Chemical Links

A chemical bond is the force that holds the atoms together in the compounds . These forces are of electromagnetic type and can be of different types and values. The energy needed to break a bond is known as bonding energy .

When a bond occurs, the atoms do not change. For example, when forming water (H 2 O), hydrogens H are still hydrogens and oxygen O is always oxygen. It is the electrons of the hydrogens that are shared with the oxygen.

How are chemical bonds formed?

Nature always tends to reach the state of least energy. Noble gases are the elements that have their full valence electron energy layer; therefore, these elements are very stable and not very reactive. Thus, the tendency of the elements to have a complete valence energy layer is the force that promotes the formation of chemical bonds.

The elements can accept, cede or share electrons in such a way that their last energy layer has 8 electrons. This is known as the octet rule .

Example 1

The electronic configuration of potassium is:

style size 14px normal K space 1 normal s squared space 2 normal s raised to 2 space high end 2 normal p raised to 6 space 3 normal s squared space 3 normal p raised to 6 space raised end 4 normal s raised to 1 end style

According to the octet rule, potassium becomes more stable if the 4s 1 electron gives way to the last level, being as follows:

style size 14px normal K raised to more space 1 normal s squared space 2 normal s raised to 2 space high end 2 normal p raised to 6 space 3 normal s squared space 3 normal p raised to 6 space high end purpose style

On the other hand, the electronic configuration of chlorine is:

style size 14px Cl space 1 normal s squared space 2 normal s squared space 2 normal p raised to 6 space 3 normal s squared space 3 normal p raised to 5 end style

To reach the octet, it is easier if the chlorine accepts an electron, so the configuration becomes:

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Then, if they are placed in the same vessel K and Cl, the electron of K is transferred to Cl and a chemical bond is formed between K and Cl, making this compound more stable than when they are separated.

Example 2

In the case of two atoms with different affinities for electrons it is easy to understand why they join. What happens between two identical atoms? Let’s see the example of molecular oxygen O 2.

The electronic configuration of oxygen is:

style size 14px normal O space 1 s squared space 2 normal s squared space 2 normal p raised to 4 end style

Then, oxygen has 6 electrons (2s 2 2p 4 ) in its outer layer. Two atoms of oxygen have the same attraction for electrons, but when they come together they share two pairs of electrons, each with eight electrons in its outer shell. This results in the oxygen molecule O 2 being more stable than two separate oxygen atoms.

Function of the electron in the chemical bond

The main actor in the links between atoms is the electron . Recalling the structure of the atom, each atom has the same number of negatively charged electrons and positively charged protons. This gives the atom a neutral charge. However, electrons have the ability to move between atoms under certain conditions.

When an atom loses or gains an electron, it acquires an electrical charge and transforms into an ion . An atom that gave up its electron now has a positive charge and is called a cation . On the contrary, when it takes an electron, it has a negative charge and is called an anion .

Types of chemical bonds

Depending on the electronic configuration of atoms and their affinity for electrons, we have different types of links, which we will describe below.

Ionic bond

Sodium and chlorine are a classic example of the formation of ionic bonds.

An ionic bond is formed when there is electron transfer between a metal and a non-metal. For example, sodium (Na) is a metal whose outer layer has an electron. This can be easily transferred and remain as Na + cation . In contrast, chlorine (Cl), has seven electrons in its outer layer, which is why it has a greater predisposition to attract an electron and be left with eight electrons, which transforms it into the chloride anion Cl  .

If sodium and chlorine are combined in aqueous solution, their opposite charges are attracted by electrostatic forces. The compounds formed in this way are arranged in crystals.

General characteristics of ionic crystals

In ionic crystals, the cations attract several anions and, in turn, the anions join several cations.
  • In the ionic bonds, a cation and an anion participate.
  • On a macroscopic scale, ionic compounds form crystalline solids.
  • In general, they have high melting points due to the strong electrostatic and multidirectional attraction between ions of opposite sign. That is, a cation can be attached to several anions at the same time. The same goes for anions.
  • They fracture when subjected to an external force by the formation of ionic repulsion planes.
  • They do not conduct electricity in solid state.
  • They conduct electricity when they are melted, due to the presence of mobile ions.
  • They conduct electricity when they are dissociated in solution.

