Metals and Non-Metals

Chromium, a metal and sulfur, a non-metal

The periodic table can be broadly segregated into two types of elements, commonly referred to as metals and non-metals. Each of these elements have varying properties and can be found in a wide variety of places, such as bridges, buildings, roads, electric cables, cars, aircraft, mobile phones, and laptops, as well as in the oxygen we breathe and the carbon dioxide we exhale.

Over the course of this wiki, you will come to understand the physical and chemical properties of both metals and non-metals. To get the most out of this wiki please read balancing chemical reactions, chemical equilibrium, acids and bases.

Let's begin by studying the physical properties of metals.

• Electrical Conductivity: Almost all metals are good conductors of electricity, though they differ in their conducting power. Silver, for instance, is the best conductor of electricity, and copper \(\ce{(Cu)}\) and aluminium \(\ce{(Al)}\) are also good conductors of electricity. However, silver \(\ce{(Ag)}\) is highly expensive and is not used for this purpose. Instead, copper and aluminium are used. Metals conduct in a solid state.

The conductivity of metal is actually due to the electrons present in the valence shell of metal atoms which they can easily release.

Mercury \(\ce{(Hg)}\), on the other hand, is a poor conductor of electricity. In addition, lead \(\ce{(Pb)}\) is almost non-conducting. We can easily check whether a metal is a good conductor of electricity. In the case of steel, connect a steel spoon to a battery and a light bulb. If the light bulb starts glowing, then we can conclude that steel conducts electricity.

• Thermal conductivity: Metals are good conductors of heat as well, which can also be easily proven. To do this, take a thin metal, a clamp, a stand, and a burner. Set up the apparatus as shown below. Heat one end of the metal sheet for a few minutes, then touch the other end of the sheet. We can observe that the end which had not been heated by the burner is also quite hot. This activity proves that metals, generally, are good conductors of heat.

Heating Iron: Thermal Conductivity[4]

Just like electrical conductivity, metals also differ in their thermal conductivity. Silver and copper are good thermal conductors, while lead and mercury are poor thermal conductors.

• Malleability: Malleability is the property by which a metal can be beaten into sheets by a hammer or similar device. All of us have heard of and seen aluminium foil, which is the metal aluminium beaten into very fine, thin sheets.

A thin Sheet of gold: Malleability[5]

• Ductility: Ductility is the property by which a metal is drawn into thin wires. When a metal can be drawn into wire, we say the metal is ductile. Copper and aluminium are some of the best ductile metals.

• Hardness: Metals are usually hard and difficult to break, though there are exceptions. Sodium \(\ce{(Na)}\) and magnesium \(\ce{(Mg)}\) are so soft they can easily be cut with a knife. The hardness of a metal depends upon the strength of the bonds present among its atoms. Such bonds are called metallic bonds. Since the strength of the bonds differ from metal to metal, their hardness varies accordingly.

• Lustre: Lustre refers to a metal's shine or brightness. Metals like silver, gold, and platinum are well known for their lustre. This property is the reason why jewelry and ornaments are often made of these metals. The lustre is present only when the metals are fresh, and once they are exposed to air, they start to lose their shine and their surface becomes dull. These metals also lose their shine when they come in contact with water vapor.

Lusture of Pyrite: Lustre[6]

• Sonorosity: Metals are generally sonorous, which means they produce a ringing sound when struck.

Though the physical properties help us classify and differentiate metals from non-metals, we need better methods to classify them, and we do that by using their chemical properties. These properties work well with almost all elements and do not seem to have many exceptions. Let us explore a few of these properties.

Reaction with Oxygen: Most metals combine with oxygen in the following manner:

\[\ce{Metal} \ + \ \ce{Oxygen} \ \to \ \text{Metal Oxide}\]

We shall look at some examples of this kind of reaction.

\[\ce{4Na}\ + \ \ce{O}_2\ \to \ce{2Na_2O}\]

\(\ce{Note:}\) Potassium \(\ce{(K)}\) also reacts with oxygen in the same manner to form potassium oxide \(\ce{(K_2O)}\)

Both of these metals react quite violently because they are highly reactive elements and occupy top positions in the reactivity series.

