Chemistry 125 Demonstrations

Winter Quarter 2002

 

Chapter 9 - Molecular Geometry and Bonding Theories - you may wish to use some of these demonstrations as you review molecular geometry and polarity

 

a)   The VSEPR Model

1)   Use sets of colored balloons and/or ball-and-stick models to effectively illustrate VSEPR and to show the shapes of molecules; different colored balloons represent lone pairs of electrons in the balloon models, and plain rods with no ball attached (or rods with a large colored foam ball attached) represent lone pairs in the ball-and-stick models.  The ball-and-stick models can be made with generic black and white balls or with different colors representing specific elements.

2)   Start with a set of 4 balloons tied together and pop them one by one to show the geometries that arise naturally from repulsion of "electron pairs"

3)   Show cardboard models of a tetrahedron and an octahedron

4)   Display a lucite tetrahedron containing a tetrahedral ball-and-stick model to illustrate the terms "tetrahedron" and "tetrahedral"

5)     Show black and white ball-and-stick models representing linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral geometry.  You can remove white balls, leaving bare sticks to represent lone pairs, or you can substitute colored balls for white balls to represent lone pairs.  (Please specify your preference.)

6)     Show ball-and-stick models made up in different colors, with plain rods (or rods with a large colored foam ball attached) representing lone pairs, to illustrate the molecular shapes possible with each type of electron-pair geometry

A)     Linear - BeCl2

B)     Trigonal planar - BF3, O3, NO2-

C)     Tetrahedral - CH4, NH4+, SO42-, NH3, H2O, CH3OH

D)     Trigonal bipyramidal - PF5, SF4, ClF3, XeF2

E)     Octahedral - SF6, IF5, XeF4

 

b)   Polarity of Polyatomic Molecules

1)   Demonstrate the effect of a charged rod on streams of H2O and CCl4 flowing from burets; show models of the compounds to relate geometry to polarity

2)   Contrast ball-and-stick models of different compounds to relate the geometry of these molecules to their dipole moments

A)  CCl4, CH2Cl2, and CH3Cl

B)  NH3, SO3, and H2CO

C)  H2O and CO2

 

 

Chapter 11 - Intermolecular Forces, Liquids, and Solids

 

a)   A Molecular Comparison of Liquids and Solids

1)   Use two sealed lucite boxes containing BB's on the overhead projector to illustrate some of the differences between gases, liquids, and solids

2)   Use magnetized steel balls in the molecular motion demonstrator on the overhead projector first to show the random motion of particles in a gas and then to show how attractive forces cause a gas to liquefy at lower temperatures

3)     Halogens - show sealed flasks of Cl2(g), Br2(l), and I2(s)

 


b)   Intermolecular Forces - use ball-and-stick models of the compounds involved in each demonstration to relate geometry and structure to intermolecular forces

1)   Show the dependence of dipole-dipole forces on geometry by contrasting the effect of a charged rod on streams of H2O and CCl4 flowing from burets

2)   Use pairs of space-filling models of the pentane isomers to show that increased branching increases compactness, decreases polarizability, and decreases London forces with a resulting decrease in boiling point.

3)   Show pairs of ball-and-stick models of HF, H2O, and NH3, with brass rods representing lone pairs, to help illustrate hydrogen bonding.  If desired, you may also request models of H2S, PH3, and SiH4 to discuss molecules that do not exhibit hydrogen bonding.

4)   Show the effect of hydrogen-bonding between H2O and NH3 effectively with the ammonia fountain demonstration

5)   Show ball-and-stick models of C2H5OH, CH3OCH3, C3H8, and CH3CHO to compare and contrast the ability of each of these compounds to form intermolecular hydrogen bonds

6)   Show ball-and-stick models of 1-propanol and methyl ethyl ether, C3H5OH and CH3OCH2CH3, to compare and contrast the ability of an alcohol and an ether to form intermolecular hydrogen bonds

7)   Show a model of ice to point out the extensive hydrogen-bonding in its structure and discuss its effect on the density of ice compared to liquid water

8)     Pass around a wave-maker, a bottle containing vegetable oil and colored water, to show an attractive result of differences in attractive forces; then relate the phenomenon to the serious consequences of an oil spill

9)   Demonstrate the negative volume of mixing for ethanol and water using a special buret; this demo provides another way to illustrate the effects of intermolecular attractions

