US20050217432A1

US20050217432A1 - Apparatus and method for the reduction of metals

Info

Publication number: US20050217432A1
Application number: US11/134,735
Authority: US
Inventors: Linnard Griffin
Original assignee: Individual
Current assignee: Individual
Priority date: 2003-11-24
Filing date: 2005-05-20
Publication date: 2005-10-06
Also published as: EP1689915A1; US20050109162A1; EP1689915A4; WO2005052219A1

Abstract

Described is an apparatus for the production of an elemental metal from a metal-containing compound comprising a reaction medium containing ions of a first metal and a second metal, wherein the second metal is in colloidal form, and a related method.

Description

RELATED APPLICATIONS

This application is a continuation-in-part of application Ser. No. 10/995,934, filed Nov. 23, 2004, which is hereby incorporated by reference for all purposes and which claims priority to provisional applications Ser. No. 60/524,469, filed Nov. 24, 2003; Ser. No. 60/531,763, filed Dec. 22, 2003; and Ser. No. 60/531,764, filed Dec. 22, 2003.

TECHNICAL FIELD

The present invention is directed to a method and apparatus for the production of metals from metal ore.

BACKGROUND

Most metals are found in nature in their oxidized form. In order to extract these metals from their ores, it is necessary to chemically reduce these metals to their elemental form. The reduction of these metals usually requires stringent reaction conditions and therefore results in a significant cost. For example, iron, as found in nature, is generally in the oxidized form iron(III) oxide, Fe₂O₃, or iron(II) oxide, FeO, or a combination of the two: magnetite, Fe₃O₄. The reduction of Fe⁺²or Fe⁺³to yield Fe is normally carried out at very high temperatures, generally in excess of 1000° C. This reduction is commonly accomplished by reaction of the iron(III) oxide with carbon as shown in equation 1.
2Fe₂O₃+3C→4Fe+3CO₂ (1)
The very high temperature required for the reaction employed in equation 1 causes the generation of metallic iron from its oxide to be difficult to achieve, and very expensive. Likewise, aluminum, as found in nature, is generally in the oxidized form aluminum oxide, Al₂O₃. The reduction of Al⁺³to yield Al is normally carried out using a procedure called the Hall-Heroult process. In the Hall-Heroult process, aluminum oxide, Al₂O₃, is dissolved in a carbon-lined bath of molten cryolite, Na₃AlF₆. Aluminum fluoride, AlF₃, is also present to reduce the melting point of the cryolite. The reactants are then electrolyzed, and liquid aluminum is produced at the cathode. The carbon anode is oxidized and forms gaseous carbon dioxide. The net chemical reaction that describes this process is shown in equation 2. The very high temperature (about 600° C.) required for the reaction employed in equation 2 causes the generation of metallic aluminum from its oxide to be difficult to achieve, and the need for the electrical energy necessary for electrolysis causes the production of aluminum to be a very expensive process.
2Al₂O₃+3C→4Al+3CO₂ (2)
Accordingly, there exists a need for a method and apparatus for the production of metals from metal ore that requires less extreme conditions and, accordingly, can be done at a lower cost.

SUMMARY

The above-described need has been addressed by providing an apparatus for the production of an elemental metal from a metal-containing compound which comprises a solution containing ions of a first metal and a second metal, wherein the second metal is in colloidal form.
In another embodiment of the apparatus, the second metal is less reactive than the first metal.
In an additional embodiment, the second metal is more reactive than the first metal.
In an additional embodiment, the apparatus further comprises a third metal, wherein the third metal is in colloidal form.
In an additional embodiment, the third metal is more reactive than the first metal.
In an additional embodiment, the apparatus further comprises a vessel for containing the solution, wherein the vessel is inert to the solution. In an additional embodiment the vessel is configured to maintain an internal pressure greater than atmospheric.
In an additional embodiment, the second metal is silver, gold, platinum, tin, lead, copper, zinc, iron, aluminum, magnesium, beryllium, nickel or cadmium.
In an additional embodiment, the apparatus further comprises a solid comprising the first metal in contact with the solution.
In an additional embodiment, the solid comprising the first metal is a metal oxide.
In an additional embodiment, the apparatus further comprises an energy source. In an additional embodiment, the energy source supplies electric energy.
In an additional embodiment, the apparatus further comprises a cathode and an anode in electrical contact with the solution.
In an additional embodiment, the temperature of the solution is less than 500° C.
In an additional embodiment, the apparatus further comprises an elemental non-metal in contact with the solution. In an additional embodiment, the elemental non-metal is carbon.
In an additional embodiment, the apparatus further comprises ions of a salt dissolved in the solution. In an additional embodiment, a cation of the salt is higher on the electromotive series than the first metal. In an additional embodiment, the salt is aluminum sulfate, magnesium sulfate or potassium aluminum sulfate.
In an additional embodiment, the solution comprises a reducing agent. In an additional embodiment, the reducing agent is hydrogen peroxide. In an additional embodiment, the reducing agent is formic acid.

BRIEF DESCRIPTION OF THE DRAWINGS

The drawing FIGURE is a diagram of equipment which may be used in connection with one embodiment of the present invention.

