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353

Chapter 23. Chemistry of the Nonmetals Lecture Outline

Figures, Tables, and Solved Examples   Figure 23.19 White Phosphorus Figure 23.20 Red Phosphorus  Table 23.4 Oxidation States of Various Nitrogen Compounds  Figure 23.21 Hydrazine and Hydrogen  Peroxide

Terms, Concepts, Relationships, Skills 23.6 Nitrogen and Phosphorus: Essential Elements for  o Nitrogen: N2 (g), 78% of Earth’s Life Elemental forms o atmosphere  Phosphorus   White: P4 tetrahedra Red: P4 chains  Black: layered like graphite o Nitrogen compounds o Hydrides: ammonia, hydrazine, hydrogen azide o 

figure: vehicle crash test  Unnumbered Figure 23.22 Tetraphosphorus Hexaoxide and Decaoxide, P4O6 and P4O10 Table 23.5 Uses of Phosphates in the Food Industry

Oxides: nitrogen monoxide, nitrogen dioxide, dinitrogen trioxide,

o dinitrogen tetraxoxide, nitrous oxide Phosphorus compounds o Nitric acid, nitrate, nitrite o Hydride: phosphine o Halides: PX3, PX5, POX3 Oxides: P4O6, P4O10 Phosphoric acid, phosphates

 

23.7Oxygen o Elemental oxygen o o O2  Ozone, O3

Unnumbered figure: photo of ocean vents Table 23.6 Types of Oxides

Oxide Peroxide: O22 o types Oxide: O2 o Superoxide: O2  Ozone o o o Strong oxidizing agent Disinfectant Ozone layer

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Chapter 23. Chemistry of the Nonmetals Teaching Tips

Suggestions and Examples

 Misconceptions and Pitfalls

23.6 Nitrogen and Phosphorus: Essential Elements for  Life 

Both extremes of nitrogen fertilizers: the 3 oxidationoxidation state of Nstates in occur in molecules that are important as

A large amount of nitrogen occurs as the elemental gas N2 in the atmosphere. The



expanded nitrate. octet because of low‐lying d‐ Phosphorus can form molecules with an orbitals.

nitrogen oxides also are gases. 





ammonia and the +5 oxidation state of N in

The most common ions containing nitrogen are nitrate and nitrite. Nitric acid is a strong mineral acid used commercially in the fertilizer industry.

Ammonia is also a gas. Its natural supply comes from bacteria that are capable of “fixing” nitrogen: cleaving the triple bond in N2.

The allotropes of elemental phosphorus include the P4 tetrahedron of white phosphorus, the chain forming interconnecting tetrahedra in red phosphorus, and the graphite‐like structure of black phosphorus.

Phosphoric acid is used to remove scale (i.e., rust) in the steel industry, in the production of fertilizers, and as a flavour agent and preservative in soft drinks. Phosphates are widely used in detergents and in the food industry.



 23.7 Oxygen The atmospheric component gases O2 and O3 are well known, but large quantities of the most common element in Earth’s crust are  found as water and oxides of many other

Ozone in the upper atmosphere and as a disinfectant is useful, but ground‐level ozone is a pollutant.

elements.

Chemical cycles of bacterial life near some volcanic vents in the ocean are based not on oxygen but on sulfur.

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Chapter 23. Chemistry of the Nonmetals Lecture Outline

Terms, Concepts, Relationships, Skills  23.8 Sulfur:oA Dangerous but Useful Element Elemental form o o  Isolation by Frasch process Oxidation of recycled H2S Hydrogen sulfide o o Toxicity, odour Product bacteria and rotting Differentofproperties from those of  H2O o Sulfur o dioxide Natural volcanic source  Combustion product of coal and o petroleum Sulfuric acid o Most abundant chemical produced in the world

Figures, Tables, and Solved Examples   Unnumbered figure: photo of molten sulfur Figure 23.23 Quenching Liquid Sulfur  Figure 23.24 Balzac Natural Gas Plant, near Calgary, Alberta  Example 23.5 Balancing of and Assigning  Oxidation States to Sulfur Reactions Table Common Metal of Sulfides Figure23.7 23.25 Dehydration Sucrose