10) You may wish to use "the wave" done at sports events in large stadiums as a loose analogy to London dispersion forces; an article in Journal of Chemical Education (Volume 75, 1998, 1301) describes the analogy

 

c)   Some Properties of Liquids: Viscosity and Surface Tension

1)   Compare the viscosity of various liquids (such as glycerin, ethanol, cyclohexanol) on the overhead projector and relate the differences to strength of attractive forces

2)   Add ethanol to one arm of a U-tube containing water, observe different heights in the two arms due to different densities; this demo shows that diffusion of liquids is very slow

3)   Show tiny pieces of camphor moving rapidly in a dish of water on the overhead projector as they break up the surface tension of water

4)   Pour water through cheesecloth into a jar, then stretch the cheesecloth over the top of the jar and invert it - the water won't run out, due to surface tension (and atmospheric pressure)

5)   Demonstrate the disruption of the surface tension of water by detergent by sprinkling baby powder or pepper in a large dish of water on the overhead projector and then touching the water lightly with a wood stick dipped in detergent.  Alternatively, float a berry basket on the surface of a dish of water, then watch it sink when you add a drop of detergent.

6)     Demonstrate capillary action by inserting a capillary tube in a small beaker of colored water and passing it around the class

 

d)   Phase Changes

1)   Changes of state - heat a beaker of ice on a hot plate during the lecture to show the changes from solid to liquid to gas.  (A magnetic stirrer and digital thermometer can be provided if you wish to observe the temperature changes.)


2)   Immerse a sealed flask of Br2(l) in liquid nitrogen to show a change of state

3)   Pour liquid nitrogen on the floor to show a change of state

4)   Make the sublimation of dry ice "visible" by dropping a piece of dry ice in a beaker of water

5)   Pour liquid nitrogen into a beaker to demonstrate a variety of phase changes: the boiling of N2(l), the deposition of H2O(g) as H2O(s) on the outside of the beaker, and the melting of H2O(s) to H2O(l) as the beaker eventually warms up again

6)   Contrast the behavior of ice and dry ice in separate beakers as they sit at room temperature and/or as they are heated on a hot plate to demonstrate melting and sublimation

7)   Critical Temperature - use a special device mounted on a modified overhead projector to dramatically demonstrate the disappearance of the liquid phase above the critical temperature. (NOTE: There are some time problems associated with this demo.)

 

e)   Vapor Pressure

1)   Demonstrate the high vapor pressure of ethanol or methanol by placing a small amount of the liquid in a 1 gallon plastic milk carton (or a large plastic carboy, subject to availability), allowing time for some to evaporate, then pouring out the liquid, and holding a lighted splint to the mouth of the "empty" container - the vapor combustion is quite spectacular

2)   Demonstrate the high vapor pressure of diethyl ether by allowing the vapor to flow from a can down a trough to a candle, resulting in a vapor flashback fire

3)   Display sealed flasks of Br2(R) and I2(s) to show the high vapor pressure of these substances, evidenced by the colored vapor in each flask

4)   Show the effect of intermolecular forces on the vapor pressure of liquids in an apparatus which contrasts the vapor pressure of two isomers, diethyl ether and 1-butanol

5)   Show the effect of intermolecular forces on vapor pressure by letting three students make streaks of water, methanol, and acetone on the blackboard and then observing the relative rates of evaporation

6)     You may wish to support the definition of vapor pressure by introducing the concept of dynamic equilibrium by having two students pour water between two 5 gallon jars (starting with one full, one empty) until equilibrium is achieved; this demo helps students view dynamic equilibrium as a state where the rates of forward and reverse reactions are equal

 

f)    Vapor Pressure and Boiling Point

1)   Heat water to boiling at atmospheric pressure and measure its temperature

2)   Show water boiling at room temperature in a beaker in an evacuated bell jar, then put your hand in the water after boiling to convince students of its low temperature.  If you prefer to drink the water after "boiling" it, warn me so I can give you a very clean beaker and drinking water.

 

g)   Phase Diagrams

1)   Demonstrate the effect of pressure on the melting point of ice by hanging a wire weighted at both ends over a block of ice - eventually the wire passes through the ice and the weights fall, leaving the ice intact, because the ice below the wire melts due to pressure but the water above the wire refreezes as the pressure is relieved.

