7.06 Oxidation and Reduction

You Can't Have One Without the Other

Oxidation and reduction go hand in hand. One substance cannot be oxidized unless another substance is reduced. If one element gains electrons (reduction), it is because another element lost them (oxidation). An element will not lose electrons (oxidation) unless another element takes them (reduction).

Oxidation

Oxidation is defined as the LOSS of electrons in a chemical reaction. In the reaction of sodium and chlorine, neutral sodium was oxidized when it lost an electron to chlorine to make a positive Na+ ion. This term got its name because many metals lose electrons when they react with oxygen to make an ionic metal oxide compound. Gaining oxygen (after a loss of electrons) is where this term got its name, "oxidized."

Oxygen

Oxygen always has an oxidation number of -2 when combined with another element, except in peroxides, when it is -1. Examples: -H2O (O = -2) -CaO (O = -2) -H2O2 (O = -1 because peroxide)

MNO4-

Oxygen has an oxidation number of -2 in all compounds, unless its a peroxide polyatomic ion. Determine the oxidation number of the manganese atom if the total sum must add up to the charge of the polyatomic ion, -1. 4(-2) + Mn = -1 Mn = +7

Real-World Examples

(1) Photosynthesis: This reaction stores energy from the sun in plants by converting carbon dioxide and water into sugar. This is a very important redox reaction in our food chain! (2) Human Body: In humans, sugars, fats, and proteins are oxidized in redox reactions that provide the energy necessary for life. (3) Batteries: Batteries (electrochemical cells) produce electricity by separating oxidation-reduction reactions into separate chambers that are connected by wires or other conductors. This controlled environment allows us to benefit from the exchange of electrons that occurs in a spontaneous redox reaction, by using the flow of electrons to power our laptops, MP3 players, and other battery-powered devices. (4) Antioxidants: In the human body, redox reactions are believed to play a role in the onset of heart disease and aging. Antioxidants are substances that may protect cells from the damage caused by losing electrons to unstable molecules known as free radicals. Antioxidants interact with and stabilize free radicals to prevent the damage that some oxidation-reduction reactions may cause inside your body. Examples of antioxidants in your food include beta-carotene, lycopene, and vitamins C, E, and A. (5) Rusting: This is a redox reaction in which oxygen oxidizes (takes electrons from) iron to form iron (III) oxide, otherwise known as rust. 4Fe (s) + 3O2 (g) → 2Fe2O3 (s) Iron is a very strong metal, whereas rust is a brittle ionic compound that flakes off of old cars and nails. What a difference a few electrons and bonds can make! (6) Combustion reactions: These reactions that provide most of the energy to power our society involve redox. The burning of methane gas is an example: CH4 + 2O2 → CO2 + 2H2O In this reaction, electrons are transferred from carbon to oxygen. However, you probably noticed that there are no ions or ionic compounds in this reaction. So, how do chemists determine if electrons have been exchanged in a reaction that does not involve ions? To answer this question, we need to review what we know about electronegativity and bond polarity.

Electronegativity and Oxidation Numbers

According to the concept of electronegativity, every atom has a certain amount of attraction for the electrons shared in a covalent bond. Some atoms have a stronger attraction for these shared electrons than others. Picture a covalent bond as a sort of tug-of-war where both atoms are competing for the shared electrons in the bond, much like two children fighting over the same toy. Just as the stronger child can pull the toy closer to herself than a weaker child in a tug-of-war, the atom with the higher electronegativity has a stronger "pull" on the electrons than the atom with the lower electronegativity. When two nonmetals compete for the electrons that are "shared" in a covalent bond, the electrons spend more of their time closer to the more electronegative element. This uneven sharing of electrons in a covalent bond means that the more electronegative element becomes more "negative" while the less electronegative element becomes more "positive." Even though the atoms involved in a covalent molecule do not have charges like an ion, we assign them positive and negative oxidation numbers to represent this "uneven sharing" that occurs. The following rules for assigning oxidation numbers will be useful when you need to determine if a reaction involves oxidation and reduction.