Common uses 23.9 Halogens: Reactive Elements with High



Electronegativity  o Elemental forms  because of o Not naturally occurring high electronegativity o o F2from o xidation of F  Cl2 from electrolysis Interhalogen compounds o A = large, B = small (B = F for n = 5, 7) ABn for n = 1, 3, 5, 7  o

  

Table 23.8 Selected Properties of the Halogens Example 23.6 Determining the Oxidation State of Halogens Table 23.9 Comparison of Halogen X—X Bond Energy Example 23.7 Identifying Changes in Oxidation States

Halogen oxides o o Often explosive Chlorine oxides Often strong oxidizing agents

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356

Chapter 23. Chemistry of the Nonmetals Teaching Tips

Suggestions and Examples  23.8 Sulfur: A Dangerous but Useful Element Elemental sulfur is isolated directly by the Frasch process or by oxidation of H2S that is recycled from other processes. Hydrogen  sulfide is toxic, so it must be removed or

 Misconceptions and Pitfalls Hydrogen sulfide looks as if it should resemble water based on their formulas.

However, nearly everything about these two compounds, including bond angle, toxicity, and physical properties, is different.

recycled anyway. 

Hydrogen sulfide has the characteristic smell of rotting eggs. It can be smelled at about 20 ppb, but the sense of smell is desensitized after exposure.

Sulfur dioxide is an environmental pollutant that arises from combustion of sulfur‐ containing coal and petroleum products.

Large quantities of sulfur dioxide are also produced from volcanic activity. 

23.9 Halogens: Reactive Elements with High  Electronegativity Elemental halogens can be formed by electrolysis of the halide salts.

Like the group 1 and 2 metals, none of the elemental forms of the halogens occur naturally because of their high reactivity.

Halogen oxides are often explosive, and most of the known ones of chlorine are strong oxidizing agents.

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357

Chapter 23. Chemistry of the Nonmetals Additional Problem for Predicting Silicate Structures (Examples 23.3 and 23.4)

Predict the silicate structure for the mineral talc Mg3Si4O10(OH)2 that serves as the basis for talcum powder, and show that the formula is charge neutral.

Si to O Ratio Si:O ratio is 4:10 or 2:5.

Determine the ratio of Si to O in the formula.

The other two oxygens occur as hydroxides and are not part of the silicate structures. (That situation is indicated by the OH in the parentheses.)

Match Ratio to List Match the Si:O ratio to the type of silicate in Table 23.2.

The 2:5 ratio is indicated for phyllosilicate sheets with a formula unit of SiO2 25.

Anion Charge SiO2 25: 2 @ 2 = 4

Determine the total anion charge.

OH: 2 @ 1 = 2 Total anionic charge = 6

Cation Charge Mg2+: 3 @ +2 = +6

Determine the total cation charge and show that it matches the total anion charge.

Total cationic charge = +6

The anionic and cationic charges balance, so the formula is charge neutral.

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358

Chapter 23. Chemistry of the Nonmetals Additional Problem for Determining the Oxidation State of Halogens in Compounds (Example 23.6)

Calculate the oxidation state for Cl in the following compounds:

a) Cl2O6 b) ClF3 c) HClO4

Oxidation states:

Solution a

oxygen: 3 @ 2 = 6 chlorine: 2 @ ?? = +6

Chlorine must be +3 oxidation state.

Oxidation states:

Solution b

fluorine: 3 @ 1 = 3 chlorine: 1 @ ?? = +3

Chlorine must be +3 oxidation state.

Oxidation states:

Solution c

hydrogen: 1 @ +1 = +1 oxygen: 4 @ 2 = 8 chlorine: 1 @ ?? = +7

Chlorine must be +7 oxidation state.

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359

Chapter 24. Metals and Metallurgy Student Objectives

24.1Vanadium: A Problem and an Opportunity 

Know that petroleum contains trace components other than the hydrocarbons that are the target fuel.



Know that petroleum contains vanadium compounds as contaminants at levels that suggest that they were part of some extinct species’ biological systems.