2)     Contrast the behavior of ice and dry ice in separate beakers as they sit at room temperature and/or as they are heated on a hot plate


3)   Display some pieces of dry ice when discussing the phase diagram of CO2 or add a few pieces of dry ice to water or tie them inside a rubber glove to demonstrate the sublimation of CO2(s) more dramatically

4)   Demonstrate the existence of three phases of CO2 at the triple point by adding crushed dry ice to a length of heavy tygon tubing fitted with a pressure gauge and a screw clamp

 

h)   Structures of Solids

1)   Amorphous Solids and Crystalline Solids - contrast a piece of charcoal, a large crystal of a mineral, and a polished quartz crystal (SiO2)

2)   Unit Cells

A)  Show ball-and-stick models of simple cubic, BCC, and FCC unit cells, as well as space-filling models and large extended lattice models of BCC and FCC if desired

B)  Stack several space-filling FCC cubes or BCC cubes to show how unit cells form an extended lattice

3)     Display unit cell models of NaCl or CsCl

4)     Close Packing of Spheres - illustrate hexagonal and cubic close-packed structures with layers of styrofoam balls and/or illustrate the ABA or ABC layering with a special set of superimposable overhead transparencies

5)   X-ray Diffraction by Crystals - use a He-Ne laser and a slide containing eight different arrays of dots to simulate x-ray diffraction experiments

 

i)    Bonding in Solids

1)   Molecular solids

A)     Show models of the CO2(s) and H2O(s) lattices

B)     Show models of benzene and toluene to illustrate the effect of structure and symmetry of molecules on their melting and boiling points

2)   Covalent network solids - show large models of C (graphite) and C (diamond); if desired, you may also request a model of C60 (buckminsterfullerene), subject to availability from Jim Coe

3)   Ionic solids

A)  Display models of NaCl, CsCl, ZnS, CaF2, and TiO2 so students can learn to count the number of ions in a unit cell

B)  Contrast a space-filling model of CsCl with an open lattice model - Cs+ ions occupy the cubic holes in the simple cubic arrangement of Cl- ions

C)  Contrast a space-filling model of NaCl with an open lattice model - Na+ ions occupy octahedral holes in the face-centered cubic arrangement of Cl- ions

D)  Use a model of ZnS to show that Zn2+ ions occupy alternate tetrahedral holes in the face-centered cubic arrangement of S2- ions

E)  Use a model of CaF2 to show that F- ions occupy all the tetrahedral holes in the face-centered cubic arrangement of Ca2+ ions

F)  Display a model of CaCO3

4)   Metallic solids - show models of the Cu, Mg, and Fe lattices (space-filling models showing the cubic close-packing of Cu and the hexagonal close-packing of Mg are also available if desired)

 

Chapter 13 - Properties of Solutions

 

a)   The Solution Process

1)     Add a few crystals of KMnO4 or another crystalline solid to water on the overhead projector so students can observe the process of dissolution


2)   Energy Changes and Solution Formation - demonstrate dramatic differences in heats of solution by dissolving NH4NO3(s) in water in a Ziploc bag to make an instant "cold pack" and CaCl2(s) in water to make an instant "hot pack", then pass the bags around the class

 

b)   Saturated Solutions and Solubility

1)   Add about 50 g NaCl to 100 mL H2O in increments to contrast unsaturated and saturated solutions

2)   Supersaturated Solutions - add a tiny crystal of NaC2H3O2 to a 2 L flask of a supersaturated solution to cause NaC2H3O2*3 H2O to crystallize out leaving almost no liquid - this demonstration is beautiful and dramatic, also the reaction is exothermic; an alternate approach is to pour the solution slowly over a crystal of NaC2H3O2 to build up a column of solid NaC2H3O2*3 H2O

 

c)   Factors Affecting Solubility: Solute-Solvent Interactions

1)   Show the solubility of NH3(g) in H2O due to hydrogen-bonding very effectively with the ammonia fountain demonstration

2)   Mix ethanol and colored water in a specially modified buret to demonstrate their miscibility and negative volume of mixing due to hydrogen-bonding

3)   Add ethanol to colored water in one beaker and add hexane to colored water in another beaker to demonstrate miscibility and immiscibility due to differences in intermolecular forces