If we look at magnesium, it occupies a lower position in the reactivity series, hence, it doesn't react very vigorously. It only reacts with oxygen upon heating. It burns with a bright flame to form magnesium oxide \(\ce{(Mg_2O)}\). Here is the reaction:

\[\ce{2Mg}\ + \ \ce{O_2}\ \mathrel{\mathop{\longrightarrow}^{\mathrm{heat}}} \ce{2MgO}\]

Reaction with water: The reactivity of metals with water is also linked with their positions in the reactivity series. As a result of reaction of a metal with water, metal hydroxide is formed along with the release of hydrogen gas.

The reactions take place in the following manner:

\[\ce{Metal} \ + \ce{Water} \ \to \text{Metal Hydroxide} + \ce{Hydrogen}\]

Let's look at some examples.

\[\ce{2K} \ + \ce{2H_2O}\ \to \ce{2KOH}\ + \ce{2H_2}\]

\[\ce{2Na} \ + \ce{2H_2O}\ \to \ce{2NaOH}\ + \ce{2H_2}\]

Note: There exists an intermediate form when the metal forms a metal oxide first. However, when the metal oxide is further made to react with water, the product formed will be a metal hydroxide.

Reaction with Acids: Generally, metals react with acids to produce hydrogen in the following manner.

\[\ce{Metal} \ + \ \ce{Acid}\ \rightarrow\ \ce{salt} \ + \ \ce{Hydrogen}\]

However, not all metals can displace hydrogen from acid to form a salt. This is because metals which are lower than hydrogen in the Reactivity Series cannot replace it. As a result, they are unable to follow the above reaction. But for now, let's see some metals which do react with acids.

\[\ce{2Na} \ + \ \ce{2HCl} \ \rightarrow \ \ce{2NaCl} \ + \ \ce{H2}\]

Here, as \(\ce{Na}\) is well above hydrogen in the activity series, it easily replaces hydrogen to form \(\ce{NaCl}\). Let's have a look at some other examples:

\[\begin{align} \ce{Mg} \ + \ \ce{H2SO4} &\rightarrow \ \ce{MgSO4} \ + \ \ce{H2}\\ \ce{Ca} \ + \ \ce{2HCl} &\rightarrow \ \ce{CaCl2} \ + \ \ce{H2}\\ \end{align}\]

However, metals like copper, mercury, and silver cannot replace hydrogen, because they occur well below hydrogen in the reactivity series.

Reaction with Solutions of other Metals: Again, these reactions involve the reactivity series, where a metal which is higher in the series can displace a metal which is lower. Here's the general formula:

\[\text{Metal A} \ + \ \text{Salt Solution of Metal B} \rightarrow \text{Salt Solution of Metal A} \ + \ \text{Metal B}\]

provided that the reactivity of Metal A is greater that the reactivity of Metal B. These reactions help us in determining the reactivity of an element, as the displaced element will be proven to be weaker in terms of reactivity than the element which displaced it. Let us see some examples now:

One very common reaction of this kind is the \(\ce{Fe-Cu}\) reaction.

\[\ce{Fe} \ + \ce{CuSO4} \rightarrow \ce{FeSO4} \ + \ \ce{Cu}\]

In each of these examples, the reactions follow according to the Reactivity Series. And in each reaction, the color of the solution changes as the reaction takes place, here are a few more:

\[\begin{align} \ce{Zn} \ + \ce{CuSO4} &\rightarrow \ce{ZnSO4} \ + \ \ce{Cu}\\ \ce{Cu} \ + \ce{2AgNO3} &\rightarrow \ce{Cu(NO3)2} \ + \ \ce{2Ag}\\ \end{align}\]

Now, let us glance through some physical properties of non-metals. Physical properties alone are not good for distinguishing non-metals due to exceptions in almost every physical property. Later, we shall study in more detail chemical properties, which is the best way to differentiate between metals and non-metals.

• State of existence: Non-metals usually exist in the three states of matter. However, most of them exist in gaseous form. Non-metals like nitrogen, oxygen, carbon dioxide, argon, neon, helium, krypton, chlorine, and fluorine are the ones which constitute the air in our surroundings.

• Hardness: Out of all the non-metals, only solids are expected to be hard. Sulphur and phosphorus are quite soft, but diamond is very hard. Diamond is also probably the hardest substance presently known.

Diamond: An exception for Hardness [7]

• Lustre: Non-metals usually have no shine since they have no loosely attached electrons which are responsible for lustre. However, there are exceptions. Out of the non-metals, Diamond and iodine have lustrous nature.

