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Consider your experimental results from part A of this lab. Suppose your strongest reducing agent were added to your strongest oxidizing agent. (Use the lowest possible coefficients. Omit states-of-matter from your answers.)

(a) Write the half-reaction for your strongest reducing agent.

Mg â Mg²⁺ + 2e^-

(b) Write the half-reaction for your strongest oxidizing agent.

MnO₄^- + 8H⁺ + 5e^- â Mn²⁺ + 4H₂O

(c) Note the number of electrons in each half reaction.

In order to balance the number of electrons lost and gained, the oxidation half-reaction must be multiplied by ---Select--- one, two, three, four, five,

and the reduction half-reaction must be multiplied by ---Select--- one. two. three. four. five.

(d) Write the net redox reaction.

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Introduction An important type of reaction mechanism is the transfer of electrons from one species to another effecting a change in oxidation state. This gain or loss of electrons is referred to as reduction and oxidation, respectively. For example, consider the two balanced, chemical equations (1 & 2) that follow Cu (s) D Cu (aq) 2e (1) Cl2(aq) 2e D 2Cl (aq) (2) The first equation (1) shows elemental copper metal with a oxidation state of zero forming aqueous copper(Il) ions having an oxidation state of +2. This gain in oxidation number, 0 to +2, is called oxidation. The second equation (2) show elemental chlorine having an oxidation state of zero forming chloride ions in aqueous solution with an oxidation state of This decrease in oxidation number, 0 to -1, is called reduction. In any chemical reaction, the electrons gained or lost are conserved, i.e.. there is no net gain or loss of electrons. In fact, the two reactions illustrated above could occur simultaneously with the electrons lost by copper, gained by the chlorine in aqueous solution. Since the copper provides the electrons that subsequently reduce the chlorine copper is referred to as the reducing agent. Likewise, since the chlorine is accepting the electrons from copper, which subsequently oxidizes the copper, chlorine is referred to as he oxidizing agent. The two reactions above (1 & 2) are called half-reactions, as much as they each constitute one-half of the total reaction Equation 1 is the oxidation half reaction showing Cu(s) losing two electrons. In Equation 2, Cla(aq) gains two electrons and is the reduction half-reaction. If these two half-reaction are added together the result is the total oxidation-reduction reaction or "redox" reaction. Adding the two equations yields (3). Cu(s) Cl2(aq) D Cu (aq) 2CI (aq) (3) Whether or not the above reaction is spontaneous proceeds from reactants to products without input of energy into the system, may be determined by observing the reaction mixture if the reactants and products differ greatly in their appearance. Copper metal is a bright copper color and chlorine in aqueous solution will appear yellow. The Cu ion may turn the liquid part of the mixture blue, however, the Cris colorless and cannot be observed directly. Copper, like many of the transition metals, exhibits a variety of oxidation states depending upon the chemical environment in which it is found. In a strong reducing environment copper may exist as the free metal, Cu(s). An example of such a situation is illustrated in the following chemical equation (4). 4Cuo(s) CH4(g) D 4Cu(s) CO2(g) 2H2O(l) (4)