4)   Add acetone to a saturated solution of CuSO4(aq) causing CuSO4(s) to crystallize out - the solubility of CuSO4 decreases as the polarity of the solvent is decreased

5)   Demonstrate the maxim that "like dissolves like" by contrasting the solubility of I2(s) and CuCl2(s) in both water and hexane in large test tubes

6)   Pass around a wave-maker, a bottle containing vegetable oil and colored water, to demonstrate an attractive plaything resulting from the immiscibility of two liquids; then relate the phenomenon to the serious consequences of an oil spill

 

d)   Factors Affecting Solubility: Pressure Effects - open a bottle of Club Soda or 7-up, listening for the hiss of CO2 escaping as the seal is broken, then watch as bubbles of CO2 are released from solution once the partial pressure of CO2 over the liquid is diminished

 

e)   Factors Affecting Solubility: Temperature Effects

1)   Heat two flasks, one containing saturated Ca(C2H3O2)2 and the other saturated KNO3; the calcium acetate will crystallize out as the potassium nitrate dissolves

2)   Cool a hot colorless solution of PbI2(aq) to produce a glittering shower of yellow PbI2(s) crystals in a large flask.  The light bulb conductivity tester can be used to observe the changes in conductivity as the solution cools, if desired - the bulb gets dimmer but never goes out completely, since PbI2 is slightly soluble even at room temperature.

 

f)    Colligative Properties: Lowering the Vapor Pressure - two large filter flasks are connected with a manometer and have equal pressures initially - each flask contains 100 mL H2O, and one flask contains a small container of NaCl; tip the container of NaCl(s) over in the flask to show the decrease in vapor pressure of a solution

 

g)   Colligative Properties: Osmosis - small dialysis bags containing equimolar solutions of C3H7OH and Ca(C2H3O2)2 are attached to long glass tubes; immerse the bags in distilled water to illustrate osmosis and to show that osmotic pressure depends on the number of particles in a solution


h)   Colloids

1)     Demonstrate the Tyndall effect and simulate a sunset on the overhead projector by reacting Na2S2O3 with HCl to produce a colloidal suspension of sulfur

2)     Shine a He-Ne laser (or a laser pen) in the darkened classroom to highlight the chalk dust suspended in air

 

Chapter 14 - Chemical Kinetics

 

a)   Introduction to Factors Affecting Reaction Rates - on the overhead projector or in beakers, use the C2O42-/MnO4- and Fe2+/MnO4- reactions to briefly introduce one or more of the four factors affecting rates - nature of the reactants, their concentration, temperature, and catalysis

 

b)   The Dependence of Rate on Concentration

1)   Perform the iodine clock reaction with three different initial concentrations of IO3-

2)   Contrast the rate of combustion of methane-filled soap bubbles with the rate of combustion of a 1:2 mixture of methane and oxygen in a balloon; note that activation energy in the form of a flame is required to initiate both reactions.  If desired, a third balloon containing a 1:2 mixture of methane and air can also be provided.

3)   Contrast the burning of steel wool in air and in a flask charged with O2(g)

 

c)   The Effect of Surface Area on Rate

1)   Contrast the rate of combustion of a pile of lycopodium powder versus the fine powder squirted into a candle flame or versus the exploding can (combustion of the powder in a sealed paint can blows the top off)

2)   Contrast the results of holding a strip of iron in a burner versus squirting powdered iron into the burner flame

3)   Contrast the results of reacting powdered iron or an iron nail with dilute sulfuric acid on the overhead projector or in beakers

4)     Contrast the rate of combustion of a small amount of ethanol in a watchglass with the rate of combustion of ethanol vapor and air in a milk carton

 

d)   Temperature and Rate

1)   Have three students add Alka-seltzer tablets to flasks containing water at different temperatures and quickly seal the flasks with stoppers fitted with balloons, which will inflate at different rates

2)   Immerse Cyalume light sticks in hot and cold water to show variation in rates depending on temperature (subject to availability of light sticks)

3)     Contrast the rate of oxidation of sucrose in the body (as shown by eating 2 M&M's) with the rate of oxidation of sucrose by KClO3 (as shown by dropping 2 M&M's into molten KClO3); body temperature is approximately 37oC, the melting point of KClO3 is 368oC

 

e)   The Collision Model and Activation Energy

1)   Use ball-and-stick models (such as NO, Cl2, NOCl, and CI- or H2, I2, and 2 HI) to show how orientation of molecules influences the effectiveness of a collision

2)     Use large wooden space-filling models of H2 and I2 to show how orientation of molecules influences the effectiveness of a collision.  These models can come apart and be reassembled to give 2 HI models.

