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 written to entice anyone interested in conducting their own experiments at home, with or without any advanced background in chemistry. The activities described her 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 to relate 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.
NOTE: Demonstration instructions are copied verbatim from the book to avoid mistakes and inaccuracies in the steps.
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.
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 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 I2out 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.
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