JOY OF CHEMISTRY (COMPLETE)

Cobb, Cathy and Fetterolf, Monty L.  The Joy of Chemistry: The Amazing Science of Familiar Things.  Amherst, New York: Prometheus Books, 2005.

This book is 393 pages in length (paperback version published by Prometheus Books).

This book is written to entice anyone interested in conducting their own experiments at home, with or without any advanced background in chemistry.  The activities described here make use of materials and chemicals that can be purchased in regular grocery stores and hardware stores.  As the author said, the experiments can be done in the kitchen, garage, or outside making use of an “adult” chemistry set derived from the kitchen or the tool box.  Unlike many chemistry sets for sale, a bonus of the book is that it gives some background information and explanations in addition to the step-by-step procedure.  In addition, many chapters contain copious detailed historical information as in Chapter 1 recounting the discovery of the electron.

The book starts off with some safety tips, important even though the chemicals and equipment used are of the household type.  It then provides a shopping list of these chemicals and materials.  Lastly, it provides some instructions on how to make up solutions that cannot be purchased using household substances:  for making copper (II) sulfate, iron (III) acetate, and purple cabbage indicator.

In its introductory demonstration, it gives instructions on how the reader might carry out his or her first “bang and splat” activities.  For the bang activity (bottle rocket), the author describes how to generate carbon dioxide gas fast and furious in a loosely corked plastic bottle to generate enough pressure to turn the cork into a projectile (must be done outside).  Carbon dioxide is generated through the reaction of baking soda or sodium bicarbonate and vinegar.  The splat activity involved making a semi-liquid homogeneous mixture of cornstarch and water.  The resulting “oobleck” (in reference to a Dr. Seuss book) has properties of both solid and liquid depending on what forces it is subjected to (fast or slow pulling forces).

Each chapter is preceded by the description of a demonstration or activity making use of the chemicals and materials in the shopping list.  The activities are chosen to stimulate the reader to think about how their observations of physical and/or chemical changes are explained by chemical principles and theories of atomic behavior and chemical reactivity.

Chapter 1 was preceded by the Water Witch activity which demonstrates how charged particles from the hair can be stripped by a plastic object.  When this plastic object is held near a stream of water, the water appears to bend.  In Chapter 1, water’s polar nature is explained by the presence and distribution of electrons.  The chapter is then followed by a further example to illustrate how the newly learned information about atomic and molecular structure and the existence of charged particles make photocopiers work.

Each of the chapter summaries that follow illustrates this sequence of demonstration, chapter explanation, and example.

PART I

INTRODUCTION: Theory, Octaves, and Scales

·         “Not unlike music and literature, chemistry is described in terms of its elements and has a theory based on fundamental principles.”

·         “Chemists think in octaves too” is a reference to the octet rule that governs electron arrangement and behavior within an atom.

Demonstration 1: Water Witch

Take a plastic spoon and rub it in your hair or on a sweater until the spoon acquires a static charge, as evidenced by the attraction of the spoon for the hair or fibers on the sweater.  Turn on a faucet so that there is a very thin stream of water.  Hold the spoon close to the water and you will see a stream of water bend. 

[My variation:  Have two burets, one with water the other with a nonpolar liquid.]

What happened?  Electron transfer from hair or cloth to plastic material which has a stronger attraction for electrons causes the plastic to acquire a negative charge.

[TRY A QUICKLY YANKED ADHESIVE TAPE AS WELL]

CHAPTER 1:  Electrons and Atoms, Elephants and Fleas

In this chapter the author describes the smallness of the atoms and sub-atomic particles and presents the difficulties of our inability to use our sensual and tactile abilities to study it.  Instead, scientists have to look for secondary effects and infer their causes.  A good quote from Rutherford to Chadwick illustrating this point:  “How could you find the Invisible Man in Piccadilly Circus?...[B]y the reactions of those he pushed aside.”

Some notes on the discovery of the electron:

J. J. Thomson is well-known for having first detected the existence of this particle whose amount of charge he was able to measure by the strength of the magnetic field used to bend a beam of these particles.  “By consensus, electrons were assigned a negative charge.”

It was Jean Perrin however in 1909 that gave a more definitive evidence for the existence of atoms.  He observed the motion of pollen particles in water (subsequently named Brownian motion after the botanist Robert Brown) and attributed it to collisions with moving atoms.  “His observations convinced the scientific community of the validity of the atomic model.”

Discovery of the proton:

In his famous experiment, Rutherford in 1910 fired alpha particles at an ultrathin gold foil and observed some of them bouncing back “as if you had fired a 15-inch shell at a piece of tissue paper and it came back and hit you”.  Rutherford had discovered the nucleus of the atom.  Henry Moseley was the one, however, to provide the experimental evidence for the existence of the positively charged particles, protons.

Discovery of the neutron:

Because of its neutral charge, it took some time to discover the presence of neutrons, but Chadwick was able to confirm their existence in 1932 by measuring the “rebound of certain radiation from nitrogen and helium and found it corresponded to a neutral particle with about the same mass as a proton”.

For Example: Protons and Photocopiers

Photocopiers make based on the principle of attraction between material and static charges.  Light passing through the paper causes electrons from a photoconductive material to be ejected.  These ejected electrons then combine with the static charge on the surface of a drum and light and dark images on the paper being copied are then “imprinted” on the drum through the combination of the photoconductively ejected electron and charge on the drum.  As the drum rotates against the toner, the regions of the drum that still have static charge attract the toner material on the drum surface and this toner material on the drum is then transferred to paper with static charge which attracts the toner material away from the drum.  Heat is used to fix the toner material on the paper which comes out as the “copy”.

Demonstration 2: Coppers and Robbers (copied verbatim from book)

(USE IN 1A/1B FOR CHAPTER 4 REDOX REACTIONS AND ACTIVITY OF METALS)

Put on your safety glasses and protective gloves. For the first part of this demonstration, take two clear plastic cups and add tap water to a depth of one inch (2.5 centimeters). Cut one-inch (2.5-centimeter) strands from the aluminum wire and the copper wire. Remove the insulation on the wires, if there is any. Place the aluminum wire in one cup and the copper wire in the other. Using a plastic spoon, add a half teaspoon (2 milliliters) of the lye specified in "Shopping List and Solutions"-the crystal drain cleaner without added aluminum metal. Swirl the cups gently to mix. There will be some small bubbles as the lye dissolves.

After about fifteen or twenty seconds, small streams of tiny bubbles should be seen coming from the aluminum wire but not from the copper wire. Copper and aluminum wire seem similar. Both are used in electrical wiring. But the aluminum visibly dissolves in this caustic solution, and the copper wire does not.

The bubbles you observed rising from the dissolving solving aluminum are hydrogen, which explains why thin strips of aluminum oil are added to some lye-based drain cleaners: to provide agitation to help break up clogs.

Keep your safety glasses on for the second part of this demonstration. Take the copper sulfate solution prepared as outlined in "Shopping List and Solutions" and pour it into a plastic cup to a quarter-inch (0.5-centimeter) or less depth. Place a steel nail and a galvanized nail in the cup, taking care that they do not touch. Both the steel nail and the galvanized nail acquire a copper coating in the copper solution, but the steel nail has a more rapid reaction. Moreover, its coating of copper will look more like copper than the coating on the galvanized nail. Steel nails are mostly iron, and galvanized nails are steel nails that have a corrosion-resisting zinc coating. Both the nails perform the physical function of fastening boards, but their chemical makeup, as determined by the positions of iron and zinc on the periodic table, accounts for the difference in chemical behaviors seen here.

CHAPTER 2: Periodically Speaking

This chapter gives a very brief introduction to and history of the periodic table.  The general layout is explained and it describes the arrangement of the elements in the table according to mass and the number of protons.  It explains how each element can be identified by the unique number of protons it has (which is the same as the atomic number) and defines the number of electrons as being equal to that of the number of protons in a neutral atom.

The electrons’ arrangement into shells and orbitals is also described and explained.  The relationship between elements in the same vertical column is described as being based on having a “similar condition of occupancy of their outmost shell” and defines the term valence.

The diagonal line in the periodic table represents the division between metals and nonmetals, with the elements straddling the diagonal referred to as metalloids, having properties of both metals and nonmetals.

It defines salt as a compound between a metal and nonmetal.

It then goes on to describe electron configuration, the filling of shells, and how this leads to the formation of distinct ions through the addition or removal of electrons.  The proportion of atoms between a metal and a nonmetal in a salt is explained by what ions are formed based on the atom aims to have a completely filled shell giving as examples the formation of KCl and CaCl₂.  It also briefly uses the same principle of filling shells to explain why there are two hydrogen atoms for one oxygen atom in H₂O and the formation of stable diatomic molecules of some elements like oxygen and nitrogen.

The abundance of nitrogen in the atmosphere, compared to oxygen for instance, is explained on the basis of its less chemically active nature.  Yet, despite this, nitrogen can be found in many compounds of salts, amino acids, smog, laughing gas, TNT, etc.  This versatility of nitrogen is attributed to its ability to achieve a filled shell in more than one way.

For Example: Elements of Diversity

The existence of different types of atoms is explained based on their formation just after the big bang.  The first atoms to be formed after the Big Bang were the lighter hydrogen and helium resulting from the condensation of debris from the Big Bang.  These lighter atoms fused together to form larger nuclei when they further condensed under compression from their own mass due to gravity.  See Figure 1.2.7 “The abundance of elements in the sun” for a distribution of the different elements in the sun.

Carbon nuclei formed after hydrogen and helium because anything in between is too unstable to form.  A carbon nucleus can be formed from three helium nuclei.

Figure 1.2.8 shows the distribution of elements in the earth’s crust, similar to the sun but with much fewer hydrogen and helium atoms which are light.  Aluminum and silicon are well represented than the heavier iron and nickel because the lighter atoms “sifted” upward while the heavier ones “sifted” toward the interior.

In Figure 1.2.9, the abundance of elements in the human body is shown.  Oxygen is the most abundant [I can only guess in terms of atom count?] followed by carbon, hydrogen, nitrogen, calcium, phosphorus, potassium, sulfur, sodium, and chlorine.

Demonstration 3: Stop-and-Go Chemistry (copied verbatim from book)

(USE IN 1A FOR CHAPTER 4 REDOX REACTIONS AND ACTIVITY OF METALS)

Don the safety glasses. Ladle two tablespoons (30 milliliters) of iron acetate solution (the vinegar solution with dissolved steel wool described in the "Shopping List and Solutions") onto a plastic-coated coated paper plate. Add a teaspoon (5 milliliters) of household ammonia. The resulting solution should turn from a light orange-brown to a dark green. Now drop in a teaspoon (5 milliliters) of hydrogen peroxide. The resulting solution is a deep blood-clot red. What happened? Iron, like nitrogen, can tolerate different numbers of electrons associated with its nucleus. The oxidation state of an element in a given environment-be it a free element, an ion, or in a compound-is an expression of the element's electron load in that environment. In the current demonstration, iron is changing between two different oxidation states: the ferric ion (Fe³⁺, or iron with a positive three charge, which means it is missing three electrons) and the ferrous ion (Fe²⁺, or iron with a positive two charge, which means it is missing two electrons). The agent causing the change in the iron is the peroxide.

ferrous iron + hydrogen peroxide à  ferric iron + water + oxygen

Iron is changed from the ferrous ion to the ferric ion because the peroxide attracts the electrons away from the iron. Hydrogen peroxide uses the extra electrons to change into water and oxygen gas, which bubbles out of the solution during the reaction. The compound formed between ferrous iron and ammonia is green. When the peroxide changes the ferrous iron to the ferric iron, the ammonia, still in solution and unchanged, forms a red compound with ferric iron. Reactions such as these, where elements gain or lose electrons, are called reduction (if they gain electrons) or oxidation (if they lose electrons). Because the reactions must happen in tandem, the combined result is called a redox reaction, a class of reactions of current and historical importance.