3)   Ignite a balloon containing H2 and O2; note that activation energy in the form of a flame is required to initiate this reaction


4)     Use the SN2 machine to demonstrate the transition state in the SN2 reaction between CH3Cl and I-.   Before and after models of  CH3Cl, I-, CH3I, and Cl- can also be provided if desired.

 

f)    Reaction Mechanisms

1)   Introduce the mystery of mechanisms with the Briggs-Rauscher oscillating reaction or the blue bottle reaction

2)   To show the formation of a reaction intermediate, add FeCl3(aq) to Na2S2O3(aq) on the overhead projector; a black intermediate of FeS2O3+ forms and then disappears, leaving colloidal sulfur as the final product

3)   To show the unlikelihood of a termolecular collision, give colored foam balls (or racquet balls) to three students and challenge them to throw the balls so that all three collide simultaneously

 

g)   Catalysis

1)   Demonstrate the catalysis of the H2O2 decomposition of NaK-tartrate with Co2+; adding Co2+ turns the solution pink, but the solution turns dark green as it begins to react vigorously, then the pink color is restored showing regeneration of the catalyst; at this point, use the solution containing the regenerated catalyst to catalyze the same reaction in a second beaker

2)   Use MnO2 to catalyze the decomposition of 30% H2O2, producing a large cloud of water vapor, in the Genie in a Bottle demonstration

3)     Demonstrate the decomposition of 30% H2O2 in the presence of dishwashing liquid, with KI as a catalyst, producing "elephant toothpaste"

 

 

Chapter 15 - Chemical Equilibrium

 

a)   The Concept of Equilibrium - introduce the concept of equilibrium as a dynamic state where the rates of forward and reverse reactions are equal by having two students pour water between two 5 gallon jars (starting with one full, one empty) until equilibrium is achieved

 

b)      The Equilibrium Constant - to demonstrate the law of mass action or shifting the equilibrium position, the demonstration with the two 5 gallon jars can be extended by offering the two students cups of two different sizes

 

c)   Le Châtelier=s Principle: Changes in Reactant or Product Concentration

1)   Apply stress to an Fe3+ + SCN- W FeSCN2+ system in five different ways to show the equilibrium shifts accompanying changes in the concentration of reactants

2)   Show the pH dependence of the CrO42-/Cr2O72+ system

3)   Demonstrate the effects of concentration changes on the Co(H2O)62+/CoCl42- equilibrium

 

d)   Le Châtelier's Principle: Changes in Volume and Pressure - illustrate the effects of pressure changes on the NO2/N2O4 system in a syringe

 

e)   Le Châtelier=s Principle: Changes in Temperature

1)   Immerse sealed tubes of NO2/N2O4 in hot and cold water to show how temperature shifts the equilibrium position and to show the reversibility of the shift; red-brown NO2 predominates at high temperatures and colorless N2O4 at lower temperatures


2)   Demonstrate the temperature dependence of the Co(H2O)62+/CoCl42- equilibrium as shown on page 580 of the text

 

 

Chapter 16 - Acid-Base Equilibria

 

a)   Bronsted-Lowry Acids and Bases

1)   Toss a racquet ball labeled H+ to a student, then hold up your hand so the student throws the ball back; now ask "Do you know what you've done? . . .You've made an acid of yourself!"  This should indelibly imprint the Bronsted-Lowry concept of an acid as a proton donor in your students' minds.