DETAILED DESCRIPTION

The drawing FIGURE shows equipment which may be used in one embodiment of the present invention. A reaction vessel 102 contains a reaction medium 104. The reaction medium 104 preferably comprises water and an acid, and preferably has a pH less than 7, although other reaction media may be used, including other solvents or non-liquid media such as gelatinous or gaseous media. A cathode 106 and an anode 108 are preferably in electrical contact with reaction medium 104. Cathode 106 is preferably in the form of a disk made of carbon, but metallic materials such as lead and iron may also be used. Cathode 106 is preferably positioned on or near a bottom 107 of vessel 102. However, cathode 106 may generally be any shape and may be positioned anywhere that is in contact with reaction medium 104 and not in direct contact with anode 108. Cathode 106 may be made of any material which is inert or of lower reactivity than the metal being reduced and is electrically conductive. Anode 108 is preferably in the form of a rod made of carbon, but metallic materials such as lead and iron may also be used. Anode 108 is preferably positioned to extend into reaction medium 104 through a top surface of reaction medium 104. However, anode 108 may generally be any shape and may be positioned anywhere in contact with reaction medium 104 and not in direct contact with cathode 106. Anode 108 may be made of any material which is inert or of lower reactivity than the metal being reduced and is electrically conductive.
Vessel 102 also preferably contains an elemental non-metal 109 in contact with reaction medium 104. The elemental non-metal is preferable in solid form, for the sake of convenience, though a gaseous or liquid form should also work. Most preferably, elemental non-metal 109 is carbon, since carbon is relatively abundant and the byproducts produced in the resulting reactions (described below) are not toxic or environmentally harmful.
An electrical potential source 110 is electrically connected to cathode 106 and anode 108 and provides an electrical potential between cathode 106 and anode 108. Electrical potential source 110 preferably provides a direct current potential of the approximate order of or greater than 12 volts. The invention described herein has been performed using as little as 1 volt and as great as 12 volts. It has been found that higher voltages increase the overall reaction rate. As an alternative to providing electrical energy, other forms of energy, such as thermal energy, or light or other electromagnetic radiation energy, may be provided to reaction medium 104. Additionally, ambient thermal energy may be used, but the necessary reactions (discussed below) will occur at a slower rate. However, unlike prior art methods of metal production, the necessary reactions do not require that the reaction medium 104 temperature be extremely high. The necessary reactions have been observed to occur at a reasonable rate at temperatures of about 75° C. when the energy source 110 provides electrical energy.
Reaction medium 104 contains ions (not shown) of a metal which is to be reduced to its elemental state. Preferably, the ion source is a metal-containing compound 112 in the solid state which is in contact with reaction medium 104. Most preferably, the metal-containing compound 112 is a metal ore found in nature, such as iron(III) oxide or aluminum oxide. Alternatively, the metal-containing compound 112 may be a derivative of a metal ore, such as aluminum hydroxide, Al(OH)₃. Additionally, the metal ions may be from a salt, either in a solid or dissolved form, a metal oxide or other ion source.
Reaction medium 104 also preferably contains a dissolved salt (not shown). The salt preferably comprises a metal cation that is higher on the electromotive series of metals than the metal of the metal-containing compound which is being reduced. The salts that have been found to be most effective are aluminum sulfate, Al₂(SO₄)₃, magnesium sulfate, MgSO₄, and potassium aluminum sulfate, KAl(SO₄)₂. However, in theory, any salt could have a similar effect.
Reaction medium 104 also contains a suspended colloidal catalyst (not shown). Most metals can be produced in a colloidal state in a liquid. A colloid is a material composed of very small particles of one substance that are dispersed (suspended), but not dissolved, in a liquid. Thus, colloidal particles do not settle out of a liquid even though they exist in the solid state. A colloid of any particular metal is then a very small particle of that metal suspended in a liquid. These suspended particles of metal may exist in the solid (metallic) form or in the ionic form, or as a mixture of the two. The very small size of the particles of these metals results in a very large effective surface area for the metal. This very large effective surface area for the metal can cause the surface reactions of the metal to increase dramatically when it comes into contact with other atoms or molecules. The colloidal metals used in the experiments described below were obtained using an apparatus for producing colloidal silver in water sold by CS Prosystems of San Antonio, Tex. The website of CS Prosystems is www.csprosystems.com. Based on materials provided by the manufacturer, the particles of a metal in the colloidal dispersions used in the experiments described below are believed to range in size between 0.001 and 0.01 microns. In such a solution of colloidal metals, the concentrations of the metals are believed to be between about 5 to 20 parts per million with the remainder being water.
Alternative to using a catalyst in colloidal form, it may be possible to use a catalyst in another form that offers a high surface-area to volume ratio, such as a porous solid or colloid-polymer nanocomposite. In general, any of the catalysts may be in any form with an effective surface area preferably of at least 298,000,000 m²per cubic meter of catalyst metal, although smaller surface area ratios may also work.
When a colloidal metal ion is treated with an oxidized metal, a voltaic oxidation-reduction will take place. The oxidized metal can be any compound where the metal is in a cationic form. Preferably, the oxidized metal will be the metal ore as found in nature. For many metals, this is the metal oxide (Me_xO_y). Equations 3 and 4 are believed to represent the oxidation and reduction reactions that occur with respect to the colloidal metal. Equations 5 and 6 are believed to represent the oxidation and reduction reactions that occur with the inclusion of elemental non-metal 109, represented by the letter “Z”. The process proceeds most successfully when elemental non-metal 109, Z, is either carbon or sulfur, but any non-metal may theoretically be employed.
The colloidal metal, M, can in principle be any metal, but it has been found that equations 3 and 4, or equations 5 and 6, work most efficiently when the colloidal metal has a higher (more positive) reduction potential than Me. Thus, equations 3 and 4 and equations 5 and 6 proceed most efficiently when the colloidal metal is as low as possible on the electromotive series of metals. Consequently, any colloidal metal will be successful, but the reactions illustrated in equations 3 and 4 and equations 5 and 6 proceed most quickly with colloidal silver ion, due to the high reduction potential of silver. When silver, for example, is employed as the colloidal metal ion in equations 3 and 4, or in equations 5 and 6, the pair of reactions is found to take place quite readily. The voltaic reaction produces a positive voltage as the oxidation and reduction reactions indicated take place. This positive voltage can be used to supply the energy required for other chemical processes. In fact, the voltage produced can even be used to supply an over potential for reactions employing equations 3 and 4, or equations 5 and 6, taking place in another reaction vessel. Thus, this electrochemical process can theoretically be made to take place more quickly without the supply of an external source of energy, if at least two of these reactions are performed in series. The addition of an external source of energy, such as thermal energy, electrical energy or light, or other electromagnetic radiation, will further enhance the reaction rate.