CHAPTER 3: Reasons, Reactions, and Redox

[Use in 30A and 1A]

This chapter describes and explains the following chemical terms and phenomena:

• Metals are normally found bound to either chlorine, sulfur, and most frequently to oxygen.  In the smelting process, the pure metal is extracted by heating it with something that has more affinity for the nonmetal.  Carbon in charcoal is the most common substance used for this process.
• Atoms lose electrons and are oxidized in the process of oxidation.  Oxidation was the term used for this process because in the smelting process making use of charcoal to reduce the metal, the carbon combines with a nonmetal which is usually oxygen.  Although these terms arose from the word oxygen, oxidation does not always require the participation of oxygen.
• To demonstrate this, the author suggests taking a paper clip and dipping it in a copper (II) sulfate solution.  Copper coats the paper clip as copper ions in solution take away electrons from the metal of the paper clip [aluminum or iron?].
• When atoms get reduced, they gain electrons in the process called reduction.
• Origin of the term reduction:  The term reduction arose from the use of redox in the smelting of iron.  When metallic iron is smelted from its ore (iron (II) and (III) oxide), the smelted iron has a lower or reduced weight (from the removal of oxygen).

For Example: Firefighters and the Chemistry of Combustion

• Combustion proceeds through a redox reaction.  Combustion is the process by which the carbon atoms and the hydrogen atoms are oxidized as they combine with oxygen.  Oxygen, in turn, is reduced from a neutral element to one that has two extra electrons.
• DEMO:  Take two glasses.  Fill one with 20 mL of water and add 2 drops of bromothymol blue (fish tank indicator).  The water should have a light green-blue color.  Hold the empty glass over a burning candle until the flame burns out.  Invert the empty glass and immediately pour the water-indicator solution into this empty glass. The color should change from green-blue to yellow.  CO₂ forms an acidic solution with water.
• The rest of the chapter was taken up by the author’s explanation of how “blow-ups” occur during a raging fire.  The decreased pressure due to consumption of oxygen from the air (20%) and the updraft caused by the heated air above the fire result in convection currents and pressure differences.  These considerable pressure differences cause a whirl of fire to move at nearly the speed of the wind causing a blow up.

Demonstration 4: Purple Cabbage Indicator

·         This is a well-known demonstration of using extracting purple cabbage juice by boiling to use as an acid-base indicator:

·         The author describes the colors resulting when three different indicators are added to the same four household chemicals:

Swimming-pool phenol red indicator

water                reddish orange

vinegar            yellow

ammonia           reddish violet

baking soda      pink

Fish-tank bromothymol blue indicator

water                greenish blue

vinegar                         pale yellow

ammonia           pale blue

baking soda      blue

Purple-cabbage indicator

water                purple

vinegar                         shocking pink

ammonia           green

baking soda      teal blue

CHAPTER 4: The Basic Stuff

·         In cake-making, mixing sour cream (containing acid) and baking soda (sodium bicarbonate, a base) causes the formation of gaseous carbon dioxide which causes cake dough to rise to give it a fluffy and light texture.

·         Baking powder contains tartaric acid which reactions with baking soda to form gaseous carbon dioxide.  Baking soda is a leavening agent.

·         Thos chapter explains the chemical difference between acids and bases.  It talks about neutralization reaction and how this process “negates” the properties of one substance by another such as in the neutralization reaction between acids and bases.

·         It gives examples of acid-base reactions that can be observed in everyday life, in particular in cooking as described above and the dissolving of limestone and marble when exposed to acid rain.

·         It then describes the role of buffers in the body to protect it against dramatic changes in pH due to ingested acids and bases or through uncontrolled respiration.  To demonstrate the action of buffers, the author suggests using whole milk:

·         Take 4 small clear containers.  To two of them, add 120 mL each of whole milk.  To the other two, add 120 mL of water.  Add several drops of the phenol red indicator from the swimming pool test kit to each container.  Add a drop of baking soda solution to one water sample and one milk sample.  The baking soda and water solution should turn a shocking pink color and the milk should stay pale yellow.

·         The chapter then proceeds to explain the role of carbon dioxide in helping buffer the blood.  CO₂ combines with water to make carbonic acid.  Another buffering component is bicarbonate, a base produced by the kidneys.  These two work to neutralize any added acid or base to the blood to prevent the pH from changing too much.

·         The pH scale is explained as a quantitative way to express the level of acidity in a given substance.  The nonlinear correlation of acidity and alkalinity to pH values is explained:  a solution that has a pH of 3 is ten times more acidic than a solution that has a pH of 4.

For Example: Swimming Pool pH

Chlorine is added to water in swimming pools to control the growth of algae, bacteria, and mosquitoes.  The bleach normally used is sodium hypochlorite.  Hypochlorite is a basic substance.  Bleach works best against bacteria when the water is acidic.  Acidic water, however, can be irritating to the skin and corrosive to metals.  To control the acidity, swimming pool water is buffered by the weak base hypochlorite from the bleach and a weak acid.  Too much base can result in the formation of “scale” or carbonate precipitates from calcium and magnesium ions, particularly prevalent in hard water.

Demonstration 5: Blue Blob, Black Ink

·         In this demonstration activity, the formation of a blue blob (copper (II) carbonate) is observed when copper (II) sulfate solution is added to sodium bicarbonate.

·         When 5 mL of hydrogen peroxide is added to 60 mL of  iron acetate solution, a reddish brown solution results as the ferrous ion is oxidized to the ferric ion.  Further addition of 120 mL of cold, brewed brown tea, a mushy black precipitate forms which was historically used as black ink.

CHAPTER 5: Chemical Partners

·         In this chapter, precipitation reaction is described and explained.  To explain why some ions have stronger attraction for each than others, the author starts by explaining the concept of electronegativity.

·         The authors also repeats for the readers the explanation for the observed pattern of electronegativity across from left to right (increasing) and down a period (decreasing) alluding to the phenomenon of reduced effective nuclear charge due to shielding by inner shell electrons. 

·         Fluorine is the most electronegative element because of its size and the resulting high charge density that results from the arrangement of its electrons (only two electrons shielding the 7 in shell 2 from a nuclear charge of 9 protons) and their proximity to the nucleus.  Hydrogen, even though it has a more naked proton has an electronegativity only comparable to that of phosphorus.  Its lower electronegativity is attributed to the fact that an added electron will have to be concentrated in such a small volume (the authors use the analogy of carrying a 20-lb dog food bag with its weight distributed through more volume than a 20-lb cannonball).

·         Molecular polarity results due to differences in electronegativity.  In bonded atoms, the more electronegative element generally pulls the electron density closer to it resulting in an uneven distribution of electron density throughout the molecule with one side more electron-heavy than another side.

·         In H2O, oxygen is more electronegative than H so electron density accumulates at that end.  This allows this negative end of the water molecule to pull the positive ion and the positive end (the H atoms) to pull on a negative ion causing the ions to separate and the salt to dissolve. 

·         Second column metals form positive two charge ions.  These are harder to pull (higher charge density) from the negative ion so their salts tend to be less soluble.

·         To illustrate the power of the polar molecules of water to dissolve salts, the author suggests comparing the solubility of table salt in water and in salad oil or mineral oil.

·         Other substances may affect the solubility of a salt in water.  Calcium carbonate in chalk does not dissolve very well in water but sodium bicarbonate does.  When vinegar is added, however, calcium carbonate dissolves as it reacts with the acid.

·         Examples of precipitation:

o   Gout is a condition in which uric acid precipitates collect in the joints.

o   Gallstones and kidney stones are precipitates.

o   Soap scum

o   Precipitation can be useful for removing toxic metals in water.  It can be precipitate using the appropriate anion and the solid can then be filtered from the solution.

For Example: Hard Water, Soft Water

·         Soap scum formation is even more common in hard water that contains more calcium and magnesium ions.  Phosphates were added to detergent so that the calcium preferentially combines with it rather than the soap ions.  Phosphates however led to abundant growth of organisms that clogged waste streams and so were eventually removed from detergent.

·         Water softeners are used to remove excess calcium and magnesium ions in hard water.  The hard water is streamed through a material saturated with sodium salts.  The sodium salts dissolve in the water while the calcium precipitates with the negative ions left behind.

Demonstration 6: Bond. Chemical Bond.  (copied verbatim from book)

[USE IN 1A AND 30A]

·         Take one of the clear plastic bottles suggested in the "Shopping List and Solutions," put on your safety glasses, and pour in about two inches (5 centimeters) of mineral oil. Add about two inches (5 centimeters) of copper sulfate solution. Put the cap on the bottle and shake it. The mixture should separate into two layers with the copper solution layer remaining bright blue and the mineral oil layer remaining clear.

·         Add a quarter teaspoon (1 milliliter) of tincture of iodine. Put the cap on and shake. When you stop shaking, the two layers should separate again, but this time the mineral oil layer should appear violet. Tincture of iodine contains both I^- and I₂ but mostly a complex of the two, which is brown. The copper ions in the copper sulfate oxidize I- to the molecular form I₂, which ruins the complex. The molecule I₂is violet and soluble in oils such as mineral oil.'

·         Now add a teaspoon (5 milliliters) of household ammonia. Put the cap on and shake. This time when the layers separate, the copper layer should be a dark royal blue because the copper is now forming strong partnerships ships with ammonia and it is no longer available to the iodine.

·         Some iodine reverts to I-, the complex re-forms, and the mineral oil clears up. You may need to add several portions of ammonia, but eventually the mineral oil clears. The molecule I₂ dissolves in the mineral oil, but the ion I- does not. The ion I- dissolves in water but not in mineral oil. Water, a polar molecule, is attracted to the ion I-, builds a cage around it, and pulls I- into the water solution (otherwise known as aqueous solution). The molecule I₂ has no polar end to attract water molecules, so the water molecules are more attracted to each other and squeeze I₂ out of solution. The mineral oil accepts I₂ because mineral oil is a nonpolar molecule and has no cage-building building propensities. Earlier we said that salts were formed when positive ions are attracted to negative ions and that these materials form a bond. Now we need to add that this type of bond is, specifically, an ionic bond. This special designation is necessary because not all materials are salts and not all chemical bonds are ionic. The bonding that occurs in non-salty salty materials, such as I₂, is called covalent. The differences between the types of bonding and the materials they produce deserve a chapter of their own, such as the one that follows.

