1220 – General Chemistry II

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Chapter 11: Liquids and Intermolecular Forces

• Comparing the Vapor Pressure of Two Liquids
  • Show the effect of intermolecular forces on the vapor pressure of liquids by contrasting the vapor pressures of two isomers, diethyl ether and 1-butanol
  • Comparing the Vapor Pressure of Two Liquids.pdf
• Effect of Pressure on the Melting Point of Ice
  • Hang a wire weighted at both ends over a cylinder of ice; eventually the wire passes through the ice and the weights fall, leaving the ice intact. The ice below the wire melts due to pressure from the weights, and the water above the wire refreezes as the pressure is relieved.
  • Effect of Pressure on MP of ice.pdf
• Halogens
  • Display flasks containing the halogens chlorine, bromine, and iodine.
  • On request, you can order a special flask of bromine that can be frozen in liquid nitrogen
  • Halogens
• Negative Volume of Mixing
  • Because of hydrogen bonding between water and absolute ethanol, when two identical volumes are mixed, the total volume is less than the sum of the two volumes.
  • Negative Volume of Mixing.pdf
• Polarity and Geometry
  • Show the dependence of dipole-dipole forces on geometry by contrasting the effect of a charged rod on streams of H₂O and “CCl₄” (actually hexane) flowing from burets.
  • Polarity and Geometry.pdf
• Triple Point Demo
  • Demonstrate the existence of three phases of CO2 at the triple point by adding crushed dry ice to a clear acrylic tube fitted with a pressure gauge and a release valve.
  • Triple Point Demo.pdf
• Viscosities of Liquids
  • Compare the viscosity of various liquids with a viscosity apparatus and relate the differences to strength of attractive forces.
  • Viscosities of Liquids
• Vapor Pressure Lowering of Solutions
  • Raoult’s Law states that the vapor pressure of a solution is lower than the vapor pressure of the pure solvent.  Use a manometer to compare the vapor pressure of water to that of a sodium sulfate (Na₂SO₄) solution
  • Vapor Pressure Lowering of Solutions.pdf

 

Chapter 13: Properties of Solutions

• Conductivity Testers
  • Sugar and Salt – Use two conductivity testers with light bulbs to contrast the conductivity of d-H₂O, sugar solution, and NaCl (aq).
  • Strong and Weak Acids and Bases – Use two conductivity testers with light bulbs to contrast the conductivity of weak and strong electrolytes: acetic acid and HCl (aq), and/or NH₃ (aq) and NaOH (aq)
  • Conductivity Testers 1.pdf
• Like Dissolves Like
  • Contrast the solubility of I₂(s) and CuCl₂(s) in both water and hexane in large test tubes
  • Like Dissolves Like.pdf
• Miscible and Immiscible Liquids
  • Mix ethanol and colored water in one beaker and hexane and colored water in another to demonstrate miscibility and immiscibility due to differences in the intermolecular forces of alcohols as the size of the alkyl group increases.  The demo can be repeated with butanol and water, to show that, as carbon chain length increases, polarity of alcohols decreases.
  • Miscible and Immiscible.pdf
• Negative Volume of Mixing
  • Mix ethanol and colored water in a graduated cylinder to demonstrate their miscibility and negative volume of mixing due to hydrogen bonding
  • Negative Volume of Mixing.pdf
• Osmotic Pressure
  • small dialysis bags containing equimolar solutions of isopropanol and CaCl₂ 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 solution.
  • Osmotic Pressure.pdf
• Solubility and Temperature
  • Heat two flasks, one containing saturated Ca(C₂H₃O₂)₂ and the other saturated KNO₃; the calcium acetate will crystallize out and the potassium nitrate dissolves
  • Solubilty and Temperature.pdf
• Tyndall Effect
  • Demonstrate the Tyndall effect and simulate a sunset on the overhead projector by reacting Na₂S₂O₃ with HCl to produce a colloidal suspension of sulfur.
  • Tyndall Effect.pdf
• Vapor Pressure Lowering of Solutions
  • Demonstrate the lower vapor pressure of solutions by knocking over a medicine cup of salt in a sealed manometer containing water.
  • Vapor Pressure Lowering of Solutions.pdf

 

