Which Equation Could Be Used To Find Mã¢ë†â E In â–â³efg?

Electrolysis

Electrolysis involves passing an electrical electric current through either a molten salt or an ionic solution.  The ions are "forced" to undergo either oxidation (at the anode) or reduction (at the cathode).  Nearly electrolysis problems are really stoichiometry issues with the addition of an corporeality of electrical electric current.  The quantities of substances produced or consumed by the electrolysis procedure is dependent upon the following:

• electric current measured in amperes or amps
• time measured in seconds
• the number of electrons required to produce or consume i mole of the substance

• Amps, time, Coulombs, Faradays, and moles of electrons
• Computing the quantity of substance produced or consumed
• Calculating the time required
• Computing the current required

Amps, Time, Coulombs, Faradays, and Moles of Electrons

Three equations relate these quantities:

• amperes x time = Coulombs
• 96,485 coulombs = 1 Faraday
• 1 Faraday = 1 mole of electrons

The thought procedure for interconverting between amperes and moles of electrons is:

amps & time Coulombs Faradays moles of electrons

Utilise of these equations are illustrated in the following sections.

Top

Computing the Quantity of Substance Produced or Consumed

To determine the quantity of substance either produced or consumed during electrolysis given the time a known current flowed::

• Write the counterbalanced half-reactions involved.
• Summate the number of moles of electrons that were transferred.
• Summate the number of moles of substance that was produced/consumed at the electrode.
• Convert the moles of substance to desired units of measure.

Example:   A twoscore.0 amp current flowed through molten fe(Iii) chloride for 10.0 hours (36,000 s).  Determine the mass of iron and the volume of chlorine gas (measured at 25^oC and 1 atm) that is produced during this time.

• Write the half-reactions that take place at the anode and at the cathode.

anode (oxidation):  ii Cl^- Cl_two(g) + ii e^-

cathode (reduction)  Fe³⁺ + iii e^-  Fe(s)

• Calculate the number of moles of electrons.



• Calculate the moles of iron and of chlorine produced using the number of moles of electrons calculated and the stoichiometries from the balanced one-half-reactions.  According to the equations, three moles of electrons produce i mole of iron and two moles of electrons produce ane mole of chlorine gas.



• Calculate the mass of iron using the molar mass and calculate the book of chlorine gas using the ideal gas police (PV = nRT).

Top

Computing the Time Required

To determine the quantity of time required to produce a known quantity of a substance given the corporeality of electric current that flowed:

• Find the quantity of substance produced/consumed in moles.
• Write the counterbalanced half-reaction involved.
• Calculate the number of moles of electrons required.
• Convert the moles of electrons into coulombs.
• Calculate the fourth dimension required.

Example: How long must a 20.0 amp current flow through a solution of ZnSO_four in club to produce 25.00 g of Zn metal.

• Convert the mass of Zn produced into moles using the tooth mass of Zn.



• Write the one-half-reaction for the product of Zn at the cathode.

Zn²⁺(aq) + 2 east^- Zn(s)

• Calculate the moles of e^- required to produce the moles of Zn using the stoichiometry of the balanced half-reaction.   Co-ordinate to the equation 2 moles of electrons will produce one mole of zinc.



• Catechumen the moles of electrons into coulombs of charge using Faraday'southward constant.



• Calculate the time using the electric current and the coulombs of charge.

Acme

Calculating the Electric current Required

To determine the corporeality of current necessary to produce a known quantity of substance in a given amount of time:

• Find the quantity of substance produced/or consumed in moles.
• Write the equation for the one-half-reaction taking place.
• Calculate the number of moles of electrons required.
• Convert the moles of electrons into coulombs of charge.
• Calculate the electric current required.

Example: What current is required to produce 400.0 50 of hydrogen gas, measured at STP, from the electrolysis of water in 1 hour (3600 s)?

• Summate the number of moles of H₂. (Remember, at STP, 1 mole of any gas occupies 22.4 L.)



• Write the equation for the half-reaction that takes place.

Hydrogen is produced during the reduction of water at the cathode.  The equation for this half-reaction is:

4 due east^- + four H₂O(fifty) 2 H₂(k) + 4 OH^-(aq)

• Calculate the number of moles of electrons. According to the stoichiometry of the equation, four mole of east^- are required to produce two moles of hydrogen gas, or 2 moles of east^-'southward for every ane mole of hydrogen gas.



• Convert the moles of electrons into coulombs of charge.



• Calculate the current required.

Top

Source: https://www.chem.purdue.edu/gchelp/howtosolveit/Electrochem/Electrolysis.htm

Posted by: keenanmaked1947.blogspot.com

Share this post