The 1827 Christmas Lectures
of Michael Faraday

Demonstration of Lecture #6 Today

Metals -- Oxides -- Earths -- Fixed Alkalies -- Salts

Teacher Presentation

Metals

Metals can be distinguished at first sight from almost all other substances in your world by their brilliant luster. The next properties you will notice is their opacity, even when the piece of metal is extremely thin, and by their high density.

• We have here a large number of common everyday objects made of metal. If you will think of all the objects you have used so far today that are made of metal - door handles, water faucets, knives, forks and spoons, keys, zippers, snaps, chains, earrings, etc. - you will realize that you are already familiar with many metals.

• Luster is one of the first properties that everyone notices about metals. Here we have pieces of iron, zinc, copper, tin, silver, and lead. All of them are as shiny as a mirror. (You know, of course, that the shiny surface of a mirror is a thin layer of metal?)

• No matter how thin you make a piece of metal, and they can be made into extraordinarily thin sheets, they remain virtually opaque. This first piece of metal is a sheet of iron. This is some aluminum foil. Notice how much thinner the foil is than the iron sheet. This is a piece of gold leaf. It is so thin that it can float like a feather in a current of air. However, not even the gold leaf will allow light to pass through.

• The property of metals that we use when we roll or hammer them into thin sheets is called malleability. This sheet of iron is half a millimetre thick. The piece of gold leaf, however, is 9 x 10^-5 millimetre thick or 9 one-hundred-thousandths of a millimetre thick!

• Metals are also ductile, they can be pulled into finer and finer wires. If you'll look at the TV monitor, you will see a number of copper wires, big and small. Also remember the wires you have seen in other situations - bundles of thin copper wires in electric cords, the tiny copper wires on computer chips, the small iron wires in steel wool, etc. Metals are the most malleable and ductile of all substances.

• Since metals are malleable and ductile, they are generally not brittle. This substance is silicon. It has the luster of metals, but you will notice that it isn't in a flat sheet. That might be because it isn't malleable. Watch what happens as we try to flatten it into a sheet. Silicon is not malleable, rather it is brittle.

• Metals feel hot or cold to our hand compared to other substances. Think of taking a metal ice cube tray out of the freezer compared to doing the same thing with a plastic one. Or using your bare hand to take a metal cookie sheet out of the oven compared to a wooden pizza board. We have here a beaker of boiling water with a metal rod and a wooden rod standing in it. You will notice that we also have a beaker with ice water and a metal rod and wooden rod in it. How many of you will volunteer to come up here and pick up each rod (as Bob is doing)?! The reason why we are careful when we touch a metal object is the ability of metals to conduct heat. When we touch a cold metal object, it rapidly conducts heat from our hand. A metal will just as rapidly conduct heat from hot surroundings to our hand.

• At room temperature, all the metals but one are solid. The exception is mercury, which you can see on the TV screen.

Metal Oxides

Faraday divided the oxides of metals into four groups: the alkalies, the earths, the acids, and all the other oxides. We begin our look at the metal oxides with some of the oxides that are not alkalies, earths, or acids.

Metals can burn! Some burn rapidly, giving off heat and light. Some react with oxygen in the air to produce an oxide (or tarnish) layer the instant they are exposed to the air. Other metals can be made to burn, if the conditions are right.

• If you'll observe the TV monitor screen again. The experimenter has a piece of potassium (another metal) on the glass. He slices some off which exposes a shiny surface of pure potassium. The potassium immediately reacts with the oxygen in the air to form a layer of tarnish over the shiny surface. The tarnish is a mixture of K₂O₂ and KO₂.

• Please keep your attention on the TV screen. The experimenter has a piece of potassium again. This time he is going to drop it into some water in a beaker. The potassium rapidly reacts with the water, releasing hydrogen which immediately catches fire.

• Bob has a pinch of iron filings in his hand. Observe what happens when he sprinkles them in the candle flame. Is the iron burning? The product you will observe is black. An oxide is formed, black iron(II) oxide.