CHAPTER 6:  The Tie that Binds the Chemicals that Bond

·         In this chapter, the authors explain the difference between ionic, covalent, and metallic bonds.

·         They emphasize that most bonds have some ionic, covalent, and even metallic character and that each bond is difficult to categorize uniquely as one of the three.

·         The periodic table can be used to predict what elements will form ionic compounds: a metal and a nonmetal.

·         When two nonmetallic elements share electrons, a covalent bond forms.

·         Analogy for what chemical bonds are:  attractions that can “be seen as negative electrons running interference between positively charged nuclei”.  In a chemical bond, the electrons between the two atoms attract both nuclei of the atoms and act as the glue.

·         Bonding orbitals normally contain electron density between the two nuclei.  When the electron density is outside and away from the nuclei, the bonding is reduced and these electrons occupy an antibonding orbital.

·         When electrons in a bond absorb energy from the light, for example, they can be promoted to an antibonding orbital which may cause the bond to break.

·         Bonding in solids:

·         Electrons in a bonding orbital between pairs of nuclei form a valence band.

·         A conduction band is an expansive orbital that extends throughout the molecule.  Electrons occupying conduction bands can move from one nucleus to another.

·         Substances that have a lot of electrons in the conduction band are conductors while those with no electrons in the conduction band are insulators.

For Example: Semiconductors

·         Semiconductor material generally has no electrons in the conduction band.  However, because of the small energy gap between the valence band and the conduction band, excitation of the electrons with a relatively small amount of energy can cause the material to conduct.  This makes semiconductors valuable in that their conductivity can be controlled by certain conditions imposed on the material.

·         Material can be converted into semiconductors by doping.  For example, silicon is not an intrinsic or natural semiconductor.  Phosphorus and aluminum are used as dopants because they are on either side of silicon in the periodic table, phosphorus has one more electron and aluminum has one fewer electron.  The extra electron in phosphorus which occupies a conduction band can be made mobile causing conduction in the material when hooked up to a battery.  Doping with aluminum creates a hole which similarly drives the movement of electrons under an electric field.

·         When doped with atoms that donate electrons, silicon is referred to as an n-type semiconductor.  When doped with atoms that have fewer electrons, silicon is called a p-type semiconductor.

·         When an n-type material is attached to a p-type material, a diode is formed.  A diode is device that allows the flow of current in only one direction: from the n-type to the p-type when the negative electrode is attached to the n-type side.    When the negative electrode is attached to the p-type side, the holes are sandwiched by negative charges and there is no net flow.

·         When an npn or pnp arrangement is constructed, conductivity can be switched on by either adding electrons to the p to turn npn to nnn or taking away electrons from the n to turn pnp to ppp.   In this way, “a small current can control a large current, switching it on and off with a little electron flow”.  These semiconductor switches are called transistors.  These transistors can then be turned on (1) or off (O) and using Boolean logic (AND, OR, NOT) and if/then logic, computers can then be programmed (switch on and off, etc.) to execute certain compounds.  The example given by the authors is:

·         “If the switch is on for the letter A, AND the switch is on for the printer, then the printer will print the letter A.”

Demonstration 7: Gemstone Chemistry (copied verbatim from book)

Into each large bag, place one level plastic teaspoonful (5 milliliters) of cream of tartar and one-half cup of hydrogen peroxide (60 milliliters). Mix these ingredients well by kneading the outsides of each bag with your fingers. Do not seal the bags yet.

Into each small sandwich-sized bag, pour a teaspoon (5 milliliters) of copper sulfate solution (prepared as described in the "Shopping List and Solutions"). Seal these bags completely, and place one into each of the two large bags.

In one of the big bags, open the little bag from the outside, without opening the big bag. Pour the copper solution from the little bag into the cream of tartar-hydrogen peroxide mixture in the big bag. The reaction begins slowly, but within ten seconds it should be fizzing away. The dramatic sky-blue copper tartarate eventually gives way to a lime-green mixture of the tartarate compound with orange copper oxide.

The reaction generates enough heat that you can observe condensation on the inside of the big bag and feel warmth on the outside. The impressive colors produced by the reaction are reminiscent of gemstones stones for a very good reason: the colors of gemstones such as azurite and turquoise derive from copper salts. Various copper salts appear as dashes and streaks of color in polished stone.

CHAPTER 7: Sticking to Principles

·         Two important chemical laws are described in this chapter:  law of conservation of mass and the law of definite proportions.

·         Law of definite proportions:  the formula for any one material is set and invariable.

·         Isomers: compounds that have the same number and type of elements but arranged in a different order. Some interesting isomers:

o   HONC is fulminic acid, used to make explosives (thus, one says, “to fulminate” or to verbally rage)

o   HOCN is cyanic acid, used to make poisonous cyanates

o   HCNO is isocyanic acid, starting material for making organic materials

·         Chemical reaction equations, when balanced, follow the law of conservation of mass, important in making chemicals in a lab but critical when manufacturing truckloads of sulfuric acid.

For Example: Engineering – Chemical , That is

·         Most of the chemical industry manufacture the following 4 chemicals: sulfuric acid, phosphoric acid, sodium hydroxide, and sodium chloride.

·         Sulfuric acid is produced in largest volume in the world:

o   Used to be called oil of vitriol because of its oily, syrupy appearance.

o   Will attack plastic, wood, skin, mucous membrane, and most metals in concentrated form but can be stored in glass.

o   Used in car batteries

o   Largest single use is fertilizers, especially in making superphosphate, a sulfuric acid/phosphate-containing rock.

·         Sodium hydroxide or lye:

·         Used in the production of soap, textiles, petroleum products, dye, detergent, and paper

·         Produced by running an electric current through a mixture of water and NaCl through the reaction 2H₂O + 2NaCl à 2NaOH + Cl₂ + H₂

·         Chemical engineers have to worry about balancing coefficients to be able to predict how much gases will need to be contained when produced.  According to the author, a ton of NaOH produced is accompanied by 1 ¾ tons of gases.  For every cubic foot of NaOH, about 2,000 cubic feet of gases are generated.

Demonstration 8: Layer Upon Layer (copied verbatim from book)

After putting on safety glasses, take a small clear glass and pour in glycerin (suggested for purchase in the "Shopping List and Solutions") until the glass is about an inch (2.5 centimeters) full. Take some water dyed with food coloring and gently pour about the same height onto the glycerin. You should observe two distinct layers in the glass. Add a layer of canola oil to this and observe the three distinct layers.

Stir the mixture. The colored water will mix with the glycerin, but the canola oil will still settle into a separate layer once agitation has ceased. It may take several minutes for the layers to separate.

Add a layer of liquid soap to the solution and watch it settle between the glycerin-water layer and the canola-oil layer. Now, when the combined layers are mixed (gently so as not to cause bubbles), the soap will cause the canola oil to blend with the glycerin. The result is a uniform mixture. The mixture may eventually reseparate into different layers, but it takes longer than when the detergent wasn't present.

The solutions initially separated because of a difference in intermolecular attraction, a type of intermolecular force. These attractions occur between molecules, as distinguished from intramolecular forces-chemical bonds-which bind the nuclei within a molecule.

·         “While chemical bonds determine the structure of a single molecule, it is intermolecular forces that will determine how this molecule will behave around other molecules.”  The behavior of molecules with respect to the other molecules is what determines their macroscopic properties.

CHAPTER 8: Slipping and Sliding, Intermolecularly

·         This chapter begins by discussing the differences between intra- and intermolecular bonds.  One particular passage I like is, “Breaking a chemical bond creates a new material that reacts differently with other chemicals.  Breaking an intermolecular attraction does not change the chemical formula or chemical reactivity.  But intermolecular forces can determine how quickly a reaction will occur – or if it can ever occur.  And intermolecular forces determine whether a material will be a liquid, a solid, or a gas at a given temperature…”

·         Another is, “If it were not for intermolecular attractions, our bodies would vaporize into gases, and if it were not for intermolecular repulsions, we would collapse into unimpressive puddles”.

·         Quantitatively understanding, measuring, and predicting the intermolecular forces in molecules require a very challenging analysis.  Molecular shapes are not static and they are always in constant motion relative to themselves and other molecules.

·         The author lists the following as the three types of attractions between molecules:  positive ion – negative ion attraction, dipole – dipole attraction, and dispersion forces.

·         These intermolecular forces manifest themselves in phenomena like friction, cohesion (including surface tension), adhesion, capillary action, solubility, and soap action.

·         Adhesion is the attraction between different materials due to intermolecular attractive forces and the force on which the adage “like dissolves like” depend.  It is the basis for capillary action.  [DEMO: PAPER TOWEL SOAKS UP WATER BETTER THAN IT SOAKS UP OIL]

·         Applications: 

·         antistatic dryer sheets add a very thin layer of oil to the clothes being dried to prevent stripping of electrons.  When placed in a pan with baked-on grease, they help remove the grease because of the nonpolar oily substance on their surface.

·         Cooking oil helps remove stuck-on hard bandages because they help dissolve the nonpolar glue.

For Example: The Physics of the Flush Toilet

Flushing the toilet works by siphoning water over a U-shaped hill.  The difference in pressure between the filled bowl and the empty pipe pushes the water over the top of the U-hill.  The intermolecular forces between water molecules are what keep the water flowing over the hill until all the water has been siphoned down the pipe.

Demonstration 9: Concentrating on Color, Salty Dog

·         In the first part of this demonstration, two different amounts of copper (II) sulfate are added to two cups containing the same amount of water.  The intensity of color in each cup is observed and correlated with the concentration.

·         In the second activity, a pinch of salt is added to each of two cups, one containing just water and the other containing a saturated solution of sodium bicarbonate or baking soda.  The saturated solution is created by adding enough baking soda to water until no more dissolves.  The liquid portion (saturated with baking soda) is then decanted and separated from any undissolved baking soda powder.  Upon addition of a pinch of salt with stirring to each, it is observed that the salt dissolves completely in the pure water liquid but turn the saturated solution cloudy.  This phenomenon is called salting out.  The less soluble sodium bicarbonate precipitates out of solution, replaced by the more soluble sodium chloride table salt in solution.  [USE IN 1B ON SOLUBILITY?]  This process, called salting out, is used in industrial applications to selectively move substances in and out of solution.   It has also been used in soap making.  Large amounts of salt are added to a soap solution to force solid soap precipitate out.

·         A third activity suggested is to demonstrate the “precipitation” or more like the gellification of liquid soap:  Sprinkle table salt on a 1/8 – ¼ inch layer of liquid soap in a glass resulting in a gelatinous solid.

CHAPTER 9: Concentration – On Being Alone Together

·         Many everyday reactions take place in solution: in cooking, in the oceans, in the rain, in the body, in swimming pools, in the atmosphere, in water sanitation ponds, etc.  Being dissolved in solution allows the reacting particles to come together more efficiently because of their mobility and separated state.