Cathode (Reduction)
4M⁺+4e ⁻→4M (3)
Anode (Oxidation)
2H₂O→4H⁺+O₂+4e ⁻ (4)
Cathode (Reduction)
4M⁺+4e ⁻→4M (5)
Anode (Oxidation)
2H₂O+Z→4H⁺+ZO₂+4e ⁻ (6)
It is believed that the oxidation-reduction reaction represented by reactions 5 and 6 occur faster and more easily than the oxidation-reduction reaction represented by equations 3 and 4 due to the thermodynamic stability of the non-metallic oxide, ZO₂.
The net result of the oxidation and reductions shown in equations 5 and 6 is equation 7, which results in the production of a colloidal metal in its elemental state plus a non-metallic oxide plus acid. $\begin{matrix} 4 M^{+} + 4 e^{-} -> 4 M + & (5) \\ 2 H_{2} O + Z \to 4 H^{+} + Z O_{2} + 4 e^{-} = & (6) \\ 4 M^{+} + 2 H_{2} O + Z \to 4 M + 4 H^{+} + {Z O}_{2} & (7) \end{matrix}$
In the absence of the non-metal, Z, the net result of the oxidation and reductions shown in equations 3 and 4 is equation 7A, which is believed to result in the production of a colloidal metal in its elemental state, plus elemental oxygen, plus acid. $\begin{matrix} 4 M^{+} + 4 e^{-} \to 4 M + & (3) \\ 2 H_{2} O \to 4 H^{+} + O_{2} + 4 e^{-} = & (4) \\ 4 M^{+} + 2 H_{2} O \to 4 M + 4 H^{+} + O_{2} & (7 A) \end{matrix}$
The colloidal elemental metal that has been produced is believed to undergo reaction with the metal ion of the substance that contains the oxidized form of the metal, which will be represented as M_e ⁺. M_e ⁺ can represent the oxidized form of any metal, which can be present in any oxidation state. Equation 8 illustrates this reaction where the oxidized form of the metal, M_e, is an oxide, but in reality can be any compound that contains the metal, M_e, in its oxidized form.
4M+2M_eO+2H₂O→4M⁺+2M_e+4OH⁻ (8)
The reaction illustrated by equation 8 will take place most efficiently when the colloidal metal, M, is more reactive than the metal, M_e. That is, the reaction in equation 8 will proceed most efficiently when the colloidal metal, M, is above the metal, M_e, on the electromotive series of metals. The hydroxide ion produced in equation 8 will react with the hydrogen ion produced in equation 7, or in equation 7A, to produce water as indicated in equation 9.
4H⁺+4OH⁻→4H₂O (9)
Since the acid produced in the electrochemical reaction depicted in equations 5 and 6 is neutralized by the base produced in the thermal reaction represented by equation 8, the entire reaction system remains at a pH close to 7 throughout. The combining of equations 5, 6, 8, and 9 results in the net process illustrated by equation 10, which represents the production of the elemental metal, M_e, produced by a reduction reaction, and the formation of an oxide of a non-metal, ZO₂, produced by an oxidation reaction. $\begin{matrix} 4 M^{+} + 4 e^{-} \to 4 M + & (5) \\ 2 H_{2} O + Z \to 4 H^{+} + Z O_{2} + 4 e^{-} + & (6) \\ 4 M + 2 M_{e} O + 2 H_{2} O \to 4 M^{+} + 2 M_{e} + 4 O H^{-} + & (8) \\ 4 H^{+} + 4 O H^{-} \to 4 H_{2} O = & (9) \\ 2 M_{e} O + Z \to 2 M_{e} + Z O_{2} & (10) \end{matrix}$
In the absence of the non-metal, Z, the combining of equations 3, 4, 8, and 9 results in the net process illustrated by equation 10A, which represents the production of the elemental metal, M_e, produced by a reduction reaction, and the formation of elemental oxygen, produced by an oxidation reaction. $\begin{matrix} 4 M^{+} + 4 e^{-} \to 4 M + & (3) \\ 2 H_{2} O + \to 4 H^{+} + O_{2} + 4 e^{-} + & (4) \\ 4 M + 2 M_{e} O + 2 H_{2} O \to 4 M^{+} + 2 M_{e} + 4 O H^{-} + & (8) \\ 4 H^{+} + 4 O H^{-} \to 4 H_{2} O = & (9) \\ 2 M_{e} O \to 2 M_{e} + O_{2} & (10 A) \end{matrix}$
The reactions shown in equations 3 and 4, or in equations 5 and 6, seem to occur best when the colloidal metal, M, is as low as possible on the electromotive series of metals (less reactive); however, the reaction depicted by equation 8 takes place most efficiently when the colloidal metal, M, is as high as possible on the electromotive series of metals. The net reaction, which is illustrated by equation 10, or by equation 10A, is merely the sum of equations 3, 4, 8, and 9 or of equations 5, 6, 8, and 9, and could, in fact, be maximally facilitated by either colloidal metals of high activity or by colloidal metals of low activity. The relative importance of the reaction illustrated by equations 3 and 4, or by equations 5 and 6, compared to the reaction shown in equation 8 determines the characteristics of the colloidal metal that would best assist the net reaction in equation 10 or in equation 10A. It has been found that the net reaction indicated in equation 10 or in equation 10A proceeds at a maximal rate when the colloidal metal is of maximum activity; that is, when the colloidal metal is as high as possible on the electromotive series of metals (more reactive). It has been found that the more reactive colloidal metals, such as, but not limited to, colloidal aluminum ion or colloidal magnesium ion, produce the most facile reduction processes for the reduction of cationic metals. It is also believed, although not yet shown analytically, that the overall reaction may proceed even more favorably when two colloidal metals are used, especially where one is higher (more reactive) and one lower (less reactive) on the electromotive series than the metal being reduced.
In addition, it has also been found that the inclusion of a small amount (such as 10 wt %) of a salt leads to a rate increase in the reaction represented by equation 10, or equation 10A. The salt has its maximal effect when it includes a cation of a metal of higher activity than M_e; that is, one that is higher (more reactive) than M_eon the electromotive series of metals. The salts that have been found to be most effective are aluminum sulfate, Al₂(SO₄)₃, magnesium sulfate MgSO₄, and potassium aluminum sulfate, KAl(SO₄)₂; however, in theory, any salt could potentially have a similar effect.
Thus, under ambient thermal conditions, the oxide of any metal can be converted to its metallic elemental state, with the concurrent formation of elemental oxygen or the oxide of a non-metal. It is believed that the thermal stability of the oxide of the non-metal, ZO₂, lowers the endothermicity of the process, and allows the reduction of the oxidized metal to proceed at lower temperatures, when the non-metal Z is used. The supplying of additional energy leads to an acceleration of the reaction rate for the process. When additional energy is supplied, it can be supplied in the manner of thermal energy, electrical energy, radiant energy, or the like. When the additional energy supplied is in the form of thermal energy, one is limited by the boiling point of the solvent. In aqueous systems, this would provide a maximum temperature of 100° C. Under pressures higher than one atmosphere, however, temperatures higher than 100° C. could be obtained, and would provide an even more enhanced reaction rate. It has been found that the increase in reaction rate is most significant when additional energy is supplied in the form of electrical energy.
An alternative to the above involves the introduction of a reducing agent into reaction medium 104. Hydrogen peroxide has been found to be an effective reducing agent for this process. With the addition of hydrogen peroxide to reaction medium 104, equations 5 and 6 are replaced by equations 11 and 12, or equations 3 and 4 are replaced by equations 11A and 12A.
Cathode (Reduction)
2M⁺+2e ⁻→2M (11)
Anode (Oxidation)