• If you'll watch the TV monitor screen again, you'll see the experimenter ``burning'' copper or forming oxide layers on a sheet of copper. If you observe closely, you'll see that there are at least two oxides formed. At first you see this purple shade on the copper sheet. Then, as the copper gets hotter, it becomes black. The purplish-red layer was copper (I) oxide; the black was copper (II) oxide.

• We have here other samples of pairs of oxides formed by some metals. Each has properties different from those of the other oxide of that metal. Each, however, has a definite and constant composition. The samples we have are black iron (II) oxide and red iron (III) oxide, green vanadium (III) oxide and yellow vanadium (V) oxide, together with purplish-red copper (I) oxide and black copper (II) oxide.

• Bob has here an overhead transparency showing the names and formulas of the oxides we have just shown you. Note that each oxide has its own mole ratio.

Metal Oxides - Alkalies

We are now going to show reactions of alkali compounds that Faraday considered oxides. One example of this is found in our next demonstration. What he called zinc oxide is actually zinc hydroxide. Faraday and his contemporaries often spoke more or less interchangeably of acids/bases and their anhydrides. This is a reaction between a metal hydroxide and an acid to form a salt.

• In the process of working out Faraday's demonstration showing zinc hydroxide's reaction with sulfuric acid, Mary happened upon this nifty reaction. She begins with solid zinc carbonate and reacts it with sodium hydroxide to form a solution of sodium zincate, Na₂Zn(OH)₄. This ion forms because of zinc hydroxide's amphoteric nature. She then reacts this with sulfuric acid by carefully floating the sulfuric acid solution on top of the denser base solution. At the interface between the base and acid, a gel-like precipitate of zinc (II) hydroxide, Zn(OH)₂ forms. The indicator phenol red shows that the bottom solution is basic, the precipitate in the middle is a weak base, and the top, which contains zinc sulfate in solution, is neutral.

• When a solution of a copper salt reacts with an alkali, you produce a lovely precipitate. Notice that the copper sulfate solution and the potassium hydroxide solution we start with are perfectly transparent. However, when we mix these two . . . .

• Other precipitates form when potassium chromate is mixed with solutions of such metals as silver, mercury and lead. Eric is combining solutions of potassium chromate and lead acetate. Notice the beautiful color of the precipitate that forms. For a long time such insoluble metal compounds were used as pigments in artists' colors and other paints. Now we tend to use less poisonous compounds.

You will remember that we had a reaction a while ago in which potassium reacted with water. The solution formed in that reaction is a base when tested with indicators. (It also feels soapy on the skin and will attack skin and other proteins.) The compound in this solution, potassium hydroxide or ``potash'', will also neutralize acids.

• Eric is neutralizing potash, potassium hydroxide, with hydrochloric acid. You will know when the solution is neutral by the change in color from red-purple to colorless, since he is using phenolphthalein as the indicator.

Metallic Oxides - Earths

The last two groups of metal oxides in Faraday's scheme are the acids and the earths. Some examples of earths are silica (silicon dioxide), alumina (aluminum oxide), lime (calcium oxide), and magnesia (magnesium oxide). The last group of oxides, the acids, were actually non-metal (or amphoteric metal) oxides. (Arsenious oxide is one example.) Faraday classified certain non-metals as metals because of some of their physical properties.

Demonstrations

GENERAL INFORMATION:

Purpose	To demonstrate the properties and reactions of oxides. To classify the oxides on the basis of their properties and reactions.
Materials	sodium oxide, water, phenolphthalein, calcium oxide, calcium carbonate, lime water, 1M sodium hydroxide, sand, alumina, 1M hydrochloric acid, iron oxide, candle, bromothymol blue, 1M potassium thiocyanate, beakers, stirring rods, Erlenmeyer flask with 1 holed rubber stopper to fit, bent glass tubing (90 degrees), Petri dish, rubber tubing, test tubes, matches.
Procedure	outlined in each demonstration
Hazards	All acid and base solutions should be handled with care. Neutralize acids and bases before disposal. Use goggles and aprons.
Equations	Na2O(s) + H2O(l) --> 2 NaOH(aq) CaO(s) + H2O(l)(r) Ca(OH)2(aq) CaCO3(s) + heat --> CaO(s) + CO2(g) CO2(g) + Ca(OH)2(aq) --> CaCO3(s) + H2O(l) CaCO3(s) + CO2(g) + H2O(l) --> Ca(HCO3)2(aq) Al2O3(s) + 6HCl(aq) --> 2AlCl3(aq) + 3H2O(l) Al2O3(s) + OH-(aq) --> 2 Al(OH)4-(aq) Fe2O3(s) + HCl(aq) --> FeCl3(aq) + H2O(l) Fe3+(aq) + SCN-(aq) --> FeSCN2+(aq) paraffin(s) + O2(g) --> CO2(g) + H2O(l) CO2(g) + H2O(l) --> CO2(aq) {H2CO3}