·         Solutions can be made out of mixing solid substances:

o   Bronze = alloy of copper and tin

o   Brass = alloy of copper and zinc

o   Steel = alloy of carbon and iron

·         Metabolic reactions in the body take place in solution. Some of these metabolic reactants are gases.  The solubility of gases in a liquid like water depend on temperature and pressure:

o   The higher the temperature, the lower the solubility of the gas.

o   The higher the pressure of the gas (specifically the same gas) above the surface of the liquid, the higher the solubility of the same gas in the liquid.

·         The dependence of solubility on pressure is a critical factor in how gas exchange takes place during metabolism, as described by the author: “To carry out metabolic functions, the body depends on the solubility of gases in blood, which, in turn, depends on the pressure of that gas and not any other.  So the solubility of oxygen in blood depends on the pressure of oxygen in inhaled air, and the solubility of oxygen in the cells depends on the pressure of oxygen in the cell.  As metabolism takes place in the cell, oxygen is used up and carbon dioxide is produced.  When oxygen-rich blood reaches the cell, oxygen passes from the blood – where the pressure is relatively high – to the cell where it is relatively low.  Carbon dioxide passes from the cell where its pressure is high to the blood where its pressure is low”.

·         Application of gas solubility to diving:  Divers who go down quite some depth underwater experience increased pressure on their lungs due to the water.  They need to breathe in pressurized air, at a pressure equal to or higher than the external pressure for air to enter the lungs.  At higher pressures of nitrogen in the lungs, more nitrogen will dissolve in the blood.  It is non-toxic.  However, if the diver ascends too fast, the rapid decrease in pressure may cause nitrogen gas to bubble out of the blood as it experiences reduced pressure.  This may cause severe pain that the resulting condition is sometimes referred to as the “bends” as sufferers are known to bend over in pain.  Helium is sometimes used as a substitute for nitrogen because helium is less soluble than nitrogen.  [Nitrox is also used which contains of a higher oxygen to nitrogen mole ratio (order of few percent) to reduce the nitrogen saturation (lower nitrogen content, lower pressure, reduce probability of narcossis) and the possibility of nitrogen bubbling out and causing the bends.]

·         Some everyday measures of concentration:

·         Proof:  a 76-proof whiskey has a 38% v/v alcohol content

·         Humidity:  measure of the amount of water vapor in the air.  Relative humidity:  amount of water in the air relative to the maximum amount of water that the air can hold at the same temperature and pressure conditions.

·         Grams per milliliter of salt or sugar in water in some common medicines

·         Moles/ liter: used by chemists.

·         The author also refers to density as a measure of concentration.

For Example: What Body, What Aroma, What Chemistry

·         The qualities of wine are measured by its aroma, bouquet, nose, astringency, finish, texture, and body.

·         Body is associate with the “perceived density of the wine in the mouth and depends in good part on the amount of alcohol in the wine”.  The amount of alcohol is dependent on the amount of sugar that is fermented.  To predict the amount of alcohol that will be produced by fermentation, the sugar content of must (the beginning mixture from grapes used in wine making) is measured by determining its density.

Demonstration 10: The Soda Bottle Crunch

·         Prepare a tub of ice water.  Fill a plastic soda with hot tap water and immerse in the ice water for several seconds.  Pour out the tap water, quickly replace the cap, and place back in the ice water.  Within a few seconds, the bottle implodes and crumbles under the weight of the atmospheric pressure.

·         Adding hot tap water heats the wall and the air inside the bottle.  Upon immersing in cold water, the air cools and condenses and the pressure inside decreases.  When the hot water is dumped out and the bottle is recapped and placed back in the ice water, the residual gas is further cooled causing decreased pressure which allows the atmosphere to crush the bottle.

CHAPTER 10: It’s a Gas

·         The word gas was coined by Johannes van Helmont in the mid-1600’s, probably from the word chaos.

·         This chapter describes the behavior of gas through 4 metrics:  temperature, pressure, volume, and amount and relates them to atomic and molecular behavior through the kinetic theory of gases.

·         Knowledge of and predictability of gas behavior are important to chemical engineers, deep sea divers, wine makers, etc.

·         We are not crushed by atmospheric pressure because we the pressure exerted on us is equally distributed on all sides.

·         Boyle’s law or the inverse relationship between pressure and gas can be demonstrated by pushing the plunger on a basting syringe (with tip closed) containing gas:  as the volume gets smaller, it becomes harder to push on the plunger because of the increased pressure of the trapped gas inside the syringe.

·          An interesting historical note is the misnaming of Charles’ law.  The direct relationship between temperature and volume was actually first determined by Joseph Louis Gay-Lussac.

·         Fahrenheit and Celsius temperature needed to be rescaled because both include negative temperatures in their domains.  Lord Kelvin proposed a scale that has as a minimum absolute zero correlated with zero volume.  The Fahrenheit and Celsius values of absolute zero temperature or 0 Kelvin are extrapolated to the point that the temperature at which volume becomes zero:  -459.67 F and -273.15 C.

·         The ideal gas equation was also mentioned and the term ideal was explained (the system is idealized to that in which the particles experience no intermolecular forces). Gas behavior will deviate from that predicted from the ideal gas equation away from a range of pressure and temperature values.

For Example:  Gutsy Gas

·         Flatulence is a mixture of gases produced from bacterial metabolism of complex sugars (e.g. cellulose in beans and peas).  These product gases themselves, carbon dioxide, methane, and hydrogen, are odorless but do get contaminated with the more odorous sulfur-containing molecules as they travel through the intestinal tract.

Demonstration 11: It’s in the Air

·         Just like solutions, the gas state also provides the mobility and means of contact to make reactions happen faster.

·         Diffusion is the term used to refer to the mobility of gas molecules or atoms in the gas state.

·         In this demonstration, a small amount of household ammonia solution is placed in a bowl and diffusion is demonstrated by the smell reaching a person in the far corner of a room.

·         Application to biology: Insects rely on diffusion to detect gaseous molecules of communication.  The male gypsy moth can detect a burst of a few hundred molecules of a sex pheromone from a female gypsy moth from three miles away.

·         Demonstration of rate of diffusion of two gases with different masses:

·         Equal volumes of household ammonia and vinegar are placed in two separate glasses. A piece of paper towel with a few drops of phenol red is placed over the vinegar.  After about 3 minutes, the red color changes to a yellow color as gaseous acetic acid molecules make contact with the phenol red in the paper towel.  The same paper towel is then draped over the glass with ammonia.  The towel turns red again but much more quickly.  [In this case however, the ammonia has been sitting there for a while and so the space above the liquid must be saturated with ammonia molecules after waiting three minutes for the vinegar.  Two things to try:  do the ammonia first then the vinegar and/or prepare the ammonia fresh from the bottle before draping the yellowed paper towel over it.]

CHAPTER 11:  When Gases Put on Airs

·         In solutions, solvent cages and viscosity due to intermolecular forces may slow particles from making contact.  But, the same attractions may also slow their pulling away from each other which can facilitate reactions as well.

·         In the liquid phase, particles are in contact longer than they do in the gas phase.  Moreover, because gas phase collisions are quite energetic, the particles fly apart as soon as they collide leaving little time for contact to facilitate any reaction.

·         Nevertheless, many gas phase reactions are explosive, implying very fast reaction rates, and these reactions generally need a spark to get started.  A primary initiator is the production of radicals when a spark is applied to gaseous molecules.

·         Free radicals formed by spark increase the probability of a successful collision leading to a reaction.

·         In an internal combustion engine, gasoline fuel is sprayed as a fine mist which then mixes with air.  This gasoline vapor – air mixture is then compressed and sparked to initiate the combustion. The temperature during this event is so high that the nitrogen and oxygen molecules in the mixture react to form NOx’s.

·         Gas-phase explosions usually proceed through a chain reaction initiated by radical formation.  The degree of explosiveness depends on the pressure of the gas mixtures:

o   At low pressures, the explosion fizzles out because the chain reaction cannot be sustained by the low density of reacting molecules.

o   At moderate pressures, the chain reaction can be sustained such that the production of heat and gas can be maintained at high rates for an explosive event.

o   At high pressures, the reaction may become less explosive again because of quenching in which the radicals lose energy as they collide with more molecules in a high density environment.

·         Solid-to-gas reactions can also produce explosions as long as high energy gas molecules are produced at very high rates causing the gas to expand rapidly and/or produce very high pressures.  The rapid conversion of solid nitrates, carbon and sulfur in gunpowder to gaseous NOx’s, carbon dioxide, and sulfur oxides results in a very fast energetic expansion that can project a bullet at speeds around 1000 feet/second.  The rapid decomposition of solid sodium azide to gaseous nitrogen upon sparking is what causes the explosive inflation of air bags.

·         Because of the dispersed nature of atoms and molecules in gaseous substances, there are certain steps required to facilitate reaction between these particles:  the pressure has to be high, the temperature has to be high, and, in many cases, a catalyst is used.  A textbook example of a reaction requiring these facilitators is the Haber process used in the production of ammonia from gaseous nitrogen and hydrogen.  Haber’s very important contribution is the discovery of the catalyst: iron and iron oxides.  Many solid catalysts increase the rate of reactions by facilitating the correct orientations between the reacting molecules.

·         In the Fisher – Tropsch process, a heterogeneous mixture forms the reactants and the reaction is facilitated by a catalyst as well.  In this process solid carbon in the form of coal reacts with steam (gasified water) to form syn gas (or synthesis gas) mixture of carbon monoxide and hydrogen.  The hydrogen product can be enriched by further reacting the carbon monoxide with steam to produce more hydrogen and carbon dioxide in what is called the water-shift reaction.

o   C(s) + H₂O(l) à CO(g) + H₂(g)               coal gasification

o   CO(g) + H₂O(l) àH₂(g) + CO₂(g)          water-shift reaction

o   The CO and hydrogen can be further reacted to form methanol and long-chain hydrocarbons.

For Example: Hardwired for Respiration

·         Respiration is the process by which organisms exchange gases with their environment.  In plant respiration, plants take up carbon dioxide and release oxygen.  In animal respiration, organisms take up oxygen and release carbon dioxide.

·         The author states that plants need oxygen as well and the reason for this is best explained by batteries.

·         Just like flowing water carries energy that can turn a water turbine (or a flow of air, wind, carries energy to turn a wind turbine), electrons flowing in a current carry energy that can power our electronics.  But, in order for there to be a flow, a source and a receptor both must be present.  In a battery, one reaction is acting as the source of electrons and another reaction acts as the receptor for the electrons which is what drives the flow of electrons.  Batteries can store energy because the two reactions driving the flow of electrons can be separated to prevent any flow. Connecting the two electrodes of a battery provide a conductive way to connect the two reaction sites which then initiates the flow of electrons and thus energy.

·         [USE IN 30A/B AND 1A/B REDOX DISCUSSION]  Plants are autotrophs (make their own food).  This food is in the form of carbohydrates produced with the liberation of oxygen when carbon dioxide and water react.  To metabolize these carbohydrates, plants need oxygen as well.  Oxygen is the final electron receptor that drives the transfer of electrons in a chain of coupled reactions collectively referred to as the electron transport chain.  Just like in a battery, this flow of electrons is what provides the energy.

