Redox, which is shorthand for reduction-oxidation, is a type of chemical reaction classified by the net transfer of electrons. In this reaction, one molecule loses electrons, called oxidation, and the other molecule gains electrons, called reduction.
To help you differentiate between the two, remember the phrase 'OIL-RIG', which stands for 'oxidation is losing, reduction is gaining'. The molecule that is oxidized is called the reducing agent because it reduces the other reactant. Similarly, the molecule that is reduced is called the oxidizing agent because it oxidizes the other molecule.
Now that we have the terminology sorted out, let's look at an example of a redox reaction, the formation of the mineral magnesium oxide. During the reaction, each magnesium atom loses two electrons. Thus, magnesium is oxidized. Each oxygen atom gains two electrons; thus, oxygen is reduced.
However, not all reactions are redox reactions. For example, the reaction of calcium oxide with carbon dioxide to form calcium carbonate is not a redox reaction. So, how can we identify a redox reaction?
To do this, we track the oxidation number of each element as it goes from reactant to product. The oxidation number is the hypothetical charge that an atom would have if its bonds to different elements were ionic, meaning that the electrons are assigned to the more electronegative atom. The sum of the oxidation numbers in a molecule equals its overall charge.
Let's look back at magnesium oxide. It is a neutral compound, so the sum of the oxidation numbers for magnesium and oxygen equals zero. Magnesium can give two electrons, so its oxidation number is plus two. Oxygen can accept two electrons, so its oxidation number is minus two.
How about the reaction? Pure neutral elemental compounds have an oxidation number of zero. Thus, both magnesium and oxygen start with oxidation numbers of zero. Both magnesium and oxygen's oxidation numbers changed during the reaction, so this is a redox reaction.
Now, let's look at the calcium carbonate reaction we saw earlier. Both reactants are neutral, so the sum of the oxidation numbers for both compounds is zero. As we saw with magnesium oxide, calcium has an oxidation number of plus two and oxygen minus two. Then, the carbon in the carbon dioxide molecule has an oxidation number of plus four and each oxygen minus two.
How about the product? Calcium is plus two and carbon plus four, just like in the reactants. Each oxygen is minus two, totaling minus six, with the net oxidation number zero. Since none of the oxidation numbers have changed, this is not a redox reaction.
Now let's introduce the four types of redox reactions. The first is a single displacement reaction, where one atom displaces another. You'll see this in a thermite reaction, where one metal is reduced, and the other metal is oxidized.
The next type is a combustion reaction, which occurs between a fuel and oxidant to form oxidized products and heat. You see this in the lab during the combustion of methane with oxygen when using a Bunsen burner.
The third is a synthesis reaction, where two reactants combine to form one product, like in the synthesis of ammonia, where nitrogen is combined with hydrogen to form ammonia.
Finally, the fourth type is a decomposition reaction, where a reactant absorbs enough energy to break its bonds to form smaller compounds. This is what happens with fireworks, where potassium chlorate decomposes to potassium chloride and oxygen after heating.
In this lab, you'll perform and identify various types of redox reactions that transform solid copper to copper oxide, and then back again to solid copper.
Oxidation and Reduction
Some chemical reactions can be classified as reduction-oxidation reactions, or redox reactions. Oxidation is the process of matter, like an atom or ionic molecule, losing one or more electrons, and reduction is the process of the matter gaining one or more electrons.
Oxidation States
Each atom in a molecule has its own oxidation state or oxidation number. The oxidation state describes how oxidized a molecule is relative to its free elemental form. The oxidation state is expressed as the charge that an atom would have if each of its bonds to other elements were purely ionic. This means that the electrons in the bond are assigned to the more electronegative atom. The oxidation state of an atom in its free elemental form is defined as 0.
There are a few rules that are followed to determine oxidation state. Elements in Group I and Group II typically have oxidation states of +1 and +2, respectively. Hydrogen and oxygen typically have oxidation states of +1 and -2, respectively, and halogens usually have an oxidation state of -1. In addition, the oxidation states of the atoms in a molecule always add up to the charge on the molecule. Thus, the oxidation state of an atom not listed above can often be deduced. For example, consider carbon dioxide (CO2), which is a neutral molecule. If each of the two oxygen molecules contributes -2, carbon’s oxidation state must be +4 to cancel out the -4 from the oxygens.