Chapter 14: Chemical Kinetics

• Alka Seltzer at Three Temperatures
  • 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
  • Alka Seltzer at Three Temperatures.pdf
• Catalysis of Reaction
  • Demonstrate the catalysis of the H₂O₂ decomposition of NaK-tartrate with Co²+.  Adding Co²+ turns the solution pink, but the solution quickly turns dark green as it begins to react vigorously. At the end of the reaction, the pink color is restored showing regeneration of the catalyst
  • Catalysis of Reaction.pdf
• Combustion of Candy
  • Contrast the rate of oxidation of sucrose in the body (by eating some candy) with the oxidation of sucrose by KClO₃ (as shown by dropping some candy into molten KClO₃, producing steam and a lavender flame.  Body temperature is ~37C, and the melting point of KClO₃ is 368C
  • Combustion of Candy.pdf
• Combustion of Ethanol (Vapor and Liquid)
  • Compare the combustion of ethanol in a small dish to the combustion of ethanol vapors.  When liquids are flammable, their vapors are explosive!
  • Compustion of Ethanol.pdf
• Elephant Toothpaste
  • Demonstrate the decomposition of 30% H2O2 in the presence of dishwashing liquid and KI, producing an upsurge of steaming foam.
• Genie in a Bottle
  • Use MnO2 to catalyze the decomposition of 30% H2O2, producing a large cloud of hot water vapor. The heat generated is intense enough to shrink the 2 L bottle used for the demo
• Iodine Clock
  • Perform the iodine clock reaction with three different initial concentrations of IO₃–
  • Iodine Clock.pdf
• Light Sticks
  • Immerse light sticks in hot and cold water to show variation in rates depending on temperature
  • Light Sticks.pdf
• Lycopodium
  • Contrast the rate of combustion of a pile of lycopodium powder versus the exploding paint can
  •  Blow lycopodium powder into a candle flame
• Reaction Intermediates
  • Add FeCl₃ (aq) to Na₂S₂O₃ (aq) on the overhead projector, giving rise to the black intermediate FeS₂O₃⁺, which forms and disappears, leaving colloidal sulfur as the final product.
  • Reaction Intermediates.pdf

 

Chapter 15: Chemical Equilibrium

• Cobalt Complexes and Temperature
  • Demonstrate effects of concentration and temperature  changes on the Co(H₂O)₆²+/CoCl₄²– equilibrium
  • Cobalt Complexes and Temperature.pdf
• Dynamic Equilibrium
  • Introduce the concept of dynamic equilibrium by having two students bail beads between two clear boxes, one empty, one full, until equilibrium is achieved. This demo helps students view dynamic equilibrium as a state where the rates of forward and reverse reactions are equal, not where the amounts of reactants and products are equal.
  • Dynamic Equilibrium.pdf
• Effect of Temperature on NO₂ ↔ N₂O₄ Equilibrium
  • 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
  • Effect of Temperature on NO2 <-> N2O4 Equilibrium.pdf
• Le Chatelier’s Principle: Iron(III) Thiocyanate Equilibria
  • Apply stress to the Fe³+ + SCN- → FeSCN²+ system in five different ways to show the equilibrium shifts accompanying changes in the concentration of reactants
  • Le Chateliers Principle.pdf

 

Chapter 16: Acid-Base Equilibria

• Acidic and Basic Oxides
  • Dissolve several oxides (CaO, ZnO, CO₂, P₄O₁₀) in water containing universal indicator to show a range of basic and acidic oxides
  • Acidic and Basic Oxides.pdf
• Acidity and Basicity of Salts (Formerly Hydrolysis of Salts)
  • Salts
  • Nitrates
  • Dissolve various salts in water containing Yamada indicator to demonstrate their acidity or alkalinity in solution.
  • Acidity and Basicity of Salts.pdf
• Buffer vs Acid
  • Compare the acid-base reaction of calcium carbonate with both acetic acid and an acetic acid/sodium acetate buffer.
  • Buffer vs. Acid.pdf
• Conductivity Tester Demos
  • Strong and Weak Acids and Bases – Contrast the extent of ionization in weak and strong acids and bases using the lightbulb conductivity apparatus.
  • Conductivity Testers 1.pdf
• Titration of a Triprotic Acid
  • Titrate H₃PO₄ with NaOH, using a combination of indicators to show removal of the first and second proton.
  • Titration of a Triprotic Acid.pdf

 

Chapter 17: Additional Aspects of Chemical Equilibrium

• Amphoteric Hydroxides: Al³⁺ and Fe³⁺
  • add NaOH to samples of Al(NO₃)₃ and Fe(NO₃)₃ to form insoluble metal hydroxides, then add HNO₃ and more NaOH to different samples of each to identify which metal hydroxides are amphoteric.
  • Amphoteric Hydroxides.pdf
• Buffer Capacity Demo
  • Contrast the buffer capacity of water, 1 M CH₃COOH/NaCH₃COO, and 0.1 M CH₃COOH/NaCH₃COO by adding increments of 6 M HCl to each in the presence of an indicator
  • Buffer Demo.pdf
• Common Ion Effect – HCl/NaCl
  •  add concentrated HCl (aq) to saturated NaCl(aq) to cause precipitation of NaCl (s)
  • Common Ion Effect.pdf
• Complex Ion Formation
  • Show color change associated with formation of complex ions. In a tall graduated cylinder of Cu²⁺ or Ni²⁺ aqueous solutions, carefully add 6 M NH₃ to create a layering effect of [M(H₂O)₆²⁺ / M(OH)₂ / [M(NH₃)₄²⁺]
  • Complex Ion Formation.pdf