Metals and Their Oxides

Properties of Metals

Metals can be distinguished from other elements by their brilliant luster. Generally, metals have a greater density than non-metal substances.

Lab. 1

Procedure	Display a large number of metal objects on your demonstration table or a side table. Be sure to have a number of objects that are familiar to your audience.

Lab. 2

Procedure	Demonstrate the luster of some sample metals by simply holding them up. Again, include as many common objects as you can. Examples might be a stainless steel knife, silver spoon, gold ring or chain, some aluminum foil, a copper sheet, etc. Be sure all tarnish or oxidation is removed.

Lab. 3

Procedure	Demonstrate opacity by placing in succession a thin metal sheet, a piece of metal foil, and then a piece of gold leaf on the overhead projector and noting that the light is blocked.

Lab. 4

Procedure	Demonstrate the malleability of metals by noting the relative thickness of each of the three samples. If you have gold leaf, mention that gold can be rolled and hammered into a sheet as thin as 3.54 X 10-6 inch or 9.01 x 10-6 centimeter.

Lab. 5

Procedure	Wires illustrate that a metal has the property of ductility. A metal can be pulled into wires as fine or finer than human hair. Only metals have the properties of malleability and ductility. Polymers, however, are often ductile and can be spun into fibers.

Lab. 6

Procedure	Metals are malleable and ductile, not brittle. To demonstrate this property, strike samples of antimony and silicon with a hammer. Compare their behavior to that of lead, copper, and zinc. Antimony and silicon are metalloids with some properties of metals and some of non-metals. They have a metallic luster but are brittle.

Lab. 7

Procedure	Another property of metals is that they have high tensile strength, i. e., they resist being pulled apart. Demonstrate this by suspending increasingly heavy weights from a fine wire.

Lab. 8

Procedure	Have a poster or picture of skyscrapers with some under construction. Call attention to the girders. Discuss the weight that is supported by a girder made of steel.

Lab. 9

Procedure

Another characteristic of metals is their ability to conduct both heat and electricity. Metals will more readily conduct heat to or from your hand than will non-metals. This can be demonstrated by having a beaker of ice-water and a beaker of boiling water on the demonstration table with the end of a metal rod and a wooden rod of similar diameter immersed in each. Invite a member of the audience to feel the metal and wood in each beaker and report the sensation.

CAUTION

Steam and hot water can cause burns.

Lab. 10

Procedure	The common metals, the ones that most people recognize, such as iron, copper, silver, gold, have a higher density than most other substances. On the demonstration table, display samples having fairly uniform volume. Pass them around the class having students compare the mass of the samples with metallic luster to similar size samples of rock, marble, wood, or other available non-metals.

Lab. 11

Procedure

Another way to compare the density of metals to non-metals is to use two large graduated cylinders, one almost full of water, the other almost full of colorless syrup. Drop pieces of two or three metals into each cylinder. Samples of lead, copper, zinc, iron, and tin work well. Drop in samples of hard plastic, marble, and wood for comparison. Have students note the rate at which the samples fall through the liquids. Although the metals with which we are most familiar are dense, there are some metals such as potassium and sodium that are less dense than water. This will be shown later in this lecture.

Lab. 12

Procedure

At room temperature, all the metals but one are solid. The exception, of course, is mercury. You can, however, melt metals and allow them to crystallize into solids again as they cool. Show crystals of lead, zinc, brass, and other metals available. You can find crystals of zinc on galvanized substances like buckets and sheet metal. Old brass doorknobs show crystals on their surfaces where chemicals from repeated handling have etched them.