H₂O₂+Z→2H⁺+ZO₂+2e ⁻ (12)
Cathode (Reduction)
2M⁺+2e ⁻→2M (11A)
Anode (Oxidation)
H₂O₂→2H⁺+O₂+2e ⁻ (12A)
Due to the fact that hydrogen peroxide has a larger (less negative) oxidation potential than water, as indicated in the comparison of equations 4A and 4B, the oxidation-reduction reaction resulting from equations 11 and 12 takes place at an enhanced rate compared to the oxidation-reduction reaction indicated by equations 5 and 6. Likewise, the oxidation-reduction reaction resulting from equations 11A and 12A takes place at an enhanced rate compared to the oxidation-reduction reaction indicated by equations 3 and 4.
H₂O→4H⁺+O₂+4e ⁻ ε⁰=−1.229 Volts (4A)
H₂O₂→2H⁺+O₂+2e ⁻ ε⁰=−0.695 Volts (4B)
The net result of the oxidation and reductions shown in equations 11 and 12 will be equation 13, which results in the production of a colloidal metal in its elemental state plus a non-metallic oxide plus acid $\begin{matrix} 2 M^{+} + 2 e^{-} \to 2 M + & (11) \\ H_{2} O_{2} + Z \to 2 H^{+} + Z O_{2} + 2 e^{-} = & (12) \\ 2 M^{+} + H_{2} O_{2} + Z \to 2 M + 2 H^{+} + Z O_{2} & (13) \end{matrix}$
Likewise, the net result of the oxidation and reductions shown in equations 11A and 12A will be equation 13A, which results in the production of a colloidal metal in its elemental state, plus elemental oxygen, plus acid. $\begin{matrix} 2 M^{+} + 2 e^{-} \to 2 M + & (11 A) \\ H_{2} O_{2} \to 2 H^{+} + O_{2} + 2 = & (12 A) \\ 2 M^{+} + H_{2} O_{2} \to 2 M + 2 H^{+} + O_{2} & (13 A) \end{matrix}$
The colloidal elemental metal that has been produced is believed to undergo reaction with the metal ion of the substance that contains the metal to be reduced, which will be represented as M_e ⁺. M_e ⁺ can represent the oxidized form of any metal, which can be present in any oxidation state. Equation 14 illustrates this reaction where the oxidized form of the metal, M_e, is an oxide, but in reality can be any compound that contains the metal, M_e, in its oxidized form.
2M+M_eO+H₂O→2M⁺+M_e+2OH⁻ (14)
The reaction illustrated by equation 14 will take place most efficiently when the colloidal metal, M, is more reactive than the metal, M_e. That is, the reaction in equation 14 will proceed most efficiently when the colloidal metal, M, is above the metal, M_e, on the electromotive series of metals. The hydroxide ion produced in equation 14 will react with the hydrogen ion produced in equation 13, or in equation 13A, to produce water as indicated in equation 15.
2H⁺+2OH⁻→2H₂O (15)
Since the acid produced in the electrochemical reaction depicted in equations 11 and 12, or in equations 11A and 12A, is neutralized by the base produced in the thermal reaction represented by equation 14, the entire reaction system remains at a pH close to 7 throughout. The combining of equations 11, 12, 14, and 15 results in the net process illustrated by equation 16, which represents the production of the elemental metal, M_e, produced by a reduction reaction, and the formation of an oxide of a non-metal, ZO₂, produced by an oxidation reaction. $\begin{matrix} 2 M^{+} + 2 e^{-} \to 2 M + & (11) \\ H_{2} O_{2} + Z \to 2 H^{+} + Z O_{2} + 2 e^{-} + & (12) \\ 2 M + M_{e} O + H_{2} O \to 2 M^{+} + M_{e} + 2 O H^{-} + & (14) \\ 2 H^{+} + 2 O H^{-} \to 2 H_{2} O = & (15) \\ H_{2} O_{2} + M_{e} O + Z \to M_{e} + Z O_{2} + H_{2} O & (16) \end{matrix}$
If the non-metal, Z, is not used, the combining of equations 11A, 12A, 14, and 15 results in the net process illustrated by equation 16A, which represents the production of the elemental metal, M_e, produced by a reduction reaction, and the formation of elemental oxygen, produced by an oxidation reaction. $\begin{matrix} 2 M^{+} + 2 e^{-} \to 2 M + & (11 A) \\ H_{2} O_{2} \to 2 H^{+} + O_{2} + 2 + & (12 A) \\ 2 M + M_{e} O + H_{2} O \to 2 M^{+} + M_{e} + 2 O H^{-} + & (14) \\ 2 H^{+} + 2 O H^{-} \to 2 H_{2} O = & (15) \\ H_{2} O_{2} + M_{e} O \to M_{e} + O_{2} + H_{2} O & (16 A) \end{matrix}$
The reactions shown in equations 11 and 12, or in equations 11A and 12A, seem to occur best when the colloidal metal, M, is as low as possible on the electromotive series of metals. However, the reaction depicted by equation 14 takes place most efficiently when the colloidal metal, M, is as high as possible on the electromotive series of metals. The net reaction illustrated by equation 16, which is merely the sum of equations 11, 12, 14, and 15, or by equation 16A, which is merely the sum of equations 11A, 12A, 14, and 15, could, in fact, be maximally facilitated by either colloidal metals of higher activity or by colloidal metals of lower activity than the metal being reduced. The relative importance of the reaction illustrated by equations 11 and 12, or by equations 11A and 12A, compared to the reaction shown in equation 14, determines the characteristics of the colloidal metal that would best assist the net reaction in equation 16, or in equation 16A. It has been found that the net reaction indicated in equation 16, or in equation 16A, proceeds at a maximal rate when the colloidal metal is of maximum activity; that is, when the colloidal metal is as high as possible on the electromotive series of metals. It has been found that the more reactive colloidal metals, such as, but not limited to, colloidal aluminum ion or colloidal magnesium ion, produce the most facile reduction processes for the reduction of cationic metals. It is also believed, although not yet shown analytically, that the overall reaction may proceed even more favorably when two colloidal metals are used, especially where one is higher and one lower on the electromotive series than the metal being reduced.
In addition, it has also been found that the inclusion of a small amount of a salt leads to a rate increase in the reaction represented by equation 16, or by equation 16A. The salt has been found to have a maximal effect when it includes a cation of a metal of higher activity than the metal being reduced; that is, one that is higher (more reactive) on the electromotive series of metals. The salts that have been found most effective are aluminum sulfate, Al₂(SO₄)₃, magnesium sulfate MgSO₄, and potassium aluminum sulfate, KAl(SO₄)₂; however, in theory, any salt could potentially have a similar effect.
Thus, under ambient thermal conditions, the oxide of any metal can be treated with hydrogen peroxide and a non-metal, and can be converted to its metallic elemental state, with the concurrent formation of the oxide of a non-metal and water. Since the oxidation of hydrogen peroxide (equation 12 or equation 12A) is more favorable than the oxidation of water (equation 6 or equation 4), the rate of metal reduction should be significantly increased when hydrogen peroxide is used in the place of water. This must be balanced by the fact that hydrogen peroxide is a more costly reagent to supply. In those cases where the rate of the metal reduction is the most critical factor, the use of hydrogen peroxide will offer a significant advantage. It is still believed that the thermal stability of the oxide of the non-metal, ZO₂, lowers the endothermicity of the process, and allows the reduction of the oxidized metal to proceed at reasonable temperatures. The supplying of additional energy leads to an acceleration of the reaction rate for the process. When additional energy is supplied, it can be supplied in the manner of thermal energy, electrical energy, radiant energy or the like. When the additional energy supplied is in the form of thermal energy, one is limited by the boiling point of the solvent. In aqueous systems, this would provide a maximum temperature of 100° C. Under pressures higher than one atmosphere, however, temperatures higher than 100° C. could be obtained, and would provide an even more enhanced reaction rate. It has been found that the increase in reaction rate is most significant when additional energy is supplied in the form of electrical energy.
