How To Draw A Ionic Bond

Learning Objectives

• State the octet dominion.
• Ascertain ionic bond.
• Draw Lewis structures for ionic compounds.

In Section four.7, nosotros demonstrated that ions are formed by losing electrons to make cations, or by gaining electrons to grade anions. The astute reader may have noticed something: many of the ions that grade have eight electrons in their valence shell. Either atoms proceeds plenty electrons to have eight electrons in the valence shell and become the accordingly charged anion, or they lose the electrons in their original valence trounce; the lower crush, now the valence shell, has 8 electrons in information technology, so the atom becomes positively charged. For whatever reason, having viii electrons in a valence shell is a particularly energetically stable arrangement of electrons. Theoctet ruleexplains the favorable trend of atoms having 8 electrons in their valence shell. When atoms class compounds, the octet rule is non always satisfied for all atoms at all times, but information technology is a very good dominion of thumb for understanding the kinds of bonding arrangements that atoms tin can brand.

It is not impossible to violate the octet rule. Consider sodium: in its elemental form, information technology has ane 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 calculation even more than energy to the ion, to make the Na ²⁺ ion. Nevertheless, that requires much more energy than is unremarkably available in chemical reactions, so sodium stops at a 1+ accuse after losing a single electron. It turns out that the Na ⁺ ion has a complete octet in its new valence shell, the n = 2 trounce, 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 take these Lewis electron dot diagrams and electron configurations:

\[\mathbf{Na\, \cdot }\; \; \; \; \; \; \; \; \; \; \mathbf{\cdot }\mathbf{\ddot{\underset{.\: .}Cl}}\mathbf{\: :}\]

\[\left [ Ne \correct ]3s^{1}\; \; \; \; \left [ Ne \right ]3s^{2}3p^{5}\]

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 cantlet:

\[\mathbf{Na\, \cdot }\curvearrowright \mathbf{\cdot }\mathbf{\ddot{\underset{.\: .}Cl}}\mathbf{\: :}\]

resulting in 2 ions—the Na ⁺ ion and the Cl ⁻ ion:

\[\mathbf{Na}^{+}\; \; \; \; \; \; \; \; \mathbf{:}\mathbf{\ddot{\underset{.\: .}Cl}}\mathbf{\: :}^{-}\]

\[\left [ Ne \right ]\; \; \; \; \; \left [ Ne \right ]3s^{2}3p^{6}\]

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

\[\mathbf{Na}^{+}\; + \; \mathbf{:}\mathbf{\ddot{\underset{.\: .}Cl}}\mathbf{\: :}^{-}\rightarrow Na^{+}Cl^{-}\; \; or\; \; NaCl\]

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

In electron transfer, the number of electrons lost must equal the number of electrons gained. Nosotros 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 2 ions each take octets equally their valence shell, and the ii oppositely charged particles attract, making an ionic bond:

\[\mathbf{Mg\,}^{2+}\; + \; \left[\mathbf{:}\mathbf{\ddot{\underset{.\: .}O}}\mathbf{\: :}\correct]^{2-}\; \; \; \; \; Mg^{2+}O^{2-}\; or\; MgO\]

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

What nearly when an Na cantlet interacts with an O cantlet? The O cantlet needs ii electrons to complete its valence octet, but the Na cantlet supplies only 1 electron:

\[\mathbf{Na\, \cdot }\curvearrowright \mathbf{\cdot }\mathbf{\ddot{\underset{.}O}}\mathbf{\: :}\]

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 iii ions attract each other to give an overall neutral-charged ionic compound, which nosotros write every bit Na ₂ O. 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 thing besides.

Example \(\PageIndex{1}\): Synthesis of Calcium Chloride from Elements

With arrows, illustrate the transfer of electrons to course calcium chloride from \(Ca\) atoms and \(Cl\) atoms.

Solution

A \(Ca\) atom has two valence electrons, while a \(Cl\) atom has seven electrons. A \(Cl\) atom needs only one more to complete its octet, while \(Ca\) atoms accept ii electrons to lose. Thus we need 2 \(Cl\) atoms to accept the 2 electrons from one \(Ca\) atom. The transfer process looks as follows:

The oppositely charged ions attract each other to make CaCl ₂ .

Exercise \(\PageIndex{ane}\)

With arrows, illustrate the transfer of electrons to form potassium sulfide from \(M\) atoms and \(S\) atoms.

Reply

Summary

• The tendency to course species that have eight electrons in the valence shell is chosen the octet rule.
• The attraction of oppositely charged ions caused by electron transfer is called an ionic bond.
• The strength of ionic bonding depends on the magnitude of the charges and the sizes of the ions.

Contributions & Attributions

This page was constructed from content via the following contributor(s) and edited (topically or extensively) by the LibreTexts development team to meet platform style, presentation, and quality:

• Marisa Alviar-Agnew (Sacramento City Higher)

• Henry Agnew (UC Davis)

Source: https://chem.libretexts.org/Courses/College_of_Marin/CHEM_114%3A_Introductory_Chemistry/10%3A_Chemical_Bonding/10.03%3A_Lewis_Structures_of_Ionic_Compounds-_Electrons_Transferred

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