Adding Them Up

All the oxidation numbers (multiplied by each of their subscripts) in a compound must add up to the total charge of that compound or ion. This means we can use subtraction to solve for any elements not listed in the previous oxidation number rules.

CASO4

Ca2+ ion = +2 || O = -2 Determine what the oxidation number of S must be to have the entire compound add up to zero. +2 + 4(-2) + S = 0 S = 8 - 2 S = +6

CH4 (g) + 2O2 (g) ----> CO2 (g) + 2H2O (g)

Combustion of Methane: In methane, hydrogen gets an oxidation number of +1 because that is its oxidation number in any covalent compound. Carbon's oxidation number in methane must be -4, to balance out the four hydrogen atoms (each -1) to give an overall sum of zero for the oxidation numbers of this neutral compound. The diatomic oxygen molecule is oxygen's standard state as a neutral element, so the oxidation number of each oxygen atom in this diatomic molecule is zero. In carbon dioxide, each oxygen atom has an oxidation number of -2, because that is oxygen's oxidation number when in a compound with another element. That gives a total negative oxidation number of negative four (2 × -2) from the oxygen atoms, which means that the oxidation number of the carbon atom must be +4. The oxygen atom in water must have an oxidation number of -2, because that is the oxidation number of oxygen when in a compound with another element. That makes the oxidation number of each hydrogen atom +1, which makes sense, because that is supposed to be the oxidation number of hydrogen when in a covalent compound with another element. Carbon's oxidation number changes from -4 in methane to +4 in carbon dioxide. This increase in oxidation number shows that carbon was oxidized by losing negatively charged electrons. Oxygen's oxidation number changes from 0 in the elemental form to -2 in both carbon dioxide and water. This decrease in oxidation number shows that oxygen was reduced by gaining negative electrons.

H2CO3

Each hydrogen atom has an oxidation number of +1, and each oxygen atom has an oxidation number of -2. Multiply these oxidation numbers by the subscripts on the appropriate elements, and then determine what the oxidation number of carbon must be for the total to add up to zero. 2(+1) + 3(-2) + C = 0 C = +4

Covalent Compounds

For covalent compounds, pretend the compound is ionic with the more electronegative element forming the negative ion (anion). For example: Fluorine is always -1 in a compound, oxygen is almost always -2, and hydrogen is +1 in covalent compounds. Examples: -CCl4 C = +4, Cl = -1 -NH3 N = -3, H = +1

Oxidizing and Reducing Agents

In a redox reaction, one reactant is always an oxidizing agent and another reactant is the reducing agent. Oxidizing Agent: A reactant that causes another substance to be oxidized (not being oxidized itself). The oxidizing agent is the reactant that is reduced. Reducing Agent: A reactant that causes another substance to be reduced (not being reduced itself). The reducing agent is the reactant that is oxidized. This may seem backward because the oxidizing agent is being reduced and the reducing agent is being oxidized, but remember that if you are an agent of something, you cause that "something" to happen to someone else.

Monatomic Ions

Monatomic ions have an oxidation number equal to their charge as an ion (when alone or when in a compound). Examples: -NaCl Na = +1, Cl = -1 -CaBr2 Ca = +2, Br = -1 -Ag+ Ag = +1 -Al3+ Al = +3

Neutral Elements

Neutral elements in their standard state (not in a compound) always have an oxidation number of zero. Examples (all have an oxidation number of zero): -O2 -Na -F2 -S8 -P4

PO4^3+

O is -2 in all compounds (unless it's a peroxide polyatomic ion). So, 4 x -2 = -8 total from the oxygen atoms. If the entire polyatomic ion must add up to its charge, -3, what is the oxidation number of the P atom? 4(-2) + P = -3 P= -3-(-8) P= +5

Memory Tool

OIL RIG Oxidation Is Loss (of electrons) Reduction Is Gain (of electrons) LEO the lion says GER Losing Electrons is Oxidation Gaining Electrons is Reduction

Reduction

Reduction is defined as the GAIN of electrons in a chemical reaction. In the reaction above, neutral chlorine got reduced when it gained electrons from the sodium to make a negative Cl- ion. When an atom or particle gains a negative electron, its overall charge is lowered, or "reduced." This is where this term comes from.