Know that metallurgy includes all processes associated with mining, separating, and refining metals to their pure forms and then combining metals into alloys that have different properties 24.2The General Properties and Natural Distribution of Metals  

Know the common properties of metal, such as conductors of electricity and heat, malleability, and ductility.

Know make up the crust and core of the Earth. crystalline inorganic solids. An ore is a Know that that metals minerals are homogeneous, naturally occurring, rock that contains a high concentration of a specific mineral.

24.3Metallurgical Processes    

Know that extractive metallurgy involves separating the metal‐containing components and nonmetal‐containing components. Crushed ore can be separated by air in an industrial cyclone or by solutions of water, oil and detergents.

Know that in pyrometallurgy, heat is used to extract metal from a mineral.

Know that hydrometallurgy uses an aqueous solution to extract metals from their ores.

Know to convert a metal salt into Know that that electrometallurgy powder metallurgyuses useselectrolysis micron‐sized metal high pressure andthe heatmetal. to form the desired compound at temperatures below melting point but high enough for the particles to fuse.

24.4Metal Structures and Alloys 

Know that structures of metals can be described as the closest packing of spheres into crystal types: face‐centred cubic, body‐centred cubic, and hexagonal closest packed.



Know that an alloy is a metallic material that contains more than one type of element in solid solutions or in specific ratios.

  

Know substitution alloys, the twoatmetals form a miscible solution.

Know that that in solid compositions that melt temperatures lower orsolid higher than either of the two pure metals are common.

Know that some alloys exhibit phase regions with limited solubility of one of the components and various mixtures can have different crystal structures.

Know that interstitial alloys form when a second metal can fit in the hole made by the other metal atoms in closest‐packed crystals.

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Chapter 24. Metals and Metallurgy 24.5 Sources, Properties, and Products of Some of the 3d Transition Metals 



Know that titanium is the fourth most abundant metal on Earth; it occurs mainly as oxides, is isolated by conversion of the oxides to intermediate TiCl4 and then reduction, is a strong and light metal used in propellers and jet engine parts, and has its largest use in the form of TiO2, a white powder used in paints.



Know that chromium compounds are brightly coloured, the oxide ore is reduced to form the meta,; an alloy with steel called stainless steel resists rusting, and the ions chromate and dichromate are





 

strong oxidizing agents used with organic molecules.

Know that manganese occurs as mineral ores, is reduced to the metal by Al or Na, is used to alloy with steel, as permanganate is used as an oxidizing agent, and as MnO2 is an additive to glass used to decolourize it.

Know that cobalt is found in ores of other metals and sulfide minerals, is ferromagnetic and important for the production of magnets, is an additive to steel, as salts it gives brilliant blue colour to pigments and inks, and is part of vitamin B12.

Know that copper can be found in its elemental form and thus has been used as a metal for many years; it has high conductivity, is used to manufacture pipes and metal surfaces on roofs, and is Section Summaries formed into bronze and brass.

Know is used as an alloy with steel to and formitstainless Know that that nickel zinc comes from ores, it forms alloys, is used assteel. an anticorrosion coating.  Lecture Outline  Terms, Relationships, Skills Figures,Concepts, Tables, and Solved Examples  Teaching Tips  Suggestions andand Examples Misconceptions Pitfalls

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Chapter 24. Metals and Metallurgy Lecture Outline

Terms, Concepts, Relationships, Skills 24.1Vanadium: A Problem and an Opportunity o Petroleum contaminants o Vanadium compounds Ancient organism biology based  on vanadium Metallurgy 24.2 The General Properties and Natural Distribution of Metals  o Properties o Opaque o Conductors of heat and o electricity  Malleable Ductile Minerals and ores 24.3 Metallurgical Processes  o Isolation steps o o Mine Separate  o Refine  Extractive metallurgy Pyrometallurgy   Calcination Roasting Smelting   Flux o Slag  Hydrometallurgy  Leaching o  Example: gold  Electrometallurgy Aluminum: Hall process o Copper Powder metallurgy

Figures, Tables, and Solved Examples  Intro figure: pieces of vanadium metal

  

    

.