A further alternative to the above involves the introduction of a different reducing agent into reaction medium 104. Formic acid has been found to be an effective reducing agent for this process. With the addition of formic acid to reaction medium 104, equations 3 and 4 are replaced by equations 17 and 18.
Cathode (Reduction)
2M⁺+2e ⁻→2M (17)
Anode (Oxidation)
CH₂O₂→2H⁺+CO₂+2e ⁻ (18)
Due to the fact that formic acid has a larger (in fact, positive) oxidation potential than water, or than hydrogen peroxide, as indicated in the comparison of equations 4A, 4B and 4C, the oxidation-reduction reaction resulting from equations 17 and 18 takes place at an enhanced rate compared to the oxidation-reduction reaction indicated by equations 3 and 4, or by equations 11A and 12A.
H₂O→4H⁺+O₂+4e ⁻ ε⁰=−1.229 Volts (4A)
H₂O₂→2H⁺+O₂+2e ⁻ ε⁰=−0.695 Volts (4B)
CH₂O₂→2H⁺+CO₂+2e ⁻ ε⁰=0.199 Volts (4C)
The net result of the oxidation and reductions shown in equations 17 and 18 will be equation 19, which results in the production of a colloidal metal in its elemental state plus carbon dioxide plus acid $\begin{matrix} 2 M^{+} + 2 e^{-} \to 2 M + & (17) \\ C H_{2} O_{2} \to 2 H^{+} + C O_{2} + 2 e^{-} = & (18) \\ 2 M^{+} + C H_{2} O_{2} \to 2 M + 2 H^{+} + C O_{2} & (19) \end{matrix}$
The colloidal elemental metal that has been produced is believed to undergo reaction with the metal ion of the substance that contains the metal to be reduced, which will be represented as M_e ⁺. M_e ⁺ can represent the oxidized form of any metal, which can be present in any oxidation state. Equation 14 illustrates this reaction where the oxidized form of the metal, M_e, is an oxide, but in reality can be any compound that contains the metal, M_e, in its oxidized form.
2M+M_eO+H₂O→2M⁺+M_e+2OH⁻ (14)
The reaction illustrated by equation 14 will take place most efficiently when the colloidal metal, M, is more reactive than the metal, M_e. That is, the reaction in equation 14 will proceed most efficiently when the colloidal metal, M, is above the metal, M_e, on the electromotive series of metals. The hydroxide ion produced in equation 14 will react with the hydrogen ion produced in equation to produce water as indicated in equation 15.
2H⁺+2OH⁻→2H₂O (15)
Since the acid produced in the electrochemical reaction depicted in equations 17 and 18 is neutralized by the base produced in the thermal reaction represented by equation 14, the entire reaction system remains at a pH close to 7 throughout. The combining of equations 17, 18, 14, and 15 results in the net process illustrated by equation 20, which represents the production of the elemental metal, M_e, produced by a reduction reaction, and the formation of a carbon dioxide, produced by an oxidation reaction. $\begin{matrix} 2 M^{+} + 2 e^{-} \to 2 M + & (17) \\ C H_{2} O_{2} \to 2 H^{+} + C O_{2} + 2 e^{-} + & (18) \\ 2 M + M_{e} O + H_{2} O \to 2 M^{+} + M_{e} + 2 O H^{-} + & (14) \\ 2 H^{+} + 2 O H^{-} \to 2 H_{2} O = & (15) \\ C H_{2} O_{2} + M_{e} O \to M_{e} + C O_{2} + H_{2} O & (20) \end{matrix}$
The reactions shown in equations 17 and 18 seem to occur best when the colloidal metal, M, is as low as possible on the electromotive series of metals. However, the reaction depicted by equation 14 takes place most efficiently when the colloidal metal, M, is as high as possible on the electromotive series of metals. The net reaction illustrated by equation 20, which is merely the sum of equations 17, 18, 14, and 15, could, in fact, be maximally facilitated by either colloidal metals of higher activity or by colloidal metals of lower activity than the metal being reduced. The relative importance of the reaction illustrated by equations 17 and 18, compared to the reaction shown in equation 14, determines the characteristics of the colloidal metal that would best assist the net reaction in equation 20. It has been found that the net reaction indicated in equation 20 proceeds at a maximal rate when the colloidal metal is of maximum activity; that is, when the colloidal metal is as high as possible on the electromotive series of metals. It has been found that the more reactive colloidal metals, such as, but not limited to, colloidal aluminum ion or colloidal magnesium ion, produce the most facile reduction processes for the reduction of cationic metals. It is also believed, although not yet shown analytically, that the overall reaction may proceed even more favorably when two colloidal metals are used, especially where one is higher and one lower on the electromotive series than the metal being reduced.
In addition, it has also been found that the inclusion of a small amount of a salt leads to a rate increase in the reaction represented by equation 20. The salt has been found to have a maximal effect when it includes a cation of a metal of higher activity than the metal being reduced; that is, one that is higher (more reactive) on the electromotive series of metals. The salts that have been found most effective are aluminum sulfate, Al₂(SO₄)₃, magnesium sulfate, MgSO₄, and potassium aluminum sulfate, KAl(SO₄)₂; however, in theory, any salt could potentially have a similar effect.
Thus, under ambient thermal conditions, the oxide of any metal can be treated with formic acid, and can be converted to its metallic elemental state, with the concurrent formation of carbon dioxide and water. Since the oxidation of formic acid (equation 18) is more favorable than the oxidation of water (equation 4), or the oxidation of hydrogen peroxide (equation 12A), the rate of metal reduction should be significantly increased when formic acid is used in the place of water, or in the place of hydrogen peroxide. This must be balanced by the fact that formic acid, while less costly than hydrogen peroxide, is a more costly reagent to supply than is water. In those cases where the rate of the metal reduction is the most critical factor, the use of formic acid will offer a significant advantage. It is believed that the thermal stability of the oxide of the carbon dioxide that is formed lowers the endothermicity of the process, and allows the reduction of the oxidized metal to proceed at reasonable temperatures. The supplying of additional energy leads to an acceleration of the reaction rate for the process. When additional energy is supplied, it can be supplied in the manner of thermal energy, electrical energy, radiant energy, or the like. When the additional energy supplied is in the form of thermal energy, one is limited by the boiling point of the solvent. In aqueous systems, this would provide a maximum temperature of 100° C. Under pressures higher than one atmosphere, however, temperatures higher than 100° C. could be obtained, and would provide an even more enhanced reaction rate. It has been found that the increase in reaction rate is most significant when additional energy is supplied in the form of electrical energy.
Finally, while all equations depicted here involve the use of just a single metal, M_e, it has been shown that all of the reactions discussed herein can be carried out using an elemental metal in addition to the oxidized form of the metal being reduced, M_e ⁺. It has been shown, in fact, that in some cases the use of multiple metals results in a significant increase in the rate, as well as a significant increase in the yield of metal reduction. In experiment #13, experiment #16 and experiment #17, for example, elemental iron in the form of small iron nuggets is used to aid the reduction of oxidized aluminum, Al⁺³. In each of these cases, a sizable yield of metallic aluminum results from a completely thermal, non-electrolytic process requiring, at most, the input of only a small amount of thermal energy. It is not clear at this point what causes these impressive enhancements in the rate of the process as well as in the yield of reduced metal that results. It is possible that the elemental metal takes part in the reaction mechanism to provide a more complicated mechanism having a greater number of steps, but a lower net activation barrier. Another possibility is that the elemental metal might provide a surface where the reduced metal, M_e, could reform more efficiently. Whatever the actual explanation is, the results from experiments #13, #16 and #17 very clearly demonstrate that the effects resulting from the addition of an elemental metal different from M_ecan be quite significant in its effect upon metal reduction.
