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Content objective:

What are we learning and why are we learning this? Content, procedures, or skills.

Vocabulary objective

Tier II: High frequency words used across content areas. Key to understanding directions & relationships, and for making inferences.

Tier III: Low frequency, domain specific terms.

Building on what we already know

Make connections to prior knowledge. This is where we build from.


Electrochemistry is the study of chemical reactions that involve electrons moving from one substance to another. This movement of electrons is called electricity.

Myths and Facts

Myth: Redox reactions are different from other chemical reactions, like combustion, decomposition, replacement, etc.

Fact: Single-replacement rxns and combustion rxns are always redox. Other types are sometimes redox.

Reduction, Oxidation, and Oxidation numbers

Electrochemical reactions involve the transfer of electrons. Mass and charge are conserved when balancing these reactions, but you need to know which atoms are oxidized, and which atoms are reduced during the reaction. Oxidation numbers are used to keep track of how many electrons are lost or gained by each atom. {Anne Marie Helmenstine}

The oxidation state (or oxidation number) is an indicator of the degree of oxidation (loss of electrons) of an atom in a chemical compound.

An oxidation state may be negative, zero or positive.

Conceptually, the oxidation state is the hypothetical charge that an atom would have – if all bonds to atoms of different elements were 100% ionic, with no covalent component. This is never exactly true for real bonds. {Wikipedia}

Rules for assigning oxidation numbers

As a general rule, shared electrons are assumed to belong to the more-electronegative atom in each bond.

More specific rules are provided by the following guidelines.

1. The atoms in a pure element have an oxidation number of zero. For example, the atoms in pure sodium, Na, oxygen, O2, phosphorus, P4, and sulfur, S8, all have oxidation numbers of zero.

2. The more-electronegative element in a binary molecular compound is assigned the number equal to the negative charge it would have as an anion. The less-electronegative atom is assigned the number equal to the positive charge it would have as a cation.

3. Fluorine has an oxidation number of −1 in all of its compounds because it is the most electronegative element.

4. Oxygen has an oxidation number of −2 in almost all compounds. Exceptions include when it is in peroxides, such as H2O2, in which its oxidation number is −1, and when it is in compounds with halogens, such as OF2, in which its oxidation number is +2.

5. Hydrogen has an oxidation number of +1 in all compounds containing elements that are more-electronegative than it; it has an oxidation number of −1 in compounds with metals.

6. The algebraic sum of the oxidation numbers of all atoms in a neutral compound is equal to zero.

7. The algebraic sum of the oxidation numbers of all atoms in a polyatomic ion is equal to the charge of the ion.

8. Although rules 1 through 7 apply to covalently bonded atoms, oxidation numbers can also be assigned to atoms in ionic compounds.

A monatomic ion has an oxidation number equal to the charge of the ion. For example, the ions Na+,Ca2+, and Cl− have oxidation numbers of +1, +2, and −1, respectively.

Let’s examine the assignment of oxidation numbers to the atoms in two molecular compounds, hydrogen fluoride, HF, and water, H2O.

HF – the bond is polar, with a partial negative charge on the fluorine atom and a partial positive charge on the hydrogen atom.

If HF were an ionic compound in which an electron was fully transferred to the fluorine atom, H would have a 1+ charge and F would have a 1− charge.

Thus, the oxidation numbers of H and F in hydrogen fluoride are +1 and −1, respectively.

H2O – The O atom is more electronegative than the H atoms.

If H2O were an ionic compound, then the oxygen atom would have a charge of 2− and the H atoms would each have a charge of 1+.

The oxidation numbers of H and O in water are therefore +1 and −2, respectively.

Because the sum of the oxidation numbers [of the atoms in a compound] must satisfy rule 6 or 7 (above) it is often possible to assign oxidation numbers even when they are not known.

Modern Chemistry, Holt, Rinehart and Winston.

Modern Chemistry, Holt, Rinehart and Winston.


Modern Chemistry, Holt, Rinehart and Winston.

Modern Chemistry, Holt, Rinehart and Winston.

Also see:  http://www.chemguide.co.uk/inorganic/redox/oxidnstates.html

Redox reactions


Balancing redox reactions


Applications of redox

* Bleach and other oxidizing cleaners

* Antioxidants in our foods – they react with free radicals, and reduce their effects

* Some redox rxns produce light: chemical luminescence

** Luminol used to detect blood

** Cyalume is used in toy light sticks

** Luciferins are a class of small-molecule substrates that are oxidized in the presence of the enzyme luciferase to produce oxyluciferin and energy in the form of light.

* Rocket fuel combusts in a redox reaction

** 2H2 (l) + O2 (l) -> 2H2O (gas) + energy

** Solid rocket boosters on space shuttle
10Al(s) + 6NH4ClO4(s) 4Al2O3(s) + 2AlCl3(s) + 12H2O(l) + 3N2(g)

Corrosion of iron

Voltaic cells and batteries

Alessandro Volta Presenting His Great Invention to Napoleon Bonaparte Nov 1800

Industrial production of substances – obtaining alkali metals

Ore refining – liberating aluminum metal from aluminum ore
in an electrolytic cell running at around 1,000 oC.

2 AL2O3 + 3C -> 4 AL + 3 CO2


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