Chemistry PPT 9: Redox Reactions and Electrochemistry

Vocabulary terms

Oxidation

Reactions in which the oxidation number (or charge) of the atoms or ions increases or becomes less negative

Oxidized

Any substance that undergoes an increase in oxidation number

Reducing agent

Any particle or substance that can cause the gain of electrons: the substance being oxidized is the reducing agent (giving up electrons; electron donor)

Reduction

A reaction in which the oxidation number (or charge) of atoms or ions decreases or becomes more negative

Reduced

A substance that undergoes a decrease in oxidation number

Oxidizing agent

Any particle or substance that can cause the loss of electrons: the substance being reduced is the oxidizing agent (gaining electrons, electron acceptor)

If oxidation occurs during a chemical reaction then

reduction must occur simultaneously

The amount of oxidation must equal

the amount of reduction

The total increase in oxidation numbers of the substance being oxidized must equal

the total decrease in oxidation numbers of the substance being reduced

Reactants and products in redox reactions may include:

monatomic ions, elements, polyatomic ions, and molecular compounds containing elements having more than one oxidation state

Oxidation half reactions:

• change from a free element to a positive ion

• Change from an ion with a lower positive charge to an ion with a higher positive charge

• Change from a negatively charged ion to a free element

• Change in oxidation state involving a polyatomic ion

A loss of electrons always results in

a gain in oxidation number

Reduction

A reaction in which a particle gains electrons (resulting in a decrease in the oxidation number)

Reduction half reactions

• change from a free element to a negative ion

• Change from an ion with a higher positive charge to an ion with a lower positive charge

• Change from a positive ion to a free element

• Change in oxidation state involving an ion made up of two or more elements

A gain of electrons always results in

a decrease in oxidation state

The principle use of oxidation numbers is

balancing redox reactions

Conservation of mass and energy

In balancing redox reactions, both mass (# of atoms) and energy (charge) must be conserved (same on both sides of the equation)

Steps of Balancing Redox Reactions

• Write the oxidation and reduction half reactions and balance for mass and charge

• Equalize electrons and add half reactions

• Transfer coefficients into main equation

Electrochemistry

The study of the interchange of chemical and electrical energy

Redox reactions, like all chemical reactions involve

energy changes

These reactions also involve the transfer of electrons (can occur in the form of electrical energy rather than heat)

Michael Faraday

Discovered that the water solutions of certain substances called electrolytes conduct an electrical current

Electrolyte

Any substance which dissolves in water to form a solution that will conduct an electric current (ions, ionic substances)

Electrolytes may be classified as

strong or weak

Non-electrolytes

Solutions that do not conduct electricity

Strong electrolyte

Solutions in which the dissolved substance (solute) is present entirely as ions

Weak electrolyte

A solute that yields a relatively low concentration of ions in solution

Dissociation

The separation of ions that occurs when an ionic substance dissolves

Ionization

Process by which a covalent/molecular substance forms ions in a solution. Water molecules must pull the ions apart due to the stronger covalent bonds

Types of electrochemical cells

• Voltaic (Galvanic) cell: spontaneous (generates electricity by itself)

• Electrolytic Cell: non-spontaneous (need outside power source)

Electrochemical cells

A system of electrodes and electrolytes in which a spontaneous or non-spontaneous redox reaction occurs

Spontaneous redox reaction

A redox reaction which will occur on its own as written (cell potential is positive). Voltaic or Galvanic cells

Non-spontaneous redox reaction

A redox reaction which will not occur without an external source of power (cell potential is negative) Electrolytic cells

Electrode

An electrical conductor (metal strip) used to establish contact with a non-metallic part of the circuit (usually an electrolyte)

Electrolyte

a liquid, paste, or gel that serves to conduct charge by moving ions in the cell

Anode

the electrode at which oxidation occurs

Cathode

the electrode at which reduction occurs

Half-Cell

a single electrode immersed in a solution of its ions

Salt bridge

A device (porous disk or bride/U-tube containing inert electrolytic solution) placed between the cells which maintains electrical neutrality by allowing ions to migrate between the cells

External circuit

The part of the cell where charge is conducted as a current of moving electrons

Standard Electrode Reduction Potential

The measurement, in volts, of the tendency for a half reaction to occur as a reduction half reaction

Cell potential

The measure of the driving force, in volts, for an electrochemical reaction

Voltaic (Galvanic) cell

Redox reactions are spontaneous and chemical energy is converted into electrical energy. The cell potential is positive and the anode is the negative electrode

Formula

E-cell = E-cathode - E-anode

Electrolytic Cell

Cell in which an external electric current is required to drive a non-spontaneous redox reaction. The cell potential is negative and the anode is the positive electrode

Calculating Overall Cell Potential

• Assign oxidation numbers and determine what is being oxidized and reduced

• Write out the half reactions

• From the table of standard reduction potentials, write down the cell potential for each half reaction changing the sign of the oxidation half reaction

• Add the half reactions together to obtain the overall cell potential

• Note: coefficients do not affect the magnitude of the cell potentials

EXTRA SLIDES TO LOOK AT/ MEMORIZE

Slide 40 - starred slide: Quick Comparison of Electrochemical Cells

Slide 51 and 53 - identifying parts of the diagrams

REMEMBER THAT

OXIDATION IS ALWAYS IN ANODES AND
REDUCTION IS ALWAYS IN CATHODES

so that when you do the cell potential math (cathode minus anode) you know what to subtract from what based off the half-reactions