Experiments #13, #16 and #17 each result in a sample of iron that has become plated with elemental aluminum. Since this technology should be valid for pairs of metals other than iron and aluminum, what could result is a general method for the plating of surfaces with a metal without the use of electrolysis.
Experimental Results:
Several experiments have been conducted using combinations of embodiments of the technology described above. The results of those experiments are given below:
Experiment #1 Summary:
An experiment was conducted using 150 mL of iron (III) chloride in an aqueous solution (commonly used as an etching solution, purchased from Radio Shack) as the starting materials. Initially, 10 mL of 93% concentration sulfuric acid (H₂SO₄) was added to the solution, at which point no reaction occurred. About 50 mL of colloidal magnesium and 80 mL of colloidal lead, each at a concentration believed to be about 20 ppm, were then added, at which point a chemical reaction began and the bubbling of gases was evident at ambient temperature. The production of gas accelerated when the solution was heated to a temperature of about 65° C. The product gas was captured in soap bubbles and the bubbles were then ignited. The observed ignition of the gaseous product was typical for a mixture of hydrogen and oxygen.
Since, it is believed, the production of hydrogen gas could only be produced with a concurrent oxidation of iron, it is evident that the iron (III) had to be initially reduced before it could be oxidized, thereby providing strong evidence of the reduction reaction. This experiment has subsequently been repeated with hydrochloric acid (HCl) instead of sulfuric acid, with similar results.
Experiment #2 Summary:
An experiment was conducted using 100 grams of Fe₃O₄(this sample was found to contain roughly equal amounts of Fe₂O₃and FeO plus a small amount of elemental carbon), 50 mL of 5% H₂SO₄, plus 40 mL of colloidal magnesium and 40 mL of colloidal lead in water. Immediately a stream of gas was evolved that was identified as carbon dioxide by gas chromatography. The mixture was then heated to a temperature of 90° C. for a period of about three hours. At this point, the stream of gas being evolved was again analyzed by gas chromatography. This gaseous mixture was found to contain 40% hydrogen and 60% carbon dioxide. Since, it is believed, the production of hydrogen gas could only be produced with a concurrent oxidation of iron, it is evident that the iron had to be initially reduced before it could be oxidized, thereby providing strong evidence of the reduction reaction.
Experiment #3 Summary:
An experiment was conducted using 5 g of Al₂(SO₄)₃.18H₂O plus 40 mL of colloidal magnesium and 40 mL of colloidal lead in water. Upon being heated to about 75° C., a stream of gas, presumed to be elemental oxygen, was produced that did not ignite, and also did not extinguish a flame. After 45 minutes of heating, the gas was found to ignite very slightly when it was exposed to a flame, indicative of the production of a small amount of elemental hydrogen. Since, it is believed, the production of hydrogen gas could only be produced with a concurrent oxidation of aluminum, it is evident that the aluminum had to be initially reduced before it could be oxidized, thereby providing strong evidence of the reduction reaction.
Experiment #4 Summary:
An experiment was conducted using 5 g of Fe₂O₃plus 40 mL of colloidal magnesium, 40 mL of colloidal lead and 80 mL of 3% aqueous H₂O₂. Almost immediately a small amount of a gaseous product was produced. As the temperature was increased, over a period of ten minutes, the yield of gas increased with a maximum yield of gas being realized at the maximum temperature of about 75° C. The product gas was found to contain a substantial amount of hydrogen, based upon the manner in which it ignited when a flame was applied. Since, it is believed, the production of hydrogen gas could only be produced with a concurrent oxidation of iron, it is evident that the iron had to be initially reduced before it could be oxidized, thereby providing strong evidence of the reduction reaction.
Experiment #5 Summary:
An experiment was conducted using 5 g of Al(OH)₃plus 40 mL of colloidal magnesium, 40 mL of colloidal lead and 80 mL of 3% aqueous H₂O₂. Almost immediately a small amount of a gaseous product was produced. As the temperature was increased, over a period of ten minutes, the yield of gas increased with a maximum yield of gas being realized at the maximum temperature of about 75° C. The product gas was found to contain a substantial amount of hydrogen, based upon the manner in which it ignited when a flame was applied. Since, it is believed, the production of hydrogen gas could only be produced with a concurrent oxidation of aluminum, it is evident that the aluminum had to be initially reduced before it could be oxidized, thereby providing strong evidence of the reduction reaction.
Experiment #6 Summary:
An experiment was conducted using 5 g of Al₂(SO₄)₃.18H₂O plus 40 mL of colloidal magnesium, 40 mL of colloidal lead and 80 mL of 3% aqueous H₂O₂. Almost immediately a small amount of a gaseous product was produced. As the temperature was increased, over a period of ten minutes, the yield of gas increased with a maximum yield of gas being realized between the temperatures of 50° C. and 75° C. The product gas was found to contain a substantial amount of hydrogen, based upon the manner in which it ignited when a flame was applied. Since, it is believed, the production of hydrogen gas could only be produced with a concurrent oxidation of aluminum, it is evident that the aluminum had to be initially reduced before it could be oxidized, thereby providing strong evidence of the reduction reaction.
Experiment #7 Summary:
An experiment was conducted using 5 g of Fe₂O₃plus 40 mL of colloidal magnesium, 40 mL of colloidal lead and 1 gram of elemental carbon in water. The mixture was heated to a temperature of about 90° C. for a period of 72 hours. A metallic-like material was produced and collected that reacted with sulfuric acid to produce an ignitable gas presumed to be hydrogen gas. The metallic material is believed to be elemental iron.
Experiment #8 Summary:
An experiment was conducted using 5 g of Al(OH)₃plus 40 mL of colloidal magnesium, 40 mL of colloidal lead and 1 gram of elemental carbon in water. The mixture was heated to a temperature of about 90° C. for a period of 72 hours. A metallic-like material was produced and collected that reacted with sulfuric acid to produce an ignitable gas presumed to be hydrogen gas. The metallic material is believed to be elemental aluminum.
Experiment #9 Summary:
An experiment was conducted using 5 g of Fe₂O₃, 40 mL of colloidal magnesium and 40 mL of colloidal lead in water. A 12 volt, 10 amp power source was then applied for a period of 5 minutes to a pair of lead electrodes that had been introduced into the solution. A metallic-like material that was produced and found on the bottom of the apparatus was collected. The metallic material reacted with sulfuric acid to produce an ignitable gas presumed to be hydrogen gas. The metallic material has been tentatively identified as elemental iron.
