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CHEM 116 Lecture Notes - Lecture 20: Standard Electrode Potential, Standard Hydrogen Electrode, Electrochemical Cell

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beigewoodchuck326

4 Aug 2016

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Electrons flow away from negative charge and towards positive charge. Oxidation corresponds to an increase in oxidation state. Reduction corresponds to a decrease in oxidation state. Procedure for half-reaction method of balancing aqueous redox equations in acidic. Balance all elements other than h and o. 20. 3: voltaic (or galvanic) cells: generating electricity from spontaneous chemical reactions. Electrical current: the flow of electric charge: redox reactions have the potential to generate electrical current. Electrochemical cell: where the generation of electricity through redox reactions occur. Voltaic (or galvanic) cell: electrochemical cell that produces electrical currents from a spontaneous chemical reaction. Electrolytic cell: another type of electrochemical cell, consumes electrical current to drive a nonspontaneous chemical reaction. Electrodes: conductive surfaces through which electrons can enter or leave the half-cells. Potential difference: is a measure of the difference in potential energy (j) per unit of charge (coulombs)

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Related Questions

Having trouble with my Pre-Lab questions I don't understand how to get them.

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)

I need help with questions 2-4 I can't figure them out.

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)

What is the oxidation state assigned to Mn in potassium permanganate, KMnO₄? (Recall that answers must be numbers, that is, F is -1 not 1- in HF.)

What is the oxidation state assigned to O in hydorgen peroxide, H₂O₂?

What is the oxidation state assigned to B in B(OH)₃?

What is the oxidation state assigned to N in the nitrite ion, NO₂^-?

What is the oxidation state assigned to N in NH₂^-?

What is the oxidation state assigned to Cr in zinc dichromate, ZnCr₂O₇?

What is the oxidation state assigned to Xe in XeO₂F₂?

What is the oxidation state assigned to Xe in XeO₃F₂?

Easy: Complete and balance the following redox reaction in acid solution:

Cl₂(g) + Br^-(aq) ? Cl^-(aq) + Br₂(l)

For the balanced oxidation half-reaction, how many electrons are produced? How many bromide ions are there? (Always use smallest integer coefficients.)

Choose at least one answer.

	a. 1 e^-
	b. 2 e^-
	c. 1 Br^-
	d. 2 Br^-

How many chlorine molecules are in the balanced equation?

Medium: Complete and balance the following redox reaction in acid solution:

Cu(s) + NO₃^-(aq) ? NO(g) + Cu²⁺(aq)

For the balanced reduction half-reaction, how many electrons are needed? How many H⁺ ions are there?

Choose at least one answer.

	a. 1 e^-
	b. 2 e^-
	c. 3 e^-
	d. 4 e^-
	e. 1 H⁺
	f. 2 H⁺
	g. 3 H⁺
	h. 4 H⁺

How many H⁺ ions are in the balanced equation?

(If the H⁺ ions are the left side (reactant), write it as a positive number. If they are the right side (product), write it as a negative number.)

Complete and balance the following redox reaction in acid solution:

MnO₄^-(aq) + C₂H₂O₄(aq) ? Mn²⁺(aq) + CO₂(g)

For the balanced oxidation half-reaction, how many electrons are produced? How many H⁺ ions are there?

Choose at least one answer.

	a. 1 e^-
	b. 2 e^-
	c. 3 e^-
	d. 4 e^-
	e. 1 H⁺
	f. 2 H⁺
	g. 3 H⁺
	h. 4 H⁺

How many H⁺ ions are in the balanced equation?

(If the H⁺ ions are the left side (reactant), write it as a positive number. If they are the right side (product), write it as a negative number.)

Complete and balance the following redox reaction in basic solution:

ClO^-(aq) + CrO₂^-(aq) ? Cl^-(aq) + CrO₄^2-(aq)

For the balanced oxidation half-reaction, how many electrons are produced? How many OH^- ions are there?

Hint given in the general feedback.

Choose at least one answer.

	a. 1 e^-
	b. 2 e^-
	c. 3 e^-
	d. 4 e^-
	e. 1 OH^-
	f. 2 OH^-
	g. 3 OH^-
	h. 4 OH^-

How many OH^- ions are in the balanced equation?

(If the OH^- ions are the left side (reactant), write it as a positive number. If they are the right side (product), write it as a negative number.)