Metal Oxides

Metals can burn, some more readily than others. Some, for example, burn rapidly on exposure to the air. Others burn only in special conditions.

Lab. 13

Procedure	Cut a piece of sodium to expose a fresh surface. Have a piece of aluminum foil near by to compare luster as the surface of the sodium oxidizes.

Lab. 14

Procedure

Carefully drop a split-pea size piece of sodium into a 600 mL or larger beaker which is two-thirds filled with water. Have students observe the activity being certain that they note that the sodium floats on the water and reacts instantly with it. After the reaction is over, add a few drops of an indicator to show that the liquid is now a base. Set the beaker and liquid aside to bring back later in the lecture.

CAUTION

Sodium is highly reactive. Consult MSDS before using. Substitution: Use calcium metal turnings, taking care not to touch with fingers. The effect can be enhanced by the addition of some sodium chloride crystals. As a safer alternative: USE A VIDEO DISC.

Lab. 15

Procedure	Place a small amount of finely divided zinc in the bottom of a crucible. Heat it with a Bunsen burner until it is glowing. Remove the burner and allow the zinc to burn. Have students observe the results of the burning.
CAUTION	Handle the crucible with tongs to prevent burns.

Lab. 16

Procedure	While you have a Bunsen burner flame, sprinkle a few iron filings in the flame. In both cases the metals are combining rapidly with the oxygen in the air. The compounds that are being formed are metal oxides.

Lab. 17

Procedure	Less reactive metals, such as lead and silver, will also burn under the proper conditions. The following method has been used: heat a charcoal block while blowing on its surface with a blowpipe. When the block is glowing red, drop small pieces of the metal onto the surface. Direct a stream of oxygen onto the pieces of metal.
CAUTION	THIS REQUIRES SPECIAL SAFETY CONSIDERATIONS AND SHOULD NOT BE UNDERTAKEN UNDER ORDINARY CLASSROOM CIRCUMSTANCES.

When metals tarnish or rust, they are undergoing slow oxidation. The compound formed is an oxide, just as when metals burn. When tarnishing takes place, no heat or light is apparent.

Lab. 18

Procedure	Show sheets of several metals. Cut off small pieces. Using tongs, hold in the burner flame. Remove and after the metal is cool compare the heated part with the unheated portion. Compare both of these to some rusted iron.

Lab. 19

OMIT	Safety Hazard: Lead Oxide

Lab. 20

Procedure	To show rapid oxidation safely, prepare a small jar of oxygen. Remove the oxygen generating apparatus from the area. Using tongs, hold some steel wool in a burner flame until it is glowing red. Lower the steel wool into the jar of oxygen.
CAUTION	Oxygen should be handled with care in the presence of combustibles.

Lab. 21

Procedure	Zinc foil can be substituted for the steel wool in the above demonstration. The same caution applies.

Faraday divided the metal oxides into four classes which he called: the alkalies, the earths, the acids, and the ordinary oxides. He discussed the ordinary oxides first.

Lab. 22

Procedure	Burn some steel wool in pure oxygen.
CAUTION	See Lab. 20.

Lab. 23

Procedure	Prepare rusty steel wool by leaving wet steel wool in the air. Compare the oxide that formed on the steel wool in the pure oxygen to the rust that formed on the wet steel wool that was left in the air.

Lab. 24

Procedure	Using tongs, hold a sheet of copper so that the edge is in a Bunsen burner flame. Remove from heat and point out the different colored oxides that have formed on the sheet indicating different oxides. The copper(II) oxide is blackish and copper(I) oxide is orange-red.

Lab. 25

Procedure	Different oxides of other transition metals can also be displayed. Stress that there are two or more oxides of each transition metal, each with a definite molar ratio.

Lab. 26

Procedure	Show a chart of the names and formulas of the transition metals and refer to the display table for samples. These oxides are solid, opaque compounds with varying degrees of luster. Most have a high density. Many are found in naturally occurring ores. They are generally insoluble.