·         In plants, the source of electrons is the light phase reaction during the first phase of photosynthesis:

o   2 H₂O à 4 H⁺ + O₂+ 4 electrons

o   This reaction provides electrons to chlorophyll.

o   Solar energy causes the electrons to move from chlorophyll and enter the electron transport chain (the wiring, according to the author) in the thylakoid membrane.

o   These electrons moving down the electron transport chain are received by nicotinamide adenine dinucleotide phosphate (NADP⁺):

o   NADP⁺ + H⁺ + 2 electrons àNADPH

o   As the electrons flow, they do work to produce a proton gradient (pushing protons from where there is less to where there is more).  In essence, the electrons have transferred their energy to these protons.

o   The proton gradient can then do work (re-release the energy) by flowing from where there is a higher concentration of them to where there is less.  This re-released energy is stored in the bonds of the molecule adenosine triphosphate (ATP). 

·         In the “dark phase” (Calvin cycle) of photosynthesis, C from carbon dioxide is fixed (converted into a less volatile form) into sugars and complex carbohydrates with the release of oxygen:

o   6 CO₂ + 6 H₂O àC₆H₁₂O₆ + 6 CO₂

o   NADPH and ATP are used to power this reaction (no net ATP consumption however)

·         In cellular respiration, the sugars are broken down to produce energy stored in ATP (happens in both plants and animals).  This works very much like the light phase except, NADH is the electron source and not water.  The electron acceptor is oxygen:

o   O₂ + 4 electrons + 4 H⁺ à2 H₂O

Demonstration 12: How Does Your Garden Grow?

In this demonstration activity, laundry bluing, household ammonia, table salt, and distilled water are used to grow crystals.  [Through a web search, I found out what laundry bluing is:  “Mrs. Stewart’s ® Bluing is a colloidal suspension of extremely minute particles of blue powder ferric hexacyanoferrate, Fe₄[Fe(CN)₆]₃.]

CHAPTER 12: Crystal-Clear Chemistry

·         This chapter focuses on solids and crystals and the three-dimensional aspect of molecules and crystals.

·         One of the important clues as to the importance of the three-dimensional shapes of molecules relate to the discovery of isomers.  The author gives as an example grain alcohol (ethanol) and dimethyl ether as examples of isomers that have different properties.  Both of these compounds have the chemical formula C₂H₆O. Grain alcohol is a liquid and dimethyl ether is a gas at room temperature.  Another example of two substances that have the same molecular formula but different properties are glycogen and cellulose, both polymers of glucose.  Glycogen can be broken down to provide energy but humans cannot digest cellulose.

·         The author describes what governs molecular shapes, electron –pair repulsive forces, using water as an example to illustrate how lone pair – lone pair, lone pair – bonding pair, and bonding pair – bonding pair repulsive forces all act to determine the positions of the atoms in three dimensional space.

·         In bulk substance, the three dimensional shape of crystals is determined by the molecular shape and the “trade-off between ionic attractions, covalent bonding, and intermolecular forces”.

·         Types of Crystals:

o   Crystals of ionic compounds are held together by ionic attractions. 

o   Crystals of molecular substances are held together by intermolecular forces.

o   Crystals of covalently bonded solids are held together by covalent bonds, e.g. quartz, diamond, and graphite [I believe this is what is referred to as giant molecule.]

o   Ionic compounds that contain polyatomic ions are held together by a combination of ionic and covalent bonds.

·         An old method for weighing an atom described by the author:  equal volume crystals of sulfate salts containing different ions are weighed.  It is assumed that each volume contains the same number of sulfate units and therefore the difference in mass can be attributed to the difference in mass of the different metal atoms.  X-ray crystallography has vastly improved the analysis of the atom arrangement and therefore numbers of atom in a given volume of crystal.

·         Some solids have an amorphous structure (e.g. glass).  Crystalline solids can be distinguished from amorphous solids by their melting points:  amorphous solids do not have a distinct melting point unlike crystalline structures.

·         Surface chemistry: Practical considerations include controlling what sticks, how it sticks, and how it can be unstuck. Chemical bonding and/or intermolecular forces cause two objects to adhere.  The word glue is related to the word gluten because some glues can be made from animal and wheat proteins.

·         Oil can be used to remove substances adhering on a surface because it can form a film that can change the surface property [essentially disrupt the intermolecular forces between the two objects adhered to each other]. Oil can be categorized as a type of surfactant because of this.

·         Soap is another type of surfactant.  It can also insinuate itself between two things are adhered to each other, working as a lubricant.  Machine part lubricants are also surfactants because they can disrupt attractions between bare metal parts.

·         Porous surfaces have special properties because their actual surface area is much larger than the apparent surface area.  Substances that adhere to porous surfaces are harder to remove as a result and prone to staining, like teeth surface.  Teeth whitening makes use of surfactants that can hold mild bleaching agents against the porous teeth surface.

·         Another interesting surface chemistry is the work of solid state catalysts. Solid state catalysts help to orient reacting molecules so the probability of a reaction upon collision is higher.

For Example: Diamonds are (Almost) Forever

·         Both gold and diamond are highly valued substances, both valued for “their scarcity as well as their utility”.

·         Gold is durable and resistant to corrosion and is therefore useful as “dental prostheses and electrical contacts”.

·         Diamond is durable, hard, and inert which make it useful as a cutting tool, drill bit, and abrasive.

·         One method for making artificial diamond is to use microwave radiation to remove H atoms from low density collection of methane molecules.  The C atoms then settle out on a substrate, bonding to each other to form a diamond crystal.  According to the author, this type of diamond is indistinguishable from natural diamond.

Demonstration 13: Hot Packs and Cold Packs

·         The dissolving of citric acid in water is an endothermic process. To demonstrate this, mix 10 mL or two teaspoons of sour salt (citric acid) and 1 tablespoon of water.  Note the cooling. Sodium bicarbonate or baking soda would also work but not as much cooling.

·         When lye (sodium hydroxide) is added to water, the process is exothermic (quite!).  To demonstrate this, fill a glass with 240 mL of water and slowly add about half a plastic spoon of lye crystals to the water.  Note the heat generated.

CHAPTER 13: When Matters Heat Up

·         Understanding thermodynamics is well-worth the effort because it “embodies a powerful theory that enables chemists to predict whether a chemical reaction will take place, the extent to which it will take place, and the amount of energy that will be generated or consumed in the process – all before mixing chemicals together.”

·         The mechanics of thermodynamic theory boils down to three concepts: energy, entropy, and free energy.

·         The concept of energy in matter arises from the molecular model that particles of matter are in constant motion: vibration, translation, and rotation. When energy is added to a substance, the energy is distributed among these various modes.

·         The internal energy is the sum of the average energies of translation, rotation, and vibration.

·         Temperature is a measure of the internal energy of a substance.  However, of the three modes, the temperature measured is due only to energy of the translational motion of particles hitting the thermometer.

·         At boiling point, any energy added goes only into overcoming intermolecular forces between the particles so it does to register as an increase in internal energy.

·         Energy will flow from an object at higher energy to an object at lower energy.  This is the type of energy that we call heat.

·         Another type of energy transfer is work.

·         When chemical changes are a source of energy to do work, the energy comes from the breaking and formation of chemical bonds.

·         The amount of internal energy can be changed by transfer through heat.

·         While most processes proceed to a lower energy state with accompanying release of energy, some processes do occur that leads to a higher energy state (and energy is absorbed).

·         The probability that a system will undergo change spontaneously is governed not just by energy change but also by entropy.

·         The author defines entropy as the “tendency of systems to move toward a state of maximum order”.  Entropy is “a real driver for physical processes and does indeed weigh in as a determinant in the fate of the universe”.

·         Systems tend toward the most disordered state simply because that is the most probable state.

·         Free energy is the energy that “results from the difference between the energy gained or lost in a chemical reaction and the amount borrowed or expended because of entropy”.  Free energy is that which left over, “free to do work”.

·         A process that absorbs energy can still occur spontaneously if there is a gain in entropy.

For Example: Refrigeration on the Road

·         In a refrigerator, the cooling process is driven by the evaporation of the liquid refrigerant which then expands as it absorbs heat to cool down the inside of the refrigerator.  The expanded gas is compressed and condensed back to the liquid state using an electric motor (heat is dumped out in the back of the refrigerator as the gas condenses and is compressed back to the liquid state).  In the author’s words, “the compression is driven by an electric motor and the expansion is driven by entropy”.

·         In an automobile refrigerator, the thermoelectric Peltier effect is in operation.  Two different semiconductors with different conductivities are connected.  When electrons flow from one semiconductor with lower conductivity to the other with higher conductivity, energy is released.  In the opposite direction, energy is absorbed.  The energy released by one flow is used by the electrons flowing in the uphill direction.  Energy absorption causes cooling and this heat can then be dumped outside the insulated area that is being cooled down.

Demonstration 14: Zoned-out Ice

·         In this demonstration, a drop of colored indicator is allowed to mix with water and the mixture is frozen.  Upon freezing, it is observed that all the indicator substance has been pushed toward the center.  The color changes involved are:  the indicator is yellow-orange when it comes out of the dispenser, turns green when added to tap water which is slightly acidic, and turns back to yellow as it separates from the tap water and reconcentrated back to its unmixed form.

·         In industry, zone refining is used to purify metals.  The molten metal rod is cooled down and solidifies starting from one end pushing impurities toward the other end.  This end can then be cut off.

CHAPTER 14: A Whole New Phase

·         Phase changes occur as heat flows in and out of a substance.

·         In addition to temperature, pressure plays a role in determining when a phase change occurs. Water can be made to boil at a lower temperature at higher elevations where the atmospheric pressure is lower.

·         This can be observed through the following activity:  fill a basting syringe (or any plastic syringe I am guessing) with rubbing alcohol. Seal the tip.  Wrap a warm cloth around it and pull the plunger back a little to create a vacuum inside (keep the tip sealed).  The alcohol is observed to boil due to the reduced pressure above and the slightly higher temperature from the warm cloth.

·         Dry ice (solid CO₂) is literally dry because, unlike ice that has some liquid on it, CO₂ sublimes and skips the liquid phase at normal atmospheric pressures.

·         In the late 1800’s, Josiah Willard Gibbs created phase diagrams to show all the pressure and temperature values at which a phase change may occur for a given substance.  At these temperatures and pressures (represented as points on the curve), two phases exist in equilibrium with each other (they are equally likely to exist) so some amount are in one phase and some amount are in another phase [how much of each?  Equilibrium constant determines this ratio].  Outside and in between these curves, the substance is either completely in the solid, liquid, or gas phase.

·         In this chapter, the author discusses two textbook examples of phase diagrams: one for water and one for carbon dioxide.

·         “A phase is any uniform and stable state of matter, pure or mixture”.  Pure metals can exist simply in its solid, liquid, or gas phase.  An alloy on the other hand can exist in more than three phases: solid Cu – solid Zn, solid Cu – liquid Zn, liquid Cu – solid Zn, liquid Cu – liquid Zn, etc.  Which of these is the stable phase depends on the temperature, pressure, and relative amounts of the two metals in the alloy.