Lustrous Iodine: An exception for Lustre[8]

• Electrical and Thermal Conductivity: Non-metals, in general, are quite poor conductors of heat and electricity. Graphite is an exception here.In fact, it is a very good conductor of electricity.

• Malleability: Non-metals cannot be beaten into thin sheets in the way that metals can. There is a weak force of attraction. As a result, they are quite brittle. Sulphur is a brittle element. If it is hammered, it would break into pieces.

Sulfur powder[3]

• Ductility: Non-metals cannot be drawn into thin wires. They would break, whether being beaten into thin sheets or drawn into wires.

• Sonorosity: Non-metals are not sonorous. They produce no ringing sound when struck on their surface.

Again, to classify metals and non-metals in a more orderly fashion, we deal with their chemical properties. The chemical properties work well with almost all elements and do not seem to have many exceptions. We will discuss some of the main chemical properties of non-metals in this section.

Reaction with Oxygen: In the presence of oxygen, when non-metals are heated they react and form non-metal oxides. These oxides are generally acidic in nature. In general, they react in the following manner:

\[\text{Non-Metal} \ + \ \text{Oxygen} \mathrel{\mathop{\longrightarrow}^{\mathrm{heat}}} \text{Non-Metal Oxide}\]

Here are some examples, one of the most common one being:

\[\ce C \ + \ \ce{O2} \mathrel{\mathop{\longrightarrow}^{\mathrm{heat}}} \text{CO2}\]

Some other examples are:

\[\begin{align} \ce{S8} \ + \ \ce{8O2} &\mathrel{\mathop{\longrightarrow}^{\mathrm{heat}}} \text{8SO2}\\ \ce{P4} \ + \ \ce{5O2} &\mathrel{\mathop{\longrightarrow}^{\mathrm{heat}}} \text{P4O10}\\ \end{align}\]

Reaction with Acids: Non-metals, being electronegative, do not react with dilute acids as they cannot displace hydrogen easily. However, they react with concentrated acids when they are heated. Here's the general formula:

\[\text{Non-Metal} \ + \ \ce{Acid} \mathrel{\mathop{\longrightarrow}^{\mathrm{heat}}} \text{Salt} \ + \ \text{Oxide} \ + \ \text{Water}\]

However, the salt may not necessarily be formed in all the cases.

The following is one of the classic examples in which the salt is not formed:

\[\ce{S} \ + \ \ce{2H2SO4} \mathrel{\mathop{\longrightarrow}^{\mathrm{heat}}} \ce{3SO2} \ + \ \ce{2H2O}\]

Other common examples being:

\[\begin{align} \ce{P} \ + \ \ce{5HNO3} &\mathrel{\mathop{\longrightarrow}^{\mathrm{heat}}} \ce{H3PO4} \ + \ \ce{5NO2} \ + \ \text{2H2O}\\ \ce{2P} \ + \ \ce{5H2SO4} &\mathrel{\mathop{\longrightarrow}^{\mathrm{heat}}} \ce{2H3PO4} \ + \ \ce{5SO2} \ + \ \ce{2H2O} \end{align}\]

Reaction with Chlorine: Non-metals react with chlorine atoms upon heating to form their respective chlorides. Sulphur \(\ce{(S)}\) and phosphorus \(\ce{(P)}\) react in this way. Let's have a look at some reactions to have a better idea of this.

\[\text{Non-Metal} \ + \ \ce{Chlorine} \mathrel{\mathop{\longrightarrow}^{\mathrm{heat}}} \text{Metal Chloride}\]

\[\begin{align} \ce{2P}\ + \ce{5Cl_2}\ &\mathrel{\mathop{\longrightarrow}^{\mathrm{heat}}} \ce{2PCl_5}\\ \ce{2S}\ + \ce{Cl_2}\ &\mathrel{\mathop{\longrightarrow}^{\mathrm{heat}}} \ce{S_2Cl_2} \end{align}\]

• Reaction with hydrogen: Non-metals form their respective hydrides when they react with hydrogen. These hydrides are non-electrolytic and exist as gases at room temperature. This is the general formula:

\[\text{Non-Metal} \ + \ \ce{Hydrogen} \mathrel{\mathop{\longrightarrow}^{\mathrm{heat}}} \text{Metal Hydride}\]

