Lesson Video: Properties and Reactions of Iron and Iron Oxides Chemistry

In this video, we will learn how to describe the properties and reactions of iron and its oxides.

18:00

Video Transcript

In this video, we will learn about the properties of iron and its oxides. And we will look at some of their reactions. Let’s start with iron.

Iron is a really important metal to the economy and is used in many applications. Iron is hardly ever used as a pure metal but is usually alloyed with other elements, for example, with carbon in steel. About five percent of the Earth’s crust is composed of this transition element. Iron is found in ores in the crust. And the important ores from which we extract iron metal include hematite, limonite, magnetite, which is a magnetic ore, and siderite.

Now that we know a bit about iron, let’s have a look at some of its properties. Pure iron is relatively soft and low in hardness. We tend to think of iron as strong and hard. But remember, it is almost always alloyed with other elements to improve its strength and other properties. Like other metals, iron is fairly malleable and can be hammered into flat sheets. Iron can be drawn into thin wires, and so we say it is ductile.

Iron is magnetic. It interacts with and is attracted to a magnetic field. Iron is a permanent magnet, meaning it has a magnetic field all on its own. It melts at 1538 degrees Celsius, which is much higher than the melting point of many other pure metals and much higher than the melting point of many alloys. How does its density compare? Iron’s density is 7.87 grams per cubic centimeter and is much more dense than aluminum, for example, whose density is 2.7 grams per cubic centimeter, but less dense than copper.

Now let’s discuss iron’s chemical properties and how it reacts. How an element reacts depends on its electronic structure or electronic configuration. This is iron’s electronic configuration. Iron loses electrons when it reacts with other substances. The common oxidation states it forms when it reacts are plus two and plus three. Here is the electronic configuration of Fe2+. The two 4s electrons have been lost. And here is the electronic configuration of Fe3+. The two 4s electrons and one 3d electron have been lost. Often, Fe2+ ions in solution make the solution pale green, whereas Fe3+ ions in solution often appear orange brown or red brown.

Now let’s clear some space to look at the chemical properties or reactions of iron in some more depth. When iron is heated till red hot, it reacts with oxygen. Black magnetic Fe3O4 is the product. This oxygen can be pure oxygen or from dry air. With gaseous water or water in its vapor form, red-hot iron also produces this magnetic iron oxide as well as hydrogen gas. With nonmetals, such as chlorine and sulfur, energy input is needed for these reactions. And the corresponding chloride or sulfide is produced.

With chlorine, the reaction is two Fe solid plus three Cl2 gas being heated to give two FeCl3 solid. Iron here has an oxidation state of plus three. The common name of this product is ferric chloride. In the case of reaction with sulfur, the equation is Fe solid plus S solid being heated to give FeS solid. The oxidation state of iron in the product is plus two. The common name of this product is ferrous sulfide. The suffixes -ic and -ous are often used with these oxidation states.

Iron can react with dilute mineral acids. The general equation is metal plus dilute acid gives salt and hydrogen gas. With dilute hydrochloric acid, the corresponding chloride salt is produced, with iron in the plus two state, which is ferrous chloride. With dilute sulfuric acid, again, the iron product is in the plus two oxidation state, and this is called ferrous sulfate, or the IUPAC name iron(II) sulfate.

However, when pure iron reacts with hot concentrated sulfuric acid, slightly different products are formed. Two different iron salts are produced, one with a plus two oxidation state and the other with a plus three oxidation state. The reaction equation is three Fe solid plus eight H2SO4 liquid being heated to give FeSO4 aqueous plus Fe2(SO4)3 aqueous plus four SO2 gas plus eight H2O gas. Now why there are two different iron products with hot concentrated sulfuric acid but only one iron product with dilute sulfuric acid is because of the oxidizing nature of sulfuric acid. Iron is oxidized to Fe2+ by sulfuric acid. But with hot concentrated sulfuric acid, the acid is so oxidizing that some of the Fe2+ can be further oxidized to Fe3+.

Let’s look at one last reaction of pure iron. When pure iron reacts with concentrated nitric acid, a layer forms on the surface of the iron. The surface layer is inactive or unreactive. The reaction produces a thin oxide layer, which coats the surface of the metal. This layer prevents the metal underneath from reacting further. This process of treating a metal surface, usually with an acid, to make it inactive or unreactive is called passivation. We say the metal surface was rendered or made passive. We won’t look at the specific chemical equation for this reaction. Passivation is beneficial to prevent or slow down rusting.

So far, we’ve learned about iron and its properties and looked at some of its reactions. Let’s now turn our attention to the oxides of iron. The three oxides of iron we will discuss are FeO, whose IUPAC name is iron(II) oxide and common name ferrous oxide; Fe2O3, whose IUPAC name is iron(III) oxide and common name ferric oxide; and Fe3O4, which is known as iron(II, III) oxide and sometimes called ferrosoferric oxide.

We know iron in ferrous oxide has an oxidation state of plus two. This is a black solid often found in powder form. It is insoluble in water, and the mineral ore in which it is found is called wüstite. Wüstite isn’t as economically important as other ores of iron. And this is because most of it is found in the mantle of the Earth and not the crust. The mantle of the Earth is much deeper and is therefore more difficult to access.