Lab. 27

Procedure	Stir samples of two or three of the oxides from Lab. 25 into bottles containing water. Note that the oxides do not dissolve. Most metallic oxides are insoluble in water. Some oxides dissolve and exhibit acidic properties in solution.

Lab. 28

Procedure

Arsenic(V) oxide in water forms arsenic acid. Testing with an indicator such as litmus will confirm this. arsenic(V) oxide + water --> meta-arsenic acid As2O5 + 4H2O --> 2H3AsO4. 1/2 H2O These oxides were called by Faraday metallic acids. They are metallic oxides with high oxygen content, for example Mn2O7 and CrO3. The metallic oxides which are not acidic (groups I and II) have low oxygen content, for example Na2O and K2O. These combine with acids to form salts many of which are of great importance.

Lab. 29

Procedure	Slowly add zinc oxide (ZnO) to dilute sulfuric acid. Using an indicator, test for the neutralization of the acid. zinc oxide +sulfuric acid --> zinc sulfate + water ZnO + H2SO4 --> ZnSO4 + H2O

Lab. 30

Procedure	Show samples of salts such as: zinc sulfate, ZnSO4, formerly called white vitriol; copper(II) sulfate, CuSO4, formerly called blue vitriol; and iron(II) sulfate, FeSO4, formerly called green vitriol. These sulfates are examples of metallic salts and not only the oxide, but the metal also may be recovered from them by proper means.

Lab. 31

Procedure	A solution of copper(II) sulfate forms a precipitate when reacted with aqueous potassium hydroxide. The dried precipitate is copper(II) oxide. Copper(II) sulfate + Potassium hydroxide --> Copper(II) hydroxide + potassium sulfate CuSO4 + 2KOH --> Cu(OH)2 + K2SO4 Copper(II) hydroxide --> Copper(II) oxide + water Cu(OH)2 --> CuO + H2O

Lab. 32

Procedure	A solution of copper(II) sulphate reacts with metallic iron yielding copper metal. copper(II) sulfate + iron --> iron(II) sulfate + copper CuSO4 + Fe --> FeSO4 + Cu

Oxides that form acidic solutions combine with basic oxides to neutralize them.

Lab. 33

Procedure	Show examples of salts such as potassium chromate or potassium dichromate. When oxides that form acidic solutions are united to alkaline oxides they generally form insoluble salts, often possessing beautiful colors.

Lab. 34

Procedure	Silver(I) chromate and mercury(II) chromate are red. Lead(II) chromate is yellow. Show samples. A great many pigments are compounded from such substances.

If we now examine the metals which are most reactive, we find that some possess very extraordinary properties. The metal potassium for instance, reacts vigorously with water.

Lab. 35

Procedure	A solution of potassium hydroxide, when tested with turmeric paper or any indicator, proves to be alkaline like ammonia. The solution also corrodes the skin, feels slippery, and neutralizes ordinary acids.

Lab. 36

Procedure	Show neutralization of an acid solution with the potassium hydroxide, using an indicator. Suggestion: Use universal indicator to show range of pH changes. If the resulting solution is evaporated, it leaves a solid, white potassium salt.

Lab. 37

Procedure	Show an alkaline compound of potassium, solid KOH. This substance has no smell and is not volatile like ammonia. It can be obtained from wood ashes; hence, the term alkali which is derived from the Arabic for ``from ashes''.

Lab. 38

Procedure	Moisten ashes of charcoal and test with indicator paper. The presence of alkaline material is readily demonstrated.

Lab. 39

Procedure	Burn a piece of wood, moisten the ash, and test with an indicator. A basic reaction is shown.

Lab. 40

OMIT

When a base is combined with an acid, it forms a salt; for example, combining with nitric acid produces a nitrate salt.

Lab. 41

Procedure	Show some specimens of nitrates; for example, potassium nitrate. Burning sea weed produces the alkaline compound, sodium hydroxide. It has properties very similar to potassium hydroxide. The metal sodium is a product of its decomposition.

Faraday divided the compounds of metals with oxygen into four categories: alkalies, earths, oxides, and acids. The earths were further subdivided into alkaline earths and non-alkaline earths. We will describe these groups in terms of today's chemistry. What Faraday called the alkalies corresponds to oxides of our group IA (1), the alkali metals.

