Chemistry Unit 3

Topics: Electrochemistry, Redox, Galvanic cell Pages: 11 (2349 words) Published: March 3, 2013
Chapter 9 – Electric Cells
9.1 – Oxidation and Reduction
* The term reduction came to be associated with producing metals from their compounds.

* Ex. Fe2O3(s) + 3CO(g) 2Fe(s) + 3CO2

* Another substance, called a reducing agent causes or promotes the reduction of a metal compound to an elemental metal. In this example, it is CO.

* Corrosion, including the rusting of metals, is now understood to be similar to combustion.

* Reactions of substances with oxygen, whether they are combustion, burning, or corrosion, came to be called oxidation.

* Oxidation is a process in which metals are converted to compounds by most non-metals.

* This does not always occur with oxygen, as other gases such as Cl or Br are also used.
* Ex. 2Mg(s) + O2(g) 2MgO(s) OR Cu(s) + Br2(g) CuBr2(s)

* A substance that causes the oxidation of metals to produce a metal compound is called an oxidizing agent. In this example, it is O2 and Br2.

Electron Transfer Theory
* A redox reaction is a combination of 2 parts, each called a half-reaction.

* In Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g)

* In the above reaction, the Zn(s) are converted to zinc ions in solution, Zn2+(aq).

* Therefore: Zn(s) Zn2+(aq) + 2e-

* Hydrogen ions in the solutions gain electrons and are converted to H2 gas. Thus becoming: 2H+(aq) + 2e- H2(g)

* According to modern theory, the gain of electrons is called reduction.

* According to modern theory, the loss of electrons is called oxidation.

* The total number of electrons gained in a reaction must equal the total # of electrons lost.

Oxidation States
* Metals and monatomic anions tend to lose electrons (become oxidized), whereas non-metals and monatomic cations tend to gain electrons (become reduced).

* An oxidation number is a positive or negative number corresponding to the apparent charge that an atom in a molecule or ion would have if the electron pairs in covalent bonds belonged entirely to the more electronegative atom.

* Oxidation numbers are written with sign first and number after. Ex. -1, +2

* Therefore, the sum of the oxidation numbers in a compound or ion must equal the total charge; zero for neutral compounds and ion charge for ions.

* Sample Problem: What is the oxidation number of carbon in methane, CH4?

CH4 = x + 4(+1) H has a charge of +1
X+4 = 0 Neutral molecule
x= -4

* Sample Problem: What is the oxidation number of sulfur in sodium sulfate?

Na2SO4 = 2(+1) + x + 4(-2) O has a charge of -2, Na has +1 2 +x -8 = 0 Neutral molecule
x= +6

* The sum of the oxidation # for a compound is 0.

* The sum of the oxidation # for a polyatomic ion equals the charge on the ion.

Oxidation Numbers and Redox Reactions
* If the oxidation number of an atom or ion changes during a chemical reaction, then an electron transfer (redox) has occurred.

* An increase in the oxidation number is defined as an oxidation

* A decrease in the oxidation number is a reduction.

* A reaction in which all oxidation numbers remain the same is not a redox reaction. 9.2 – Balancing Redox Equations
Oxidation Number Method
* One way of recognizing a redox reaction is to assign oxidation numbers to each atom/ion.

* Then, look for any changes in the oxidation numbers.

* In terms of oxidation numbers, this means that the changes in oxidation numbers must be balanced between the reactants and the products.

* The total increase in oxidation # for a particular atom/ion must equal the total decrease in oxidation # of another atom/ion.

* Sample Problem: Look Pg. 665 - 668

* Procedure for balancing redox equations using oxidation numbers:

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