Table 24.1 Thermal Conductivity and Electrical Resistivity of Several Metals Unnumbered figure: salt production by desalinization Unnumbered figure: vein of gold metal

Figure 24.1 Separation by Air Figure 24.2 Separation by Using a Solution Figure 24.3 The Hall Process Figure 24.4 Copper Electrolysis Cell Unnumbered figure: different metal products

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Chapter 24. Metals and Metallurgy Teaching Tips

Suggestions and Examples 24.1 Vanadium: A Problem and an

 Misconceptions and Pitfalls Vanadium may only be present in trace amounts, but the process of converting petroleum to fuel products or burning the fuels that contain it may concentrate the samples.

 Opportunity Use vanadium as an example to understand the pros and cons of certain sources and uses of different metals.

24.2 The General Properties and Natural Distribution of Metals 

 Most metals do not occur naturally in their elemental form but must undergo reduction of the cationic form from ores that have been concentrated.

The distribution of metals in nature requires an understanding of the structure and composition of ores and minerals.

24.3 Metallurgical Processes

 



Metallurgical processes have been developed over time to make the

Metallurgy takes advantage of basic chemical principles to concentrate an element or mineral or to reduce metal cations to the elemental form.

isolation of metals more efficient.

The technique used for a particular metal is likely to depend on the nature and quantity of an elemental metal or its ore.

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Chapter 24. Metals and Metallurgy Lecture Outline

Terms, Concepts, Relationships, Skills 24.4Metal Structures and Alloys o Closest o packing of spheres o Face‐centred cubic Body‐centred cubic  Alloys oHexagonal closest packed  Substitutional  Miscible solid 

o

solutions

Alloys with limited  solubility Interstitial  Phase diagrams Octahedral Tetrahedralholes holes

24.5 Sources, Properties, and Products of Some of the 3d Transition Metals   o Titanium o Sources (ores, minerals), uses Chromium  Sources (ores, minerals), uses, o alloys Manganese  Sources (ores, minerals), uses, alloys Cobalt o Nickel o Sources (ores, minerals), uses,  Zinc o alloys 

Tables, and Solved Examples Figures, Table 24.2 Thefigure: Crystal Structures of the 3d Elements Unnumbered face‐centred structure of  crystallized copper Unnumbered figure: structure of nickel copper alloy Figure 24.5 Cu–Ni Phase Diagram Figure 24.6 Cr–V Phase Examples 24.1and 24.2 Diagram Determining Alloy

   

Compositions from a Phase Diagram

24.7 Cr–Ni Diagram  Figure Examples 24.3 andPhase 24.4 Alloy Compositions in a  Solid Solution with Limited Solubility Figure 24.8 Octahedral Holes in Closest Packed Crystals



ViewHole of aninOctahedral Hole  Figure Figure 24.9 24.10AADifferent Tetrahedral a Closest Packed  Crystal  Table 24.3 Formulas of Several Interstitial Alloys

  Table 24.4 The Colours of Various Chromium  Compounds Figure 24.11 Chromium Compounds  Unnumbered figure: tetrahedral arrangements of chromate and dichromate ions  Unnumbered figure: ignition of ammonium dichromate



Unnumbered figure: varied applications of copper metal Unnumbered figure : zinc phosphate adhering to steel

Sources (ores, minerals), uses, alloys Sources (ores, minerals), uses, alloys

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Chapter 24. Metals and Metallurgy Teaching Tips

Suggestions and Examples 24.4Metal Structures and Alloys 



24.5 the 3

Misconceptions and Pitfalls  The formulation of alloys takes advantage of knowledge of both microscopic properties and macroscopic properties.

Knowledge of crystal lattices takes on a practical approach. Substitutional alloys require knowledge of the size of a potential component and melting information as a function of the mole percent of the components. Interstitial alloys can take advantage of the holes in the closest packed crystals.

Conceptual Connection 24.1 Interstitial Alloys



Sources Properties and Products of Some of d Transition Metals 

The descriptive chemistry of seven metals is outlined: the isolation of the metal depends on the primary source while the practical uses depend on the microscopic properties of the metal.

Seven metals are compared and contrasted in terms of isolation, practical applications, and their compounds.

