What happens to the electrons in an ionic bond?

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Ionic bonding

Ionic bonding is a type of chemical bonding that involves the electrostatic attraction between oppositely charged ions, or between two atoms with sharply different electronegativities, and is the primary interaction occurring in ionic compounds. It is one of the main types of bonding along with covalent bonding and metallic bonding. Ions are atoms (or groups of atoms) with an electrostatic charge.

Atoms that gain electrons make negatively charged ions (called anions). Atoms that lose electrons make positively charged ions (called cations). This transfer of electrons is known as electrovalence in contrast to covalence. In the simplest case, the cation is a metal atom and the anion is a nonmetal atom, but these ions can be of a more complex nature,

e.g. molecular ions like NH⁺
₄ or SO²⁻
₄. In simpler words, an ionic bond results from the transfer of electrons from a metal to a non-metal in order to obtain a full valence shell for both atoms.

It is important to recognize that clean ionic bonding — in which one atom or molecule completely transfers an electron to another — cannot exist: all ionic compounds have some degree of covalent bonding, or electron sharing.

Thus, the term “ionic bonding” is given when the ionic character is greater than the covalent character – that is, a bond in which a large electronegativity difference exists between the two atoms, causing the bonding to be more polar (ionic) than in covalent bonding where electrons are shared more equally. Bonds with partially ionic and partially covalent character are called polar covalent bonds.

Ionic compounds conduct electricity when molten or in solution, typically not when solid. Ionic compounds generally have a high melting point, depending on the charge of the ions they consist of. The higher the charges the stronger the cohesive forces and the higher the melting point. They also tend to be soluble in water; the stronger the cohesive forces, the lower the solubility.

Electron Transfer: Ionic Bonds

In the section “Lewis Electron Dot Diagrams”, we saw how ions are formed by losing electrons to make cations or by gaining electrons to form anions. The astute reader may have noticed something: Many of the ions that form have eight electrons in their valence shell. Either atoms gain enough electrons to have eight electrons in the valence shell and become the appropriately charged anion, or they lose the electrons in their original valence shell.

The lower shell, now the valence shell, has eight electrons in it, so the atom becomes positively charged. For whatever reason, having eight electrons in a valence shell is a particularly energetically stable arrangement of electrons. The trend that atoms like to have eight electrons in their valence shell is called the octet rule.

When atoms form compounds, the octet rule is not always satisfied for all atoms at all times, but it is a very good rule of thumb for understanding the kinds of bonding arrangements that atoms can make.

It is not impossible to violate the octet rule. Consider sodium: in its elemental form, it has one valence electron and is stable. It is rather reactive, however, and does not require a lot of energy to remove that electron to make the Na⁺ ion. We could remove another electron by adding even more energy to the ion, to make the Na²⁺ ion.

However, that requires much more energy than is normally available in chemical reactions, so sodium stops at a 1+ charge after losing a single electron. It turns out that the Na⁺ ion has a complete octet in its new valence shell, the n = 2 shell, which satisfies the octet rule. The octet rule is a result of trends in energies and is useful in explaining why atoms form the ions that they do.

Now consider an Na atom in the presence of a Cl atom. The two atoms have these Lewis electron dot diagrams and electron configurations:

 

 

For the Na atom to obtain an octet, it must lose an electron; for the Cl atom to gain an octet, it must gain an electron. An electron transfers from the Na atom to the Cl atom:

Resulting in two ions — the Na⁺ ion and the Cl⁻ ion:

Both species now have complete octets, and the electron shells are energetically stable. From basic physics, we know that opposite charges attract. This is what happens to the Na⁺ and Cl⁻ ions:

Where we have written the final formula (the formula for sodium chloride) as per the convention for ionic compounds, without listing the charges explicitly. The attraction between oppositely charged ions is called an ionic bond, and it is one of the main types of chemical bonds in chemistry. Ionic bonds are caused by electrons transferring from one atom to another.

In electron transfer, the number of electrons lost must equal the number of electrons gained. We saw this in the formation of NaCl. A similar process occurs between Mg atoms and O atoms, except in this case two electrons are transferred:

The two ions each have octets as their valence shell, and the two oppositely charged particles attract, making an ionic bond:

Remember, in the final formula for the ionic compound, we do not write the charges on the ions.

What about when an Na atom interacts with an O atom? The O atom needs two electrons to complete its valence octet, but the Na atom supplies only one electron:

The O atom still does not have an octet of electrons. What we need is a second Na atom to donate a second electron to the O atom:

These three ions attract each other to give an overall neutral-charged ionic compound, which we write as Na₂O. The need for the number of electrons lost being equal to the number of electrons gained explains why ionic compounds have the ratio of cations to anions that they do. This is required by the law of conservation of matter as well.

What happens to electrons when an ionic bond forms?

 

 

Lewis Dot Structures:

Lewis dot structures are a shorthand to represent the valence electrons of an atom. The structures are written as the element symbol surrounded by dots that represent the valence electrons. The Lewis structures for the elements in the first two periods of the periodic table are shown below.

lewis_H Lewis Dot Structures lewis_He

lewis_Li

lewis_Be lewis_B lewis_C lewis structure-nitrogen lewis_O lewis_F lewis_Ne Lewis structures can also be used to show bonding between atoms. The bonding electrons are placed between the atoms and can be represented by a pair of dots or a dash (each dash represents one pair of electrons, or one bond). Lewis structures for H2 and O2 are shown below.

H2 H:H or H-H

O2 lewis structure – oxygen3lewis structure – oxygen3 lewis structure – oxygen2

Quartz, the Most Common Type of Ionic Bonding

Now, one of the most common types of ionic bonding in nature links silicon, which is element 14, and oxygen, element 8. On the periodic table, you’ll see that silicon, element 14, then wants to lose 4 electrons, become a +4  ion, while oxygen, element 8, wants to gain two electrons and become a -2 ion, forming a strong ionic bond between one silicon and two oxygens which is SiO₂.

This ionic Si-O bond is particularly strong because it links a +4 ion with a -2 ion. Now, remember Coulomb’s law which describes the force, the electrostatic force between two charges.

The force is proportional to the product of the charge in the first object and the charge in the second object. In silicon-oxygen bonding, the relative strength is 4 times 2, or 8. So, for a given separation of ions, the silicon-oxygen bond is eight times stronger, making a very strong and resistant material.

Strong Si-O bonds are extremely common in nature. They form, for example, the mineral quartz, SiO₂, which forms beautiful crystals. These are the power crystals you may have heard about. But quartz forms beach sand, and sand grains are typically grains of quartz. Si-O bonds also form window glass.

 

 

 

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