Examples of ionic compounds

Many of the ionic compounds are precious stones such as fluorite or calcium fluoride CaF 2 . Calcium chloride CaCl 2 is an ionic compound used mainly to prevent ice formation and as a dehumidifier. Magnesium bromide MgBr 2 is used as an accelerator of chemical reactions.

Carbon C shares its four valence electrons with four hydrogens H and forms methane CH 4 .

A covalent bond is established between two atoms when they share electrons. The electrons are not fixed, they move between the two atoms depending on the electronegativity of each atom, that is, the attraction of electrons in atoms.

Polar covalent bond

Electrons that share oxygen and hydrogen are more attracted to oxygen.

When substances with different capacity to attract electrons form a covalent bond, they are said to be polar. For example: in the hydrogen sulfide molecule HS, sulfur S is more electronegative than hydrogen, therefore, the electrons they share will be closer to the sulfur.

Another example of polar covalent bonding is found in the bond between carbon and CF fluorine. Both share electrons, but because fluorine attracts more electrons, they create an electric dipole in which the fluoride side is more negative and the carbon side more positive.

In the formation of a polar covalent bond we do not speak of anions or cations; the atom with the highest electronegativity remains with a negative partial electric charge:

style size 14px bold delta raised to bold minus end style

The atom with the lowest electronegativity remains with a positive partial charge:

style size 14px bold delta raised to bold more end style

Non-polar covalent bond

Formation of a non-polar covalent bond between two chlorine atoms (Cl).

When substances with a similar capacity to attract electrons form a bond, it is said that it is non-polar, since electrons are shared equally among atoms.

For example: the union between carbons in the molecule of ethane C 2 H 6 is nonpolar, because between the two carbons the attraction by electrons is equal.

Depending on the number of electrons that are shared, you can have a single, double or triple covalent bond. Next, we explain each one.

Simple covalent bond

A simple covalent bond occurs when only a pair of electrons are shared. It is represented as a line between two atoms. For example, the oxygen molecule:

style size 14px normal O subscript 2 arrow double right normal OR less normal OR end style

Double covalent bond

The covalent bond double this type of covalent bond, there are four electrons shared between atoms. They are represented by two parallel lines between the two atoms. This union is stronger than the simple covalent bond. For example, ethene:

style size 14px normal C subscript 2 normal H subscript 4 arrow double right normal H subscript 2 normal C equal CH subscript 2 end style

Triple covalent bond

A triple bond means that you are sharing six electrons between two atoms. It is represented by three parallel lines between the elements. For example, the nitrogen molecule:

style size 14px normal N subscript 2 arrow double right normal N identical normal N end style

Characteristics of covalent compounds

  • The electrons are shared between two or more atoms. These joints frequently occur between similar elements or between non-metals.
  • They can form molecules, unlike ionic crystals.
  • The molecules formed are neutral.
  • They can not conduct electricity.
  • When dissolved they do not produce charged particles.
  • When the molecules of these substances are held together by weak intermolecular forces, they have low melting points, so they are gases or liquids at room temperature.
  • Covalent solids with multidirectional forces have high melting points (diamond, graphite, silica) and form reticular or periodic solids.

The majority of organic compounds, where carbon is a primordial element, are characterized by the presence of covalent bonds. Molecules such as adrenaline, methane CH 4 and glucose C6H 12 O 6 are formed by covalent bonds.

Carbon monoxide CO, a toxic gas, is also a covalent compound.

Key concepts to remember in chemical bonds

Electronegativity : the ability of an atom to attract valence electrons.

Valence electrons: the electrons that can form the bond are the valence electrons. These are the electrons that are in the outermost layer of energy of an atom.

Outstanding scientists in chemical bonds

In 1858, the German chemist Friedrich August Kekulé (1829-1896) was the first to define the ability of the atom of an element to join atoms of other elements. So he predicted that carbon was tetravalent, which meant it could join four other atoms.

Another German chemist, Richard Abegg (1869-1910) discovered that noble gases (which do not bind to other atoms) possessed 8 valence electrons. He then suggested that the atoms join to reach the configuration of the noble gases, that is, with 8 electrons in their outer layer.

For its part, the American chemist Gilbert Newton Lewis (1875-1946) discovered that, in covalent bonds, electrons are shared between atoms. It is also attributed to Lewis how to represent electrons with points around the chemical symbol of the atom.

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