Experiment #10 Summary:
An experiment was conducted using 5 g of Al(OH)₃plus 40 mL of colloidal silver and about 0.1 g sodium hydroxide in water. A 12 volt, 10 amp power source was then applied for a period of about thirty minutes to an iron anode and a carbon cathode that had been introduced into the solution. After about five minutes, the solution was titrated to a pH of about 7 using H₂SO₄. A metallic-like material that was produced and found attached to the anode was collected. The metallic material reacted with sulfuric acid to produce an ignitable gas presumed to be hydrogen gas. The metallic material has been tentatively identified as elemental aluminum. An X-Ray Photoelectric Spectrum (XPS) was taken of this material that indicates the presence of some elemental aluminum in this material.
Experiment #11 Summary:
An experiment was conducted using an initial solution comprising 10 mL of 93% concentration H₂SO₄and 30 mL of 35% concentration HCl that was reacted with 25 g of Al₂(SO₄)₃.18H₂O plus 80 mL of colloidal lead. Over a period of thirty minutes, the reaction mixture was heated on a hot plate, and the temperature increased to a value of 75° C. Over this period, a small amount of an ignitable gas presumed to be hydrogen gas was produced. Since, it is believed, the production of hydrogen gas could only be produced with a concurrent oxidation of aluminum, it is evident that the aluminum had to be initially reduced before it could be oxidized, thereby providing strong evidence of the reduction reaction.
Experiment #12 Summary:
An experiment was conducted using an initial solution comprising 10 mL of 93% concentration H₂SO₄and 30 mL of 35% concentration HCl that was reacted with 25 g of Al₂(SO₄)₃.18H₂O plus 80 mL of colloidal lead, and 80 mL of 3% aqueous H₂O₂. Over a period of thirty minutes, the reaction mixture was heated on a hot plate, and the temperature increased to a value of 75° C. Over this period, an impressive amount of an ignitable gas presumed to be hydrogen gas was produced. The rate of gas formation at this point was measured to be 80 mL per minute, or 4.8 L per hour. Since, it is believed, the production of hydrogen gas could only be produced with a concurrent oxidation of aluminum, it is evident that the aluminum had to be initially reduced before it could be oxidized, thereby providing strong evidence of the reduction reaction.
Experiment #13 Summary:
An experiment was conducted using an initial solution comprising 80 mL of 90% aqueous formic acid, CH₂O₂, that was reacted with 25 g of Al₂(SO₄)₃.18H₂O plus 80 mL of colloidal lead. Over a period of thirty minutes, the reaction mixture was heated on a hot plate, and the temperature increased to a value of 75° C. Over this period, a small amount of an ignitable gas presumed to be hydrogen gas was produced. There was then added to the reaction mixture 80 g of metallic iron. The rate of gas formation was found to increase drastically to a measured rate of 80 mL per minute or 4.8 L per hour. After an additional hour of gas formation at a temperature of 75° C., the reaction mixture was allowed to cool to a temperature of 20° C. The iron metal was examined and was found to have metallic aluminum plated over the iron. The aluminum was identified by its reaction with water in an aqueous solution of NaOH to produce an ignitable gas presumed to be hydrogen.
Experiment #14 Summary:
An experiment was conducted using an initial solution comprising 20 mL of 90% aqueous formic acid, CH₂O₂, that was reacted with 20 mL of colloidal lead and one gram of metallic aluminum. Over a period of thirty minutes, the reaction mixture was heated on a hot plate, and the temperature increased to a value of 75° C. No significant amounts of gases were emitted and, in fact, no noticeable chemical reaction of any kind was observed to occur under these conditions. Since metallic aluminum does not react significantly under these reaction conditions, it would be likely that if elemental aluminum were to be produced under similar conditions, the aluminum could be isolated.
Experiment #15 Summary:
An experiment was conducted using an initial solution comprising 40 mL of 90% aqueous formic acid, CH₂O₂, that was reacted with 40 mL of colloidal lead and 40 g of metallic iron. Over a period of thirty minutes, the reaction mixture was heated on a hot plate, and the temperature increased to a value of 75° C. Over this period, a moderate amount of gas was produced. There was then added to the reaction mixture log of Al₂(SO₄)₃.18H₂O, and the rate of gas formation was found to increase significantly. The reaction mixture was allowed to cool to a temperature of 20° C. The iron metal was examined and found to have metallic aluminum plated over the iron. The aluminum was identified by its reaction with water in an aqueous solution of NaOH to produce an ignitable gas presumed to be hydrogen.
Experiment #16 Summary:
An experiment was conducted using an initial solution comprising 80 mL of colloidal lead that was reacted with 30 g of Al₂(SO₄)₃.18H₂O plus 1 g of metallic iron. The reaction mixture was heated using a hot plate to a temperature of 75° C., and this temperature was maintained over a period of 24 hours. The reaction mixture was allowed to cool to a temperature of 20° C. The iron metal was examined and was found to have metallic aluminum plated over the iron. The aluminum was identified by its reaction with water in an aqueous solution of NaOH to produce an ignitable gas presumed to be hydrogen. A different iron nugget from the same reaction mixture was then studied using Energy Dispersive X-Ray Spectroscopy (EDX), and was found to contain approximately 9% elemental aluminum by mass.
Experiment #17 Summary:
An experiment was conducted using an initial solution comprising 40 mL of 90% aqueous formic acid, CH₂O₂, that was reacted with 40 mL colloidal lead plus 25 g of Al₂(SO₄)₃˜18H₂O plus 9 g of metallic iron. The reaction mixture was maintained at an ambient temperature of approximately 20° C. over a period of 48 hours. The iron metal was examined and was found to have metallic aluminum plated over the iron. The aluminum was identified by its reaction with water in an aqueous solution of NaOH to produce an ignitable gas presumed to be hydrogen.
It is believed the experimental results described above demonstrate the potential value of the inventions described herein. However, the results, calculations and conclusions are based on the theoretical reaction mechanisms that are described above and that are believed to accurately characterize the reactions involved in these experiments. However, if it is discovered that the theoretical reaction mechanisms used to rationalize the experimental findings, or the calculations based thereon are in error, the inventions described herein nevertheless are valid and valuable.
The embodiments shown and described above are exemplary. Many details are often found in the art and, therefore, many such details are neither shown nor described. It is not claimed that all of the details, parts, elements, or steps described and shown were invented herein. Even though numerous characteristics and advantages of the present invention have been described in the drawings and accompanying text, the description is illustrative only, and changes may be made in the detail, especially in matters of shape, size, and arrangement of the parts within the principles of the invention to the full extent indicated by the broad meaning of the terms of the attached claims.
The restrictive description and drawings of the specific examples above do not point out what an infringement of this patent would be, but are to provide at least one explanation of how to use and make the invention. The limits of the invention and the bounds of the patent protection are measured by and defined in the following claims.