·         [USE PHASE DIAGRAM FOR THE SOLUBILITY OF SUGAR IN WATER , Figure 1.14.4]

·         Application to hollandaise sauce: Hollandaise sauce is a suspension of butter in egg which pours out as a one-phase mixture at warm temperatures.  When it cools down, the egg and butter separate.  The hot hollandaise mixture is in a metastable stable (a stable state above the lowest energy state). 

·         Metastable State:  The author gives a bowling ball balanced on top of step ladder as an analogy to a metastable state.  The bowling ball is stable but not at its lowest energy state.  Metal alloys can also form in a metastable state if cooled at a fast rate with not enough time for the metal atoms to position themselves at the lowest energy state. 

·         “A system that is proceeding to its most stable state so slowly that it is virtually at rest can also be considered metastable.”

For Example: When Good Pizza Goes Bad

·         Many familiar household substances we use are in a metastable state:  medicinal emulsions and suspensions, shampoos, salad dressings, sauces, catsup, mayonnaise, mustard, and cheese.

·         Cheese is a metastable mixture of oil, fats, water, calcium compounds, and milk containing proteins.

·         Aging involves the action of bacteria in coagulating the components of cheese initiate by enzymes.  Well-aged cheese keeps better because they have already gone through the action of bacteria.

·         The metastable state in which cheese is produced and sold causes it to change texture when subjected to wide temperature ranges.

·         Mozzarella is used for pizza topping because it melts uniformly and smoothly.  It also melts quickly.  In this context, mozzarella can be considered a eutectic mixture, a single-melting-point mixture (just like solder is).

Demonstration 15: All Things Being Equal

A demonstration of the equilibrium condition and Le Chatelier’s principle using household chemicals:

A solution of baking soda (NaHCO₃) is added to a solution of copper (II) sulfate dropwise.  A white precipitate forms as a product:

Cu²⁺+ CO₃^2- à CuCO₃(s)

Upon addition of vinegar containing acetic acid, the carbonate ions decrease as they react with acetic acid:

2 HC₂H₃O₂ + CO₃^2- àH₂O + CO₂ + 2 C₂H₃O₂^-

As the CO₃^2- concentration goes down, the CuCO₃precipitation reaction above shifts left (according to Le Chatelier’s principle) and some of the CuCO₃ precipitate dissolves.

Cu²⁺+ CO₃^2- ß CuCO₃(s)

When some baking solution is added back, more CuCO₃ white precipitate appears.  Addition of the carbonate ions (from the sodium bicarbonate) causes the precipitation reaction to shift right to make more CuCO₃ precipitate:

Cu²⁺+ CO₃^2- à CuCO₃(s)

CHAPTER 15: Equilibrium – Chemistry’s Two – Way Street

·         In this chapter, the author addresses in a simplified way the following concepts:

o   the reversibility of reactions, equilibrium condition and mixture

o   the dynamic nature of equilibrium

o   reactions at equilibrium can respond to any changes in amounts or condition (as governed by Le Chatelier’s Principle)

·         After the reaction has reached equilibrium, the reactants and products in the reaction vessel constitute an equilibrium mixture.  For some reactions, the equilibrium mixture may contain a lot of products and very little reactants left (e.g. gunpowder reaction), a lot of reactants left and very few products formed (ammonia synthesis from nitrogen and hydrogen), or something in between.

·         The author describes quite a complex analogy of a trade show to try to explain the “why” of reactants or products dominating at equilibtium.

·         At equilibrium (in a closed system), reactant and product particles will distribute themselves so that the overall energy of the system is minimized and the entropy maximized.  The trade-off is between lowered energy and maximized entropy until a compromise is reached by way of ZERO free energy.  [USE IN 1A]

·         The author ends this chapter with the following statement: “An equilibrium always carries the day. Equilibrium may be held at bay for a time, and systems may even enter a metastable state, but ultimately, equilibrium rules.  Chemical reactions are able to respond to and adjust to lower energy and maximize entropy because chemical reactions are reversible and dynamic.

For Example: Breathing Out and Breathing In

·         The author uses the binding of O₂ to hemoglobin to form a hemoglobin – oxygen complex to illustrate the role that equilibrium plays in biological processes:

·         Hb + O₂ ⇌ HbO₂

·         HbO₂ formation is favored when there is a high concentration of O₂:  in the lungs (the reaction shifts right)

·         The release of O₂ from HbO₂(reaction shifting left) is favored when there is little O₂ present: in the cells where they have been consumed.

·         At higher elevations where O₂concentration is low, the body produces more Hb to push the reaction forward and capture more O₂

Demonstration 16: Antifreeze and Antiboil

The four activities below demonstrate the colligative properties of a solution:

Freezing Point Depression
Fill two containers with equal amounts of water.  Into one of them, add enough salt until the solution is saturated.  Place both in the freezer.  After a day or two, observe that only the pure water has frozen while the salt solution has remained liquid.

Boiling Point Depression

Heat water from the tap.  Measure the temperature of the boiling water.  Now heat vinegar until it boils.  Measure the temperature of boiling.  The vinegar’s boiling point should be a few degrees higher than the water.

Vapor Pressure Lowering

Crush three aspirin tablets.  Add the crushed aspirin to one of two cups.  Add a half a cup of rubbing alcohol into each cup; make sure the two volumes are the same.  Swirl the aspirin-alcohol cup to mix it (not all solid may dissolve).  Let the two cups sit side-by-side for about an hour.  After about an hour, observe that the liquid level in the cup that has only alcohol will have gone down more than in the cup with the alcohol-aspirin mixture due to slower evaporation rate.

Osmosis

Add a raisin to pure water.  The raisin will plump up after about an hour as water diffuses from where there is more of it to where there is less of it (the raisin has a much higher concentration of solute).  In another experiment, add a thin slice of potato to each of two cups, one containing just water and the other containing a salt brine (on tablespoon of salt in ½ cup of water).  Observe the difference between the two slices. The one in the salt brine should become limp as water leaves it to go to where there is less of it and a higher solute concentration (salt brine).

CHAPTER 16: Colligative Properties – Strength in Numbers

·         Colligative properties are physical properties involved with phase changes.  The magnitude of colligative properties depend only on the number of particles dissolved (atoms, ions, molecules) and not on their identity.

·         Adding salt on an icy road is an example of the result of freezing point depression.  Ice is in equilibrium with liquid water.  When salt is added to ice, it dissolves in the liquid portion making a solution.  The ice starts to melt because as mixing occurs, more water enter the liquid phase.

·         The same solute added to water in antifreeze (to lower the freezing point of water) also works as an anti-boil as it elevates the boiling point of water.

·         The vapor pressure of water is lowered when non-volatile solutes are added to it.  This colligative property of a solution of lowered vapor pressure is also dependent just on the number of dissolved particles and not the identity.

·         Osmosis is the tendency of solvent to flow where there is a higher concentration of solutes.  Examples given include the making of pickles (cucumber in a salt brine), the puckering of skin as skin cells swell when water from the bath diffuses through the semipermeable membrane of the cells.

·         From a thermodynamic point of view, entropy is thought to drive these properties to maintain or achieve a more dilute solution (higher disorder, higher entropy).

For Example: Kidneys and Chemistry

·         “Colligative properties of solution figure into the body chemistry in a big way because the body is just one big salt solution partitioned into cells by semipermeable membrane.”

·         The flow of substances in and out of cells and in and out of blood as it gets filtered through the kidneys is governed in large part by control of the net diffusion of solvent and solutes through osmosis and dialysis and other mechanisms. 

·         Kidneys use filtration to remove water and small molecules by applying pressure.  Transporter proteins on the kidney membranes recognize and reabsorb molecules such as sugars and proteins that are accidentally pushed out of the blood.  These transporter proteins do not recognize non-human made substances so drugs are eliminated.  Their dosage has to be timed just right to maintain the needed amount in the blood.

·         The kidneys also help maintain a proper blood pH by regulating the amount of carbonate and bicarbonate ions.  Carbon dioxide is also produced by the kidneys.

·         The antidiuretic hormone ADH helps control the permeability of the cell membranes and regulate osmosis. When the blood has high levels of sodium, the pituitary gland releases ADH so that more water is retained in the blood to keep it dilute.

Demonstration 17: Kick it up a Notch

This next activity demonstrates the effect of concentration and temperature on the rate of a reaction (kinetics).

Effect of concentration:

Pour a half a cup of bleach each in two cups.  Add two drops of methyl red to one and, immediately, 4 drops in the other.  Start timing and note the time it takes for the two drops to disappear and the time for the 4 drops to disappear.

Effect of temperature:

Into one cup place 120 mL of cold tap water and into another cup the same volume of hot tap water.  Drop a table of effervescent antacid into each cup at the same time.  Observe that the fizzing is much more vigorous in the hot water indicative of a faster reaction taking place (and lower solubility of the CO₂).

CHAPTER 17: Chemical Kinetics – A Veritable Explosion

·         “Thermodynamics deals with the effects of heat and entropy on equilibrium and is useful for predicting the extent to which a reaction will occur.”

·         Thermodynamics can predict the extent to which a change will occur but not when and how fast.  Because graphite is the more thermodynamically stable, all diamond will turn into graphite but it will take years (billions?).

·         The word kinetics derives from the fact that the rate of reactions is related to the motion of particles.  Gas phase particles can travel at speeds of 500 m/s, each one colliding 7 billion times every second due to the density of particles in the atmosphere.

·         The collision theory of kinetics is used to understand and explain the different conditions that affect reaction rate.  For reactions to occur, reactant particles must collide and must do it in the correct orientation.

·         Factors/conditions that can be controlled to affect reaction rates:

o   Concentration

o   State of reactants:  Add powdered chalk (calcium carbonate) versus a stick of chalk to vinegar to demonstrate the difference in rate. Liquid gasoline burns whereas gasoline vapor explodes.  The presence of surfaces can also impede or facilitate a reaction by affecting the orientation of reactants in contact with the surface.

o   Temperature:  The vast majority of reactions occur faster as the temperature increases.  Mammals respond to injury by raising the body temperature to increase the rate of immune response reactions and repair reactions.  Note, however, that heat can affect the equilibrium position depending on whether the reaction is endothermic or exothermic.

o   Presence of a catalyst:  A catalyst functions to lower the activation energy required such that the reaction can occur at a faster rate.  Enzymes are biological catalysts.  Enzymes in dietary supplements help digest dairy products and roughage.  Enzymes in detergents help breakdown stains faster. Catalytic converters in cars increase the rate that NOx’s and unburned hydrocarbons are converted to nitrogen gas, oxygen gas, water, and carbon dioxide.

·         Some catalytic mechanisms are known but not many others.  Many enzymes, for instance, lower the activation energy and accelerates reactions by helping orient reacting molecules correctly.

·         The mechanism for why copper works as a catalyst but not steel wool in the decomposition of hydrogen peroxide is unknown.

For Example:  When Your Car Comes Knocking

·         Explosions arise from very fast reactions, so fast that the “surroundings don’t have time to accommodate the change”.