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Introduction An important type of reaction mechanism is the transfer of electrons from one species to another effecting a change in oxidation state. This gain or loss of electrons is referred to as reduction and oxidation, respectively. For example, consider the two balanced, chemical equations (1 & 2) that follow Cu (s) D Cu (aq) 2e (1) Cl2(aq) 2e D 2Cl (aq) (2) The first equation (1) shows elemental copper metal with a oxidation state of zero forming aqueous copper(Il) ions having an oxidation state of +2. This gain in oxidation number, 0 to +2, is called oxidation. The second equation (2) show elemental chlorine having an oxidation state of zero forming chloride ions in aqueous solution with an oxidation state of This decrease in oxidation number, 0 to -1, is called reduction. In any chemical reaction, the electrons gained or lost are conserved, i.e.. there is no net gain or loss of electrons. In fact, the two reactions illustrated above could occur simultaneously with the electrons lost by copper, gained by the chlorine in aqueous solution. Since the copper provides the electrons that subsequently reduce the chlorine copper is referred to as the reducing agent. Likewise, since the chlorine is accepting the electrons from copper, which subsequently oxidizes the copper, chlorine is referred to as he oxidizing agent. The two reactions above (1 & 2) are called half-reactions, as much as they each constitute one-half of the total reaction Equation 1 is the oxidation half reaction showing Cu(s) losing two electrons. In Equation 2, Cla(aq) gains two electrons and is the reduction half-reaction. If these two half-reaction are added together the result is the total oxidation-reduction reaction or "redox" reaction. Adding the two equations yields (3). Cu(s) Cl2(aq) D Cu (aq) 2CI (aq) (3) Whether or not the above reaction is spontaneous proceeds from reactants to products without input of energy into the system, may be determined by observing the reaction mixture if the reactants and products differ greatly in their appearance. Copper metal is a bright copper color and chlorine in aqueous solution will appear yellow. The Cu ion may turn the liquid part of the mixture blue, however, the Cris colorless and cannot be observed directly. Copper, like many of the transition metals, exhibits a variety of oxidation states depending upon the chemical environment in which it is found. In a strong reducing environment copper may exist as the free metal, Cu(s). An example of such a situation is illustrated in the following chemical equation (4). 4Cuo(s) CH4(g) D 4Cu(s) CO2(g) 2H2O(l) (4)

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11 Oxidation-Reduction Reactions Introduction We encounter oxi uction reactions every day Examples include the combustion of coal natural gas, oil, and gasoline, the operation of automobile's battery and the sains by a Although there are a variety of oxidation-reduction reactions .when you compare a number of you can some common featu Each these of simultaneous oxidation and reduction because electrons are lost by one atom and gained by another one. Moreover, the reactants include an agent and a reducing agent always lost agent and gained by the oxidizing agent (Ebbing/Gammon, Section 4.5) Purpose You will begin by preparing aqueous solutions of three halogens: Cl2, Bra, and Iz. You will prepare these substances in solution from stoichiometric quantities of certain oxidizing and agents. By observing the effect ofthese halogens on the corresponding halides (CI,Br, and you will be able the halogens according to their oxidizing strengths and the halides according to their reducing strengths. Your subsequent observations on the reactions of two halides, Br and I with two common oxidizing agents will enable you to rank the latter according to their oxidizing strengths. You will also find that the products obtained from an oxidation-reduction reaction in an acidic solution can differ markedly from those obtained in a basic solution. Finally, and most important, you will write a balanced equation for each reaction that you have observed The Halogens and the Halides You will deal with halogens and halides in each part ofthis experiment. A brief description of some of their properties will enable you to understand some of your observations more easily. The halogens are found in Group VIIA of the periodic table. The members of this group are fluorine (F) chlorine (CI, bromine (B), iodine (i) and astatine these elements belong to the same 12, group, have many similar properties. For example, their elemental diatomic (F2.cl. Br2, they Atal, and each of them is an oxidizing agent. This experiment uses only Cl, Bra, and Iz, however, because F is such a strong oxidizing agent that special conditions are required for its study and because astatine is radioactive. Aqueous solutions ofcl, be prepared either by dissolving these elements in water or by generating them in Br, reactions. Aqueous solutions of these halogens are solution as products of certain oxidation-reduction respectively. Each called chlorine water, brominee and iodine thickness solution. has its own characteristic color, depending on the concentration and red-brown to water is colorless to yellow; bromine water is yellow to red-brown; and iodine water is brown These halogens are also soluble in cyclohexane (C6H12, a substance that is itselfinsoluble in water. layer When cyclohexane is added to an aqueous solution of a halogen, two layers are formed. The upper is cyclohexane, because it is insoluble and less dense than water. A portion of the halogen passes from the aqueous layer to the upper cyclohexane layer (a process called extraction). A characteristic color is then imparted to this layer. The color ofa halogen in cyclohexane differs somewhat from its color in an aqueous solution, as you will discover during the experiment. be scanned, copied or duplicated or possed toa publicly accessible webwie, iawhole or in Cengage Leaming All Rights Resened c 2013

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