Claims

1. An apparatus for the production of an elemental metal from a metal-containing compound, comprising:

a reaction medium containing ions of a first metal; and

a second metal, wherein the second metal is in colloidal form.

2. The apparatus of claim 1 further comprising a reaction vessel for containing the reaction medium.

3. The apparatus of claim 2, wherein the reaction vessel is configured to maintain an internal pressure above atmospheric.

4. The apparatus of claim 1, wherein the reaction medium comprises formic acid.

5. The apparatus of claim 4, wherein the second metal is less reactive than the first metal.

6. The apparatus of claim 4, wherein the second metal is more reactive than the first metal.

7. The apparatus of claim 4, further comprising a third metal, wherein the third metal is in colloidal form.

8. The apparatus of claim 7, wherein the third metal is more reactive than the first metal.

9. The apparatus of claim 7, wherein the third metal is less reactive than the first metal.

10. The apparatus of claim 4, wherein the second metal is silver, gold, platinum, tin, lead, copper, zinc, iron, aluminum, magnesium, beryllium, nickel or cadmium.

11. The apparatus of claim 4, further comprising a solid comprising the first metal in contact with the solution.

12. The apparatus of claim 11, wherein the solid comprising the first metal is a metal oxide.

13. The apparatus of claim 1, further comprising an elemental non-metal in contact with the solution.

14. The apparatus of claim 1, further comprising ions of a salt dissolved in the solution.

15. A method for the production of a metal from a metal-containing compound, comprising the steps of:

suspending a first colloidal metal in a reaction medium; and

providing a metal-containing compound in contact with the reaction medium.

16. The method of claim 15, wherein the reaction medium comprises formic acid.

17. An apparatus for the production of elemental metal from a metal-containing compound, comprising:

a reaction medium;

a first metal at least partially submerged in the reaction medium, the first metal having a surface area of at least 298,000,000 m²per cubic meter of first metal; and

a metal-containing compound in contact with the reaction medium.

18. The apparatus of claim 17 further comprising a reaction vessel for containing the reaction medium.

19. The apparatus of claim 18, wherein the reaction vessel is configured to maintain an internal pressure above atmospheric.

20. The apparatus of claim 17, wherein the reaction medium comprises formic acid.