2)   Show ball-and-stick models of Bronsted-Lowry acid and base reactants and products, such as the following:

a)   NH3 and H2O to give NH4+ and OH-

b)   H2O and H2O, to give H3O+ and OH-

 

b)   The pH Scale

1)   Add a chunk of dry ice to a 2 L cylinder containing a basic solution and universal indicator; the dry ice gradually acidifies the solution causing the color to change in the order purple, blue, green, yellow, orange

2)   Have a student wearing a white or light-colored shirt use a straw to blow into water containing universal indicator or bromthymol blue to observe the color change accompanying the reaction CO2 + H2O --> HCO3- + H+

3)   Measuring pH

a)     Use pH paper to test the pH of an acid, a base, and water

b)     Use universal indicator to test the pH of various household substances or familiar acids and bases

c)     Use a pH meter (subject to availability) to test the pH of various household substances or familiar acids and bases

 

c)   Strong Acids and Bases

1)   Contrast the extent of ionization in weak and strong acids and bases using the lightbulb conductivity apparatus

2)   Contrast the rates of reaction of 1 M HCl and 1 M HC2H3O2 with Mg metal or NaHCO3 as shown on page 611 of the text

 

d)   Weak Acids - demonstrate a method for determining Ka of acetic acid: compare the conductivities of 1 M HCl and 1 M HC2H3O2 using two light bulb conductivity devices, then test 1 M HC2H3O2 successively against 0.1 M, 0.01 M, and 0.001 M HCl; when the two bulbs glow with equal intensity, the HCl concentration tells [H+] and [C2H3O2-] for 1 M acetic acid, so Ka can be calculated

 

e)   Acid-Base Properties of Salt Solutions - demonstrate acid-base properties of various salts (i.e., NH4NO3, NaF, NaCl, Na2CO3, Fe(NO3)3, NaHSO4) by dissolving them in water containing universal indicator

 

f)    Lewis Acids and Bases

1)   Continue the racquet ball analogy from a above, this time using a ball labeled with an electron pair; whoever tosses this ball makes an acid of the person who catches it, as that person becomes an electron pair acceptor


2)   Demonstrate the reaction of HCl(g) and NH3(g) to form NH4Cl(s)

3)   Show ball-and-stick models of Lewis acid and base reactants and products, such as the following:

A)  NH3 and BF3 to give F3B:NH3

B)  AlCl3 and Cl-, to give AlCl4-

C)  SnCl4 and 2 Cl-, to give SnCl62-

D)  NH3 and HCl to give NH4+ and Cl-

4)   Hydrolysis of Metal Ions - demonstrate the acidity of various metal nitrates by dissolving them in water containing universal indicator

 

 

Chapter 17 - Additional Aspects of Aqueous Equilibria

 

a)   The Common Ion Effect - add NaC2H3O2 to a solution of acetic acid and universal indicator; the color change accompanying the change in pH shows the equilibrium shift caused by the common ion effect

 

b)   Buffered Solutions

1)   Contrast the buffer capacity of water, 1 M HC2H3O2/NaC2H3O2, and 0.1 M HC2H2O2/NaC2H3O2 by adding increments of 6 M HCl to each in the presence of a mixed indicator; the experiment can be repeated with additions of 6 M NaOH, if desired

2)   Starting with two 1 L graduated cylinders each containing 10 g CaCO3, simultaneously add 1 M acetic acid to one and 1 M acetic acid/sodium acetate to the other, and observe the relative rates of reaction

 

c)   Acid-Base Titrations - use color transparencies from the current text or previous texts to illustrate various titration curves and to show how the shape of the curve and pH at the equivalence point determine the appropriate indicator

 

 

Chapter 19 - Chemical Thermodynamics

 

a)   Spontaneous Processes - display a large box lid containing white balls on one side and colored balls on the other side, then walk around the room carrying the box to demonstrate the spontaneous mixing of the balls.  Is the "reaction" spontaneously reversible?

 

b)      Spontaneity, Enthalpy, and Entropy - two possible demonstrations show that even endothermic reactions can be spontaneous if they have a large positive ΔS

1)   Set a beaker of ice in front of the class, observe it as the lecture progresses

2)   Shake solid Ba(OH)2*8 H2O with solid NH4NO3 to produce an aqueous mixture of Ba(NO3)2(s) and NH3(aq); the reaction is endothermic enough to freeze the flask to a wet block of wood, or a digital thermometer can be used to measure the change in temperature

 

c)   A Molecular Interpretation of Entropy

1)   Use a ball-and-spring model to point out the various possible stretching, bending, and rotational motions