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Chapter 25. Transition Metals and Coordination Compounds Student Objectives

25.1The Colours Rubies and Know thatof the colours ofEmeralds gemstones are attributable to electronic transitions between the d‐orbitals of transition elements.

25.2Electron Configurations of Transition Metals Recall that the transition metals occupy the d block of the periodic table.  Recall how to write electron configurations and determine oxidation states for d‐block elements.  Recall trends in radius, ionization energy, and electronegativity for d‐block elements.

Define complex ion, ligand, coordination compound, and coordinate covalent bond.

25.3Coordination Compounds   Define chelate and chelating agent.  Recognize monodentate, bidentate, and polydentate ligands.  Recognize the common geometries of complex ions.

Write names and formulas for coordination compounds.

25.4Structure and Isomerization   Differentiate between structural isomers and stereoisomers.  Recognize and identify coordination isomers and linkage isomers.

Identify and draw geometric isomers.

Recognize and draw optical isomers.

25.5Bonding in Coordination Compounds  Explain the crystal field theory of d‐orbital splitting in octahedral and tetrahedral complexes.  State the hybridization of the metal centre using valence bond theory for particular geometries.

Understand the relationship between crystal field splitting and the colours of transition metal   complexes. between weak‐field, moderate‐field, andmagnetic strong‐field ligands. of transition metal  Know Know the the difference relationship between crystal field splitting and the properties  complexes.

Understand the difference between high‐spin low‐spincomplex. in transition metal complexes.

Predict the number of unpaired electrons in anand octahedral 25.6Applications of Coordination Compounds  Know the role of transition metals in biomolecules such as hemoglobin, cytochrome c, chlorophyll, Know that EDTA is used as a chelating agent to bind toxic heavy metals. and carbonic anhydrase.

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Chapter 25. Transition Metals and Coordination Compounds Section Summaries

 Outline Lecture  Terms, Relationships, Skills Figures,Concepts, Tables, and Solved Examples  Teaching Tips  Suggestions andand Examples Misconceptions Pitfalls

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367

Chapter 25. Transition Metals and Coordination Compounds Lecture Outline

 Figures, Tables, and Solved Examples

Terms, Concepts, Relationships, Skills 25.1The Colours of Rubies and Emeralds  Crystal field Cr3+ o In o Al2O3 (ruby, red)  In Be3Al2(SiO3)6 (emerald, green) In garnet (red) Fe2+ o In o peridot (yellow‐green) Cu2+ o  In turquoise

the shape of the dz2

Intro figure: illustration of ruby colour and  orbital Unnumbered figure: photos of ruby, emerald, garnet, peridot, and turquoise

  25.2 Electron Configurations of Transition Metals Electron o configurations Periodic o properties o Atomic radius  First ionization energy Electronegativity Oxidation states



Table 25.1 First‐Row Transition Metal Orbital Occupancy Figure 25.1 First‐Row Transition Metal Oxidation States

  25.3 Coordination Compounds o Coordination complex o Complex ion plus counterions Common ligands Monodentate: water, ammonia, o chloride, carbon monoxide, o

    

thiocyanate

Bidentate: oxalate, ethylenediamine Polydentate: ethylenediamine o tetraacetate Geometries and coordination number o o Linear o Square planar Tetrahedral  Names oOctahedral Ligands and prefixes o Metal centre o Names of metals in anionic complexes



    

.

Unnumbered figure: molecular model of [Co(NH3)6]Cl3 Unnumbered figure: Lewis acid–base reaction Unnumbered figure: bidentate ligand Table 25.2 Common Ligands Figure 25.5 Bidentate and Polydentate Ligands Coordinated to Co(III) Table 25.3 Common Geometries of Complex Ions Unnumbered table: Guidelines for Naming Complex Ions Unnumbered figure: EDTA Table 25.4 Names and Formulas of Common Ligands Table 25.5 Names of Common Metals when Found in Anionic Complex Ions Examples 25.1 and 25.2 Naming Coordination Compounds 368

Chapter 25. Transition Metals and Coordination Compounds Teaching Tips

Misconceptions and Pitfalls

Suggestions and Examples 25.1The Colours of Rubies and Emeralds The colours of gemstones, which are caused by transition metals in different mineral environments, are an introduction to crystal field theory, coordination compounds, and complex ions.