Hydrogen combines with sulphur at \(713K\) to form hydrogen sulphide in the following manner:

\[\ce{S}\ + \ce{H2}\ \mathrel{\mathop{\longrightarrow}^{\mathrm{heat}}} \ce{H2S}\]

Another common reaction is the formation of ammonia, which occurs at \(773 K\) in the presence of an iron catalyst and it is a reversible reaction:

\[\ce{N2}\ + \ce{3H2}\ \rightleftharpoons \ce{2NH3}\]

Reaction with Salt Solutions: This type of reaction is similar to the one we discussed in the chemical properties of metals. Here again, the reactions involve the reactivity series. A non-metal which is higher in the series can displace a non-metal from it's salt which is lower than it. Here's the general formula:

\[\text{Non-Metal A} \ + \ \text{Salt solution of Non-Metal B} \longrightarrow \text{Non-Metal B} \ + \ \text{Salt Solution of Non-Metal A}\]

provided that the reactivity of Metal A is greater than the reactivity of Metal B. Let's have a look at one very common example:

\[\ce{Cl2}\ + \ce{2NaBr}\ \longrightarrow \ce{2NaCl} \ + \ \ce{Br2}\]

Metals and non-metals do not have stable electronic configurations like the noble gas elements do. Hence, they will have a strong tendency to achieve the noble gas configuration.

Metal atoms have surplus electrons to give away in their valence shell. On the other hand, non-metals need electrons to fill their last valence shell.

From this, we can say that atoms of metals lose valence electrons while non-metal atoms accept electrons. When the metal atoms give away their valence electrons, they become positively charged or cations, while the non-metal atoms, when they lose electrons, become negatively charged, or in other words, anions. The opposing type of charges bring the ions closer due to electrostatic force of attraction between them, resulting in the formation of ionic compounds with an ionic bond between them.

Let's look at a few examples.

The formation of common salt or table salt \(\ce{(NaCl)}\) is shown below. The electron from the sodium atom gets transferred to the chlorine atom. Hence, an ionic bond is formed between them.

\[\ce{Na}\cdot \ + \ \cdot\ce{Cl} \rightarrow \ce{Na-Cl}\]

Here, the chlorine atom changes to \(\ce{Cl^-}\) and gets the electronic configuration \(\ce{2, 8, 8}\) of argon \(\ce{(Ar)}\). Sodium, as it changes to \(\ce{Na^+}\) ion, gets the electronic configuration \(\ce{2, 8}\) of neon \(\ce{(Ne)}\).

When they react, they will form \(\ce{NaCl}\)

\[\ce{Na^+}\ + \ce{Cl^-}\ \to \ce{NaCl}\]

Hence, the ionic compound sodium chloride \(\ce{(NaCl)}\) is formed.

• Atoms, Molecules, Elements, Compounds

• Ionic Compounds

• Acids and Bases

• Balancing Chemical Reactions

• Chemical Equilibrium

• Reactivity Series

[1] MTG Publications, Foundation Course for Chemistry Revised 2014 edition, MTG: Indian Subcontinent adaptation 2014.

[2] Image from https://en.m.wikipedia.org/wiki/Chromium#/media/File%3AChromiumcrystalsand1cm3cube.jpg under the creative commons attribution for reuse and modification.

[3] Image from https://en.m.wikipedia.org/wiki/Sulfur#/media/File%3ASulfur-sample.jpg under the creative commons attribution for reuse and modification.

[4] Image from https://en.m.wikipedia.org/wiki/Metal#/media/File%3AHot_metalwork.jpg under the creative Commons license for reuse and modification.

[5] Image from https://en.m.wikipedia.org/wiki/Ductility#/media/File%3AKanazawaGoldFactory.jpg under the creative Commons attribution for reuse and modification.

[6] Image from https://en.m.wikipedia.org/wiki/Lustre(mineralogy)#/media/File%3APyrite3.jpg under the creative Commons attribution for reuse and modification.

[7] Image from https://en.m.wikipedia.org/wiki/Diamond#/media/File%3ATheHopeDiamond-SIA.jpg under the creative Commons license for reuse and modification.

[8] Image from https://en.m.wikipedia.org/wiki/Nonmetal#/media/File%3AIodinecrystals.JPG under the creative Commons attribution for reuse and modification.