·         Combustion in an internal combustion engine occurs as a series of mini explosions occurring at a rate of several thousand per minute.

·         In a smooth coordination of the piston moving up to compress the gasoline air mixture, a spark initiating an explosion, and the explosion pushing the piston down with the greatest impact, timing is everything.  If the explosion occurs too early as the piston is still moving up, then the impact of the explosion is lessened.  If the explosion occurs too late after the piston is already in its downward movement, then it will not receive the full impact of the explosion.  The time it takes for the combustion to occur and the explosion to spread needs to be taken into account in determining when a spark is initiated to ignite the mixture.

·         Premature ignition will results in uncontrolled burning which causes “knocking” (due to uneven pressures and burning at several places at once).  This can occur if the temperature gets too hot (compression of the gasoline and air mixture will cause the temperature to increase).  A fuel with high octane rating is formulated to prevent premature ignition before the spark is applied.  The octane rating for gasoline is based on a ratio between 8-carbon atom hydrocarbon molecules and 7-carbon atom hydrocarbon molecules.

Demonstration 18: Black-Light Chemistry, Paper Clip Chemistry

Black-Light Chemistry

The activity described here is to demonstrate the effects of light on chemical reactions.

Pour a half a cup of bleach each in two cups.  Expose one of the cups to black light (UV) for 15 minutes.  “The black light should be moved to within an inch of the solution’s surface for maximum effect” and should be directly toward the surface not through the wall of the cup.  After 15 drops, add equal drops of the methyl red indicator to the two cups.  The disappearance of color in the bleach exposed to UV light should occur faster as the UC has broken down the bleach compound into more reactive species.

Paper Clip Chemistry

Pour about a half inch of copper (II) sulfate solution into a cup.  Untwist a paper clip and dangle it over the lip so that one end is immersed in the copper (II) sulfate solution.  Stand a galvanized iron in the liquid without it touching the paper clip.  Drape a piece of aluminum so that one end is in the solution.  After about an hour: the paper clip end should have a copper coating, the nail should turn black, and the aluminum should not change.  Electrons on the surface of the steel (iron) paper clip and the zinc on the galvanized surface of the nail have been transferred to the copper ions in solution converting them to metallic copper atoms. 

CHAPTER 18:  Electrons and Photons: Turning on the Light

·         Photochemistry:  Light can be a product (fireworks) or a reactant in chemical reactions (photosynthesis).

·         Electrochemistry:  Electrons can be made to flow spontaneously (batteries) and electrons can be forced to flow in the non-spontaneous direction to make a reaction happen (plating).

·         Both of these involve the flow of particles: photons and electrons.

·         Light is an oscillating electromagnetic field, when absorbed at the right frequency can impart energy to atoms.  This can cause electron rearrangement in molecules and atoms that may lead to new bonding arrangement and new substances.

·         The behavior of light as it interacts with matter can be described as either a wave or as stream of particles.

·         Another example of a reaction that produces light can be found in glow sticks.  Phosphorescence (or fluorescence?) is emitted when the substances in the glowstick react.  [See downloaded website article from University of Michigan for more details about the reactions.]

·         In fireworks, gunpowder explosion is used to heat up metals and excite electrons which then relax back down to the ground state emitting visible light of varying colors.

·         Photochemistry plays an important role in atmospheric chemistry:

o   When lightning strikes, molecular nitrogen and molecular oxygen react to form NO.  NO then reacts with O₂ to form NO₂.  NO₂ dissolves in rain and falls to the ground where bacteria can then use it to synthesize proteins and other nitrogen containing compounds.

o   High amounts of NO₂ released from automobile gas combustion, when released into the atmosphere, can breakdown into NO and O when hit with UV light.  The O can then react with O₂ and form O₃ or ozone. Ozone is beneficial if found in the higher atmosphere because of their ability to block excessive UV.  At ground level, however, excessive amounts of ozone is damaging to a lot of materials: rubber, plastic, plants, and animals.  It reacts with exhaust from cars producing irritating pollutants.

·         Electrochemistry:  the oxidation site is separated from the reduction site connected by a conducting wire through which electrons can be made to flow.  The two solutions must also be connected by a conductive medium (an electrolyte) to complete the circuit.  In car batteries, this electrolyte solution is sulfuric acid.

·         Voltaic electrochemical cells (like batteries) produce energy because electrons flow in the spontaneous direction releasing energy in the process.  Electrolytic cells require an external provider of energy to force electrons to flow in the non-spontaneous direction.

·         See the following links to rechargeable batteries:

o   http://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Voltaic_Cells/Case_Study%3A_Battery_Types/Rechargeable_Batteries

o   http://www.scientificamerican.com/article/how-do-rechargeable-that/

·         Voltaic Cell -- Lemon electrochemical set-up:  stick a zinc-coated nail at one end and a copper penny in a slit at the other end.  Use a voltmeter to measure the voltage.

·         Electrolytic Cell – Take the copper-coated nail from the lemon electrochemical set-up.  Dangle it in a glass with about an inch of freshwater pH-lowering solution [I think this is just a dilute sulfuric acid solution].  Hook the copper coated nail to the positive end of a D-cell battery.  Dangle an iron nail on the other side and hook it to the negative end of the D-cell battery.  The iron nail will eventually get plated with copper.

For Example: Photon Meets Electron – Solar Energy

Solar radiation can excite electrons in titanium oxide particles.  These excited electrons can form mini-batteries that can drive electrons through organic pollutant molecules and break them down.

Part 2

In Part 2, the authors look at what they refer to as further “explorations” in chemistry: organic chemistry, inorganic chemistry, biochemistry, and analytical chemistry.

Organic Chemistry Demonstration: Aspirin to Acid

[MIGHT TRY THIS AS A 30A LAB]

Acid-catalyzed hydrolysis of acetylsalicylic acid:  he unique feature of this demonstration activity of a common organic chemistry lab is the use of household chemicals. 

In the first step, about 120 mL of rubbing alcohol (or 70% isopropyl alcohol) are added to 10-15 tablets of aspirin (plain or buffered OK).  This mixture is then warmed in a microwave on half power for 30 seconds.  The rubbing alcohol which is polar will extract the soluble acetylsalicylic acid from the tablets separating it from the insoluble binder (mostly starch).  Crushing the insoluble parts will help maximize the amount of acetylsalicylic acid extracted.  This process takes about 15-20 minutes. 

Filter out the undissolved material by pouring the mixture through a paper towel filter pushed down through a glass as in a funnel.  The clear liquid draining through the paper towel, the filtrate, should contain the acetylsalicylic acid. 

Try not to get any of the liquid on your skin as it contains acetylsalicylic acid.  Remove the paper towel with the solid residue and dispose of in the waste container.

Fill the glass containing the extract (filtrate) with some cold tap water until about ¾ full.  As the cold water is added, acetylsalicylic acid should precipitate (crystallize) out.  Acetylsalicylic acid is not as soluble in water as it is in isopropyl alcohol.

Set-up another paper towel filter and collect the solid crystals of acetylsalicylic acid.  Use more cold water to wash out any crystals remaining in the glass and pour this over the solid on the filter.

Let dry overnight and make observations about appearance and smell.

In the next step, take a quarter of the acetylsalicylic acid (wet or dry) and put in a glass container suitable for microwaving.  Add aquarium pH – lowering solution to the solid until it is completely covered with the liquid.  Aquarium pH lowering solution is sulfuric acid that is not very concentrated.  It is used as a catalyst in the reaction with water being the other reactant.  Warm the mixture on low power in the microwave,

Gently waft some of the vapor from the warmed mixture.  You should be able to detect the smell of vinegar because acetic acid is one of the products in the hydrolysis of acetylsalicylic acid.  The other product is salicylic acid, a natural analgesic.

PART 2, CHAPTER 1: Simply Organic

·         The authors start off the chapter by pointing out that many of the materials that we are familiar are made of organic compounds including all living things.

·         Substances made from hydrocarbons may seem ubiquitous but carbon does not figure in the top 8 of elements (by amount) in the earth’s composition or even in the earth’s crust.

·         Despite its relative scarcity, carbon forms a huge variety of compounds because of its ability to form long, stable chains (concatenation).  This “talent” for concatenation is owed by carbon to the relative smallness of its atoms and their ability to form 4 stable bonds in a tetrahedral geometry that can extend these chains fluidly in 4 different directions.  In addition, C can form double and triple bonds, adding to a diverse geometric repertoire for bonded carbon atoms.

·         The stability of these C-C and C-H bonds is evident in their strength (energy): huge amounts of energy are released when they are broken during combustion reactions for instance.

·         The hydrocarbons and other families containing other elements such as O, N, S, halogens, etc are characterized by functional groups, “identifying structures” that define the physical and chemical properties of each family.

·         Oxygen-containing families like alcohols, ethers, organic acids, and esters with examples are discussed.

·         The concept of handedness or chirality is also briefly touched upon along with the broad concept of isomerism.  The chapter describes a very commonly referred to example of isomers with dramatic differences: HONC (fulminic acid, explosive), HOCN (cyanic acid, a poison), and HCNO (isocyanic acid, used in synthesis). 

·         Examples of chiral isomers given are

·         carvone: one enantiomer smells like spearmint and the other smells like caraway.

·         Limonene: one smells like lemon and the other smells like oranges

·         “In many biological systems, if the molecule does not fit, it doesn’t work.”

·         “When light encounters a chiral molecule, the light may experience a twist from the electric field of the unsymmetrical molecule”.  Enantiomers interact with polarized light in opposite ways.

·         The human body favors most of the time the L form of enantiomeric molecules.

·         The drug thalidomide, prescribed as a sedative, unfortunately came with an enantiomer that happened to be a teratogen, a compound that causes birth defects.

Inorganic Chemistry Demonstration: Entropy, Ammonia, and Amoeba

Place a puddle of a about a half inch diameter of iron (II?) acetate and about a quarter inch away a similar size puddle of household ammonia.  Use a straw or stick or glass rod to draw a line of liquid connecting the two puddles.  Observe any changes.  Do the same for a puddle pair of copper (II) sulfate and ammonia and copper (II) sulfate and baking soda solution.

PART 2, CHAPTER 2: Chemistry Rocks

·         “If organic chemistry can be described the chemistry of hydrocarbons, then inorganic chemistry can be described as the chemistry of everything else”.

·         Topics that are the “special purview” of inorganic chemistry are discussed:

o   Coordination complex compounds of transition metals*

o   Representative group chemistry

o   Radiochemistry

·         *To iron (II) acetate:

o   Add NH₃ àorange to green (coordination chemistry)

o   Add H₂O₂ à green to red (redox and coordination chemistry)

o   “When ammonia was added to the iron solution, the ammonia molecules crowded toward the metal ion, displacing some of the water molecules and acetate ions, and in the process changed the color of the compound.  When we added the peroxide, we changed the oxidation state of iron and how tightly the ligands were held.

·         *Radiator cleaner contains oxalic acid which complexes calcium ions deposited on radiator walls.  Oxalic acid molecules form a water-soluble cage around the calcium ions.