2)     Pass out large rubber bands to a number of students and have them observe the temperature changes that occur when they stretch or relax the rubber band and hold it against their upper lip; use the demonstration to illustrate the relationship between entropy and heat, as stretching the rubber band increases the order of the system


d)   Spontaneity versus Kinetics - you may wish to point out that the rate of a spontaneous process is determined by kinetics, not by the value of ΔG

1)   Use calculations to show that formic acid is thermodynamically unstable and should spontaneously decompose, then display a reagent bottle of formic acid to show that it can be kept on the shelf; although ΔG is negative, the rate of reaction is very slow

2)   Add ethanol to one arm of a tall U-tube containing water; observe the different heights due to different densities, and note that, although the mixing of the two liquids by diffusion should be a spontaneous process, it is also an extremely slow process

 

e)   Gibbs Free Energy and the Equilibrium Constant - hold both ends of a long ribbon so that it assumes the shape of the free energy curve for the course of a spontaneous reaction; vary the height of your hands to show how different values of ΔGo affect the position of equilibrium

 

Chapter 20 - Electrochemistry

 

a)   Oxidation-Reduction Reactions - perform one or more of the following demonstrations, then review redox terminology

1)   React Zn and S to produce ZnS

2)   Burn a piece of Mg ribbon in air

3)   Immerse a strip of Cu in ZnSO4(aq) and a strip of Zn in CuSO4(aq)

4)   Immerse a strip of Al in CuCl2(aq) and a strip of Cu in Al(NO3)3(aq)

5)   Immerse a strip of Zn in Pb(NO3)2(aq) and a strip of Pb in Zn(NO3)2(aq)

6)   Immerse an Fe nail in CuSO4(aq) and a strip of Cu in FeSO4(aq)

7)     Immerse a coil of Cu wire in AgNO3(aq) and a Ag wire in CuSO4(aq)

8)     Burn Mg in dry ice (CO2) to produce MgO and elemental C

9)     Add mossy Zn to HCl(aq); if desired, show a second redox reaction by using a lighted splint to ignite the H2(g) generated by the reaction of Zn and HCl

 

b)   Balancing Oxidation-Reduction Reactions - use one of these demonstrations to show that H+ or OH- is necessary to balance some redox reactions; when the two reactants in 1 and 2 are mixed, the expected product will not appear until acid is added

1)   Reduce pink MnO4- with NO2- in aqueous solution to produce colorless Mn2+

2)   React I- and IO3- in aqueous solution to produce I2; the iodine appears orange in aqueous solution, but if you add dichloroethane it appears pink in the dichloroethane layer

3)   React Al(s) with aqueous base to produce Al(OH)4- (aq) and H2(g)

 

c)   Voltaic Cells - show the spontaneous reaction between Zn(s) and I2(aq) in a test tube, then harness the reaction by separating the Zn and I2 in a voltaic cell and use it to operate a motor to turn a propeller

 

d)   Cell EMF

1)   Set up a Zn|Zn2+||Cu2+|Cu cell in a beaker or on the overhead projector, then use the Fluke Multimeter to get a digital reading of the voltage; this cell can also be attached to a motor to turn a propellor, if desired

2)   Set up cells with Ag|Ag+ versus Zn|Zn2+ and/or Cu|Cu2+ and measure the emf with the Fluke Multimeter; use readings from different cells to rank the various half-cell potentials


e)   Effect of Concentration on Cell EMF

1)     Set up a concentration cell in a tall beaker with 1 M Cu2+ on the bottom and 0.01 M Cu2+ on the top and copper plates immersed in the solutions as electrodes; the voltage read from the Fluke multimeter should be close to 59 mV as predicted by the Nernst equation

2)     Set up a Zn|Zn2+||Cu2+|Cu cell using 1 M ZnSO4 and 0.1 M CuSO4, then use the Fluke Multimeter to get a digital reading of the voltage, which should vary from 1.10 V as predicted by the Nernst equation

 

f)    Batteries - you may wish to pass around a disassembled 9 V battery to show that it is comprised of six 1.5 V cells

 

g)      Electrolysis

1)   Electrolyze Na2SO4(aq) in the Hoffman apparatus, which clearly shows the production of 2 H2(g) for each O2(g); the presence of bromthymol blue shows that H+ and OH- are produced at the electrodes, but the neutral color is restored when the solutions are combined at the end

2)   Electrolyze NaCl(aq) in the Hoffman apparatus - some Cl2(g) as well as O2(g) and H+(aq) are produced at the anode while H2(g) and OH+(aq) are produced at the cathode, so the final solution is basic. (Unfortunately the Cl2(g) produced bleaches the indicator, so the final pH cannot be evaluated visually.)