Ferric oxide is a red-brown solid. It is the main component of rust. It is also insoluble in water. It is widespread in nature in the ore hematite. Hematite is a reddish-black color. This ore is the main source of iron for the steel-making industry and so is a very important ore. Because of the lovely red-brown color of this iron oxide, it is often used in paint pigments. And we know that the oxidation state of iron in this compound is plus three.

Now the last oxide of iron that we will look at is interesting. In Fe3O4, the iron ions have different oxidation states: plus two and plus three. It is often considered to be an Fe2+ oxide and an Fe3+ oxide mixed together as a combined compound. Sometimes its formula is written as FeO.Fe2O3. But we just simplify it and say Fe3O4. This oxide is a black solid and is used as a black pigment. It is strongly magnetic and is found in nature in the ore magnetite, which happens to be the most common iron ore.

We now know a bit about the iron oxides as well as some of their physical properties. Now let’s have a look at their chemical reactions. We will also discuss how they are prepared. Let’s start with iron(II) oxide. How is it synthesized? FeO is prepared by the thermal decomposition of iron(II) oxalate in an inert atmosphere, in other words in the absence of air. This is the structure of iron(II) oxalate, and the byproducts are carbon monoxide and carbon dioxide gases.

FeO can also be prepared when other oxides of iron are reduced. Let’s have a look. In this reaction, hydrogen gas is used as the reducing agent to convert iron in its plus three oxidation state to a plus two oxidation state. Sometimes the reducing agent that is used is carbon monoxide. The reaction equation for the reduction of iron in its three plus state to a two plus state is Fe2O3 solid plus H2 gas being heated to give two FeO solid plus H2O gas.

Now let’s have a look at this reaction. Again, hydrogen gas is being used as the reducing agent. But in this case, the starting oxide is the iron(II, III) oxide. And we get the same product: iron(II) oxide or ferrous oxide. The reaction equation is identical to the previous one except that we have a different starting oxide, and so the balancing is different. So these are three ways to prepare iron(II) oxide.

Now let’s have a look at some reactions of iron(II) oxide. When iron(II) oxide is in the presence of oxygen gas and the reaction system is heated, Fe2O3 or iron(III) oxide is produced. The reaction equation is four ferrous oxides reacting with oxygen under heat to give two ferric oxides. Oxygen causes the oxidation of iron from a plus two oxidation state to a plus three oxidation state. Iron(II) oxide can react with dilute mineral acids according to the following general equation: metal oxide plus dilute acid gives salt and water. For example, iron(II) oxide can react with dilute sulfuric acid. The corresponding sulfate salt is produced as well as liquid water.

Now let’s go on to the next oxide, Fe2O3. To prepare iron(III) oxide, an iron(III) plus salt, such as iron(III) chloride, in solution is reacted with a hydroxide, such as ammonium hydroxide, also known as ammonia solution. Iron(III) hydroxide comes out of solution as a precipitate. If this precipitate is then heated, it decomposes to form the desired product and water in the gaseous form or water vapor.

Another way to prepare iron(III) oxide is to heat iron(II) sulfate, which decomposes to give ferric oxide as well as two oxides of sulfur, sulfur dioxide gas and sulfur trioxide gas. We’ll have a look at one quick reaction of iron(III) oxide, namely, its reaction with hot concentrated mineral acid. In this example, let’s take sulfuric acid. The iron(III) oxide and sulfuric acid react together under hot conditions and produce the salt iron(III) sulfate as well as water. The reaction equation is Fe2O3 solid plus three H2SO4 liquid being heated to give Fe2(SO4)3 aqueous plus three H2O gas.

We’ve looked at many equations so far for the preparation and reaction of the first two oxides of iron, namely, iron(II) oxide and iron(III) oxide. Let’s have a look at just three more equations, and those are for the last oxide, the iron(II, III) oxide. Besides extracting it from its ore, how can we prepare the iron(II, III) oxide?

We can synthesize this oxide by reacting Fe2O3 or iron(III) oxide with carbon monoxide gas. The reaction system needs to be heated to up to nearly 300 degrees Celsius. Gaseous carbon dioxide is produced as well as our desired iron oxide. You might be wondering if this oxide could not more easily be produced by just reacting iron with oxygen from the air. And yes it can, but we will not look at that reaction equation here.

The last two equations that we will look at in this video are the reactions of the iron(II, III) oxide. The first one is its reaction with hot concentrated sulfuric acid and the last reaction with oxygen. Here is the reaction equation with hot concentrated sulfuric acid. I won’t read the whole reaction equation out again this time. But notice that there are two different products containing iron. And the iron ions in these two salts both have different oxidation states and both contain sulfate ions because sulfuric acid was one of the reactants.

Let’s look at the last reaction. When the iron(II, III) oxide reacts with oxygen and the system is heated, ferric oxide is formed. Notice that the iron ions that were in the three plus state remain in the three plus state. But oxygen gas causes the oxidation of those iron ions in the two plus state, two or three plus state. So all the iron ends up in a three plus state.

Now it’s time to sum up everything we have learnt. We learnt about some of the physical properties of pure iron, for example, its relative softness, its magnetic abilities, and its high melting point. We also learnt about three oxides of iron — the iron(II) oxide, the iron(III) oxide, and the iron(II, III) three oxide — and that the last two are found in important ores. We looked at many reactions, too many to list here, but saw that many of the reactions of iron and its oxides require heat energy to occur and that iron and its oxides react with acids to produce iron salts.

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