For a more general approach, draw the Lewis structure of the molecule, identify the bonds between different atoms, and assign each bond to the more electronegative atom. Then, count the number of electrons on each atom, with each bond contributing two electrons. Subtract the number of electrons that are currently on the atom from the standard number of valence electrons for that atom to get the oxidation number.
Consider carbon dioxide again. Each oxygen has two lone pairs of electrons and is connected to the central carbon by a double bond. Oxygen is more electronegative than carbon, so each C=O bond, which accounts for four electrons, is assigned to its oxygen. Thus, each oxygen is assigned a total of eight electrons (four from the lone pairs and four from the double bond), and carbon is assigned none. The default number of valence electrons for oxygen is six, so the oxidation number for each oxygen is 6 – 8 = -2. The default number of valence electrons for carbon is four, so the oxidation number for carbon is 4 – 0 = +4.
Redox Reactions
Not all chemical reactions are classified as a redox reaction. A redox reaction is any reaction in which there is a change in an atom's oxidation state. Thus, to check whether a reaction is a redox reaction, determine the oxidation states of each atom in the reactants and products and look for any changes.
Many redox reactions involve a transfer of electrons directly from one molecule or atom to another. In those reactions, if a molecule, or atom, gains an electron, another molecule, or atom, must lose an electron. One simple way to remember the definitions of oxidation and reduction is through the phrase OIL-RIG, which stands for: Oxidation Is Losing – Reduction Is Gaining.
The molecule gaining an electron is being reduced, but it is called an oxidant or oxidizing agent because it is oxidizing the other molecule. Similarly, the molecule that loses an electron is being oxidized, but it is called a reductant or reducing agent because it reduces the other molecule.
There are four major reaction types that typically involve redox processes.
- Single Displacement Reaction: An atom displaces another atom that is part of a compound, and replaces it.
- Combustion Reaction: A compound is reduced by a strong oxidant, typically oxygen gas. Combustion reactions that occur between hydrocarbons and organic compounds typically produce carbon dioxide and water.
- Synthesis Reaction: Two reactants form a single product.
- Decomposition Reaction: A single reactant breaks into two or more products.
References
1. Harris, D. C. (2015). Quantitative Chemical Analysis. New York, NY: W. H. Freeman and Company.
Redox, which is shorthand for reduction-oxidation, is a type of chemical reaction classified by the net transfer of electrons. In this reaction, one molecule loses electrons, called oxidation, and the other molecule gains electrons, called reduction.
To help you differentiate between the two, remember the phrase 'OIL-RIG', which stands for 'oxidation is losing, reduction is gaining'. The molecule that is oxidized is called the reducing agent because it reduces the other reactant. Similarly, the molecule that is reduced is called the oxidizing agent because it oxidizes the other molecule.
Now that we have the terminology sorted out, let's look at an example of a redox reaction, the formation of the mineral magnesium oxide. During the reaction, each magnesium atom loses two electrons. Thus, magnesium is oxidized. Each oxygen atom gains two electrons; thus, oxygen is reduced.
However, not all reactions are redox reactions. For example, the reaction of calcium oxide with carbon dioxide to form calcium carbonate is not a redox reaction. So, how can we identify a redox reaction?
To do this, we track the oxidation number of each element as it goes from reactant to product. The oxidation number is the hypothetical charge that an atom would have if its bonds to different elements were ionic, meaning that the electrons are assigned to the more electronegative atom. The sum of the oxidation numbers in a molecule equals its overall charge.
Let's look back at magnesium oxide. It is a neutral compound, so the sum of the oxidation numbers for magnesium and oxygen equals zero. Magnesium can give two electrons, so its oxidation number is plus two. Oxygen can accept two electrons, so its oxidation number is minus two.
How about the reaction? Pure neutral elemental compounds have an oxidation number of zero. Thus, both magnesium and oxygen start with oxidation numbers of zero. Both magnesium and oxygen's oxidation numbers changed during the reaction, so this is a redox reaction.
Now, let's look at the calcium carbonate reaction we saw earlier. Both reactants are neutral, so the sum of the oxidation numbers for both compounds is zero. As we saw with magnesium oxide, calcium has an oxidation number of plus two and oxygen minus two. Then, the carbon in the carbon dioxide molecule has an oxidation number of plus four and each oxygen minus two.