3)     Electrolyze one of the following in a special cell on the overhead projector; production of H+ or OH- at the electrodes can be detected with an indicator

A)  Electrolysis of CuSO4(aq) or CuCl2(aq) - Cu plates out on the cathode, H+(aq) and O2(g) or H+(aq), O2(g), and Cl2(g) are produced at the anode

B)  Electrolysis of SnCl2(aq) in HCl(aq) - produces H2(g) and crystals of Sn(s) at the cathode and Cl2(g) at the anode

 

8)      Quantitative Aspects of Electrolysis - perform the electrolysis of SnCl2(aq) in HCl(aq) on the overhead as described in g3B above, contrasting the rate of production of Sn(s) using one 9-volt battery versus the rate using two 9-volt batteries in parallel

 

 

Chapter 12 - Modern Materials

(This chapter is not part of the standard syllabus for Chemistry 122, but the demonstrations are provided in case you decide to incorporate some of the material into the course.)

 

a)   Liquid Crystals - demonstrate the color changes initiated by temperature changes on two Mylar sheets covered with black ink and Parker Liquid Crystals; placing a hand on the film or "writing" on it with a damp Kimwipe are two ways to bring about color changes

 

b)   Polymers

1)     You may wish to show the 29-minute World of Chemistry Video "The Age of Polymers" (UTS #3614-2nd program); UTS requires a MINIMUM of 48 hours notice

2)     Make a styrofoam (expanded polystyrene) cup disappear by placing it in a dish of acetone

3)   Nylon - demonstrate the polymerization of hexamethylenediamine with sebacoyl chloride to produce the polyamide, Nylon 6-10


4)   Cross-linking Polymers

A)  See how much water you can add to a super-absorbent disposable diaper, then cut open another diaper to show the super-absorbent powder, Water Lock J-550, which is polysodium acrylate cross-linked with starch; the original polymer results from multiple addition reactions of the alkene functional groups of acrylic acid molecules.  (One diaper holds 1 L of water!)

B)  Take time out to make "Slime", a cross-linked gel produced by mixing solutions of polyvinyl alcohol and borax; use this demo to relate concepts such as polymers and hydrogen-bonding to a commercial product students are familiar with.  (This demonstration requires 24-36 hours notice to prepare the PVA solution.)

 

c)      Ceramics - demonstrate the Meissner effect, the levitation of a tiny magnet above a superconductor

 

 

Chapter 18 - Chemistry of the Environment

(This chapter is not part of the standard syllabus for Chemistry 122, but the demonstrations are provided in case you decide to incorporate some of the material into the course.)

 

a)      Sulfur Compounds and Acid Rain - burn S8 in air enriched with O2 to produce SO2(g), then dissolve the SO2 in H2O containing an indicator to show that an acid is produced (SO2 + H2O --> H2SO3);  this demonstration shows how acid rain results from burning high sulfur coal

 

b)      Nitrogen Oxides and Photochemical Smog - You may wish to show all or selected parts of a 9-minute videotape titled "Oxides of Nitrogen" in lecture or in lab. PLEASE NOTE: showing the tape in classrooms other than 1000 MP and 1015 MP requires a MINIMUM of 48 hours notice. Demonstrations on the tape include the following:

1)   Reaction of Cu with dilute HNO3 to produce NO (about 1 min.)

2)   Reaction of NO with O2 to produce NO2 (about 30 sec.)

3)   Reaction of Cu with concentrated HNO3 to produce NO2 (about 40 sec.)

4) Thermal decomposition of Pb(NO3)2 to produce NO2 (about 50 sec.)

5)   Three redox reactions of NO2 with KI, KMnO4, and H2O (about 3 min.)

6)     Catalytic conversion of NH3 to NO (about 1.5 min.) followed by highlights of equations in the Ostwald process (about 1 min.)

 

 

 

 

 

 

 

 

 

 

 

 

NOTE:  This listing was prepared by Mary H. Bailey to accompany the eighth edition of Chemistry - The Central Science, by Brown, LeMay, and Bursten.