  25.2 Electron Configuration of Transition Metals 

transition metals fill the 4s before the 3d Electron thebefirst‐row unless allconfigurations the 3d orbitalsofcan filled evenly

The electron configurations of transition metals follow the rules and guidelines set



(e.g., Cr and Cu). remove electrons from the 4s before the 3d.

When forming cations of transition metals,

forth previously in Section 8.3.

Oxidation states are likely to be somewhat confusing because of the variety of values some elements can assume. In general, the first‐row transition metals can have oxidation states of +2 and +3. Elements in the centre of the row have more possible states than the ones on the ends. For example, Sc is only +3, Zn is only +2, but Mn can be any state from +2 to +7.



 25.3 Coordination Compounds Students may be confused about the organization of metal cations, ligands, and additional properties for complex ions. Table 25.2 lists the common ligands, and Table 25.3 lists the shapes and coordination numbers.  Likewise, one can often predict the geometry

the metal are found in the d orbitals. For this In complexes, all of the valence electrons on reason, complexes of Ir+ are isoelectronic with complexes of Pt2+ and Au3+.

given the formula of a complex ion. 

Students may struggle with predicting the formulas of complex ions. For example, what does Fe2+ form when reacted with ammonia?

Naming complex ions follows some basic guidelines. Often, coordination chemistry relies more on structures than on names.

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Chapter 25. Transition Metals and Coordination Compounds Lecture Outline

Figures, Tables, and Solved Examples

Terms, Concepts, Relationships, Skills 25.4Structure and Isomerization o Structural isomers o o  Coordination isomers Linkage isomers Stereoisomers  Cis –trans Geometric isomers  Fac–mer o

   

Examples 25.3 and 25.4 Identifying and  Drawing Stereoisomers  Figure 25.6 Stereoisomerism in [Co(en)3]3+ Example 25.5 Recognizing and Drawing

Optical isomers

Enantiomers

 25.5 Bonding o in Coordination Compounds Valence o bond theory o Geometry Hybridization  o Orbitals d theory Crystal o field orbital orientation   Octahedral field  Splitting pattern  Strong field field  Weak Colours and crystal field 

o o

Table 25.6 Ligands Capable of Linkage Figure 25.4 Cis–trans Isomerism Isomerization Figure Figure 25.3 25.5 Linkage Fac–MerIsomers Isomerism in Co(NH3)3Cl3



Figure 25.7 Relative Positions of d Orbitals



Figure 25.8 d Orbital Splitting in an and Ligands in an Octahedral Complex

     

splitting energy



Spectrochemical series   Magnetism High‐spin Low‐spin  Tetrahedral field  Splitting pattern Square planar field Splitting pattern



Octahedral Field Figure 25.9 Colours of Complex Ions Figure 25.10 The Colour Wheel Figure 25.11 The Colour and Absorption Spectrum of [Ti(H2O)6]3+ Example 25.6 Calculating Crystal Field Splitting Energy Unnumbered figures: orbital diagrams of low‐ and high‐spin Fe2+ complexes Examples 25.7 and 25.8 and Figure Splitting of d d High‐ Orbitals byLow‐ Figure 25.12 25.13 Splitting of Orbitals by aa Tetrahedral Ligand Geometry Spin Octahedral Complexes Square Planar Ligand Geometry

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370

Chapter 25. Transition Metals and Coordination Compounds Teaching Tips

Suggestions and Examples 25.4Structure and Isomerization 

 

and Pitfalls  Misconceptions Tetrahedral complexes cannot exhibit cis– trans isomerism, and square planar complexes cannot exhibit optical isomerism.

Some ligands can be attached in two distinct ways, giving rise to linkage isomers.

Models are likely to be helpful for seeing or Fac and merthe require three ligands the in visualizing arrangement of theof ligands octahedral structures. same kind in an octahedron.

Visualizing optical isomerism in octahedral complexes is much more difficult than in tetrahedral complexes. Models are essential for the beginning student. 