·         The chapter discusses families of representative elements and their general properties. 

·         The first and the second groups are the alkali and alkaline earth metals respectively.

·         The third and fourth groups are the semiconductors.  Combinations of tin, lead, indium, and aluminum have been used in making semiconductors.

·         “The importance of recycling aluminum metal does not arise from its rarity but rather from the amount of energy and materials that must be expended to claim it from the rocks and purify it.”

·         In a halogen lamp, the tungsten filament heats up and a small amount of it vaporizes.  Halogen gas is added to prevent complete vaporization of the tungsten filament, extending the life span of the bulb.

·         Most of the radioactive material are found in the rare earth elements group (lanthanides and actinides).

·         Carbon-14 dating is limited to material that is less than 50 years old (carbon – 14 is constantly being produced by cosmic ray bombardment of nitrogen.  Carbon – 14 is taken up by living organism in which its concentration stays constant until the organism dies.  At death, the carbon-14 starts to decrease in concentration (new C-14 is not being absorbed) at a rate depending on its half-life.  When compared with the amount of C-14 in a living organism, the time passed at time of death can be approximated (also based on decay rate compared to current decay rates).

Biochemistry Demonstration: Fat, Flatulence, and Bean Soup

·         Cornstarch mixed in water (a quarter of a teaspoon in 240 mL of water) is “digested” by a tablet containing alpha-galactosidase (crush before adding).  Prepare another cornstarch-water mixture without the enzyme.  Leave for about an hour.  Add tincture of iodine to each one.  Note the blue-violet color in the mixture without the enzyme and the same brown yellow color of iodine in the mixture with an enzyme

CHAPTER 3: THE BODY OF CHEMISTRY MEETS THE CHEMISTRY OF THE BODY

·         Plastics and many biological macromolecules are both polymeric in nature.

·         Carbohydrates: long-chain polysaccharides form the essential materials of starch, cellulose, and glycogen.

·         Lipids are biomolecules that are not soluble in water. Two common types are fatty acids and steroids. 

·         Fatty acids store energy, help protect and insulate organs, form insulation on nerve fibers, and form the basic component of cell membrane.  Arachidonic acid is a fatty acid that is released when a cell is cut or injured.  It is converted into a compound that stimulates nerve cells (for pain sensation), rushes blood to the site (causing inflammation), and raises the temperature of the body to speed up the repair process (causing fever).

·         Cholesterol is a steroid.  It is used in the synthesis of hormones and the cell membrane.  Too much of it ends up in the blood vessel walls and can cause constriction of blood flow.

·         Proteins are polymers composed of amino acids.  They have many functions including structural, regulatory, catalytic, etc.  Enzymes, catalytic biomolecules, allow our bodies to respond to needs as fast as necessary.  Enzymes can be turned off or turned on or broken down when not in use so that the amino acids can be used somewhere else.

·         Nucleic acids like DNA are polymers as well, each monomer containing a base, a phosphate, and a sugar molecule.  DNA contains information for how to make proteins.

·         “Our rapid and cursory survey of biochemistry, combined with our previous discussions of biochemical systems, shows that in life all our chemical principles come into play: acid-base reactions, redox reaction chemical bonding, intermolecular forces, concentration, solids, solubility, gas phase, kinetics, phase transition, thermodynamics, equilibrium, entropy, etc.

·         Even though biological systems seem so “organized and ordered”, entropy plays a big role in its workings.  Cells are “irregular and messy” and even DNA contain “jumping genes and junk genes” that appear to have no role but simply to cause confusion or perhaps variety.  “The irregularity means that no two biological systems will ever be completely alike.”

Analytical Chemistry Demonstration: The Proof is in the Prints

·         Developing and lifting fingerprints:  Place a glass jar with fingerprints inside a plastic tub and seal closed.  Squirt some superglue onto a piece of aluminum foil shaped into a bowl or square pan (the larger the tub, the greater the amount of superglue needed).  Make sure the superglue contains cyanoacrylate.  Place this inside the plastic tub containing the glass jar.  Leave overnight.  The fingerprint should be more visible, tomorrow, and can be lifted off with a tape.  Sticking the tape onto a black background will make the gray lines of the fingerprint show up more prominently.

·         Testing for the presence of blood can be done with hydrogen peroxide.  Blood contains hemoglobin that can catalyze the decomposition of hydrogen peroxide into water and oxygen.  Take a red dye and add it to a little bit of water in a cup.  In another cup, squeeze blood from a beef container absorbent pad.  Add hydrogen peroxide to both.  The red dye water should turn clear as the red dye gets oxidized by the hydrogen peroxide.  The hydrogen peroxide in the blood should start to bubble vigorously as the decomposition process is catalyzed by the real blood.

CHAPTER 4: Chemist as Analyst

·         “In analytical chemistry, all the tools of chemistry and all the talents of chemists are brought to bear on two questions: what is it and how much of it do we have?”

·         “The job of the analytical chemist is to find – or invent – the tool or tools necessary to determine the quantity of a given material or the nature of unknown materials.”

·         “Analytical chemists must be able to defend their results by reporting parameters such as confidence limits, instrument noise levels, and significant figures.”

·         In anything type of analysis of material, the first step is an attempt to separate the components of a mixture by one or more of the following:

o   Filtration

o   Solvent extraction

o   Chromatography (gas, ion, high-performance liquid, etc.

·         The next problem to address is concentration.  This can be optimized by:

o   Evaporation

o   Precipitation

o   Centrifuging

·         Polymerase chain reaction for DNA

·         The next step is analysis:

·         GCMS

o   Spectroscopy is the general name for any technique that has a spectrum as its final result.  A spectrum has two components: the property or substance being tracked and its quantity.

o   Good analogy for how different mass ions can be separated in a mass spectrometer:  throwing bags of various sizes with the same effort and having them end up at different spots.

o   In mass spectrometry, molecules are hit by energetic electrons which may cause them to ionize and/or fragment into pieces depending on relative bond strengths.  These fragments are hurtled through a magnetic field and because of their different mass to charge ratio arrive at the detector at the different times, spots, and abundance.  The same molecules are likely to fragment in the same manner under the same conditions and therefore produce the same spectra.  The first peak to look for is the parent peak, usually the highest mass peak but not necessarily the most abundant.  Identities of the parts and pieces can be deduced from the mass to charge ratio, presence of isotopes, and perhaps abundance.  Mass spectrometry is a destructive technique.

·         Light spectroscopy takes advantage of the differences in the way different light interacts with different matter.  Types of light spectroscopy include ultraviolet, visible, infrared, microwave.

·         For inorganic salts, precipitation is a way to detect the presence of ions in an unknown sample.

·         Absorption spectroscopy is a destructive technique for inorganic substances but it is a destructive one as the sample is atomized (turned into a fine mist) and sprayed into a flame, just in the right place.  Flames have interesting chemistry: some regions are hotter than others.  Some regions promote oxidation while others reduction.  The way that substances in the atomized sample changes the quality of the flame and how much is determined by the substance in the spray and its amount.

·         In emission spectroscopy, a lamp containing the same ion being analyzed emits light.  The amount of the ion present then absorbs the light emitted by the lamp and this decrease in intensity is detected as proportional to the amount of ion in the flame.

·         Simple chromatography demo:  crush and cut up three green, fresh tree leaves and place in the bottom of a drinking glass. Add about 3 inches of acetone (from fingernail polish remover) which should just cover the leaves. Let it sit for several minutes to let the acetone extract compounds from the leaf.  Cut a strip of paper towel and dangle it over the glass until the bottom just touches the liquid (stick to a pencil and hang).  After several hours, there should be at least two distinct bands of color as the acetone travels up the paper towel by capillary action.  One band might be green and the other bands may be yellow, orange, or red. (In the fall, the chlorophyll is gone so the reds, oranges, and yellows start to show through).

Demonstration for the Future: Medicine Cabinet Chemistry

·         Go to the medicine cabinet and get a “synthetic mimic of the extract of Salix willow bark, modified to buffer the acidic effect”, aka aspirin.

CHAPTER 5: Harry, Hogwarts, and Folk Pharmacopoeias – Mysteries in the Past and Magic in the Future

·         In our increasing abilities to scour nature’s millions of different compounds or synthesize a seemingly infinite repertoire of atomic combinations for potential drugs, the need for rational drug design becomes even more crucial.

·         Rational drug design process provides a couple of choices:

·         Combinatorial approach: try a variation of a compound that has worked in the past.  The author gives as an example drugs that have bacteria have become resistant to can be modified and tested for renewed efficacy.  The disadvantage of the combinatorial approach is it is only sampling “soil near its roots” and therefore a potentially more effective but completely different drug may be missed.

·         Study of folk medicines: remedies that may arise from millions of tests from millions of years of human-plant interactions.  Aspirin and quinine were discovered this way as well as using iodine-infused seaweed for curing goiter and lime juice to treat scurvy.

·         Phytochemistry – study of the chemistry of plants

·         Why do plants carry so much potential for substances that may cure us?

·         Plants share the same chemical space (take up carbon dioxide and produce oxygen, the opposite for humans; ovaries become fruits so animals can eat them and carry the seeds away).

·         Plants protect themselves from the same scourges from which we protect ourselves: bacteria, fungi, viruses.

·          ”History tells us that aconite and belladonna were once used together to produce the hallucination of flying.' Why would plants produce chemicals with these peculiar properties? Do plants have headaches? Do they have trouble sleeping? Why have they evolved chemicals with analgesic, sedative, or psychoactive properties? Again, it could be just an accident. Plants evolved chemicals to kill predators, but a toxin that would kill an insect might just tantalize or sedate a human. The difference between a sedative and a poison is often a matter of dosage. The ingestion of toxins below lethal levels for the sake of euphoric effects has become known as intoxication, as a nod to its toxic origins.”

·          Why did plants evolve analgesic, sedative, or psychoactive chemicals?  Plants produce many of these chemicals to ward off or kill predators.  While these amounts can kill animals, these toxins may have only a milder effect as an analgesic or sedative to humans.  “The difference between a sedative and a poison is often a matter of dosage.”  “The ingestion of toxins below lethal levels for the sake of euphoric effects has become known as intoxication.”

·         Pharmacognosy is the study of natural product drug discovery and development.  It involves investigation of ethno-botanical claims (“folk medicines”) but also looks to natural sources. Aside: To protect their secret medicinal potions, healers in the European Middle ages, conceal the identifiable taste of aromatic herbs by mixing it with foul-tasting ingredients or they would use smoke to conceal the steps.  Waving a stick was thought to distract the patient so they would not see what goes on in the brew.

·         Besides plants, of course, there the oceans, the polar ice caps, underground, other planets, and the moon to be looking for the possibility of “chemical cornucopia”.  The question is, where does one start?

Epilogue

·         “There is so much work to be done, and so many reasons to do it.  The work is far from finished. And the joy never ends.”

I celebrate myself, and sing myself,

And what I assume you shall assume,

For every atom belonging to me as good belongs to you.

Walt Whitman, Leaves of Grass, 1855.

END OF BOOK