Lab. 42

Procedure	Dissolve sodium oxide in water. It is soluble as sodium hydroxide.

Lab. 43

Procedure	Test the resulting solution with phenolphthalein. The indicator turns pink. These oxides are called basic anhydrides since they form basic solutions. Faraday's alkaline earths correspond to group IIA (2), alkaline earth oxides.

Lab. 44

Procedure	Attempt to dissolve calcium oxide in water. It is only slightly soluble.

Lab. 45

Procedure	Test solution with phenolphthalein. The indicator turns pink.

Group II oxides are also basic anhydrides.

Lab. 46

Procedure	Heat calcium carbonate (marble chips) in an Erlenmeyer flask with a one-holed stopper through which a bent piece of glass has been placed. Attach rubber tubing to the glass tube and bubble the gas through a solution of limewater in a test tube. Calcium carbonate precipitates.

Lab. 47

Procedure	Continue heating the calcium carbonate but direct the gas through a solution of sodium hydroxide. No precipitate forms. The results observed illustrate the difference between the solubilities of group I and group II carbonates.
NOTE	Excess carbon dioxide causes the formation of calcium hydrogen carbonate which is soluble.

Faraday's non-alkaline earths correspond to the modern metalloid classification.

Lab. 47

Procedure	Add water to sand, silicon dioxide. It does not dissolve.

Lab. 48

Procedure	Add water to alumina, aluminum oxide. It does not dissolve.

Lab. 49

Procedure	Divide each of the suspensions into two parts. Add acid to a portion of each. If freshly prepared, the aluminum oxide dissolves while the silicon dioxide does not dissolve.

Lab. 50

Procedure	Add base to the remaining portions. The aluminum oxide dissolves, the silicon dioxide does not. Since aluminum oxide dissolves in both acidic and basic solutions, aluminum is called amphoteric.
NOTE	Gallium, germanium, and indium, which are also metalloids, were not known in Faraday's time.

Faraday's ordinary oxides correspond to the oxides of transition metals and the heavy metals.

Lab. 51

Procedure	Add water to rust. It does not dissolve.

Lab. 52

Procedure	Add acid to another sample of rust. It dissolves.

Lab. 53

Procedure	Add base to a sample of rust. No change is observed. Iron(III) oxide is not amphoteric. It dissolves in hydrochloric acid to form a characteristic yellow solution, but does not dissolve in base.

Lab. 54

Procedure	Add a solution of thiocyanate ions (as KSCN) to the acidified iron(III) chloride solution. The resulting red complex is a test for the presence of iron(III) ions.

Faraday makes reference to ``metallic acids'' such as the anhydride of arsenious acid. We can expand this category to include all those oxides which yield an acidic solution when dissolved in water. These compounds are oxides of non-metals. They are referred to as acid anhydrides. NOTE: Faraday's example (arsenic oxide) would be more properly classified as a metalloid compound today.

Lab. 55

Procedure

To a few milliliters of water in a Petri dish add a few drops of bromothymol blue. Place a candle in the center of the dish. Light the candle. Cover it with a beaker. Allow the candle to burn until extinguished. Shake the water without breaking the seal to dissolve the carbon dioxide. The bromothymol blue color changes shows that the solution is acidic. The carbon dioxide produced by the combustion of the candle wax dissolves in the water, producing a solution of carbonic acid.

Student activities

To further explore the ideas of this lecture students can:

• design a poster to advertise this lecture.

• write a review of the lecture as if written for publication in a newspaper of the time.

• research and present information about the development of the pH scale and the development of the use of pH measurement.

• prepare concept maps using the terms from the lecture.

• research to find elements that could be used as substitutes for the ones that Faraday used in his lecture.

• write a critique of the safety of Faraday's experiments, using Flynn catalogue and MSDS sheets for sources.

Authors

Julianne Shepelavy, Michael Sixtus, Eric Stelter, Anne Stowe, Susana Suarez, Gail Thompson, Kathleen Thompson, Robert Van Milligan and Alice Veyvoda.

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