 

Chapter 19: Chemical Thermodynamics

• Money to Burn
  • Soak a dollar bill in a water-alcohol mixture and then light it with a match; the high specific heat of water keeps the combustion temperature low enough to prevent burning the bill
  • Money to Burn.pdf
• An Endothermic Reaction
  • 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 piece of cardboard.  Alternatively, a digital thermometer can be used to record the temperature change
  • Endothermic Reaction.pdf

 

Chapter 20: Electrochemistry

• 9V battery cut in half (photos)
  • Cross section of a 9V battery
• Balancing a Redox Reaction
  • Reduce pink MnO₄^– with NO₂^– in aqueous solution to produce colorless Mn²⁺. The reaction requires H+ ions, just like in the balanced reaction.
  • Balancing a Redox Reaction.pdf
• Concentration Cell
  • Set up a concentration cell with 1 M Cu2+ on the bottom and 0.01 M Cu2+ on the top with copper plates immersed in the solutions as electrodes; the voltage read from a multimeter should be close to 59 mV as predicted by the Nernst equation
• Copper-Zinc Voltaic Cell
  • Demonstrate a copper/zinc voltaic cell turning a motor to show that a spontaneous reaction can be harnessed to do work.
  • Copper Zinc voltaic cell.pdf
• Electrolysis of Water
  • Electrolyze water (dilute Na₂SO₄ solution with indicator) in the Hoffman apparatus to decompose it into its component elements, hydrogen and oxygen
  • Electrolysis of Water.pdf
• Metal Redox Reactions
  • Copper and Zinc – Immerse a strip of Cu in ZnSO₄ (aq) and compare to a strip of Zn in CuSO₄ (aq).
  • Copper Star Oxidation – Immerse copper wire in AgNO₃ (aq). This is best done on the document camera over a lengthy period of time so students can observe the continuing reaction.
  • Metal Redox Reactions.pdf
• Multiple Oxidation States of Vanadium
  • Shake a solution of ammonium meta-vanadate with a Zn-Hg amalgam to reduce the vanadium from +5 to +4 to +3 to +2 with different colors
    at each stage
  • Multiple Oxidation States of Vanadium.pdf

 

Chapter 21: Nuclear Chemistry

• Detection of Radioactivity
  • Use a Geiger counter to demonstrate the radioactivity (or lack thereof) of several substances, including NaI, NaCl and uranium salts. A sheet of lead is provided to display the ability of lead to block radiation.
  • Fiesta ware and Uranium plates also available
  • Detection of Radioactivity.pdf

Chapter 23: Transition Metals and Coordination Chemistry

• Cobalt and Nickel Ammine Complexes
  • Show the dependence of color on both the metal ion and its oxidation state. Add concentrated ammonia to Ni²⁺ and Co²⁺ solutions to show different colors with the same ligand. Next, shake some of the resulting [Co(NH₃)₆]²⁺ complex with O₂ to shift the oxidation state from Co²⁺ to Co³⁺, to show different colors with different oxidation states.
  • Ni-Co ammine complex.pdf
• Complex Ion Formation Ni and Cu
  • Show color change associated with formation of complex ions. In a tall graduated cylinder of Cu²⁺ or Ni²⁺ aqueous solutions, carefully add 6 M NH₃ to create a layering effect of [M(H₂O)₆²⁺ / M(OH)₂ / [M(NH₃)₄²⁺]
  • Complex Ion Formation.pdf
• Multiple Oxidation States of Vanadium
  • Shake a solution of ammonium meta-vanadate
    with a Zn-Hg amalgam to reduce the vanadium from +5 to +4 to +3 to +2 with different colors
    at each stage
  • Multiple Oxidation States of Vanadium.pdf
• Paramagnetic and Diamagnetic Salts
  • Provide experimental evidence of electron spin by
    bringing a powerful magnet close to suspended test tubes of MnSO₄, FeSO₄, NiSO₄, and
    ZnSO₄ to show the different responses due to different numbers of unpaired electrons
  • Paramagnetic and Diamagnetic Salts.pdf
• Polarizing Filters and Limonene
  • Place small beakers of (R)-(+)-limonene and (S)-(-)-limonene between two polaroid sheets on the overhead projector to show the equal but
    opposite rotation of plane-polarized light by these enantiomers.
  • Polarizing Filters and Limonene.pdf
• Precipitates and Complexes of Nickel
  • Add different amounts of ethylenediamine to beakers of Ni²⁺ to contrast the colors of Ni(H₂O)₆²⁺ and the Ni²⁺ chelate complexes with one, two, and three ethylenediamine molecules.
  • Precipitates and Complexes of Nickel.pdf