CHM 1046

CHM 1046
General Chemistry II
Dr. Michael Blaber

Aqueous Reactions

Introduction to Oxidation-Reduction Reactions

Many metals, when placed in an acid solution, will react chemically to produce hydrogen gas. For example, zinc metal in an aqueous solution of hydrochloric acid will react as follows:

Some troubled youths are actually aware of this reaction: they have been known to put swimming pool acid (hydrochloric acid, commonly called muriatic acid) into a wine bottle (preferably a large Gallo jug), dump in galvanized nails, quickly put a balloon over the top of the bottle, and watch the balloon fill with hydrogen (the balloon is then ignited with much rejoicing).

What is going on in this type of reaction?

The zinc starts out as a solid metal, i.e. the neutral (uncharged) elemental form of zinc. The zinc ends up as an ion (Zn²⁺)in an ionic compound involving chloride ion
The hydrogen starts out as a hydrogen ion (H⁺) from the acid, and ends up in the neutral (covalently bonded) form of diatomic hydrogen
Chloride ion (Cl^-) is a spectator ion in this reaction

In order for these ionic changes to occur, electrons leave the metal (producing zinc cations) and join the hydrogen ions (producing neutral diatomic hydrogen).

When an atom has become more positively charged (by LOSING electrons) chemists say that it has been OXIDIZED

In the above reaction, the zinc metal has become oxidized (producing Zn²⁺ cations)

When an atom becomes more negatively charged (by GAINING electrons) chemists say that is has been REDUCED

In the above reaction, the hydrogen ions have been reduced (to produce neutral diatomic hydrogen gas)

Oxidation of the metal refers to the fact that this type of reaction was actually first characterized by studying the reaction of metals with oxygen:

Many metals react with oxygen in the air to form metal oxides

In this reaction the neutral metal loses electrons to the oxygen, forming a cation (Ca²⁺ in this case). The metal cation participates in forming an ionic solid with the oxide (O^2-) anion. Thus, the metal becomes oxidized

When one substance in a reaction loses electrons, another substance in the reaction must gain them. The oxidation of a reactant is always accompanied by the reduction of another reactant in the reaction

These types of reactions are therefore called oxidation/reduction (or REDOX) reactions

In addition to oxidation by O₂, metals can be oxidized by both acids and salts

We saw an example of the oxidation of a metal (Zn) by an acid (HCl) above:

In this case, the metal is oxidized by aqueous protons to produce zinc cation, and the hydrogen is correspondingly reduced to neutral diatomic hydrogen gas

Metals can also be oxidized by aqueous solutions of various salts:

In this case solid iron can be oxidized by aqueous nickel ions to produce iron cations, and the nickel ions are correspondingly reduced to produce nickel metal

Again, whenever one substance is oxidized, another substance is reduced in a redox reaction

In the above example of the oxidation of iron by nickel:

Why doesn't the reaction go the other way around? Why doesn't iron oxidize nickel?
Can we predict which metal will be preferentially reduced and which will be oxidized when elemental or ionic forms of different metals are combined?

The different metals differ in the ease with which they can be oxidized

Some metals, e.g. Li and K, are easy to oxidize
Other metals, e.g. platinum (Pt) and gold (Au) are difficult to oxidize

In predicting oxidation reactions between different metals, compare the relative ease with which the two metals can be oxidized

Of the two metals, the one that is relatively more difficult to oxidize will preferentially exist in the neutral elemental form (i.e. be reduced)

The other metal will preferentially exist in the oxidized (cation) form (i.e. will give its electrons up to the other metal)

A ranking of the metals by the relative ease with which they can be oxidized, is known as an Activity Series. Hydrogen is also included in such lists so as to include the behavior of acids along with the metals:

At the top (the easiest to oxidize) are the alkali metals (group 1A) and the alkaline earth metals (group 2A)
At the bottom (the least likely to oxidize) are the transition metals, and in particular, the metals that are typically found in jewelery or coins are at the very bottom

The difficulty in oxidizing gold, for example, means that it prefers to exist in its uncharged, elemental form (i.e. as a metal and not as a cation or metal salt). Thus, those metals at the bottom of the Activity Series are generally unreactive and are therefore useful for use in metal items that we want to stay in the metal form (e.g. coins and jewelry). Gold, platinum, silver are often called the Noble Metals in reference to their lack of reactivity in redox reactions. Oxidation of metals during redox reactions represent what is known as corrosion. Our coins and jewelry should not corrode! Silver has a slight tendency to oxidize, or tarnish. We can restore its luster by placing it in a redox reaction where it will become reduced. Silver cleaning solutions are essentially redox reactions that promote the reduction of silver.

The Activity Series can be use to predict the outcome of reactions between metals and either metal salts or acids

Any metal in the list can be oxidized by the ions of any metal below it (or hydrogen ions if they fall below the metal in question). The corresponding ions of the other metal (or hydrogen) will become reduced to the elemental form

For example, silver is below copper in the Activity Series, therefore, copper metal will be oxidized by silver cations (resulting in the formation of copper cations and elemental silver)

Metals can react with acids to form diatomic hydrogen gas and metal cations (i.e. the oxidation of the metal and the reduction of the hydrogen ion). Which metals can participate in such reactions?

Metals that are higher than H⁺ ions on the Activity Series can be oxidized by the H⁺ ions (i.e. can react with acids to form hydrogen gas)
Metals below H⁺ ions on the Activity Series cannot be oxidized by H⁺, and therefore, will not react with acids to produce hydrogen gas (the metals below hydrogen represent the group of metals used in coins and jewelry)

Let's build a boat out of aluminum!

Aluminum is a very light and strong metal - the development of aluminum as a building material allowed the successful production of large aircraft. Use in boat hulls seemed to be an appropriate application of this metal. However, the ocean is an aqueous environment that contains a a variety of metal ions, including many that fall below aluminum in the Activity Series (e.g. manganese, iron, etc.) as well as inherent concentrations of hydrogen ions. Thus, early attempts at aluminum hulls resulted in disaster (oxidation of the elemental aluminum hulls to aluminum ions). The hulls corroded and dissolved and the ships sank. The solution? Magnesium lies above aluminum in the Activity Series, so in comparison to aluminum it will preferentially oxidize (and any aluminum ions will reduce to elemental aluminum). So, aluminum boats always have a small piece of magnesium connected to the aluminum hull. This magnesium piece will preferentially corrode (i.e. oxidize) before the aluminum hull. It is replaced before it completely oxidizes.