 25.5 Bonding in Coordination Compounds





Students often will simply memorize the splitting patterns for octahedral, tetrahedral, and square planar complexes. These patterns orientation of the d orbitals and the positions  can be simply deduced by considering the of the incoming ligands.

Students often confuse the terms weak‐ and strong‐field with high‐ and low‐spin.

The crystal field model is completely electrostatic and views all ligands as merely point negative charges. This model does not explain the spectrochemical series at all.

Strong‐ and weak‐field effects of ligands can be confirmed by the number of unpaired electrons in octahedral complexes.



Demonstrations of colour (especially cobalt complexes) and magnetism (diamagnetism, paramagnetism, and number of unpaired electrons) are easy and highly illustrative.

Conceptual Connection 25.1 Weak‐ and Strong‐Field Ligands

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371

Chapter 25. Transition Metals and Coordination Compounds Lecture Outline

Terms, Concepts, Relationships, Skills  25.6 Applications of Coordination Chemical analysis: SCN Compounds Chelating agents: EDTA   Colouring agents o Hemoglobin and cytochrome c: iron in Biomolecules o porphyrin ligand o o  Chlorophyll: magnesium in porphyrin

Tables, and Solved Examples Figures, Figure 25.14 Chemical Analysis SCN Table 25.7 Transition Metals andwith Some of  Their Functions in the Human Body      

Carbonic anhydrase: zinc Drugs and therapeutic agents Cisplatin: cis‐[Pt(NH3)2Cl2]

cytochrome c, heme Unnumbered figure: structures of Figure 25.15 Porphyrin Structures Figure 25.16 Hemoglobin Figure 25.17 Chlorophyll Figure 25.18 Carbonic Anhydrase Figure 25.19 Cisplatin

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372

Chapter 25. Transition Metals and Coordination Compounds Teaching Tips

Suggestions and Examples  25.6 Applications of Coordination Compounds The use of chelating agents to remove lead and mercury is a documented treatment when  a case of heavy metal poisoning is known or strongly suspected.

Misconceptions and Pitfalls

With an unsubstantiated belief that mercury‐ containing preservatives found in a few antiviral vaccines cause autism, some children are being subjected to EDTA treatment. What are some of the potential dangers or hazards of chelation therapy?

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373

Chapter 25. Transition Metals and Coordination Compounds Additional Problem for Writing Electron Configurations for Transition Metals

Write the ground state electron configuration for Nb and Nb3+.

Identify the noble gas Identify the noble gas that precedes the element and put it in square brackets.

Nb is atomic number 41 in period 5 and Group 15.

Complete noble gas shell = [Kr]

Count the periods Count down the periods to determine the outer principal quantum level—this is the quantum level for the s orbital. Subtract one to obtain the quantum level for the d orbital.

Nb is in the fifth period:

[Kr] 5s4d

Count the electrons Nb has five more electrons than [Kr]:

Count across the row to see how many electrons are in the neutral atom and fill the orbitals accordingly.

[Kr] 5s25d3 For an ion, remove the required number of electrons, first from the s and then from the d orbitals.

Nb3+ has lost three electrons relative to Nb:

[Kr] 5s04d2 or [Kr] 4d2

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374

Chapter 25. Transition Metals and Coordination Compounds Additional Problem for High- and Low-Spin Octahedral Complexes (Examples 25.7 and 25.8)

How many unpaired electrons would you expect for the complex ion [Cu(NH)]2+ ?

36

Charge and number of d electrons Complex ion = +2

Begin by determining the charge and number of d electrons on the metal.

NH3: 6 @ +0 = 0 Cu: 1 @ +2 = +2  Cu2+

Cu: [Ar] 4s13d10 Cu2+: [Ar] 4s03d9  Cu2+ has d9 configuration.

Ligands in spectrochemical series NH3 is a strong-field ligand, so  will be relatively

Look at the spectrochemical series to determine whether the ligand is a strong-field or weak-field ligand.

large.

High-spin or low-spin?

Strong-field ligands yield low-spin configurations.

Decide if the complex will be high- or low-spin and draw the electron configuration.

Unpaired electrons This configuration has one unpaired electron.

Count the unpaired electrons.

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