How about the product? Calcium is plus two and carbon plus four, just like in the reactants. Each oxygen is minus two, totaling minus six, with the net oxidation number zero. Since none of the oxidation numbers have changed, this is not a redox reaction.
Now let's introduce the four types of redox reactions. The first is a single displacement reaction, where one atom displaces another. You'll see this in a thermite reaction, where one metal is reduced, and the other metal is oxidized.
The next type is a combustion reaction, which occurs between a fuel and oxidant to form oxidized products and heat. You see this in the lab during the combustion of methane with oxygen when using a Bunsen burner.
The third is a synthesis reaction, where two reactants combine to form one product, like in the synthesis of ammonia, where nitrogen is combined with hydrogen to form ammonia.
Finally, the fourth type is a decomposition reaction, where a reactant absorbs enough energy to break its bonds to form smaller compounds. This is what happens with fireworks, where potassium chlorate decomposes to potassium chloride and oxygen after heating.
In this lab, you'll perform and identify various types of redox reactions that transform solid copper to copper oxide, and then back again to solid copper.
Redox, which is shorthand for reduction-oxidation, is a type of chemical reaction classified by the net transfer of electrons. In this reaction, one molecule loses electrons, called oxidation, and the other molecule gains electrons, called reduction.
To help you differentiate between the two, remember the phrase 'OIL-RIG', which stands for 'oxidation is losing, reduction is gaining'. The molecule that is oxidized is called the reducing agent because it reduces the other reactant. Similarly, the molecule that is reduced is called the oxidizing agent because it oxidizes the other molecule.
Now that we have the terminology sorted out, let's look at an example of a redox reaction, the formation of the mineral magnesium oxide. During the reaction, each magnesium atom loses two electrons. Thus, magnesium is oxidized. Each oxygen atom gains two electrons; thus, oxygen is reduced.
However, not all reactions are redox reactions. For example, the reaction of calcium oxide with carbon dioxide to form calcium carbonate is not a redox reaction. So, how can we identify a redox reaction?
To do this, we track the oxidation number of each element as it goes from reactant to product. The oxidation number is the hypothetical charge that an atom would have if its bonds to different elements were ionic, meaning that the electrons are assigned to the more electronegative atom. The sum of the oxidation numbers in a molecule equals its overall charge.
Let's look back at magnesium oxide. It is a neutral compound, so the sum of the oxidation numbers for magnesium and oxygen equals zero. Magnesium can give two electrons, so its oxidation number is plus two. Oxygen can accept two electrons, so its oxidation number is minus two.
How about the reaction? Pure neutral elemental compounds have an oxidation number of zero. Thus, both magnesium and oxygen start with oxidation numbers of zero. Both magnesium and oxygen's oxidation numbers changed during the reaction, so this is a redox reaction.
Now, let's look at the calcium carbonate reaction we saw earlier. Both reactants are neutral, so the sum of the oxidation numbers for both compounds is zero. As we saw with magnesium oxide, calcium has an oxidation number of plus two and oxygen minus two. Then, the carbon in the carbon dioxide molecule has an oxidation number of plus four and each oxygen minus two.
How about the product? Calcium is plus two and carbon plus four, just like in the reactants. Each oxygen is minus two, totaling minus six, with the net oxidation number zero. Since none of the oxidation numbers have changed, this is not a redox reaction.
Now let's introduce the four types of redox reactions. The first is a single displacement reaction, where one atom displaces another. You'll see this in a thermite reaction, where one metal is reduced, and the other metal is oxidized.
The next type is a combustion reaction, which occurs between a fuel and oxidant to form oxidized products and heat. You see this in the lab during the combustion of methane with oxygen when using a Bunsen burner.
The third is a synthesis reaction, where two reactants combine to form one product, like in the synthesis of ammonia, where nitrogen is combined with hydrogen to form ammonia.
Finally, the fourth type is a decomposition reaction, where a reactant absorbs enough energy to break its bonds to form smaller compounds. This is what happens with fireworks, where potassium chlorate decomposes to potassium chloride and oxygen after heating.
In this lab, you'll perform and identify various types of redox reactions that transform solid copper to copper oxide, and then back again to solid copper.
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